1

Classification of matter

Physical separation Chemical reactions

Depending on structure or their

interatomic arrangement or States of Matter Elements are made of Atoms Chemical Bonding An element is a substance that cannot be separated into simpler substances by a chemical change. When two or more elements combine in a chemical reaction, they form a compound. A technique called electrolysis can help to distinguish between elements and compounds. During electrolysis, an electric current is passed through a substance. If the substance is a compound, it may be broken down into the separate elements that form it. A mixture is a blend of two or more pure substances. A mixture that has visibly different parts is called a heterogeneous mixture. Other mixtures do not contain visibly different parts are called homogeneous mixtures. Heterogeneous mixtures composed of liquids and solids are often separated by filtration . Homogeneous mixtures cannot be separated by filtration, but by distillation, crystallization , and chromatography. ============================================================================================

Atomic Theories

Periodic Trends and Valence Electrons (Electronic Configurations)

Shapes Bonding

Physical Properties Chemical Properties

Solubilities/ Types & Rates Melting/Boiling Points of Reactions Uses and Preparation

From heterogeneous mixtures found in nature, which has two or more phases, homogeneous matter may be obtained

Homogeneous matter is uniform throughout and has only one phase

Solutions are mixtures of at least 2 constituents, homogeneous throughout. The proportion of components can vary from one solution to another. Solutions may be

Pure substances are homogeneous throughout. They must be either elements or compounds

Elements may be grouped into rough categories based on some general chemical and physical properties

Compounds always have elements in the same constant proportion by weight. Compounds can be classified as either

gas liquid solid metals nonmetals ionic molecular

2 Classification Systems for Elements: 1. A system based on the electron configuration of the elements, in which elements are described as

(i) Noble gases – These are elements in the 8th column (Group VIIIA). They are all gases at room temperature, and have little tendency to form chemical compounds. Krypton, xenon, and radon reluctantly form compounds with fluorine; radon is radioactive. Except helium

(1 2s ), the distinguishing electron completes the p shell ( p6 ). They have low melting and boiling points.

(ii) Representative Elements– Except the noble gases; they are all the elements in the s and p areas of the Periodic Table. The distinguishing electron partially or completely fills an s subshell or partially fills a p subshells. Most of the more common elements are Representative Elements.

(iii) Transition Elements – These are all elements in the d are with the distinguishing electron in the d subshell; exhibit a wide range of chemical and physical properties; characteristically strong, hard metals with high melting points; good conductors of electricity; variable reactivity; form ions with variable charges; many react with oxygen to form oxides; some will react with solutions of strong acids to form hydrogen gas.

(iv) Inner-transition Elements – These are all elements in the f area. There is very little variance in the properties of either the 4f or 5f series of inner transition elements. Period 6 – lanthanides – rare earth elements with atomic numbers from 57 to 70. Period 7 – actinides – actinium and the 13 elements that follow it in the 7th row of the table; elements with atomic numbers 89 to 102; includes transuranic elements with atomic number 93 or greater; transuranic elements are synthetic; they follow uranium in the periodic table.

2. A system based on selected physical properties of the elements, in which elements are described as (i) Metals - luster, thermal conductivity, electrical conductivity, and malleability, all (except mercury) are solids at room temperature

( 25°C ), high density, high melting points. The majority elements (87) are metals. Group IA – alkali metals – soft, silver-coloured elements; solid at SATP; exhibit metallic properties; react violently with water to form basic solutions and liberate hydrogen gas; react with halogens to form compounds similar to sodium chloride, NACl s( ) ; stored under oil or in a vacuum to prevent reaction with air.

Group IIA – alkaline earth metals – light, very reactive; solids at SATP; exhibit metallic properties; form oxide coatings when exposed to air; react with oxygen to form oxides with the general chemical formula, MO s( ) ; all except beryllium will react with

hydrogen to form hydrides with the general chemical formula XH2 ; react with water to liberate hydrogen.

(ii) Nonmetals – absence of luster, thermal conductivity, electrical conductivity, and malleability, many are gases (only bromine is liquid at room temperature). Solid nonmetals include carbon, sulfur, and phosphorus. In general, the nonmetals have lower densities and melting points than metals. 22 elements are nonmetals. Group VIIA – halogens – may be solids, liquids, or gases at SATP; exhibit nonmetallic properties – not lustrous and nonconductors of electricity; extremely reactive, with fluorine being the most reactive; react readily with hydrogen and metals. Group VIIIA – noble gas – gases at SATP; low melting and boiling points; extremely unreactive; krypton, xenon, and radon reluctantly form compounds with fluorine; radon is radioactive.

(A) Nuclear Reactions change the composition of an atom’s nucleus – Radioactivity is the spontaneous emission of radiation

from an atom. When an atom emits one of these kinds of radiation, it is said to be undergoing a radioactive decay. Almost all the atoms have stable nuclei. Only a few atoms in nature are radioactive. And many elements have one or more isotopes called radioisotopes or radionuclides that are unstable. Why are some nuclei stable and some are not? Part of the reason has to do with the number of protons and neutrons that they contain. Not all combinations of protons and neutrons make a stable nucleus. In the nucleus, adjacent protons experience both an electrical force repelling them and an strong nucleus force attracting them. Since a neutron does not experience electrical repulsion, it adds a net attractive force to the inside of a nucleus. All stable nuclei follow a distinct pattern. For elements with atomic numbers between 1 and 20 (from hydrogen to calcium, stable nuclei have almost equal numbers of protons of protons and neutrons. Beyond 20 protons, nuclei need increasingly more neutrons than protons to be stable. When the atomic number exceeds 83 (the element bismuth), no number of neutrons is sufficient to glue the nucleus together indefinitely. All nuclei with atomic numbers greater than 83 are radioactive.

Applying the Rule of Binding Energy it is found that isotopes having mass numbers near 56 (iron ) should have the most stable nuclei. However, several isotopes of iron, all synthetic, are unstable. So scientists have developed several “rule of thumb” to aid them in picking those candidates for synthesis that stand some chance of having half-lives long enough to allow the isotope to be detected and studies.

The Band of Stability: This rule uses an array of all known stable and unstable isotopes. Each isotope is given a location in the array

according to its number of protons and neutrons, as shown in Fig. 23.8 element 83 (bismuth) is the last one to have at least one stable

NOTES FOR CHEMISTRY 11/12 – I: Matter and Properties 4

nucleus, so the array is not plotted beyond 83. The two curved lines are drawn to enclose all stable nuclei, and between these lines lies

what is called the band of stability. Not all of the isotopes in this band are stable, but those that are not have half-lives long enough to

enable detection. Any isotope not represented anywhere on the array of Fig. 23.8 quite likely has a half-life at least as short as 810−

second, too short for detection and observation. Therefore, as a rule of thumb, the search for isotopes lying at some distance from

the band of stability will likely be unsuccessful. The isotopes not plotted on the graph, those with atomic numbers above 83, tend to be

alpha emitters. The beta emitters have too many neutrons. The beta decay reduces the neutron-t0-proton (P

N ) ratio, while positron

emission has the opposite effect. Positron emitters are found only among artificial radionuclides, which have too few neutrons for

stability.

Odd-Even Rule: When the numbers of neutrons and protons in a nucleus are both even, the isotope is far more likely to be stable than

when both numbers are odd. Out of the 264 stable isotopes, only five have odd numbers of both protons and neutrons, whereas 157 have

even numbers of both. The rest have an odd number of one and an even number of the other. In Fig. 23.8, the horizontal lines on which

the dark squares most commonly occur correspond to even numbers of neutrons, and the vertical lines on which the dark squares most

often fall correspond to even numbers of protons. With even numbers, the particles are spinning in opposite directions and their spins are

said to be paired resulting in the least amount of energy, and therefore more stable.

Every radioisotope has a characteristic property called its half-life. The half-life of a radioactive substance is the time taken for half of the

original number of radioactive atoms to decay. The half-lives of radioisotopes vary considerably. Carbon-14 has a half-life of 5730 a. It

emits a beta particle as it decays. Small amounts of this isotope occur naturally in the atmosphere, where it reacts with oxygen to form

radioactive carbon dioxide. Carbon-14 is absorbed by plants when a mixture of radioactive and nonradioactive carbon dioxide is taken in

during photosynthesis. Carbon-14 then finds its way into other living organisms through the food chain. When a living organism dies, it

stops taking in material, including carbon. Radioactive decay will gradually reduce the amount of carbon-14 present in its tissues. As a

result, the ratio of radioactive carbon to nonradioactive carbon present in the organism will gradually decrease. Comparing this ratio to

the normal ratio present in a living organism provides a measure of the time elapsed since the organism’s death. The practice of

measuring the carbon-14 to carbon-12 ratio is known as carbon-14 dating.

Mathematically, radioactive decay is represented by 2

1

2

10

T

t

NN

= , where

2

1T is the half-life of the particular isotope.

The number of nuclei decaying per second is called the activity and is measured in becquerels (Bq). Activity can be described by the

equation: 2

1

2

10

T

t

AA

= and Nt

NA λ−=

∆∆= , where λ is called the decay constant of the isotope.

Representative Elements - Metals (1) Alkali Metals The alkali metals are a series of chemical elements forming Group 1 (IUPAC style) of the periodic table: lithium (Li), sodium (Na), potassium (K),

rubidium (Rb), caesium (Cs), and francium (Fr). (Hydrogen, although nominally also a member of Group 1, very rarely exhibits behaviour

comparable to the alkali metals). The alkali metals provide one of the best examples of group trends in properties in the periodic table, with

well characterized homologous behaviour down the group.

The alkali metals (i) are all highly reactive and are never found in elemental form in nature. As a result, in the laboratory they are stored under

mineral oil or paraffin oil. They also (ii) tarnish easily and (iii) have low melting points, and (iv) densities. Potassium and rubidium possess a

weak radioactive characteristic due to the presence of long duration radioactive isotopes.

The alkali metals (v) are shiny, silver-colored (caesium has a golden tinge), (vi) soft metals, which (vii-a) react readily with halogens to form ionic

salts, and (vii-b) with water to form strongly alkaline (basic) hydroxides, (vii-c) with oxygen in air to form oxide. These elements (viii) all have

one electron in their outermost shell, so the energetically preferred state of achieving a filled electron shell is to lose one electron to form a

singly charged positive ion, i.e. cation.

Hydrogen, with a solitary electron, is usually placed at the top of Group 1 of the periodic table, but it is not considered an alkali metal; rather it

exists naturally as a diatomic gas. Removal of its single electron requires considerably more energy than removal of the outer electron for the

alkali metals. As in the halogens, only one additional electron is required to fill in the outermost shell of the hydrogen atom, so hydrogen can in

some circumstances behave like a halogen, forming the negative hydride ion. Binary compounds of hydride with the alkali metals and some

transition metals have been prepared. Under extremely high pressure, such as is found at the core of Jupiter, hydrogen does become metallic

and behaves like an alkali metal.

Alkali metals have the (ix) lowest ionization potentials in their respective periods, as removing the single electron from the outermost shell

gives them the stable inert gas configuration. Their second ionization potentials are very high, as removing an electron from a species having a

noble gas configuration is very difficult.

Alkali metals are famous for their vigorous reactions with water, and these reactions become increasingly violent as one moves down the

group. The reaction with water is as follows: Alkali metal + water → Alkali metal hydroxide + hydrogen gas

With potassium as an example: 2K (s) + 2H2O (l) → 2KOH (aq) + H2 (g)

NOTES FOR CHEMISTRY 11/12 – I: Matter and Properties 5

Trends - Like other groups, the members of this family show patterns in its electron configuration, especially the outermost shells resulting

in trends in chemical behavior:

The alkali metals show a number of trends when moving down the group - for instance: decreasing electronegativity, increasing reactivity, and

decreasing melting and boiling point. Density generally increases, with the notable exception of potassium being less dense than sodium, and

the possible exception of francium being less dense than caesium.

(2) Alkaline Earth Metals

The alkaline earth metals are (i) silver colored, (ii) soft metals, (iii-a) which react readily with halogens to form ionic salts, and (iii-b) with water,

though not as rapidly as the alkali metals, to form strong alkaline (basic) hydroxides. For example, where sodium and potassium react with

water at room temperature, magnesium reacts only with steam and calcium with hot water:

Mg + 2 H2O → Mg(OH)2 + H2

Beryllium is an exception: It does not react with water or steam, and its halides are covalent. All the alkaline earth metals have two electrons

in their valence shell, so the energetically preferred state of achieving a filled electron shell is to lose two electrons to form doubly charged

positive ions.

Just like their names, they don’t differ completely. The main difference is the electron configuration, which is ns2 for alkaline earth metals and

ns1 for alkali metals. For the alkaline earth metals, there are two electrons that are available to for a metallic bond, and the nucleus contains an

additional positive charge. Also, the elements of group 2A (alkaline earth) (iv) have much higher melting points and boiling points compared to

those of group 1A (alkali metals). The alkali also have a softer and more light-weight figure whereas the (v) alkaline earth metals are much

harder and denser.

The second valence electron is very important when it comes to comparing chemical properties of the alkaline earth and the alkali metals. The

second valence electron is in the same “sublevel” as the first valence electron. Therefore, the Zeff is much greater. This means that the elements

of the group 2A contain a smaller atomic radius and much higher ionization energy than the group 1A. Even though the group 2A contains

much higher ionization energy, they still form an ionic compound with 2+ cations. Beryllium, however, behaves differently. This is due to the

fact that in order to remove two electrons from this particular atom, it requires significantly more energy. It never forms Be2+

and its bonds are

polar covalent.

Beryllium - As mentioned earlier, Be is “special”; it behaves differently. If the Be2+

ion did exist, it would polarize electron clouds that are

near it very strongly and would cause extensive orbital overlap, since Be has a high charge density. All compounds that include Be have a

covalent bond. Even the compound BeF2, which is the most ionic Be compound, has a low melting point and a low electrical conductivity when

melted.

Important reactions: Note: E = elements that act as reducing agents 1. The metals reduce halogens to form ionic halides: E(s) + X2 → EX2 (s) where X = F, Cl, Br or I 2. The metals reduce O2 to form the oxides: 2E(s) + O2 → 2EO(s) 3. The larger metals react with water to produce hydrogen gas: E(s) + 2H2O(l) → E2+

(aq) + 2OH-aq + H2 (g) where E = Ca, Sr or Ba

Representative Elements - Non-Metals The elements generally regarded as nonmetals are:

• hydrogen (H) • In Group 14: carbon (C) • In Group 15(the pnictogens): nitrogen (N), phosphorus (P) • Several elements in Group 16, the chalcogens: oxygen (O), sulfur (S), selenium (Se) • All elements in Group 17 - the halogens • All elements in Group 18 - the noble gases

There is no rigorous definition for the term "nonmetal" - it covers a general spectrum of behaviour. Common properties considered

characteristic of a nonmetal include:

• poor conductors of heat and electricity when compared to metals

• they form acidic oxides (whereas metals generally form basic oxides)

• in solid form, they are dull and brittle, rather than metals which are lustrous, ductile or malleable

• usually have lower densities than metals

• they have significantly lower melting points and boiling points than metals

• non-metals have high electronegativity

Only 18 elements in the periodic table are generally considered nonmetals, compared to over 80 metals, but nonmetals make up most of the crust, atmosphere and oceans of the earth. Bulk tissues of living organisms are composed almost entirely of nonmetals. Most nonmetals are monatomic noble gases or form diatomic molecules in their elemental state, unlike metals which (in their elemental state) do not form molecules at all. A handful of nonmetals of low atomic weights make up most of the universe due mostly to their nuclear stability and the fact that they form strong covalent bonds with each other. Moreover, most of them, such as hydrogen, oxygen, carbon, and nitrogen, can form double and triple bonds as well as single bonds. Carbon, oxygen, and nitrogen all have an unusual ability to form double bonds

NOTES FOR CHEMISTRY 11/12 – I: Matter and Properties 6

HHyyddrr ooggeenn -- Hydrogen has very low density and is highly flammable. Isotopes of hydrogen: (1) protium – H, 1 proton, makes up 99.8% of elemental hydrogen; (2) deuterium – D, 1 proton + 1 neutron; (3) tritium – T , 1 proton + 2 neutrons, radioactive and formed in nuclear reactions. HHyyddrriiddeess –– bbiinnaarryy ccoommppoouunnddss bbeettwweeeenn hhyyddrrooggeenn aanndd mmeettaallss oorr ootthheerr nnoonnmmeettaallss aarree ooff tthhrreeee ttyyppeess;;

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tthhee mmeettaall .. TThheeyy oofftteenn bbeehhaavvee aass ddoonnoorrss ooff ggaasseeoouuss 2H ..

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OOxxyyggeenn -- Oxygen is the element with atomic number 8. It is a member of the chalcogen group on the periodic table, and is a highly reactive

nonmetallic period 2 element that readily forms compounds (notably oxides) with almost all other elements. At standard temperature and

pressure two atoms of the element bind to form dioxygen, a colorless, odorless, tasteless diatomic gas with the formula O2. Oxygen is the third

most abundant element in the universe by mass after hydrogen and helium and the most abundant element by mass in the Earth's crust.

Diatomic oxygen gas constitutes 20.9% of the volume of air.

All major classes of structural molecules in living organisms, such as proteins, carbohydrates, and fats, contain oxygen, as do the major

inorganic compounds that comprise animal shells, teeth, and bone. Oxygen in the form of O2 is produced from water by cyanobacteria, algae

and plants during photosynthesis and is used in cellular respiration for all complex life. Oxygen is toxic to obligately anaerobic organisms, which

were the dominant form of early life on Earth until O2 began to accumulate in the atmosphere 2.5 billion years ago. Another form (allotrope) of

oxygen, ozone (O3), helps protect the biosphere from ultraviolet radiation with the high-altitude ozone layer, but is a pollutant near the surface

where it is a by-product of smog. At even higher low earth orbit altitudes atomic oxygen is a significant presence and a cause of erosion for

spacecraft. Oxygen is more soluble in water than nitrogen; water contains approximately 1 molecule of O2 for every 2 molecules of N2,

compared to an atmospheric ratio of approximately 1:4. The solubility of oxygen in water is temperature-dependent, and about twice as much

(14.6 mg·L−1

) dissolves at 0 °C than at 20 °C (7.6 mg·L−1

). At 25 °C and 1 standard atmosphere (101.3 kPa) of air, freshwater contains about

6.04 milliliters (mL) of oxygen per liter, whereas seawater contains about 4.95 mL per liter. At 5 °C the solubility increases to 9.0 mL (50% more

than at 25 °C) per liter for water and 7.2 mL (45% more) per liter for sea water. Oxygen condenses at 90.20 K (−182.95 °C, −297.31 °F), and

freezes at 54.36 K (−218.79 °C, −361.82 °F). Both liquid and solid O2 are clear substances with a light sky-blue color caused by absorption in the

red (in contrast with the blue color of the sky, which is due to Rayleigh scattering of blue light). High-purity liquid O2 is usually obtained by the

fractional distillation of liquefied air; Liquid oxygen may also be produced by condensation out of air, using liquid nitrogen as a coolant. It is a

highly reactive substance and must be segregated from combustible materials.

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iinnssiiddee aann aauuttoommoobbiillee ccyylliinnddeerr.. AAss ssoooonn aass tthhee nniittrrooggeenn mmoonnooxxiiddee iinn tthhee hhoott eexxhhaauusstt hhiittss tthhee oouuttssiiddee aaiirr,, iitt rreeaaccttss wwiitthh mmoorree ooxxyyggeenn ttoo

mmaakkee ssoommee nniittrrooggeewwnn ddiiooxxiiddee..

((ii)) ,, wwhhiicchh tthheenn rreeaaccttss uunnddeerr tthhee ssuunnlliigghhtt::

((iiii))

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((iiiiii))

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ppeerrooxxiiddeess aarree ssttrroonngg bblleeaacchhiinngg aanndd ddiissiinnffeeccttiinngg aaggeennttss..

NNii ttrr ooggeenn -- Nitrogen is a chemical element that has atomic number 7 and atomic mass 14.00674 u. Elemental nitrogen is a colorless,

odorless, tasteless and mostly inert diatomic gas at standard conditions, constituting 78% by volume of Earth's atmosphere. Many industrially

important compounds, such as ammonia, nitric acid, organic nitrates (propellants and explosives), and cyanides, contain nitrogen. The

extremely strong bond in elemental nitrogen dominates nitrogen chemistry, causing difficulty for both organisms and industry in converting the

N2 into useful compounds, and releasing large amounts of energy when these compounds burn or decay back into nitrogen gas. When an

animal digests the proteins and nucleic acids, it produces toxic nitrogenous wastes that must be eliminated from its body. The sources of the

nitrogen in these waste products is the (i) amino group ( ) that is removed from the amino acids of proteins and (ii) the nitrogenous bases

of nucleic acids. (a) Most aquatic animals like fish convert the amino groups into ammonia, which is highly toxic and soluble in water, and

excrete them through gills into the surrounding water as dissolved ammonium ions, ; (b) Mammals, amphibians, and sharks convert

amino groups into urea , which is much less toxic, and excrete it as urine; (c) Birds, insects, and reptiles turn the amino group

into uric acid, , a substance much less soluble in water, and store it as solid waste within the egg until the bird hatches.

NOTES FOR CHEMISTRY 11/12 – I: Matter and Properties 7

(i) Nitrogen dioxide, , is a red-brown gas that gives color to old supplies of concentrated nitric acid and to smog.

(ii) Dinitrogen monoxide,

(iii) Nitrogen monoxide,

(iv) Dinitrogen trioxide,

(v) A pair of oxides in equilibrium,

(vi) Dinitrogen pentoxide,

(vii) Nitrous acid, an unstable monoproptic acid made by adding a strong acid to a metal nitrite -

. At higher temperature:

(viii)Hydrazine, a rocket fuel.

CCaarr bboonn -- Carbon is the chemical element with atomic number 6. As a member of group 14 on the periodic table, it is nonmetallic and tetravalent—making four electrons available to form covalent chemical bonds. There are three naturally occurring isotopes, with 12C and 13C being stable, while 14C is radioactive, decaying with a half-life of about 5730 years. Carbon is one of the few elements known since antiquity. There are several allotropes of carbon of which the best known are graphite, diamond, and amorphous carbon. The physical properties of carbon vary widely with the allotropic form. For example, diamond is highly transparent, while graphite is opaque and black. Diamond is among the hardest materials known, while graphite is soft enough to form a streak on paper (hence its name, from the Greek word "to write"). Diamond has a very low electrical conductivity, while graphite is a very good conductor. Under normal conditions, diamond has the highest thermal conductivity of all known materials. All the allotropic forms are solids under normal conditions but graphite is the most thermodynamically stable.

All forms of carbon are highly stable, requiring high temperature to react even with oxygen. The most common oxidation state of

carbon in inorganic compounds is +4, while +2 is found in carbon monoxide and other transition metal carbonyl complexes. The

largest sources of inorganic carbon are limestones, dolomites and carbon dioxide, but significant quantities occur in organic

deposits of coal, peat, oil and methane clathrates [CH4•5.75H2O, also called methane hydrate, hydromethane, methane ice, fire

ice, natural gas hydrate, or gas hydrate, is a solid clathrate compound (more specifically, a clathrate hydrate) in which a large amount of

methane is trapped within a crystal structure of water, forming a solid similar to ice]. Carbon forms more compounds than any other

element, with almost ten million pure organic compounds described to date, which in turn are a tiny fraction of such compounds

that are theoretically possible under standard conditions.

Carbon is the 15th most abundant elements in the Earth's crust, and the fourth most abundant element in the universe by mass

after hydrogen, helium, and oxygen. It is present in all known life forms, and in the human body carbon is the second most

abundant element by mass (about 18.5%) after oxygen. This abundance, together with the unique diversity of organic

compounds and their unusual polymer-forming ability at the temperatures commonly encountered on Earth, make this element

the chemical basis of all known life.

CCaarrbboonn ccoommppoouunnddss ffoorrmm tthhee bbaassiiss ooff aallll kknnoowwnn lliiffee oonn EEaarrtthh,, aanndd tthhee ccaarrbboonn--nniittrrooggeenn ccyyccllee pprroovviiddeess ssoommee ooff tthhee eenneerrggyy

pprroodduucceedd bbyy tthhee SSuunn aanndd ootthheerr ssttaarrss.. AAlltthhoouugghh iitt ffoorrmmss aann eexxttrraaoorrddiinnaarryy vvaarriieettyy ooff ccoommppoouunnddss,, mmoosstt ffoorrmmss ooff ccaarrbboonn aarree

ccoommppaarraattiivveellyy uunnrreeaaccttiivvee uunnddeerr nnoorrmmaall ccoonnddiittiioonnss.. AAtt ssttaannddaarrdd tteemmppeerraattuurree aanndd pprreessssuurree,, iitt rreessiissttss aallll bbuutt tthhee ssttrroonnggeesstt

ooxxiiddiizzeerrss.. IItt ddooeess nnoott rreeaacctt wwiitthh ssuullffuurriicc aacciidd,, hhyyddrroocchhlloorriicc aacciidd,, cchhlloorriinnee oorr aannyy aallkkaalliiss..

((ii)) AAtt eelleevvaatteedd tteemmppeerraattuurreess ccaarrbboonn rreeaaccttss wwiitthh ooxxyyggeenn ttoo ffoorrmm ccaarrbboonn ooxxiiddeess..

((iiii)) IItt wwiillll rreedduuccee ssuucchh mmeettaall ooxxiiddeess aass iirroonn ooxxiiddee ttoo tthhee mmeettaall.. TThhiiss eexxootthheerrmmiicc rreeaaccttiioonn iiss uusseedd iinn tthhee iirroonn aanndd sstteeeell iinndduussttrryy

ttoo ccoonnttrrooll tthhee ccaarrbboonn ccoonntteenntt ooff sstteeeell::

FFee33OO44 ++ 44 CC((ss)) →→ 33 FFee((ss)) ++ 44 CCOO((gg))..

((iiiiii)) IItt rreeaaccttss wwiitthh ssuullffuurr ttoo ffoorrmm ccaarrbboonn ddiissuullffiiddee aanndd wwiitthh sstteeaamm iinn tthhee ccooaall--ggaass rreeaaccttiioonn::

CC((ss)) ++ HH22OO((gg)) →→ CCOO((gg)) ++ HH22((gg))..

((iivv)) CCaarrbboonn ccoommbbiinneess wwiitthh ssoommee mmeettaallss aatt hhiigghh tteemmppeerraattuurreess ttoo ffoorrmm mmeettaalllliicc ccaarrbbiiddeess,, ssuucchh aass tthhee iirroonn ccaarrbbiiddee cceemmeennttiittee

iinn sstteeeell,, aanndd ttuunnggsstteenn ccaarrbbiiddee,, wwiiddeellyy uusseedd aass aann aabbrraassiivvee aanndd ffoorr mmaakkiinngg hhaarrdd ttiippss ffoorr ccuuttttiinngg ttoooollss..

TThhee cchhiieeff iinnoorrggaanniicc ccoommppoouunnddss ooff ccaarrbboonn aarree ttwwoo ooxxiiddeess,, ccaarrbboonniicc aacciidd (( aanndd iittss ssaallttss,, ccyyaanniiddeess,, aanndd ssoommee ccaarrbbiiddeess..

CCaarrbboonn mmoonnooxxiiddee,, aa rreedduucciinngg aaggeenntt,, iiss uusseedd iinndduussttrriiaallllyy ttoo cchhaannggee cceerrttaaiinn mmeettaall oorreess ttoo tthheeiirr mmeettaallss.. CCaarrbboonn ddiiooxxiiddee,, oonnee eenndd

pprroodduucctt iinn tthhee ccoommbbuussttiioonn ooff aannyytthhiinngg ccoonnttaaiinniinngg ccaarrbboonn,, wwiillll ddiissssoollvvee iinn wwaatteerr,, aanndd ssoommee rreeaaccttss ttoo ffoorrmm uunnssttaabbllee ccaarrbboonniicc aacciidd,,

..

.. TThheerreeffoorree,, aaqquueeoouuss ccaarrbboonn ddiiooxxiiddee nneeuuttrraall iizzeess mmeettaall hhyyddrrooxxiiddeess:: AAll ll bbiiccaarrbboonnaatteess aanndd ccaarrbboonnaatteess rreeaacctt wwii tthh ssttrroonngg aacciiddss ttoo ggiivvee ccaarrbboonn ddiiooxxiiddee aanndd wwaatteerr pplluuss ssoommee ssaall tt..

NOTES FOR CHEMISTRY 11/12 – I: Matter and Properties 8

SSooddiiuumm ccyyaanniiddee rreelleeaasseess tthhee ppooiissoonnoouuss ccyyaanniiddee iioonn,, iinn wwaatteerr.. TThhiiss iioonn ccoommbbiinneess wwii tthh hhyyddrrooggeenn iioonn ttoo ll iibbeerraattee hhyyddrrooggeenn ccyyaanniiddee,, aa ppooiissoonnoouuss ggaass uusseedd iinn cceerrttaaiinn ppeesstt--ccoonnttrrooll ooppeerraattiioonnss aass wweell ll aass iinn tthhee ggaass cchhaammbbeerrss tthhaatt ssoommee ggoovveerrnnmmeennttss uussee ttoo ccaarrrryy oouutt ddeeaatthh sseenntteennccee.. AAss ttoo 22000099,, ggrraapphheennee aappppeeaarrss tthhee ssttrroonnggeesstt mmaatteerriiaall eevveerr tteesstteedd.. HHoowweevveerr,, tthhee pprroocceessss ooff sseeppaarraattiinngg ii tt ffrroomm ggrraapphhii ttee wwii ll ll rreeqquuiirree ssoommee tteecchhnnoollooggiiccaall ddeevveellooppmmeenntt bbeeffoorree ii tt iiss eeccoonnoommiiccaall eennoouugghh ttoo bbee uusseedd iinn iinndduussttrriiaall pprroocceesssseess

SSuullffuurr -- Sulfur or sulphur is the chemical element that has the atomic number 16. It is denoted with the symbol S. It is an

abundant, multivalent non-metal. Sulfur, in its native form, is a beautiful bright yellow crystalline solid consisting of molecules of

in which the atoms are joined in crown-like rings. In nature, it can be found as the pure element and as sulfide and sulfate

minerals, and some hydrogen sulfide, and in the groups in proteins. It is an essential element for life and is

found in two amino acids, cysteine and methionine. Its commercial uses are primarily in fertilizers, but it is also widely used in

black gunpowder, matches, insecticides and fungicides. Elemental sulfur crystals are commonly sought after by mineral collectors

for their brightly colored polyhedron shapes. In nonscientific contexts, it can also be referred to as brimstone.

Sulfur dissolves in solutions of metal sulfites to give the thiosulfate ion, , and in solutions of metal sulfides or ammonium

sulfide to give polysulfide ions, sulfide to give polysulfide ions, , where x varies from 2 to about 10. The polysulfide

solutions react with strong acids to give corresponding sulfanes. These are unstable yellow oils with the general formula, .

In times, sulfanes decompose to hydrogen sulphide and finely divided colloidal sulphur. When sulphur dioxide dissolves in water,

some sulfurous acid ( ) forms, a weak, unstable, diprotic acid having reasonably stable salts – hydrogen sulfites and sulfites.

These salts all react with strong acids to give sulphur dioxide plus sulfurous acid.

PPhhoosspphhoorr uuss -- Phosphorus is the chemical element that has atomic number 15. A multivalent nonmetal of the nitrogen group,

phosphorus is commonly found in inorganic phosphate rocks. Elemental phosphorus exists in two major forms in molecules of 4

atoms joined together in a tetrahedron, , - white phosphorus, very poisonous, and red phosphorus. Although the term

"phosphorescence", meaning glow after illumination, derives from phosphorus, glow of phosphorus originates from oxidation of

the white (but not red) phosphorus and should be called chemiluminescence.

Due to its high reactivity, phosphorus is never found as a free element in nature on Earth. The first form of phosphorus to be

discovered (white phosphorus, in 1669) emits a faint glow upon exposure to oxygen — hence its name given from Greek

mythology, Φωσφόρος meaning "light-bearer" (Latin Lucifer), referring to the "Morning Star", the planet Venus.

Phosphorus is a component of DNA, RNA, ATP, and also the phospholipids which form all cell membranes. It is thus an essential

element for all living cells. The most important commercial use of phosphorus-based chemicals is the production of fertilizers.

Phosphorus compounds are also widely used in explosives, nerve agents, friction matches, fireworks, pesticides, toothpaste and

detergents.

SSii ll iiccoonn -- Silicon is the most common metalloid. It is a chemical element, which has the symbol Si and atomic number 14. A

tetravalent metalloid, silicon is less reactive than its chemical analog carbon. As the eighth most common element in the

universe by mass, silicon very rarely occurs as the pure free element in nature, but is more widely distributed in dusts, planetoids

and planets as various forms of silicon dioxide (silica) or silicates. In Earth's crust, silicon is the second most abundant element

after oxygen, making up 25.7% of the crust by mass.

Silicon has many industrial uses. It is the principal component of most semiconductor devices, most importantly integrated

circuits or microchips. Silicon is widely used in semiconductors because it remains a semiconductor at higher temperatures than

the semiconductor germanium and because its native oxide is easily grown in a furnace and forms a better

semiconductor/dielectric interface than any other material.

In the form of silica and silicates, silicon forms useful glasses, cements, and ceramics. It is also a constituent of silicones, a

class-name for various synthetic plastic substances made of silicon, oxygen, carbon and hydrogen, often confused with silicon

itself.

NOTES FOR CHEMISTRY 11/12 – I: Matter and Properties 9

Silicon is an essential element in biology, although only tiny traces of it appear to be required by animals. It is much more

important to the metabolism of plants, particularly many grasses, and silicic acid (a type of silica) forms the basis of the striking

array of protective shells of the microscopic diatoms.

Notable characteristics - The outer electron orbitals (half filled subshell holding up to eight electrons) have the same structure as

in carbon and the two elements are sometimes similar chemically. Even though it is a relatively inert element, silicon still reacts

with halogens and dilute alkalis, but most acids (except for some hyper-reactive combinations of nitric acid and hydrofluoric acid)

do not affect it. Having four bonding electrons however gives it, like carbon, many opportunities to combine with other elements

or compounds under the right circumstances.

Both silicon and (in certain aspects) carbon are semiconductors, readily either donating or sharing their four outer electrons

allowing many different forms of chemical bonding. Pure silicon has a negative temperature coefficient of resistance, since the

number of free charge carriers increases with temperature. The electrical resistance of single crystal silicon significantly changes

under the application of mechanical stress due to the piezoresistive effect.

In its crystalline form, pure silicon has a gray color and a metallic luster. It is similar to glass in that it is rather strong, very brittle,

and prone to chipping.

Silicon is one of the few substances, like water/ice and gallium, which density is higher in liquid than in solid state, so it expands

when it freezes.

HHaallooggeennss The halogens are five non-metallic elements found in group 17 of the periodic table. The term "halogen" means "salt-former"

and compounds containing halogens are called "salts". All halogens have 7 electrons in their outer shells, giving them an

oxidation number of -1. The halogens exist, at room temperature, in all three states of matter: • Solid- Iodine, Astatine

• Liquid- Bromine

• Gas- Fluorine, Chlorine

Transition Metals [leave this part of note until you learn electron configuration in Grade 12 Chemistry] The term transition metal (sometimes also called a transition element) has two possible meanings:

• In the past it referred to any element in the d-block of the periodic table, which includes groups 3 to 12 on the periodic table. All

elements in the d-block are metals.

• The modern, IUPAC definition states that a transition metal is "an element whose atom has an incomplete d sub-shell, or which can

give rise to cations with an incomplete d sub-shell." Group 12 elements are not transition metals in this definition.

The name transition comes from their position in the periodic table of elements. In each of the four periods in which they occur, these

elements represent the successive addition of electrons to the d atomic orbitals of the atoms. In this way, the transition metals represent the

transition between group 2 elements and group 13 elements.

Transition metals play an important role in living organisms, and are also extremely valuable as strong, structurally useful materials.. for

example, iron is the primary building metal in our society. Chromium, a hard, silvery metal, very resistant to corrosion, is used as a protective

coating on metals and in producing alloy such as stainless steel, which contains mainly iron, chromium, and nickel. Silver and gold are used in

jewelry and coins. Cobalt is found in vitamin and iron is an important part of the hemoglobin required for oxygen transport.

(i) Classification In the d-block the atoms of the elements have between 1 and 10 d electrons.

Group 3 4 5 6 7 8 9 10 11 12

Period 4 Sc 21

Ti 22

V 23

Cr 24

Mn 25

Fe 26

Co 27

Ni 28

Cu 29

Zn 30

Period 5 Y 39

Zr 40

Nb 41

Mo 42

Tc 43

Ru 44

Rh 45

Pd 46

Ag 47

Cd 48

Period 6 La 57

Hf 72

Ta 73

W 74

Re 75

Os 76

Ir 77

Pt 78

Au 79

Hg 80

Period 7 Ac 89

Rf 104

Db 105

Sg 106

Bh 107

Hs 108

Mt 109

Ds 110

Rg 111

Uub 112

NOTES FOR CHEMISTRY 11/12 – I: Matter and Properties 10

(ii) Coloured compounds

Co(NO3)2 K2Cr2O7 K2CrO4 NiCl2 CuSO4 KmnO4

From left to right, aqueous solutions of: Co(NO3)2 (red); K2Cr2O7 (orange); K2CrO4 (yellow); NiCl2 (green); CuSO4 (blue); KmnO4 (purple).

Colour in transition metal compounds may be due to electronic transitions of two principal types.

• charge transfer transitions. An electron may jump from a predominantly ligand orbital to a predominantly metal orbital, giving rise to a ligand-to-metal charge-transfer (LMCT) transition. These can most easily occur when the metal is in a high oxidation state. The colour of chromate, dichromate and permanganate ions is due to LMCT transitions. A metal-to ligand charge transfer (MLCT) transition will be most likely when the metal is in a low oxidation state and the ligand is easily 10xidized. Mercuric iodide, HgI2, is red because of a MLCT transition. As this example shows, charge transfer transitions are not restricted to transition metals.

• d-d transitions. An electron jumps from one d-orbital to another. In complexes of the transition metals the d orbitals do not all have the same energy. The pattern of splitting of the d orbitals can be calculated using crystal field theory. The extent of the splitting depends on the particular metal, its oxidation state and the nature of the ligands. The actual energy levels are shown on Tanabe-Sugano diagrams.

In centrosymmetric complexes, such as octahedral complexes, d-d transitions are forbidden by the Laporte rule and only occur because of vibronic coupling in which a molecular vibration occurs together with a d-d transition. Tetrahedral complexes have somehat more intense colour because mixing d and p orbitals is possible when there is no centre of symmetry, so transitions are not pure d-d transitions. The molar absorptivity of bands caused by d-d transitions are relatively low, roughly in the range 5-500 M-1cm-1 (where M = mol dm-3). Some d-d transitions are spin forbidden. An example occurs in octahedral, high-spin complexes of manganese(II), which has a configuration in which all five electron has parallel spins; the colour of such complexes is much weaker than in complexes with spin-allowed transitions. In fact many compounds of manganese(II) appear almost colourless.

(iii) Oxidation states A characteristic of transition metals is that they exhibit two or more oxidation states, usually differing by one. For example, compounds of vanadium are known in all oxidation states between −1, such as [V(CO)6]

−, and +5, such as [VO3] −.

Main group elements in groups 13 to 17 also exhibit multiple oxidation states. The “common” oxidation states of those elements differ by two. For example, compounds of gallium in oxidation states +1 and +3 exist in which there is a single gallium atom. No such compound of Ga (II) is known: any such compound would have an unpaired electron and would behave as a free radical and be destroyed rapidly. However, under the right conditions dimeric compounds such as [Ga2Cl6]

2− can be made in which a Ga-Ga bond is formed from the unpaired electron on each Ga atom. Thus the main difference, regarding oxidation states, between transition elements and other elements is that oxidation states are known in which there is a single atom of the element and one or more unpaired electrons.

The maximum oxidation state in the first row transition metals is equal to the number of valence electron from titanium (+4) up to manganese (+7), but decreases in the later elements. In the second and third rows the maximum occurs with ruthenium and osmium (+8). In compounds such as [MnO4]

− and OsO4 the elements achieve a stable octet by forming four covalent bonds.

The lowest oxidation states are exhibited in such compounds as Cr(CO)6 (oxidation state zero) and [Fe(CO)4]2− (oxidation state −2) in

which the 18-electron rule is obeyed. These complexes are also covalent.

Ionic compounds are mostly formed with oxidation states +2 and +3. In aqueous solution the ions are hydrated by (usually) six water molecules arranged octahedrally.

(iv) Magnetism Transition metal compounds are paramagnetic when they have one or more unpaired d electrons. In octahedral complexes with between four and seven d electrons both high spin and low spin states are possible. Tetrahedral transition metal complexes such as [FeCl4]

2− are high spin because the crystal field splitting is small so that the energy to be gained by virtue of the electrons being in lower energy orbitals is always less that the energy needed to pair up the spins. Some compounds are diamagnetic. These include octahedral, low-spin, d6 and square-planar d8 complexes. In these cases, crystal field splitting is such that all the electrons are paired up.

Ferromagnetism occurs when individual atoms are paramagnetic and the spin vectors are aligned parallel to each other in a crystalline material. Metallic iron and the alloy alnico are examples of ferromagnetic materials involving transition metals. Antiferromagnetism is another example of a magnetic property arising from a particular alignment of individual spins in the solid state.

Inner Transition Metals The elements of the 4f series are generally called lanthanides, after the first member of the series, lanthanum (La). Similarly, the elements of the 5f series are called actinides, after the first member of that series, actinium (Ac).all the elements from lanthanum to lutetium have extremely similar properties, but their electron configurations are not perfectly regular across the series. For example the electron configuration of lanthanum (La) is [ , whereas that of the next element certium (Ce) is [ . Lanthanides are all very similar to one another because they differ principally in the number of electrons in the 4f and 5d sublevels rather than in their outer 6s sublevels. They readily lose 3 electrons to form 3+ ions. They are all rather soft, silvery

NOTES FOR CHEMISTRY 11/12 – I: Matter and Properties 11

metals. Somewhat less reactive than the heavier alkaline earth metals, they are still too reactive to be used as structural materials. Actinides are of greatest interest because of their radioactivity. All isotopes of these elements are radioactive. Only thorium(Th) and uranium (U) occur to any extent in nature. All the elements after uranium are artificial elements.

# Element Sources Uses 1 Oxygen 23% of air, 89% of mass of water, 46% of

crustal rocks Its ability to support respiration, combustion, and decay make it an essential element for life.

2 Silicon Silica ( ) As described above

3 Aluminum Bauxite ( )

4 Iron Hematite ( -rust), magnetite ( ), siderite ( ), iron pyrite ( )

5 Calcium Limestone ( )

7 Sodium 3% of seawater is salt (NaCl) Table salt (NaCl), caustic soda (NaOH), bleach (NaOCl), soda ash ( ), baking soda ( )

8 Potassium Potassium in nature occurs only in ionic salts. As such, it is found dissolved in seawater (which is 0.04% potassium by weight), and is part of many minerals.

- it's an essential component of plant nutrition and is found in most soil types - used in agriculture as a fertilizer - tomato paste, orange juice, beet greens, white beans, and bananas, are good dietary sources of potassium - used in baking powder and as a food preservative - used in dyes and stains (bright yellowish-red colour) - used in explosives and fireworks, safety matches, tanning of leather, and fly paper - used in specialized glasses, ceramics, and enamels

6 Magnesium Magnesite ( ) 9 Hydrogen The two main methods for getting hydrogen

are reforming it from another substance or separating from water

. Largest use to make ammonia.

10 Titanium But platinum & iridium are very rare. Titanium is very expensive because it is awkward to extract from its ores - for example, from rutile, TiO2. Titanium can't be extracted by reducing the ore using carbon as a cheap reducing agent. The problem is that titanium forms a carbide, TiC, if it is heated with carbon, so you don't get the pure metal that you need. The presence of the carbide makes the metal very brittle.

Titanium is corrosion resistant, very strong and has a high melting point. It has a relatively low density (about 60% that of iron). It is also the tenth most commonly occurring element in the Earth's crust. That all means that titanium should be a really important metal for all sorts of engineering applications. In fact, it is very expensive and only used for rather specialized purposes. Titanium is used, for example: • in the aerospace industry - for example in aircraft engines and air frames; • for replacement hip joints; • for pipes, etc, in the nuclear, oil and chemical industries where corrosion is

likely to occur. Nuclear Stability, Bonding Ability and Relative Abundance A few nonmetallic elements of low atomic weights make up most of the universe. Why these and not others? Nuclear stability is one factor – perhaps a far more likely factor than chemical stability. Isotopes of the lightest elements have atomic nuclei that are the most stable with respect to various modes of breakdown. But what causes oxygen and silicon to make up so much of the crust of the earth, and what makes oxygen, carbon, nitrogen, and hydrogen so prevalent in living systems? Perhaps one answer is that atoms of these (and others) nonmetallic elements can form strong covalent bonds with each other. Moreover, most of them can form double and triple bonds as well as single bonds. Hydrogen is rare in the atmosphere because it is light enough to escape Earth’s gravitational attraction. Most of the Earth’s hydrogen is combined with oxygen as water. It also occurs frequently in combination with carbon in a large variety of organic compounds called hydrocarbons. Hydrocarbons typically are found deep beneath the Earth’s surface. In part, this because carbon reacts readily with oxygen and nitrogen, which are very common in the atmosphere. Nonmetals and Pi Bonds Carbon, oxygen, and nitrogen all have an unusual ability to form double bonds. In an atom of any of these elements, the p orbitals of the electrons available for bonding are in the second shell, not too far from the atom’s nucleus. If we drop down to the third row of the periodic table to silicon, sulfur and phosphorus, the p orbitals of the bonding electrons are farther away from the nucleus in shell three. Therefore, when third-level p electrons go into pi bonds, they cannot be as effective in holding nuclei near other as when the p electrons come from the second level. Thus, pi bonds formed by the overlap of third shell p orbitals are not as strong as those made from second shell p orbitals. The oxides of carbon and silicon illustrate that the types of bonds used to hold the atoms together can depends entirely on e level in which the bonding electrons reside. Carbon dioxide is a gas at room temperature and pressure. Silicon dioxide is a solid that in its purer forms makes up quartz and quartz sand. A molecule of dioxide of carbon, a second-row element, has two double bonds, and therefore two pi-bonds, and the gas is made up of molecules. Silicon, just beneath the carbon, in the third-row, forms a dioxide without individual molecules and pi bonds are not present. The silicon-oxygen bonds are single bonds in a huge silicon-oxygen network of atoms. Noble Gas Compounds – Only the two most electronegative elements, fluorine and oxygen, give noble gas compounds, and only xenon and krypton are known to be reactive.

NOTES FOR CHEMISTRY 11/12 – I: Matter and Properties 12

/ Mendeleev’s (1869) Periodic Law stated that elements arranged in order of increasing atomic mass (total # of protons and neutrons) show a periodic recurrence of properties at regular intervals. The modern periodic law that was revised from Mendeleev’s based on H.G.J. Moseley’s (1887-1915) discovery of the atomic number (Z) stated that elements arranged in order of increasing atomic number (total # of protons) show a periodic recurrence of properties at regular intervals. In accordance with this periodic law, elements are arranged in a row-and-column table called the Periodic Table:

• A group – elements with similar chemical properties in a vertical column in the main part of the table, numbered from 1 to 18. • A period – elements, arranged in a horizontal row, whose properties change from metallic on the left to nonmetallic on the right,

numbered from 1 to 7 (top to bottom). • The “staircase line” – a zigzag line that separates metals (to the left) from nonmetals (to the right).

Periodic Trends: 1. Atomic sizes (radii) decrease from left to right across a period in the periodic table and decrease from bottom to top in a

group. In a period, each succeeding atom has an additional electron and an additional proton. The electrons are added to the

same shell and are located at a relatively constant distance from the nucleus. The increasing nuclear charge has a stronger

effect than the electrons’ repelling force. It attracts the electron clouds closer to the nucleus, thereby decreasing the overall

size. From the top each row adds an additional shell to the atom’s structure. This added outer shell of electrons must be

further from the nucleus than the preceding shell. In addition, the added distance means that the attractive force of the

nucleus is reduced. Cations are positively charged ions, and anions are negatively charged. Cations are always smaller than

their neutral atoms, and anions are always larger

2.The process of removing an electron from an atom is called ionization, and the energy needed to do this called ionization

energy. The energy needed to remove only one electron is called the first ionization energy. Removal of each successive

electron from an Atom requires an increased amount of energy because the removal of the electron results in a stronger

attraction between the nucleuses and the remaining electrons. The ionization energies increase from left to right across a

period in the periodic table, and increase from the bottom to the top of a group, creating a diagonal relationship in the table.

3. The ionization energy is one of the factors involved in determining the electronegativity, which follows the same diagonal

trend as the ionization energy does.

4. The electron affinity is defined as the energy needed to add an electron to an atom. It has the same diagonal relationship as

the ionization energy in the periodic table

1. European System: IA to VIIIA, then IB to VIIIB in sequential order of columns. 2. American System: 1A , 2A, then jump over transition metals columns, to 3A to 8A; start 1B in column 11, then 2B in column 12, then 3B to 8B back in columns 3 to 10 3. Compromised System: Arabic Numerals 1 to 13 in sequential order (not widely accepted yet).

NOTES FOR CHEMISTRY 11/12 – I: Matter and Properties 13

Summary of Periodic Trends:

1. Atomic Sizes ………...increases from right to left; from top to bottom: 2. Ionization Energy …...increases from left to right; from bottom to top:

(less likely to lose an electron; metals have lower IE, therefore they tend to lose electrons to non-Metals)

3. Electron Affinity ….. increases from left to right; from bottom to top: (more likely to gain an electron with release of energy) 4. Electronegativity ….. increases from left to right; from bottom to top:

(oxidizing agent: higher tendency to attract electrons) 5. Oxoacids Strength ….increases from left to right; from bottom to top: 6. Binary Acids Strength increases from left to right; from top to bottom:

The term ionization energy (EI) of an atom or molecule means the energy needed to remove an electron from an atom. Large atoms require low ionization energy while small atoms require high ionization energy. This quantity was formerly called ionization potential, and was at one stage measured in volts. The name "ionization energy" is now strongly preferred. In atomic physics the ionization energy is measured using the unit "electronvolt" (eV). In chemistry, the value is usually given in kJ/mol (or formerly kcal/mol). This value is strictly the "molar ionization energy" and corresponds to the energy required to remove (to infinity) one mole of electrons from one mole of gaseous atoms or molecules. However it is often just called "ionization energy". The Electron affinity of a molecule or atom is the energy change when an electron is added to the neutral atom to form a negative ion. This property can only be measured in an atom in gaseous state. X + e− → X− The electron affinity, Eea, is defined as positive when the resulting ion has a lower energy, i.e. it is an exothermic process that releases energy: Eea = Einitial − Efinal Alternatively, electron affinity is often described as the amount of energy required to detach an electron from a singly charged negative ion, i.e. the energy change for the process X− → X + e− Although Eea varies greatly across the periodic table, some patterns emerge. Generally, nonmetals have more negative Eea than metals. Atoms whose anions are more stable than neutral atoms have a greater Eea. Chlorine most strongly attracts extra electrons; mercury most weakly attracts an extra electron. The electron affinities of the noble gases have not been conclusively measured, so they may or may not have slightly negative values. Eea generally increases across a period (row) in the periodic table. This is caused by the filling of the valence shell of the atom; a group 7A atom releases more energy than a group 1A atom on gaining an electron because it obtains a filled valence shell and therefore is more stable. A trend of decreasing Eea going down the groups in the periodic table would be expected. The additional electron will be entering an orbital farther away from the nucleus, and thus would experience a lesser effective nuclear charge. However, a clear counterexample to this trend can be found in group 2A, and this trend only applies to group 1A atoms. Electron affinity follows the trend of electronegativity. Fluorine (F) has a higher electron affinity than oxygen and so on. Electronegativity, symbol χ (the Greek letter chi), is a chemical property that describes the tendency of an atom or a functional group to attract electrons (or electron density) towards itself and thus the tendency to form negative ions. An atom's electronegativity is affected by both its atomic number and the distance that its valence electrons reside from the charged nucleus.

NOTES FOR CHEMISTRY 11/12 – I: Matter and Properties 14

Electronegativity Table of the Elements

Group

Period 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18

1 H 2.1

He 0

2 Li 0.98

Be 1.57

B 2.04

C 2.55

N 3.04

O 3.44

F 3.98

Ne 0

3 Na 0.93

Mg 1.31

Al 1.61

Si 1.9

P 2.19

S 2.58

Cl 3.16

Ar 0

4 K 0.82

Ca 1

Sc 1.36

Ti 1.54

V 1.63

Cr 1.66

Mn 1.55

Fe 1.83

Co 1.88

Ni 1.91

Cu 1.9

Zn 1.65

Ga 1.81

Ge 2.01

As 2.18

Se 2.55

Br 2.96

Kr 0

5 Rb 0.82

Sr 0.95

Y 1.22

Zr 1.33

Nb 1.6

Mo 2.16

Tc 1.9

Ru 2.2

Rh 2.28

Pd 2.2

Ag 1.93

Cd 1.69

In 1.78

Sn 1.96

Sb 2.05

Te 2.1

I 2.66

Xe 2.6

6 Cs 0.79

Ba 0.89

La 1.1

Hf 1.3

Ta 1.5

W 2.36

Re 1.9

Os 2.2

Ir 2.2

Pt 2.28

Au 2.54

Hg 2

Tl 2.04

Pb 2.33

Bi 2.02

Po 2

At 2.2

Rn 0

7 Fr 0.7

Ra 0.89

Ac 1.1

Rf Db Sg Bh Hs Mt Uun Uuu Uub

Lanthanides

Ce 1.12

Pr 1.13

Nd 1.14

Pm 1.13

Sm 1.17

Eu 1.2

Gd 1.2

Tb 1.1

Dy 1.22

Ho 1.23

Er 1.24

Tm 1.25

Yb 1.1

Lu 1.27

Actinides

Th 1.3

Pa 1.5

U 1.38

Np 1.36

Pu 1.28

Am 1.3

Cm 1.3

Bk 1.3

Cf 1.3

Es 1.3

Fm 1.3

Md 1.3

No 1.3

Lr

Key: White= No data

0-.66 .66-1 1-1.33 1.33-1.66 1.66-2 2-2.33 2.33-2.66 2.66-

This table is the Pauling electronegativity scale. There are other ways of measuring electronegativity, such as the Mulliken scale and the

Allred-Rochow scale. Linus Pauling's electronegativity scale is the most common. Note that atoms toward the upper right are more

electronegative, and those to the lower left are least electronegative. Pauling did not assign electronegativities to the noble gasses because

they typically do not form covalent bonds.

In general electronegativity is the measure of an atom's ability to attract electrons to itself in a covalent bond. Because fluorine is the most

electronegative element, the electrons tend to "hang out" more toward the fluorine atom when fluorine is covalently bonded to other atoms.

Oxygen is the 2nd most electronegative element.

When you examine a periodic table, you will find that (excluding the noble gases) the electronegativity values tend to increase as you go to the

right and up. The reverse statement is that the values tend to decrease going down and to the left. This pattern will help when you are asked to

put several bonds in order from most to least ionic without using the values themselves.

Electronegativity values are useful in determining if a bond is to be classified as nonpolar covalent, polar covalent or ionic.

What you should do is look only at the two atoms in a given bond. Calculate the difference between their electronegativity values. Only the

absolute difference is important.

I. Nonpolar Covalent: This type of bond occurs when there is equal sharing (between the two atoms) of the electrons in the bond. Molecules

such as Cl2, H2 and F2 are the usual examples.

Textbooks typically use a maximum difference of 0.2 - 0.5 to indicate nonpolar covalent. Since textbooks vary, make sure to check with your

teacher for the value he/she wants. Here we will use 0.5.

NOTES FOR CHEMISTRY 11/12 – I: Matter and Properties 15

One interesting example molecule is CS2. This molecule has nonpolar bonds. Sometimes a teacher will only use diatomics as examples in lecture

and then spring CS2 as a test question. Since the electronegativities of C and S are both 2.5, you have a nonpolar bond.

II. Polar Covalent: This type of bond occurs when there is unequal sharing (between the two atoms) of the electrons in the bond. Molecules

such as NH3 and H2O are the usual examples.

The typical rule is that bonds with an electronegativity difference less than 1.6 are considered polar. (Some textbooks or web sites use 1.7.)

Obviously there is a wide range in bond polarity, with the difference in a C-Cl bond being 0.5 -- considered just barely polar -- to the difference

the H-O bonds in water being 1.4 and in H-F the difference is 1.9. This last example is about as polar as a bond can get.

III. Ionic: This type of bond occurs when there is complete transfer (between the two atoms) of the electrons in the bond. Substances such as

NaCl and MgCl2 are the usual examples.

The rule is that when the electronegativity difference is greater than 2.0, the bond is considered ionic.

So, let's review the rules:

1.If the electronegativity difference (usually called DEN) is less than 0.5, then the bond is nonpolar covalent.

2. If the DEN is between 0.5 and 1.6, the bond is considered polar covalent [examples of polar bond: C−−−−F, C−I, C−O, C−N, O−H, N−H]

3. If the DEN is greater than 2.0, then the bond is ionic. That, of course, leaves us with a problem. What about the gap between 1.6 and 2.0? So,

rule #4 is:

4. If the DEN is between 1.6 and 2.0 and if a metal is involved, then the bond is considered ionic. If only nonmetals are involved, the bond is

considered polar covalent.

Here is an example: Sodium bromide (formula = NaBr; ENNa = 0.9, ENBr = 2.8) has a DEN = 1.9. Hydrogen fluoride (formula = HF; ENH = 2.1, ENF =

4.0) has the same DEN. We use rule #4 to decide that NaBr has ionic bonds and that HF has a polar covalent bond in each HF molecule.

Polar Bonds and Electronegativity

A nonpolar covalent bond has a uniform distribution of electron charge between the bonded atoms. The simplest nonpolar covalent bonds exist in "homonuclear diatomic" molecules like H2 and Cl2. Both atoms attract the shared electrons equally. The shared electrons spend equal time on both ends of the bond and molecule. There is no permanent localized electric charge build up.

A polar bond has a unsymmetric electron cloud distribution. The electrons in the bond are not shared equally. This typically happens between two nonmetal atoms that are two or more positions apart in the periodic table. An example is Hl. The electrons in the bond spend more time around the chlorine nucleus. this makes the chlorine end more electron rich than the hydrogen end of the bond. The "arrow" indicates the direction of the electron shift. This is a polar bond.

Electronegativity, EN, is an index that tells the relative attraction an element has for electrons in a bond. Electronegativity has a high value of 4.0 for F, fluorine. The lowest electronegativity value is about 0.7 for Cs, cesium. The table below shows the nonmetals have relatively high electronegativities. The metals have relatively low electronegativities. The electronegativities follow the same trends as atomic sizes (radius). Electronegativity gets smaller with increasing distance from fluorine. Atoms that are equidistant from fluorine have similar (not identical) electronegativities. The rare gases generally are not tabulated for EN values.

Bond polarity is directly related to electronegativity difference. The greater the difference in electronegativity the more polar the bond.

When two elements are next to one another in the periodic table they have similar electronegativities. Chlorine has a value of 3.0 while

bromine has a value of 2.8. These two atoms in BrCl would have a nonpolar covalent bond.

The bond between carbon with 2.5 and nitrogen with 3.0 would be polar. The bond between chlorine, 3.0, and boron 2.0, would be polar.

When electronegativities get too different the bond is not polar, it is ionic. The nonmetal gets the electrons from the metal essentially 100% of

the time. The metal atom is stripped of its valence electrons. These atoms don't share the electrons. There is a transfer to the nonmetal.

Differences in EN are related to positions in the periodic table. The greater the separation between the elements the greater the

electronegativity difference. Look at fluorine and cesium. This is the greatest difference possible, 4.0 - 0.7 = 3.3. This combination is so different

that they form ionic bonds.

NOTE: hydrogen really fits in best at the bottom of the halogens, group 7A. The fact that the EN changes so predictably with the element

position allows us to make predictions about bonds based on periodic table location and not even have the EN values.

Example 1: Which of the following pairs of bonds would be predicted to be the most polar?

a. Cl and B b. F and C c. C and I

NOTES FOR CHEMISTRY 11/12 – I: Matter and Properties 16

Example 2:

Which of the following pairs of bonds would be predicted be the least polar?

a. Cs and F b. B and At c. H and O

Answer for Example 1:

a. Sorry if you picked a. The chlorine is one atom away from fluorine. The boron is four atoms away from fluorine. The fact that

chlorine is less electronegative than fluorine makes the B-Cl bond less polar than the B-F bond. and boron are both Cl and B

b. Great if you picked b. The fluorine and boron have the biggest difference in EN so they would have the most polar bond.

c. Sorry if you picked c. Carbon and iodine are two and three atoms away from fluorine. They have similar electronegativity,EN,

values. Actually carbon has EN = 2.5 and Iodine has EN = 2.5. the values are identical.

Answer for Example 2:

a. Sorry if you picked a. The cesium is a nonmetal and as far from fluorine as an element can be. The bond is ionic.

b. Great if you picked b. The boron and At are both four atoms away from fluorine. The EN values are boron, 2.0, and astatine, At, 2.2.

The electronegativity difference is very small. The bond is slightly polar.

c. Sorry if you picked c. Hydrogen is a bad acting element. It really has an electronegativity that would put it at the bottom of group 7A.

The bond between hydrogen and oxygen is one of the most polar covalent bonds possible. They have very different electronegativity,

EN, values. Actually hydrogen has EN = 2.1. Oxygen is second only to fluorine in the electronegativity scale. Oxygen has EN = 3.5. The

values are different by 1.4 units. The bond is very polar.

Classification of Hydrogen Compounds There are three main types of binary hydrogen compounds, organized by the relative electronegativities of hydrogen and the other element. The known binary EHn compounds are organized on the following periodic chart:

1 Saline or Ionic Hydrides The saline, or Ionic, hydrides, very strong bases, are produced when hydrogen reacts with the very electropositive elements under heat. They are compounds of Mn+ and n H–, and are, when very pure, colourless ionic compounds with structures similar to those of the ionic halides. However, the materials that are commercially available are always found to be contaminated by colloidal-size particles of the metallic element, and typically appear from light to very dark gray. They exist in the following crystalline forms:

Compounds Crystal Structure LiH, NaH, KH, RbH, CsH Rock salt (NaCl) MgH2 Rutile

CaH2, SrH2, BaH2 Distorted PbCl2

They are usually binary compounds. Binary compounds occur when there are two elements in a compound. They are also insoluble in solutions.

Ionic hydrides combine vigorously with water to produce hydrogen gas and hydroxide ions , and readily react with various compounds.

2 Metallic or Interstitial Hydrides The more electronegative, transition metals do not form saline hydrides. Instead they form a range of, typically, non-stoichiometric hydrogen derivatives that are of great theoretical and practical interest, but of little utility in chemistry in general.

The idea and basis for the non-stoichiometric chacteristic is that with metal and hydrogen bonding there is a crystal lattice that H atoms can and may fill in between the lattice while some are not a definite ordered filling. Thus it is not a fixed ratio of H atoms to the metals. Even so, metallic hydrides consist of more basic stoichiometric compounds as well. 3 Molecular Hydrogen Compounds By far the most important hydrogen compounds are the molecular compounds. Molecular hydrogen compounds form for most of the p-block elements. Unlike the halogen derivatives of the elements, there is usually only one oxidation state of each of these

NOTES FOR CHEMISTRY 11/12 – I: Matter and Properties 17

elements that forms a stable hydrogen compound. However, there are a few unstable species. Thus carbon, which has CH4 as its stable hydrogen compound, also has the unstable form CH2, the carbene molecule. Compared to carbon, for which a vast array of hydrocarbon derivatives exist, based on chains and rings of the element with hydrogen satisfying the remaining valences, the remaining elements have far less diversity. Thus nitrogen has ammonia, NH3 and hydrazine, H2NNH2 (N2H4). Oxygen has water and hydrogen peroxide. The heavier group 14 elements do have a few recognized extended structures, but nothing even close to the diversity of carbon.

These hydrides can be volatile or non-volatile. Volatile simply means being readily able to be vaporized at low temperatures. One such example of a covalent hydride is when hydrogen bonds with chlorine and forms hydrochloric acid (HCl).

The hydrides of nonmetals on the periodic table become more electronegative as you move from group 13 to 17. This means that they are less capable of donating an electron, and want to keep them because their electron orbital becomes fuller. Instead of donating a H-, they would instead donate a H+ because they are more acidic.

We can classify the molecular compounds of hydrogen into three classes, as shown in the following table that also gives the names of these compounds.

Class Group Formula Trivial name IUPAC name Electron-Deficient 13 B2H6 Diborane Diborane(6) 13 AlH3 (polymeric) Alane Alane 13 Ga2H6 (< –30°C) Gallane Gallane Electron-Precise 14 CH4 and hydrocarbons Methane Methane 14 SiH4 and silanes Silane Silane 14 GeH4 and germanes Germane Germane 14 SnH4 and stananes Stanane Stanane Electron-Rich 15 NH3 Ammonia Azane 15 PH3 Phosphine Phosphane 15 AsH3 Arsine Arsane 15 SbH3 Stibine Stibane 16 H2O Water Oxidane 16 H2S Hydrogen sulf ide Sulfane 16 H2Se Hydrogen selenide Sellane 16 H2Te Hydrogen telluride Tellane 17 HF Hydrogen fluoride Hydrogen fluoride 17 HCl Hydrogen chloride Hydrogen chloride 17 HBr Hydrogen bromide Hydrogen bromide 17 HI Hydrogen iodide Hydrogen iodide

The electron-deficient hydrides are those that cannot complete an octet of electrons around the central atom. They are chiefly the Group 13 elements, although the gas-phase only species BeH2 also fits this description. The electron precise compounds are those that have an octet of electrons, while the electron-rich elements have additional electrons belonging to the central atom that function as lone pairs. Note that although they are all electron rich, this group of compounds vary greatly in the availability of these extra electrons.

Boron Hydrides Boron can form many different types of hydrides. One of them is borane (BH3). They react violently with air and is easily oxidized. Borane occurs as a gaseous substance, and can form B2H6by two borane molecules combined with each other. Borane is not a stable compound because it does not follow a complete octet rule since it has six valence electrons. Synthesizing organic compounds require boron hydrides, therefore they are significant in our daily lives. Note that while the structures of these hydrides follow similar patterns, i.e. all the EH3 in group 15 are pyramidal , the detailed structures differ significantly. Consider the following table of bond lengths for both groups 15 and 16:

Group 15 hydrogen compounds Bond angle Group 15 hydrogen compounds Bond angle

NH3 106.6° H2O 104.5° PH3 93.8° H2S 92.1°

AsH3 91.8° H2Se 91° SbH3 91.3° H2Te 89°

Group 15: phosphine, PH3 - The electronegativity of P is 2.1, the same as hydrogen, therefore PH3 is a non-polar molecule with a boiling point similar to the equally non-polar ethane.

NOTES FOR CHEMISTRY 11/12 – I: Matter and Properties 18

Some of these changes can be explained

using molecular orbital methods. The

bending of H–E–H systems is due to a Jahn-

Teller effect that lowers the energy of what

would be a degenerate set of E π orbitals in

the linear molecule. The greater bending in

the third and subsequent periods is due to

a second-order Jahn-Teller effect, and is a

consequence of the greater similarity in the

atomic orbital energies of hydrogen and

elements such as sulfur and phosphorus,

compared to the large difference that exists

in water and ammonia.

Jahn-Teller Effect – Jahn-Teller Distortion The Jahn-Teller effect correlates electronic structure with a geometric distortion of a molecule. The original Jahn-Teller effect, also called first-order Jahn-Teller distortion, is frequently encountered in transition-metal chemistry and frequently leads to tetragonal distortions of octahedral complexes. We can distinguish between two types of Jahn- Teller distortions:

(i) Jahn-Teller distortion (first-order Jahn-Teller distortion): Any non-linear molecule with an incompletely and unevenly filled (e.g., one or three electrons in a double degenerate set of orbitals), degenerate MO level (degenerate electronic ground state) will undergo a structural distortion that will remove the degeneracy.

(ii) second-order distortion: A second-order Jahn-Teller distortion arises from the mixing of an evenly filled degenerate set of MOs with an empty MO, resulting in a structural distortion that lowers the filled molecular orbital. The original electronic state is non-degenerate. The requirements for a second-order Jahn-Teller distortion are (a) a small energy difference between these two MOs and (b) the correct symmetry of these two orbitals that allow for such a mixing/distortion.

An important property of the electron-rich hydrogen compounds is that of

hydrogen bonding. In the figure at right, the boiling points at one atmosphere

of the hydrogen compounds of groups 14 through 17 are graphed in a

comparative manner. We can take the electron-precise series from group 14 as a

reference point in that no hydrogen bonding can occur for these elements.

Indeed, the boiling-point curve for this class of compound is a smooth curve,

with increasing bp as the mass increases. Compared to these, all the

electron-rich compounds from groups 15, 16 and 17 have higher boiling points,

indicative of strong dipole forces operating in these polar molecules. However,

the unusual displacements to higher values of H2O, HF, NH3 and HCl are

indicative of additional inter-molecular forces, and these have been identified as

hydrogen bonding. Hydrogen bonds have energies considerably smaller than

those of the equivalent covalent bonds, but are still significantly strong. The

table below presents some comparisons between H- bonds and E–H

covalent bond energies, and also provides metrical data for the short

contacts that are typical for hydrogen bonding.

Hydrogen bond E––E distance, pm ∑v.d.Waals radii, pm Energy (kJ/mol) Covalent Bond Energy (kJ/mol) HS–H⋅⋅⋅SH2 370 7 S–H 363

H2N–H⋅⋅⋅NH3 294 – 315 300 17 N–H 386 HO–H⋅⋅⋅OH2 248 – 290 280 22 O–H 464 F–H⋅⋅⋅ F–H 245 – 249 270 29 F–H 565

HO–H⋅⋅⋅Cl– 295 – 310 320 55 Cl–H 428

[F⋅⋅⋅H⋅⋅⋅F]– 227 270 165 H–F 565

Group 15: Nitrogen Hydrides – Ammonia, a polar molecule, is an important nitrogen hydride that is possible due to the synthesis of nitrogen and water which is called the Haber-Bosch process. The chemical equation for this reaction is: N2(g) + 3H2(g) → 2NH3(g) In order to yield ammonia, there needs to be a catalyst to speed up the reaction, a high temperature and a high pressure. Ammonia is a reagent used in many chemistry experiments and is used as fertilizer.

Ammonia can react with sulfuric acid to produce ammonium sulfate, which is also an important fertilizer. In this reaction, ammonia acts as a base since it receives electrons while sulfuric acid gives off electrons. 2NH3(aq) + H2SO4(aq) → (NH4)2SO4(aq) Other hydrides of nitrogen include ammonium chloride, hydrazine and hydroxylamine. Ammonium chloride is widely used in dry-cell batteries and clean metals.

NOTES FOR CHEMISTRY 11/12 – I: Matter and Properties 19

Group 16: Hydrogen sulphide, H2S - The electronegativity of S is 2.5 so the S-H bond is slightly polar and there is a small increase in boiling point compared to the non-polar ethane and phosphine and has a higher boiling point than hydrogen chloride with its more polar bond. This may be due to the fact that there are two S-H bonds and the electrons are more spread out and more polarisable.

Group 17: Hydrogen chloride, HCl - Similarly, in the case of the obviously polar hydrogen chloride molecule, the presence of the permanent dipole has virtually no effect on the boiling point compared to what you might expect for an 18 electron non-polar molecule. This suggests that the intermolecular forces operating in the first three molecules all have their origin in the instantaneous dipole - induced dipole forces, despite the picture below!

attractions seem to make little difference to the bpt! Summary of Trends and Properties of Hydrogen Compounds 1. No hydrogen bond occurs for hydrides in the group 14 which includes C, Si, Ge. The boiling-point curve for this class of

compounds is a smooth curve, with increasing bp as the mass increases. These are non-polar compounds because the elements in this group has about the same electronegativity as that of hydrogen.

2. All the electron-rich compounds from groups 15, 16 and 17 have higher boiling points, indicative of strong dipole forces operating in these polar molecules because these elements are much more electronegative than hydrogen.

3. The unusual displacements to higher values of H2O, HF, and NH3, thus breaking the trend of increasing boiling points as the mass (or the # of electrons) increases, are indicative of additional inter-molecular forces, and these have been identified as hydrogen bonding.

4. The curve of increasing boiling points generally moves up from group 14 to group 17. However, the compounds of group 16 generally have higher boiling than those in group 17 because the compounds in group 16 have one more hydrogen and are more polar than those group 17 with only one hydrogen.

Practice Problems 1. Which group does this hydride belong to, based on the below ? NH3(g) is a hydride and is bonded by a union between the two elements, in which the hydrogen is attracted to a more electropositive element, and its bonding is done by sharing electrons between atoms, and or other atoms and other "sharing electron bonds". Solution: Covalent Hydride

2. Which type of hydrides describes the hydride that often bond between hydrogen and metals with alkaline earth metals and alkali metals? Solution: Ionic Hydrides

3. Can a hydride be a binary compound? Solution: Yes

4. Name one nitrogen hydride. NH3, hydrazine and hydroxylamine are all accepted.

5. True/False: Ionic hydrides react slowly with water. Solution: False

NOTES FOR CHEMISTRY 11/12 – I: Matter and Properties 20

(B) To understand Chemical Reactions, it is first necessary to have an understanding of how chemical compounds are formed:

Chemical bonds are the attractive forces that hold atoms together in more complex units. There are two types of chemical bonds: (I) Ionic Bond – results from the transfer of one or more electrons from one group of atoms to another. An ion is an atom (or

group of atoms) that is electrically charged as a result of the loss or gain of electrons. This bond model is particularly applicable for compounds containing atoms of metallic and nonmetallic elements. When elements having low ionization energies (metals) react with elements having high electron affinity (nonmetals), ionic compounds (bases and salts) form.

(II) Covalent Bond – results from the sharing of one or more pairs of electrons between atoms. This bond model is most applicable for compounds (acids) containing only atoms of nonmetal elements. Covalent bonds are formed between similar or even identical atoms.

Lewis Electron-Dot Structure – which is the shorthand system for designating the number of valence electrons, consists of an element’s symbol with one dot for each valence electron placed around the elemental symbol.

There are three important generalizations about valence electrons: 1. Representative elements in the same group of the periodic table have the same number of valence electrons. 2. The number of valence electrons for representative elements in a group is the same as the Roman numerical periodic table group number. 3. The maximum number of valence electrons for any element is eight.

The Octet Rule – In compound formation, atoms of elements lose, gain, or share electrons in such a way as to produce a noble-gas electron configuration for each of the atoms involved. All nonmetals and most representative element metals (all Group 1A, Group IIA metals, and Ag, Zn, Cd, Al, Ga) follow the octet rule and form only one type of ion. However, there are exceptions to this rule. There are many other metals that exhibit a less predictable behaviour and are able to form more than one type of ion.

Exceptions (the atoms in some molecules cannot obey the octet rule because there are either two few or too many electrons): (i) Atoms with less than an octet: Many compounds of boron and beryllium do not follow the octet rule. Example: boron trifluoride (BF3 ) – only 6 valence electrons surround the boron atom; (BeCl2) – 4 e- around Be. Boron trifluoride, a gas at

normal temperatures and pressures, reacts very energetically with molecules such as water and ammonia that have unshared electron pairs (lone pairs). In the molecules ofBF3 , each of the fluorine atom form a single bond with the boron atom in the centre. In this structure, the boron

atom has only six electrons around it. This electron deficiency causesBF3 to react vigorously with electron-rich molecules such as3NH to

form33NBFH . It is also characteristic of beryllium to form molecules where the beryllium atom is electron-deficient.

(ii) Atoms with more than an octet: Some atoms beyond the 2nd row of the periodic table – most notably phosphorus and sulfur- sometimes form bonds that give them more than an octet of electrons. The additional electrons fill the 3d orbitals of these atoms.

Examples: the sulfur atom in SF4 - 10 valence electrons around the sulfur atom. Other examples: 6SF (12 e- around S), 5PCl (10 e- around

P). (iii) Molecules with an odd number of electrons cannot follow the octet rule. Example: nitrogen monoxide (NO) has a total of 11 electrons. The

“odd” electron is usually on the nitrogen atom. So nitrogen monoxide is very unstable and very reactive. But it is a vital messenger molecule in the human body. Same situation applies to nitrogen dioxide.

(iv) The elements C, N, O, and F obey the octet rule in the vast majority of their compounds. However, there is one important exception is the oxygen molecule

2O . The octet rule of Lewis structure would be satisfied if the two oxygen atoms bond in a double bond. But experiment has

shown that oxygen is paramagnetic, meaning that it contains unpaired electrons. No simple Lewis structure can explain this Paramagnetism.

(I) Compounds that are formed by ionic bond are called ionic compounds: 1. Ionic Compounds usually contain both a metallic and a nonmetallic element. 2. The metallic element atoms lose electrons to produce positive ions and the non-metallic elements atoms gain electrons to

produce negative ions. 3. The electrons lost by the metal atoms are the same ones that are gained by the nonmetal atoms. Electron loss must always equal

electron gain. 4. The ratio in which positive metal ions and negative nonmetal ions combine is the ratio that achieves charge neutrality for the

resulting compound. 5. Metals from Groups IA, IIA, and IIIA of the periodic table form ions with charges of +1, +2, and +3, respectively. Nonmetals of

Groups VIIA, VIA, and VA of the periodic table form ions with charges of –1, -2, and –3, respectively. When the representative elements form ions they tend to achieve noble gas configuration having 8 electrons in their outer shell.

6. All ionic compounds are solids at room temperature and tend to have high melting points. 7. Their solids are generally hard and brittle. 8. In solid state, ionic compounds do not conduct electricity because the attractive forces prevent the movement of ions through the

crystal. But the liquid conducts electricity well. Ionic solids consist of positive and negative ions arranged in such a way that each ion is surrounded by the nearest neighbours of the opposite charge. Discrete molecules do not exist in an ionic solid. Therefore, the formula for these solids cannot represent the composition of a molecule of the substance. The formulas for ionic solids only represent ratios and they are used in equations and in chemical calculations such as molecular weights in the same way as the formulas for molecular species. Thus, the molecule is not the smallest unit capable of a stable existence for all pure substance. Ionic compounds without the OH--1 are also known as salts.

NOTES FOR CHEMISTRY 11/12 – I: Matter and Properties 21

(II) Compounds that are formed by covalent bond are called covalent or molecular compounds. The formation of a covalent bond always involves the process of electron sharing where the octet rule and electron-dot structures apply. The number of covalent bonds that an atom forms is equal to the number of electrons it needs to achieve a noble-gas configuration with 8 electrons in its outer shell, the valence shell. The octet rule can also be used to predict formulas in covalent compounds. A single-covalent bond is a bond where a single pair of electrons is shared between two atoms. A double-covalent covalent bond is a bond where two atoms share two pairs of electrons. A triple-covalent bond is a bond where two atoms share three pairs of electrons. A double-covalent bond is stronger than a single-covalent bond, but not twice as strong, because two electron pairs repel each other and cannot become fully concentrated between the two atoms. A triple-covalent bond is not triple in bonding strength for the same reason. Not all elements can form double- or triple-covalent bonds. There must be at least two vacancies in an atom’s valence electron shell prior to bond formation. This requirement eliminates Group VIIA elements (fluorine, chlorine, bromine, iodine) and hydrogen from participating in such bonds.

Examples: (i) Triple-covalent bond - A diatomic N 2molecule, the natural form in which nitrogen occurs in the atmosphere, is the simplest

triple-covalent bond: ••

••••••

••

••

••≡N N or N N

(ii) Triple-covalent bond - C H2 2 has a carbon-carbon triple-covalent bond and two carbon-hydrogen single bonds:

H C C H or H C C H••

••••••

•• − ≡ −

(iii) Triple-covalent bond – HCN has a heteroatomic carbon-nitrogen triple bond: H C N or H C N••

••••••

•• − ≡

(iv) Double-covalent bonds are found in numerous molecules – a common molecule is carbon dioxide (CO2 ), where two

carbon-oxygen double bonds are present: ••

••••••

••••

••

••

••

••

••

••=O C O or O = C O

A coordinate-covalent bond is a bond in which both electrons of a shared pair come from one of the two atoms involved in the bond. Examples:

(i) N O2 - ••

••••••

••

••

••••N N O in which the nitrogen-nitrogen triple bond is a normal covalent bond; the nitrogen-oxygen bond is a

coordinate- covalent bond , where both electrons are supplied by the nitrogen atom. (ii) CO - •

•••••••

••C O in which four of the six electrons in the triple carbon-oxygen bond can be considered to have come from the

oxygen atom. (iii) Another example of a coordinate-covalent bond occurs when a molecule having an incomplete valence shell reacts with a

molecule having electrons that aren’t being used in bonding. Compounds like [ 33NHBCl ] which are formed by simply joining

two smaller molecules, are called addition compounds:

The ionic and covalent models for bonding are actually closely related to each other and represent the extremes of a broad continuum of bonding patterns. Their close relationship can be explained by the concept of electronegativity, which is a measure of the relative attraction that an atom has for the shared electrons in a bond. The higher the electronegativity value for an element is, the greater the electron-attracting ability of atoms of that element for shared electrons in bonds. Electronegativity increases from left to right across a period of the periodic table (in the same direction as the atomic size decreases), and decreases from top to bottom in a group of the periodic table (in the same way as the atomic size increases). Generally, nonmetals have higher electronegativities than metals, consistent with the fact that metals tend to lose electrons and nonmetals tend to gain electrons when an ionic bond is formed. That is why an element from Group IA or Group IIA reacts with an element from the upper right hand corner of the periodic table to form predominantly ionic compound. Electronegativity values differ from element to element because of differences in (i) atom size, (ii) nuclear charge, and (iii) number of inner-shell (non-valence) electrons. 100% ionic bonding occurs when the difference in electronegativity is very high such as 3 or higher; the bond is more than 50% ionic when the difference exceeds 1.7. In a nonpolar covalent bond, there is no difference in electronegativity, so the pair of bonding electrons is shared equally. It is important to note that element hydrogen located to the far left in period 1 in the periodic table have an electronegativity of

2.1, which is between boron (2.0) and carbon (2.5), both elements of period 2.

NOTES FOR CHEMISTRY 11 – UNITS I, II. & IV: Matter, Properties, States & Stoichiometry 22

Chemical Bonding – more detailed explanation A chemical bond is an attraction between atoms that allows the formation of chemical substances that contain two or more atoms. The bond is

caused by the electromagnetic force attraction between opposite charges, either between electrons and nuclei, or as the result of a dipole

attraction. The strength of chemical bonds varies considerably; there are "strong bonds" such as covalent or ionic bonds and "weak bonds" such

as dipole-dipole interactions, the London dispersion force and hydrogen bonding.

Since opposite charges attract via a simple electromagnetic force, the negatively charged electrons orbiting the nucleus and the positively

charged protons in the nucleus attract each other. Also, an electron positioned between two nuclei will be attracted to both of them. Thus, the

most stable configuration of nuclei and electrons is one in which the electrons spend more time between nuclei, than anywhere else in space.

These electrons cause the nuclei to be attracted to each other, and this attraction results in the bond. However, this assembly cannot collapse to

a size dictated by the volumes of these individual particles. Due to the matter wave nature of electrons and their smaller mass, they occupy a

very much larger amount of volume compared with the nuclei, and this volume occupied by the electrons keeps the atomic nuclei relatively far

apart, as compared with the size of the nuclei themselves.

In general, strong chemical bonding is associated with the sharing or transfer of electrons between the participating atoms. The atoms in

molecules, crystals, metals and diatomic gases— indeed most of the physical environment around us— are held together by chemical bonds,

which dictate the structure of matter.

Examples of Lewis dot-style chemical bonds between carbon C, hydrogen H, and oxygen O. Lewis dot depictures represent an early

attempt to describe chemical bonding and are still widely used today.

Overview of Main Types of Chemical Bonds and Intramolecular Forces: In the simplest view of a so-called covalent bond, one or more electrons (often a pair of electrons) are drawn into the space between the two

atomic nuclei. Here the negatively charged electrons are attracted to the positive charges of both nuclei, instead of just their own. This

overcomes the repulsion between the two positively charged nuclei of the two atoms, and so this overwhelming attraction holds the two nuclei

in a fixed configuration of equilibrium, even though they will still vibrate at equilibrium position. In summary, covalent bonding involves sharing

of electrons in which the positively charged nuclei of two or more atoms simultaneously attract the negatively charged electrons that are being

shared. In a polar covalent bond, one or more electrons are unequally shared between two nuclei.

In a simplified view of an ionic bond, the bonding electron is not shared at all, but transferred. In this type of bond, the outer atomic orbital of

one atom has a vacancy which allows addition of one or more electrons. These newly added electrons potentially occupy a lower energy-state

(effectively closer to more nuclear charge) than they experience in a different atom. Thus, one nucleus offers a more tightly bound position to an

electron than does another nucleus, with the result that one atom may transfer an electron to the other. This transfer causes one atom to

assume a net positive charge, and the other to assume a net negative charge. The bond then results from electrostatic attraction between atoms,

and the atoms become positive or negatively charged ions.

All bonds can be explained by quantum theory, but, in practice, simplification rules allow chemists to predict the strength, directionality, and

polarity of bonds. The octet rule and VSEPR theory are two examples.

The types of strong bond differ due to the difference in electronegativity of the constituent elements. A large difference in electronegativity leads

to more polar (ionic) character in the bond.

(1) Covalent bond - Covalent bonding is a common type of bonding, in which the electronegativity difference between the bonded atoms is small

or nonexistent. Bonds within most organic compounds are described as covalent. See sigma bonds and pi bonds for LCAO-description of such

bonding.

(2) A polar covalent bond is a covalent bond with a significant ionic character. This means that the electrons are closer to one of the atoms than

the other, creating an imbalance of charge. They occur as a bond between two atoms with moderately different electronegativities, and give rise

to dipole-dipole interactions. The electronegativity of these bonds is 0.3 to 1.7 .

A coordinate covalent bond is one where both bonding electrons are from one of the atoms involved in the bond. These bonds give rise to Lewis

acids and bases. The electrons are shared roughly equally between the atoms in contrast to ionic bonding. Such bonding occurs in molecules such

as the ammonium ion (NH4+) and are shown by an arrow pointing to the Lewis acid. Also known as non-polar covalent bond, the electronegativity

of these bonds range from 0 to 0.3.

Molecules which are formed primarily from non-polar covalent bonds are often immiscible in water or other polar solvents, but

much more soluble in non-polar solvents such as hexane.

The structure of ice

is quite open. Each

water molecule has

only four nearest

neighbors with

which it interacts

by means of

hydrogen bonds

(blue dashed lines)

�

NOTES FOR CHEMISTRY 11 – UNITS I, II. & IV: Matter, Properties, States & Stoichiometry 23(3) Ionic bond - Ionic bonding is a type of electrostatic interaction between atoms which have a large electronegativity

difference. There is no precise value that distinguishes ionic from covalent bonding, but a difference of electronegativity of over

1.7 is likely to be ionic, and a difference of less than 1.7 is likely to be covalent. Ionic bonding leads to separate positive and

negative ions. Ionic charges are commonly between −3e to +3e. Ionic bonding commonly occurs in metal salts such as sodium

chloride (table salt). A typical feature of ionic bonds is that the species form into ionic crystals, in which no ion is specifically paired

with any single other ion, in a specific directional bond. Rather, each species of ion is surrounded by ions of the opposite charge,

and the spacing between it and each of the oppositely charged ions near it, is the same for all surrounding atoms of the same type.

It is thus no longer possible to associate an ion with any specific other single ionized atom near it. This is a situation unlike that in

covalent crystals, where covalent bonds between specific atoms are still discernable from the shorter distances between them, as

measured by with such techniques as X-ray diffraction.

Ionic crystals may contain a mixture of covalent and ionic species, as for example salts of complex acids, such as sodium cyanide,

NaCN. Many minerals are of this type. X-ray diffration shows that in NaCN, for example, the bonds between sodium cations (Na+)

and the cyanide anions (CN-) are ionic, with no sodium ion associated with any particular cyanide. However, the bonds between C

and N atoms in cyanide are of the covalent type, making each of the carbon and nitrogen associated with just one of its opposite

type, to which it is physically much closer than it is to other carbons or nitrogens in a sodium cyanide crystal.

When such salts dissolve into water, the ionic bonds are typically broken by the interaction with water, but the covalent bonds

continue to hold. In solution, the cyanide ions, still bound together as single CN- ions, move independently through the solution, as

do sodium ions, as Na+. These charged ions move apart because each of them are more strongly attracted to a number of water

molecules, than to each other. The attraction between ions and water molecules in such solutions is due to a type of weak

dipole-dipole type chemical bond.

(4) Metallic bond – In a metallic bond, bonding electrons are delocalized over a lattice of atoms. By contrast, in ionic compounds, the

locations of the binding electrons and their charges are static. The freely-moving or delocalization of bonding electrons leads to classical

metallic properties such as shininess (surface light reflectivity), electrical and thermal conductivity, ductility, and high tensile strength.

========================================================================================================

You may skip the (4) to (7) in the following as these are not covered in the Grade 11 Chemistry course:

(4) One- and three-electron bonds - Bonds with one or three electrons can be found in radical species, which have an odd number of electrons.

The simplest example of a 1-electron bond is found in the hydrogen molecular cation, H2+. One-electron bonds often have about half the bond

energy of a 2-electron bond, and are therefore called "half bonds". However, there are exceptions: in the case of dilithium, the bond is actually

stronger for the 1-electron Li2+ than for the 2-electron Li2. This exception can be explained in terms of hybridization and inner-shell effects.

The simplest example of three-electron bonding can be found in the helium dimer cation, He2+, and can also be considered a "half bond" because,

in molecular orbital terms, the third electron is in an anti-bonding orbital which cancels out half of the bond formed by the other two electrons.

Another example of a molecule containing a 3-electron bond, in addition to two 2-electron bonds, is nitric oxide, NO. The oxygen molecule, O2 can

also be regarded as having two 3-electron bonds and one 2-electron bond, which accounts for its paramagnetism and its formal bond order of 2.

Molecules with odd-electron bonds are usually highly reactive. These types of bond are only stable between atoms with similar electronegativities.

(5) Bent bonds - Bent bonds, also known as banana bonds, are bonds in strained or otherwise sterically hindered molecules whose binding orbitals

are forced into a banana-like form. Bent bonds are often more susceptible to reactions than ordinary bonds.

(6) 3c-2e and 3c-4e bonds - In three-center two-electron bonds ("3c-2e") three atoms share two electrons in bonding. This type of bonding occurs

in electron deficient compounds like diborane. Each such bond (2 per molecule in diborane) contains a pair of electrons which connect the boron

atoms to each other in a banana shape, with a proton (nucleus of a hydrogen atom) in the middle of the bond, sharing electrons with both boron

atoms.

Three-center four-electron bonds ("3c-4e") also exist which explain the bonding in hypervalent molecules. In certain cluster compounds, so-called

four-center two-electron bonds also have been postulated.

In certain conjugated π (pi) systems, such as benzene and other aromatic compounds (see below), and in conjugated network solids such as

graphite, the electrons in the conjugated system of π-bonds are spread over as many nuclear centers as exist in the molecule, or the network.

(7) Aromatic bond - In organic chemistry, certain configurations of electrons and orbitals infer extra stability to a molecule. This occurs when π

orbitals overlap and combine with others on different atomic centres, forming a long range bond. For a molecule to be aromatic, it must obey

Hückel's rule, where the number of π electrons fit the formula 4n + 2, where n is an integer. The bonds involved in the aromaticity are all planar.

In benzene, the prototypical aromatic compound, 18 (n = 4) bonding electrons bind 6 carbon atoms together to form a planar ring structure. The

bond "order" (average number of bonds) between the different carbon atoms may be said to be (18/6)/2=1.5, but in this case the bonds are all

identical from the chemical point of view. They may sometimes be written as single bonds alternating with double bonds, but the view of all ring

bonds as being equivalently about 1.5 bonds in strength, is much closer to truth.

In the case of heterocyclic aromatics and substituted benzenes, the electronegativity differences between different parts of the ring may dominate

the chemical behaviour of aromatic ring bonds, which otherwise are equivalent.

========================================================================================================

Intermolecular bonding - There are four basic types of bonds that can be formed between two or more (otherwise

non-associated) molecules, ions or atoms. Intermolecular forces cause molecules to be attracted or repulsed by each other. Often,

these define some of the physical characteristics (such as the melting point) of a substance.

i. A large difference in electronegativity between two bonded atoms will cause a permanent charge separation, or dipole, in a

molecule or ion. Two or more molecules or ions with permanent dipoles can interact in dipole-dipole interactions. The

NOTES FOR CHEMISTRY 11 – UNITS I, II. & IV: Matter, Properties, States & Stoichiometry 24bonding electrons in a molecule or ion will, on average, be closer to the more electronegative atom more frequently than

the less electronegative one, giving rise to partial charges on each atom, and causing electrostatic forces between molecules

or ions.

ii. A hydrogen bond is effectively a strong example of a permanent dipole. The large difference in electronegativities between

hydrogen and any of fluorine, nitrogen and oxygen, coupled with their lone pairs of electrons cause strong electrostatic forces

between molecules. Hydrogen bonds are responsible for the high boiling points of water and ammonia with respect to their

heavier analogues.

iii. The London dispersion force arises due to instantaneous dipoles in neighbouring atoms. As the negative charge of the

electron is not uniform around the whole atom, there is always a charge imbalance. This small charge will induce a

corresponding dipole in a nearby molecule; causing an attraction between the two. The electron then moves to another part

of the electron cloud and the attraction is broken.

iv. A cation-pi interaction occurs between the electron density of pi bonds above and below an aromatic ring and a cation.

Summary of Intramoleculat & Intermolecular Forces & Their Relative Strength (see also page 30): Relative

Strength Force Description/Example Picture

Strongest

Ion - Ion

• Attraction between two oppositely charged ions.

• Forms a crystal lattice structure.

• Strong (but easily dissociated when dissolved in water)

• E.g. ionic salts ( .

Hydrogen

Bonding

• Attraction between an ion and the oppositely charged side of a

polar molecule.

• Ionic salts dissolving in water

• If the molecules have O−−−−H, N−H, F−H bonds, they can form

hydrogen bonds with themselves and with water

• If the molecules contain O, N, or F atoms that are not bonded

to hydrogen atoms, they may accept hydrogen bonds from

water

Dipole – Dipole

• A special type of dipole-dipole attraction that is very strong.

• Occurs between the unshared pair of electrons of an O, N, or F

atom of one molecule with the H atom of an H-O, H-N, or H-F

bond on a nearby molecule.

• E.g. , .

Dipole - Induced

Dipole

• The attraction of the negatively charged end of one polar

molecule to the positively charged end of a nearby polar

molecule

• E.g. , .

Induced Dipole –

Induced Dipole

(London

Dispersion)

• Attraction between a polar molecule and the weak dipole it

temporarily induces on a non-polar species.

• E.g. or dissolved in water.

Weakest

Induced Dipole –

Induced Dipole

(London

Dispersion)

• The attraction between momentary induced dipoles in

non-polar molecules.

• Electrons in bond are always vibrating, thus causing

momentary, uneven distributions of charge (a non-polar

molecule becomes polar for an instant).

• At the instant when this uneven charge distribution exists, the

molecule is capable of inducing a momentary dipole in a

nearby molecule.

• A force of attraction results between the dipoles.

• The dipoles exist for just an instant, but they continue to be

formed in a random but ongoing basis.

• Dispersion forces are weak, but increase in strength as the size

of the molecule increases leading to higher boiling point

• E.g. liquid or hydrocarbons.

NOTES FOR CHEMISTRY 11 – UNITS I, II. & IV: Matter, Properties, States & Stoichiometry 25

NOTES FOR CHEMISTRY 11 – UNITS I, II. & IV: Matter, Properties, States & Stoichiometry 26

Chemical polarity Different Types of Water:

A water molecule, a commonly-used example of polarity. The two charges are present with a negative charge in the middle (red

shade), and a positive charge at the ends (blue shade).

In chemistry, polarity refers to a separation of electric charge leading to a molecule or its chemical groups having an electric dipole

or multipole moment. Polar molecules interact through dipole–dipole intermolecular forces and hydrogen bonds. Molecular

polarity is dependent on the difference in electronegativity between atoms in a compound and the asymmetry of the compound's

structure. For example, a molecule of water is polar because of the unequal sharing of its electrons in a "bent" structure, whereas

methane is considered non-polar because the carbon shares the electrons with the hydrogen atoms almost uniformly. Polarity

underlies a number of physical properties including surface tension, solubility, and melting- and boiling-points.

Hard water is water that contains dissolved metals ions (principally ) that

form insoluble compounds (precipitates) either with soap or upon heating.

Soft water is water in which the ions that cause hardness have been removed or tied up

chemically. A popular method for softening water is through a process called ion exchange

using a naturally occurring material called zeolites, which are complex sodium aluminum

silicates ( represents the hard water ions):

Deionized water is softened water in which dissolved ions (both positive and negative) have

been removed. In the production of deionized water, the hard water is first passed through

an “acidic” ion-exchange material that replaces all positive ions with ions. Then the

water is passed through a “basic” ion-exchange material that replaces all negative ions with

ions. The ions and ions now present in the water immediately react with

each other to produce additional water.

Fluoridated water is water with ions added by adding NaF in the water to reduce

dental caries. The compound hydroxyapatite ( is the major consitituent

of teeth, Ions associated with this compound, such as ion, are readily replaced by

ions. The resulting fluoride-containing material is less soluble and less reactive than the

non-fluroride material is. This decreased solubility and decreased reactivity afford greater

protection against caries. Fluoridated toothpaste is much less effective than fluoridated

drinking water in reducing the incidence of tooth decay. However, many fluorine-containing

compounds (at concentration much greater than in fluoridated drinking water) exhibit

toxic effects.

Hydrogen bonds (dashed lines) among water molecules.

The solution process for an ionic solid in water �

NOTES FOR CHEMISTRY 11 – UNITS I, II. & IV: Matter, Properties, States & Stoichiometry 27Theory - Electrons aren't always shared equally between two bonding atoms: one atom might exert more of a force on the

electron cloud than the other. This "pull" is termed electronegativity and measures the attraction for electrons a particular atom

has. The unequal sharing of electrons within a bond leads to the formation of an electric dipole: a separation of positive and

negative electric charge. Fractional charges are denoted as δ+ (delta plus) and δ− (delta minus). Atoms with high

electronegativities — such as fluorine, oxygen, and nitrogen — exert a greater pull on electrons than atoms with lower

electronegativities. In a bonding situation this can lead to unequal sharing of electrons between atoms, as electrons will spend

more time closer to the atom with the higher electronegativity.

Bonds can fall between one of two extremes — being completely non-polar or completely polar. A completely non-polar bond

occurs when the electronegativities are identical and therefore possess a difference of zero. A completely polar bond is more

correctly termed ionic bonding and occurs when the difference between electronegativities is large enough that one atom takes an

electron from the other. The terms "polar" and "non-polar" bonds usually refer to covalent bonds. To determine the polarity of a

covalent bond using numerical means, the difference between the electronegativity of the atoms is taken. If the result is between

0.4 and 1.7 then, generally, the bond is polar covalent.

Polarity of molecules - While the molecules can be described as "polar covalent", "non-polar covalent", or "ionic", it must be

noted that this is often a relative term, with one molecule simply being more polar or more non-polar than another. However, the

following properties are typical of such molecules.

A molecule is composed of one or more chemical bonds between molecular orbitals of different atoms. A molecule may be polar

either as a result of polar bonds due to differences in electronegativity as described above, or as a result of an asymmetric

arrangement of non-polar covalent bonds and non-bonding pairs of electrons known as a full molecular orbital.

Polar molecules - Examples of common household polar molecules include sugar, for instance the sucrose sugar variety. Sugars

have many polar oxygen–hydrogen (-OH) groups and are overall highly polar. Due to the polar nature of the water molecule (H2O)

itself, polar molecules are generally able to dissolve in water.

Sucrose, a sugar, has many polar -OH groups Water is a polar solvent

• Example 1. The hydrogen fluoride, HF, molecule is polar by virtue of polar covalent bonds — in the covalent bond

electrons are displaced towards the more electronegative fluorine atom.

Hydrogen fluoride: the more electronegative fluoride atom is shown in yellow. Hydrogen fluoride: red represents partially

negatively charged regions

• Example 2. In the ammonia, NH3, molecule the three N–H bonds have only a slight polarity (towards the more

electronegative nitrogen atom). However, the molecule has two lone electrons in an orbital, that points towards the

fourth apex of the approximate tetrahedron, (VSEPR). This orbital is not participating in covalent bonding; it is electron

rich which results in a powerful dipole across the whole ammonia molecule.

Ammonia: the two lone electrons are shown in yellow, the hydrogen atoms in white. Ammonia: red represents partially

negatively charged regions

• Example 2.5. In the ozone, O3, molecule the two O–O bonds are non-polar (there is no electronegativity difference

between atoms of the same element). However, the distribution of other electrons is uneven — since the central atom

has to share electrons with two other atoms, but each of the outer atoms only have to share electrons with one other

atom, the central atom is more deprived of electrons than the others (the central atom has a formal charge of +1, while

the outer atoms each have a formal charge of −1/2). Since the molecule has a bent geometry, this results in a dipole

across the whole ozone molecule.

Ozone

NOTES FOR CHEMISTRY 11 – UNITS I, II. & IV: Matter, Properties, States & Stoichiometry 28

Non-polar molecules

Diagram showing the net effect of symmetrical polar bonds (direction of yellow arrows show the migration of electrons) within boron trifluoride

cancelling out to give a net polarity of zero. δ- shows an increase in negative charge and δ+ shows an increase in positive charge.

A molecule may be non-polar either because there is (almost) no polarity in the bonds (when there is an equal sharing of electrons

between two different atoms) or because of the symmetrical arrangement of polar bonds. Examples of household non-polar

compounds include fats, oil and petrol/gasoline. Therefore (per the "oil and water" rule of thumb), most non-polar molecules are

water insoluble (hydrophobic) at room temperature. However many non-polar organic solvents, such as turpentine, are able to

dissolve polar substances. When comparing a polar and non-polar molecule with similar molar masses, the polar molecule

generally has a higher boiling point, because of the dipole–dipole interaction between their molecules. The most common form of

such an interaction is the hydrogen bond, which is also known as the H-bond.

• Example 3. In the methane molecule (CH4) the four C–H bonds are arranged tetrahedrally around the carbon atom. Each bond

has polarity (though not very strong). However, the bonds are arranged symmetrically so there is no overall dipole in the

molecule.

Methane: the bonds are arranged symmetrically so there is no overall dipole. Boron trifluoride: trigonal planar arrangement of

three polar bonds results in no overall dipole

• Example 4. The boron trifluoride molecule (BF3) has a trigonal planar arrangement of three polar bonds at 120o. This results in no

overall dipole in the molecule.

• Example 5. The oxygen molecule (O2) does not have polarity in the covalent bond because of equal electronegativity, hence

there is no polarity in the molecule.

Hybrids - Large molecules that have one end with polar groups attached and another end with non-polar groups are good

surfactants. They can aid in the formation of stable emulsions, or blends, of water and fats. Surfactants reduce the interfacial

tension between oil and water by adsorbing at the liquid–liquid interface.

Predicting molecule polarity (Summary Diagram on the following page)

• This classification table gives a good general understanding of predicting molecular dipole of some general molecular

structures, however one should not interpret it literally:

Formula Description Example

Polar AB Linear Molecules CO

HAx Molecules with a single H HF

AxOH Molecules with an OH at one end C2H5OH

OxAy Molecules with an O at one end H2O

NxAy Molecules with an N at one end NH3

Non-polar A2 Diatomic molecules of the same element O2

CxAy Most carbon compounds CO2

• Determining the point group is a useful way to predict polarity of a molecule. Generally, a molecule will not possess dipole

moment, if the individual bond dipole moments of the molecule cancel each other out. This is because dipole moments

are euclidean vector quantities with magnitude and direction, and a two equal vectors who oppose each other will cancel

out.

Any molecule with an centre of inversion ( "i" ) or a horizontal mirror plane ( "σh ") will not possess dipole moments. Likewise, a

molecule with more than one Cn axis will not possess dipole moment because dipole moments can't lie in more than one

dimension. As a consequence of that constraint, all molecules with D symmetry (Schönflies notation) will therefore not have dipole

moment because, by definition, D point groups have two or multiple Cn axis. Since C1, Cs,C∞h Cn and Cnv point groups do not have

a centre of inversion, horizontal mirror planes or multiple Cn axis, molecules in one of those point groups will have dipole moment.

Solubility - A popular aphorism used for predicting solubility is "like dissolves like". This statement indicates that a solute will dissolve

best in a solvent that has a similar chemical structure to itself. This view is simplistic, but it is a useful rule of thumb. The overall solvation

capacity of a solvent depends primarily on its polarity. For example, a very polar (hydrophilic) solute such as urea is very soluble in highly

polar water, less soluble in fairly polar methanol, and practically insoluble in non-polar solvents such as benzene. In contrast, a non-polar

NOTES FOR CHEMISTRY 11 – UNITS I, II. & IV: Matter, Properties, States & Stoichiometry 29or lipophilic solute such as naphthalene is insoluble in water, fairly soluble in methanol, and highly soluble in non-polar benzene. The

solubility is favored by entropy of mixing and depends on enthalpy of dissolution and the hydrophobic effect. Synthetic chemists often

exploit differences in solubility to separate and purify compounds from reaction mixtures, using the technique of liquid-liquid extraction.

Following the aphorism, "like dissolves like", ionic compounds dissolve in polar solvents, especially those that ionize, such as water

and ionic liquids. They are usually appreciably soluble in other polar solvents such as alcohols, acetone and dimethyl sulfoxide as

well. Ionic compounds tend not to dissolve in nonpolar solvents such as diethyl ether or petrol.

When the oppositely charged ions in the solid ionic lattice are surrounded by the opposite pole of a polar molecule, the solid ions are pulled out

of the lattice and into the liquid. When this force is more than the electrostatic attraction of the lattice, the ions become dissolved in the liquid.

Physical properties of Ionic Compounds:

Physical Properties Ionic Compounds States(at room temperature) Solid

Electrical conductivity Solid: No Liquid: Yes Aqueous: Yes (if soluble) Boiling point and Melting Point High

Solubility in water Often high Thermal conductivity Low

Physical properties of covalent compounds:

Physical Properties Covalent Compounds

States(at room temperature) Solid, liquid, gas

Electrical conductivity Usually none

Boiling point and Melting Point Variable (usually lower than ionic compounds)

Solubility in water Variable (usually lower than ionic compounds)

Thermal conductivity Usually low

NOTES FOR CHEMISTRY 11 – UNITS I, II. & IV: Matter, Properties, States & Stoichiometry 30

Covalently bonded groups of atoms are called molecules. Compounds of this type are called covalent or molecular compounds. Molecules, as well as bonds, can have polarity . Molecular polarity depends on the polarity of the bonds within a molecule and the geometry of the molecule. Molecular geometry describes the way in which atoms in a molecule are arranged in space relative to each other. Within the individual molecules the atoms are held to each other very strongly, but between neighboring molecules the attractions are very weak. So molecular compounds such as water and candle wax tend to have low melting points. Molecules are uncharged particles, so they do not conduct electricity in the solid state or when melted. Most molecular substances also will not conduct electricity when dissolved in water. Resonance: When Lewis Structures Fail – The bonding in some molecules and ions cannot be adequately described by a single Lewis structure. There are some molecules and ions for which Lewis structures do not agree with experimental measurements of bond length and bond energy. For example, the structure for −

2CHO suggests that one carbon-oxygen bond should be longer

than the other, but experiment shows that they are identical. In fact, the OC − bond lengths are about halfway between that expected for a single bond and that expected for a double bond. We get around this problem by introducing the concept of resonance, that is, the actual structure of the ion is a resonance hybrid of two contributing structures as illustrated in the following diagrams:

Extending the Lewis Theory of Bonding by Nevil Sidgwick: (1) One atom could contribute both electrons that are shared; (2) An octet of electrons around an atom may be desirable, but is not necessary in all molecules and polyatomic ions.

NOTES FOR CHEMISTRY 11 – UNITS I, II. & IV: Matter, Properties, States & Stoichiometry 31

NOTES FOR CHEMISTRY 11 – UNITS I, II. & IV: Matter, Properties, States & Stoichiometry 32 A nonpolar-covalent bond is one in which there is an equal sharing of electrons when two identical atoms (atoms of equal electronegativity) share one or more pairs of electrons. A polar-covalent bond is one in which there is unequal sharing of bonding electrons. The significance of a polar-covalent bond is that it creates partial positive and negative charges on atoms. Most chemical bonds are not 100% covalent (equal sharing) or 100% ionic (no sharing). Instead, most bonds are somewhere in between (unequal sharing) as shown in the diagram in previous page. Rule 1: When there is zero difference in electronegativity between bonded atoms, the bond is called a nonpolar-covalent bond. Rule 2: When the electronegativity difference between bonded atoms is greater than zero but less than 1.7, the bond is called a

polar- covalent bond. Rule 3: When the difference in electronegativity between bonded atoms is 1.7 or greater, the bond is called an ionic bond. The shape of a molecule and the polarity of its bonds together determine whether the molecule is polar or nonpolar:

(a) 2CO (b) 4XeF

Not every structure containing lone pairs on the central atom produces polar molecules.

NOTES FOR CHEMISTRY 11 – UNITS I, II. & IV: Matter, Properties, States & Stoichiometry 33

States of Matter

A. Solids – Many substances form crystalline solids, those with a regular arrangement of their components. The structure of a

crystalline solid can be described by choosing a small portion of the structure as a representative unit called a unit cell. There

are three main types of crystalline solids:

i) ionic solids (components are ions) such as sodium chloride – stable substances with high melting points that are held

together by the strong forces that exist between oppositely charged ions. Many chemical reactions of ionic substances

occur in water. This can sometimes cause ionic crystals to include water molecules in their solid structure thus forming a

hydrate, example: CuSO H O4 25• (copper sulfate pentahydrate).

ii) molecular solids (components are molecules) such as sucrose – substances that tend to melt at relatively low

temperatures because the intermolecular forces (dipole-dipole forces for solids with dipole moment, and London dispersion

forces for nonpolar molecules) that exist among the molecules are relatively weak.

iii) atomic solids (components are atoms) such as graphite and diamond – The properties vary greatly because of the different

ways in which the fundamental particles, the atoms, can interact with each other. For example,

a) The solids of the Group VIII elements have very low melting points because these atoms, having filled valence orbitals, cannot

form covalent bonds with each other. So the forces in these solids are the relatively weak London dispersion forces.

b) On the other hand, diamond, a form of solid carbon, is one of the hardest substances and has an extremely high melting point

(about 35000C). These properties are due to the very strong covalent carbon-carbon bonds in the crystal, which lead to a giant

molecule. Diamond having a giant molecule is not classified as a molecular solid because by convention molecular solids refer to

those with small molecules only. Rather, substances such as diamond that contain giant molecules are called covalent-network

solids (this may be viewed as the 4th

category of solids).

c) Metals represent another type of atomic solids. The shapes of most pure metals can be changed relatively easily; they are durable

and have high melting points. These properties are due to the fact that the bonding in most metals is strong but non-directional.

In addition to ionic bonds, covalent bonds, the third category of bonding is called metallic bonds, which hold atoms together in a

metallic substance. Metallic bonds are the attractive forces between fixed ions and the moving valence electrons in a metal.

Because of the strong metallic bonds, most metals are solids at room temperature, with mercury one familiar exception. Cesium

and gallium have melting points slightly above room temperature.

In the electron sea model, a metal solid is pictured as consisting of regular arrays of atoms in a “sea” of valence electrons that are

shared among the atoms in a non-directional way and that are quite mobile in the metal crystal. The mobile electrons can conduct heat

and electricity, and the atoms can move rather easily, as, for example, when the metal is hammered into a sheet or pulled into a wire.

Because of the nature of the metallic crystal, other elements can be introduced relatively easily to produce substances called

alloys, which are substances that contain a mixture of elements and have metallic properties. Two common types of alloys: 1) In a

substitutional alloy, some of the host metal atoms are replaced by other metal atoms of similar sizes: i) brass – 1/3 copper atoms

replaced by zinc atoms; ii) sterling silver – 93% silver, 7% copper; iii) pewter – 85% tin, 7% copper, 6& bismuth, & 2% antimony. 2)

An interstitial alloy is formed when some of the interstices (holes) among the closely packed metal atoms are occupied by atoms

much smaller than the host atoms: i) steel – contains carbon atoms in the “holes” of an iron crystal. Alloy steels can be viewed as being

mixed interstitial (carbon) substitutional (other metals) alloys: stainless steel with chromium and nickel atoms replacing iron atoms.

Type of Solids

1. Types of Particles

2. Forces between Particles 3. Properties 4. Examples

1. Ionic + ve and - ve ions

electrostatic attractions

hard, brittle, high melting point, poor electrical and thermal conductivity

typical salts - NaCl KBr MgSO, , 4

2. Molecular (Covalent)

atoms or molecules

hydrogen bond, dipole-dipole, London dispersion

soft, low to moderately high melting point, poor electrical and thermal conductivity

most organic compounds such as CH C H O4 12 22 11, as well

as many inorganic compounds such as SO CO H O4 2 2, ,

3. Metallic atoms metallic bond soft to hard, low to high melting point, excellent electrical and thermal conductivity, malleable, ductile

all metallic elements – for examples: Al, Cu, Na, Ag, Fe

4. Covalent- Network

atoms covalent bonds very hard, very high melting point, often poor electrical and thermal conductivity

diamond, C; graphite, C; silicon, Si; quartz, SiO2

Some substances are rigid and appear solid but do not behave like crystalline solids. These are called amorphous solids, examples: glass, rubber, and several plastics. You can think of these substances as liquids that have been cooled to such low temperatures that their viscosities have become very high.

NOTES FOR CHEMISTRY 11 – UNITS I, II. & IV: Matter, Properties, States & Stoichiometry 34

Diamond and graphite are both composed entirely of carbon, but they have different physical properties. This is due to the fact that diamond atoms in carbon are bonded tetrahedrally with single bonds while the atoms in graphite are bonded in a trigonal planar arrangement, where there is one double bond and two single bonds on each carbon. In graphite there are delocalized electrons, wheras the electrons in diamond are all part of bonds. This allows graphite to conduct electricity while diamond cannot. Also, the 3-D tetrahedral arrangement of diamond provides for strength, while the trigonal planar arrangement of graphite produces flat layers, which slip by each other easily. Thus, graphite makes a good lubricant.

Changes of State

1. Vaporization – liquid to gas. 2. Condensation – gas to liquid 3. Freezing – liquid to solid 4. Melting – solid to liquid 5. Sublimation – solid to gas, example: iodine and dry ice (solid carbon dioxide). Solids exhibit vapor pressure just as liquids

do, but generally they are much lower. Solids with high vapor pressures sublime relatively easily. Solids that do not have strong attractive forces between their particles sublime readily. Of the four categories of solids, molecular solids have the weakest forces between their particles. Familiar examples of sublimation are found in this category. The characteristic odor of naphthalene (moth balls) is due to its ease of sublimation. Other fragrant organic compounds that have this property are used in solid air fresheners and deodorizers.

6. Deposition – gas to solid. Sublimation and deposition play a role in weather conditions. Snow can disappear through sublimation even though the temperatures stay below the freezing point. Snow itself is a result of deposition.

NOTES FOR CHEMISTRY 11 – UNITS I, II. & IV: Matter, Properties, States & Stoichiometry 35

B. Gas Physical Properties common to all Gases: � Gases have mases � It is easy to compress gases � Gases filled their containers completely � Different gases can move through each other quite rapidly. The movement of one substance through another is called diffusion. � Gases exert pressure � The pressure of a gas depends on its temperature Gas properties are explained by a kinetic-molecular model that describes the behavior of the submicroscopic particles that make up a gas.

Kinetic Molecular Theory of Gases – a theory that explains the behaviour of an ideal gas. Its postulates: 1) Gases consist of tiny particles (atoms or molecules). 2) These particles are so small, compared with the distances between them, that the volume (size) of the individual

particles can be assumed to be negligible. 3) The particles are in constant random motion, colliding with the walls of the container. These collisions with the

walls cause the pressure exerted by the gas. 4) The particles are assumed not to repel or attract each other. 5) The average kinetic energy of the gas particles is directly proportional to the Kelvin temperature of the gas. 6) Gas particles exert no force on one another. In other words, attractive forces between gas particles are so weak that

the model assumes them to be zero. Unit of Pressure – mm Hg (millimeters of mercury) also called the torr. Another related unit is the standard atmosphere (atm). The SI unit for pressure is the pascal (Pa). A unit of pressure used in engineering science and for measuring tire pressure is pounds per square inch (psi).

1 atm = 760 mm Hg = 760 torr = 101,325 Pa ≈ 105 Pa; 1 atm = 14.69 psi.

Gas Laws

1. Boyle’s Law - PV k k= , is a constant at a specific temperature for a given amount of gas.

2. Charles’s Law - V bT b= , is the proportionality constant

3. Avogadro’s Law - V an a n= , is the proportionality constant, is the number of moles

4. The Ideal Gas Law ( combining above 3 Laws) - KatmLRRnRTPViiP

TnRVi

/ 08206.0 = & constant, gas universal theis )(

= )(

=

For 1 mol of an ideal gas at 0 273o C K ( ) and 1 atm, the volume of the gas given by the ideal gas law is:

VnRT

P

mol L atm K mol K

atm= = =( . )( . / )( )

..

100 0 08206 273

10022 4

L .

The condition 0 273o C K ( ) and 1 atm are called standard temperature and pressure (STP)

Gases do not obey the ideal gas law at high pressures and low temperatures. As pressure is increased to very high levels, compression becomes more and more difficult because the gas particles do have volume of their own. As the volume of the particles themselves becomes a significant proportion of the total volume of a gas, the ideal gas law begins to fail. As decrease in temperature tales away energy from the gas, the gas particles slow down. Attractive forces between gas particles, which are negligibly small when the particles are moving fast, become significant. Thus the ideal gas law also fails.

5. Dalton’s Law of Partial Pressures – The partial pressure of a gas is the pressure that the gas would exert if it were alone in the container: P P P P P P Ptotal = + +1 2 3 1 2 3, , , where are the partial pressures.

Assuming the gases behave ideally, the partial pressures can be calculated as follows:

Pn RT

VP

n RT

VP

n RT

V11

22

33= = =, ,

Then,

++=V

RTn

V

RTnnn

V

RTn

V

RTn

V

RTnP totaltotal =)+(=+ 321

321

The fact that the pressure exerted by an ideal gas is affected by the number of gas particles and is independent of the nature of the gas particles tells us two important things about ideal gases:

1) The volume of the individual gas particle (atom or molecule) must not be very important,

NOTES FOR CHEMISTRY 11 – UNITS I, II. & IV: Matter, Properties, States & Stoichiometry 362) The forces among the particles must not be very important.

Kinetic-Molecular Theory can be applied to liquids and solids as well. According to the kinetic-molecular theory, the state of a substance at room temperature depends on the strength of the attractions between its particles. Substances in which particles are held together by strong forces of attraction are solids at room temperature. Substances in which the attractive forces between particles are of moderate strength are typically liquids at room temperature. And substances in which the forces of attraction between particles are very weak are gases at room temperature.

C. Liquids

Atoms of molecules of gases and liquids are held together by covalent bonds. Liquids and solids are collectively referred to as the condensed states of matter because substances in these states have substantially higher densities. Energy requirement for the changes of state – the bonding forces that hold the atoms of a molecule together are called intramolecular forces. The forces that occur among molecules that cause them to aggregate to form a solid or a liquid are called intermolecular forces. The physical properties, such as viscosity and surface tension, of liquids are determined mainly by the nature and strength of the intermolecular forces. Only intermolecular forces are involved in changes of state. It takes energy to overcome the intermolecular forces in order for the changes of state (solid to liquid, and liquid to gas) to take place. The energy to melt 1 mol of a substance is called the molar heat of fusion. The energy to change 1 mol of liquid to its vapor is called the molar heat of vaporization.

Review of Intermolecular Forces (see pages 24, 25): i) Dipole-Dipole Forces – the forces of attraction between oppositely charged ends of polar molecules (e.g. HCl), typical about 1%

as strong as covalent or ionic bonds.

ii) Hydrogen Bonding - Particularly strong dipole-dipole forces occur between molecules in which hydrogen is bound to a highly electronegative atom, such as nitrogen, oxygen, or fluorine , thus accounting for the high boiling points of their compounds compared to those of hydrogen compounds of other elements in the same group. In this case, a proton is shared between two pairs of electrons, rather than a pair of electrons being shared between protons as in covalent bonds. 2 factors: the great polarity of the bond and the close approach of the dipoles made possible by the small size of the hydrogen atom. This force is given a special name called Hydrogen Bonding.

iii) London Dispersion Forces - the rather weak forces that exist among noble gases and nonpolar molecules caused by instantaneous dipoles, the attraction force between protons and electrons in different molecules, (as different from permanent dipoles). Dipole-Dipole forces and London Dispersion Forces are called van der Waals forces.

Viscosity is the friction or resistance to motion, that exists between the molecules of a liquid when they move past each other. Therefore the greater the attraction force among molecules, the greater the viscosity. Molecules at the surface of the liquid experience attractive forces downward, toward the inside of the liquid, and sideways, along the surface of the liquid. This is unlike the uniformly distributed attractive forces that molecules in the centre of the liquid experience. The imbalance of forces at the surface of a liquid results in a property called surface tension. The uneven forces make the surface behave as if it had a tight film stretched across it. Depending on the magnitude of the surface tension of the liquid, the film is able to support the weight of small objects. Surface tension causes small quantities of a liquid to take on spherical shapes in order to minimize surface area. Surface tension is greater with strong intermolecular forces of attraction.

Vapor Pressure – the pressure of the vapor present at equilibrium with its liquid. This pressure is determined by the intermolecular forces that act upon the molecules.

Water – Some Unusual Properties:

i) The unexpectedly high boiling point (due to its hydrogen bonds among molecules) is why it is a liquid at room temperature. Other hydrogen compounds such asNH H S HF3 2, , , are corrosive gases at room temperature.

ii) Water can absorb or release relatively large quantities of heat without large changes in temperature because it has unusually high specific heat. This is why oceans and lakes exert a moderating influence on climate.

iii) The density of the solid form of water – ice – is less than the density of its liquid form. This is due to the fact that hydrogen bonding is even more extensive in ice than it is in liquid water. Ice has a more open structure because the hydrogen bonds keep other water molecules from getting inside the hexagonal ring structure. Thus there are fewer molecules packed in a given volume, giving ice its lower density. So ice floats on water.

iv) Water has a relatively high surface tension. This property plays an important part in carrying water from the roots to the tops of the tallest trees by a phenomenon called capillary action.

v) Water has a very high heat of vaporization. This property explains the cooling effect of perspiration.

vi) Water is referred to as the universal solvent because of the polar nature of its molecules.

NOTES FOR CHEMISTRY 11 – UNITS I, II. & IV: Matter, Properties, States & Stoichiometry 37

Solutions A solution is a homogeneous mixture of two or more substances in a single physical state.

Type of Solutions Solute Solvent Example 1. Gaseous Solution Gas Gas Air (oxygen and nitrogen) 2. Liquid Solution Gas Liquid Seltzer (carbon dioxide in water) 3. Liquid Solution Liquid Liquid Antifreeze (ethylene glycol in water) 4. Liquid Solution Solid Liquid Ocean water (salt in water) 5. Aqueous Solution Liquid / Solid Water Salt (electrolyte); Sugar (non-electrolyte) 6. Solid Solution Gas Solid Charcoal filter (poisonous gases in carbon) 7. Solid Solution Liquid Solid Dental filling (mercury in silver) 8. Solid Solution (alloy) Solid Solid Sterling silver (copper in silver), an alloy

Some pairs of liquids can mix in any amount. These are said to be miscible in all proportions.

1. Solvent – the substance present in the largest amount; Solute – the other substance.

2. Aqueous solutions – solutions with water as the solvent.

3. Mass Percent = mass of solute

mass of solution×100%=

grams of solute

grams of solute + grams of solvent×100%

4. Concentration of a solution is the amount of solute in a given amount of solvent or solution. The most common measurement of concentration are the following:

a) Molarity (M) = moles of solute

liters of solution= mol

L

b) Molality (m) = moles of solute

kilograms of solvent=

mol

kg

c) Mole Fraction (X) =moles of component

total moles of solution, and X Xsolute solvent+ = 1

d) Formality = (liters) volume

weightsformula

e) H ereplaceabl of no.

massmolar WeightEquivalent = ;

solution of liters weightequivalent

solute of grams Normality

×=

How a Solution Forms?

Intermolecular forces also operate between solute and solvent particles in a solution as in pure substances. Sodium chloride dissolves in water

because the water molecules have a sufficient attraction for theNa+ and Cl− ions – enough to overcome the attraction of these two ions for one another in the crystal. Water molecules orient themselves on the surface of the NaCl crystal so that they can separate, dissociate, the ions and pull

them into solution. Once separated from their crystal, theNa+ and Cl− ions are surrounded by water molecules. The interaction between solute and solvent particles is called solvation. The interaction is called hydration when the solvent is water. The water molecules are also separated from one another to make room for the solute particles, so that the solute and solvent particles are intermingled. So the formation of solution of NaCl in water involves the breaking of attractions among solute particles, the breaking of attractions among solvent particles (endothermic process - energy absorbing), and the formation of attractions between solute and solvent particles (exothermic process – energy releasing). Whether energy, in the form of heat, is absorbed or given off in the overall process depends on the balance between the processes. So, the forming of sodium hydroxide solution is exothermic while that of ammonium nitrate (NH NO4 3 ) is endothermic. Instant cold pack used to

reduce swelling cause by an injury is equivalent to the melting a great deal of ice because when the pack is hit the breaking of attractions in forming this solution absorbs more energy than is released. The heat is absorbed from outside the pack thereby cooling the injured area. In an opposite process, a supersaturated solution ofNa S O2 2 3is used to make instant heat pack. When the pack is squeezed, a crystal ofNa S O2 2 3is

released from a small compartment in the pack. The crystal causes excess solute to come out of solution. The process during which the crystal pulls particles out of solution is exothermic. The heat pack can be recycled by placing it in boiling water to re-dissolve theNa S O2 2 3.

Solubility - The solubility is the amount of a solute that will dissolve in a specific solvent under given conditions, which is the amount of solute required to form a saturated solution, and are usually expressed in grams of solute per 100 grams of solvent at a specified temperature and pressure.

Factors affecting solubility: 1) Nature of solute and solvent: Like Dissolve Like Rule for a solid in a liquid:

Solute Polar solvent Non-polar solvent polar soluble Insoluble nonpolar insoluble Soluble Ionic (similar to polar) soluble Insoluble

NOTES FOR CHEMISTRY 11 – UNITS I, II. & IV: Matter, Properties, States & Stoichiometry 38

Some Polar and Nonpolar Substances

Polar Non-polar

Water ( OH 2 ) In general, greases, petroleum oils, vegetable oils, waxes, tars, gasoline Alcohols Methyl alcohol ( OHCH3 )

Ethyl alcohol ( OHHC 52 )

Isopropyl alcohol ( OHHC 73 )

Hexane ( 146HC )

Acetone ( OHC 63 ) Heptane ( 177HC )

Acetic acid ( 232 OHHC ) Octane ( 188HC )

Formic acid ( 2HCHO ) Carbon tetrachloride ( 4CCl ) Chloroform ( 3CHCl )

Some compounds contain both polar and nonpolar components, yet exhibit properties more like one component than the other. Cholesterol (C H O27 26 ) is such a compound. It is considered nonpolar, because it is insoluble in water and soluble in

nonpolar solvent such as fat tissue.

2) Temperature – Solutions of gases in liquids are greatly affected by changes in temperature. As the temperature increases, the kinetic energy of the solute gas becomes greater. The gas particles acquire more of a tendency to escape from the solvent. Thus as the temperature increases, the solubility of a gas in a liquid decreases. The effect of temperature changes on the solubility of solids in liquids is very different from that for gases. Generally, the solubility of a solid solute increases as the temperature increases. The fact that solubility of a solid solute changes with temperature is the key to preparing supersaturated solutions. To prepare a supersaturated solution, the solution must be heated and then excess solute added. If the solution is then cooled slowly, the extra solute will stay in the solution. Shaking or disturbing a supersaturated solution or adding a tiny crystal of the solid solute can destroy the super-saturation and cause the excess solid solute to crystallize, leaving a saturated solution.

3) Pressure (for gases) – While the solubility of solids and liquids is not appreciably affected by pressure, the solubility of a gas in a liquid is strongly influenced by pressure. According to Henry’s Law , the solubility of a gas was proportional to the partial pressure of the gas above the liquid.

Factors affecting rate of a solid solute dissolving in a solution:

1) Surface area – grinding the solute into smaller particles thereby increasing the surface increases the rate of dissolving. 2) Stirring – similar effect as grinding solute into small particles. 3) Temperature – raising the temperature of a solvent increases the rate at which a solute dissolves, because as temperature

increases, solvent particles move faster. As solvent particles move faster, more particles come into contact with the solute. Colligative Properties: Some physical properties of liquid solutions differ from those of the pure solvent. A property that depends on the concentration of solute particles but is independent of their nature is called colligative property.

1) Vapor pressure reduction – The extent to which a nonvolatile solute lowers the vapor pressure is proportional to its concentration.

2) Boiling point elevation – The difference is directly proportional to the number of solute particles per mole of solvent particles, that is, it is proportional to the molality of the solute: mKT bb =∆ , where bK is the molal boiling point elevation

constant. The value of bK depends on the solvent. This property can be used to determine molar mass:

solvent kg

solute mol=∆=b

bK

Tm , → solvent kgsolute mol ×= m →

solute of massmolar

solute mass=solute mol , →

solute mol

solute mass=massmolar .

(Note: Any of the four colligative properties can be used to determine the molar mass of an unknown substance in this fashion.)

3) Freezing point depression – Like boiling point elevation, the decrease in the freezing point ( fT∆ )is directly proportional

to the to the molality of the solute: mKT ff =∆ .

4) Osmotic pressure – The process that allows a net flow of solvent molecules from the less concentrated solution to the more concentrated solution is called osmosis. The pressure required to prevent osmosis is known as the osmotic pressure ( π ) of the solution. When two solutions with identical osmotic pressure are separated by a semi-permeable membrane, there is no osmosis, and the solutions are said to be isotonic. Fluids administered intravenously to people needing replacement of body fluids must be isotonic with body fluids because the membranes of red blood cells are semi-permeable.

NOTES FOR CHEMISTRY 11 – UNITS I, II. & IV: Matter, Properties, States & Stoichiometry 39If one solution has a lower osmotic pressure than another, it is said to be hypotonic, conversely, the other solution is called hypertonic.