Chemistry UNIT 3

Chemistry UNIT 3

PERIODICITY

3.1 The periodic table3.1 The periodic table

•3.1.1 Describe the arrangement of elements in the periodic 3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic numbertable in order of increasing atomic number

Metalloids are the elements found along the stair-step line that distinguishes metals from non-metals.

This line is drawn from between Boron and Aluminum to the border between Polonium and Astatine.

The only exception to this is Aluminum, which is classified under "Other Metals".

Metalloids

3.1.2 Distinguish between the term group and period

• Group • These are the numbers represented by Roman Numerals

at the top of the periodic table.

• This number tells us the number of valence electrons in an atom.

• Valence electrons are important because they determine the chemical reactivity of elements.

• Period: Represent the number of energy levels /Shells

3.1.1 Apply the relationship between the electron arrangement of elements and their position in the periodic table up to Z = 20

Hydrogen H 1s 

Helium He  1s2

Lithium Li  1s2 2s1

Beryllium Be  1s2 2s2

Boron B  1s2 2s2 2p1

Carbon C  1s2 2s2 2p2

Nitrogen N  1s2 2s2 2p3

Oxygen O  1s2 2s2 2p4

Neon Ne  1s2 2s2 2p6

Sodium Na  1s2 2s2 2p6 3s1

Sulphur S  1s2 2s2 2p6 3s2 3p4

Argon Ar  1s2 2s2 2p6 3s2 3p6

Potassium K  1s2 2s2 2p6 3s2 3p6 4s1

Calcium Ca  1s2 2s2 2p6 3s2 3p6 4s2

The electron configuration of the first 20 elements

3.1.4 Apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the periodic table.

• Do the following questions: 1-3 p. 75 students text book

• Question from work book. Exercise 3.1

P. 97-99

• 3.2 Physical Properties

• Elements in the same group therefore have different physical properties

• Physical properties include:

• Effective nuclear charge

• Atomic radius

• Ionic radius

• Ionization energy

• Electronegativity

• Melting point

The trend in the physical and chemical properties are governed by the effective nuclear charge.

• The nuclear charge of the atom is given by the atomic number. As you go across the periodic table the atomic number increases by one, as one proton is added to the nucleus.

• The outer electron (that determine many of the physical and the chemical properties) do not experience the full attraction of this charge (protons) as they are shielded from the nucleus and repelled by the inner electrons.

• The presence of the inner electrons reduces the attraction of the nucleus for the outer electrons.

• The effective charge experienced by the outer electron is less than th efull nuclear charge.

• As you go across a period, one proton is added to the nucleus and one electron is added to the outer electron shell. The effective charge increases with the nuclear charge as there is no change in the number of the inner electrons

elements Na Mg Al SiNuclear Charge

11 12 13 14

Electron arrangement

2,8,1 2,8,2 2,8,3 2,8,4

• Chemical property• Group I Metals (Alkali Metals)

• These metals react with water to form alkaline solutions

• They have low melting points

• Low boiling points

• Soft

• Low density

• Very reactive

3.2 Physical properties: Trends in Melting and boiling points                                                                                                            You will see that both the melting points and boiling points fall as you go down the Group.

The alkaline earth elements are metallic elements found in the second group of the periodic table. All alkaline earth elements have an oxidation number of +2, making them very reactive. Because of their reactivity, the alkaline metals are not found free in nature. The Alkaline Earth Metals are: Beryllium Magnesium Calcium Strontium Barium Radium

Group II Metals

• Extracted from oxides in the earth’s crust

• Less reactive than the alkali metals

Metalloids have properties of both metals and non-metals. Some of the metalloids, such as silicon and germanium, are semi-conductors. This means that they can carry an electrical charge under special conditions.

This property makes metalloids useful in computers and calculators

The Metalloids are:

Boron Silicon

Germanium Arsenic

Antimony Tellurium Polonium

• Properties of Metals

- They have high melting and boiling points

• Properties of Non-metals– Have low densities– Brittle

Group VII THE HALOGENS

The halogens are five non-metallic elements found in group 7 of the periodic table. The term "halogen" means "salt-former" and compounds containing halogens are called "salts".

The Halogens are: Fluorine Chlorine Bromine Iodine

Astatine

• The halogens exist, at room temperature, in all three states of matter:

• Solid- Iodine, Astatine

• Liquid- Bromine

• Gas- Fluorine, Chlorine

The Halogens are: Fluorine

Iodine Bromine

Chlorine

Astatine

The six noble gases are found in group VIII (8) of the periodic table.

These elements are very stable and do not react with anything else.

They are sometimes referred to as inert gasesHelium Neon Argon

Krypton Xenon Radon

•These gases have low melting points and boiling points.

Group VIII The Noble Gases

PERIODS

• Period Number indicates the number of occupied shells.

• Elements in the same period share a gradual change in their physical and chemical properties.

Trends in the Periodic Table

• The trends in the periodic table are

- atomic radius

- first ionisation energy

- electronegativity

- melting and boiling points

- density.

3.2.2 Trends in Atomic Radius

• ATOMIC SIZEThe radius of an atom is found from the distance

between the nuclei in a molecule of two touching atoms, and then halving that distance.

Trends in atomic radius in Periods 2 and 3

Trends in atomic radius down a group

• It is fairly obvious that the atoms get bigger as you go down groups. The reason is equally obvious - you are adding extra layers of electrons.

Trends in atomic radius across periods

• You have to ignore the noble gas at the end of each period. Because neon and argon don't form bonds.

• Leaving the noble gases out, atoms get smaller as you go across a period.

You can see that the atomic radius increases as you go down the Group.

Explaining the increase in atomic radius

The radius of an atom is governed by

• the number of layers of electrons around the nucleus

• the pull the outer electrons feel from the nucleus.

• Compare lithium and sodium:– Li – Na

In each case, the outer electron feels a net pull of 1+ from the nucleus. The positive charge on the nucleus is cut down by the negativeness of the inner electrons.

• The only factor which is going to affect the size of the atom is therefore the number of layers of inner electrons which have to be fitted in around the atom.

• The more layers of electrons you have, the more space they will take up - electrons repel each other.

• This means that the atoms are bound to get bigger as you go down the Group.

Trends in First Ionisation Energy

• First ionisation energy is the energy needed to remove the most loosely held electron from one mole of gaseous atoms.

• The state symbols - (g) - are essential. When you are talking about ionisation energies, everything must be present in the gaseous state.

• Ionisation energies are measured in kJ mol-1 (kilojoules per mole). They vary in size from 381 KJ (which you would consider very low) up to 2370 KJ (which is very high).

• All elements have a first ionisation energy.

Example: Helium (1st I.E. = 2370 kJ mol-1) Large amount of energy that is needed to remove one of its electrons. (Break a complete shell)

Patterns of first ionisation energies in the Periodic TableThe first 20 elements

• Factors affecting the size of ionisation energy- Ionisation energy is a measure of the energy needed to pull a particular electron away from the attraction of the nucleus. A high value of ionisation energy shows a high attraction between the electron and the nucleus.

The size of that attraction will be governed by:• The charge on the nucleus.

- The more protons there are in the nucleus, the more positively charged the nucleus is, and the more strongly electrons are attracted to it.

• The distance of the electron from the nucleus.- Attraction falls off very rapidly with distance. An electron close to the nucleus will be much more strongly attracted than one further away.

The patterns in periods 2 and 3

• Explaining the general trend for ionisation energies across periods 2 and 3

• The general trend is for ionisation energies to increase across a period.

• The major difference is the increasing number of protons in the nucleus as you go from lithium to neon.

• This causes greater attraction between the nucleus and the electrons and so increases the ionisation energies.

• The increasing nuclear charge also drags the outer electrons in closer to the nucleus.

• That increases ionisation energies still more as you go across the period.

Removal of successive electrons

• Second ionisation energy is defined by the equation:

• It is the energy needed to remove a second electron from each ion in 1 mole of gaseous 1+ ions to give gaseous 2+ ions.

More ionisation energies• You can then have as many successive ionisation

energies as there are electrons in the original atom.

:                                       

• The first four ionisation energies of aluminium, for example, are given by

1st I.E. = 577 kJ mol-1

2nd I.E. = 1820 kJ mol-1

3rd I.E. = 2740 kJ mol-1

4th I.E. = 11600 kJ mol-1

• In order to form an Al3+(g) ion from Al(g) you would have to supply:

577 + 1820 + 2740 = 5137 kJ mol-1

Why do successive ionisation energies get larger?

• Once you have removed the first electron you are left with a positive ion. Trying to remove a negative electron from a positive ion is going to be more difficult than removing it from a neutral atom.

• Removing an electron from a 2+ or 3+ (etc) ion is going to be progressively more difficult.

Using ionisation energies to work out which group an element is in

• A big jump between two successive ionisation energies is typical of suddenly breaking in to an inner level. (new shell)

• You can use this to work out which group of the Periodic Table an element is in from its successive ionisation energies.

• Magnesium (2,8,2) is in group 2 of the Periodic Table and has successive ionisation energies:

• Here the big jump occurs after the second ionisation energy.

• It means that there are 2 electrons which are relatively easy to remove (the 2 electrons in last shell), while the third one is much more difficult (because it comes from an inner level closer to the nucleus and with less screening).

• Silicon (2,8,4) is in group 4 of the Periodic Table and has successive ionisation energies:

:

• Decide which group an atom is in if it has successive ionisation energies?

Plotting graphs of successive ionisation energies

• Chlorine has the electronic structure 2,8,7

• This graph plots the first eight ionisation energies of chlorine.

• The green labels show which electron is being removed for each of the ionisation energies.

.

• The seventeenth ionisation energy of chlorine is nearly 400,000 kJ mol-1, and the vertical scale has to be squashed to accommodate this.

.

• Lets do some problems

Three Types of Bonding

• IONIC BONDING

• COVALENT BONDING

• METALLIC BONDING

IONIC STRUCTURES

• This type of bonding occurs only between a metal and a non metal.

Example:Ionic bonding in sodium chloride• Sodium (2,8,1) has 1 electron more than a stable

noble gas structure (2,8). If it gave away that electron it would become more stable. cation

• Chlorine (2,8,7) has 1 electron short of a stable noble gas structure (2,8,8). If it could gain an electron from somewhere it too would become more stable. anion

• The answer is obvious. If a sodium atom gives an electron to a chlorine atom, both become more stable.

.

Some other examples of ionic bondingmagnesium oxide

Covalent Bonding

• Occurs between two non metals

• Involves the sharing of electrons.

Some very simple covalent moleculesChlorine

For example, two chlorine atoms could both achieve stable structures by sharing their single unpaired

electron as in the diagram.

The two chlorine atoms are said to be joined by a covalent bond. The reason that the two chlorine atoms stick together is that the shared pair of electrons is attracted to the nucleus of both chlorine atoms.

• Hydrogen chloride

• The hydrogen has a helium structure, and the chlorine an argon structure.

For example:

Metallic Bonding

• Metals tend to have high melting points and boiling points suggesting strong bonds between the atoms. Even a metal like sodium (melting point 97.8°C) melts at a considerably higher temperature than the element (neon) which precedes it in the Periodic Table.

Metallic BondingMetallic Bonding

•atoms in metals are packed very closely in an orderly arrangement

                                                                                                                      •each atom loses its valence electrons to become a positive ion

• The electrons can move freely within these molecular orbitals, and so each electron becomes detached from its parent atom. The electrons are said to be delocalized. The metal is held together by the strong forces of attraction between the positive nuclei and the delocalized electrons.

• Metallic bonding is the electrostatic attraction between the positively charged ions and negatively charged electrons

Trends in Electronegativity

• Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons.

• It is usually measured on the Pauling scale, on which the most electronegative element (fluorine) is given an electronegativity of 4.0 and values range down to caesium and francium which are the least electronegative at 0.7.

What happens if two atoms of equal electronegativity bond together?

Consider a bond between two atoms, A and B.

•If the atoms are equally electronegative, both have the same tendency to attract the bonding pair of electrons, and so it will be found on average half way between the two atoms.

What happens if B is slightly more electronegative than A?

• B will attract the electron pair rather more than A does.

What happens if B is a lot more electronegative than A?

In this case, the electron pair is dragged right over to B's end of the bond.

A has lost control of its electron, and B has complete control over both electrons.

Patterns of electronegativity in the Periodic Table

• The most electronegative element is fluorine. If you remember that fact, everything becomes easy, because electronegativity must always increase towards fluorine in the Periodic Table.

Trends in electronegativity across a period• As you go across a period the electronegativity

increases. The chart shows electronegativities from sodium to chlorine - you have to ignore argon. It doesn't have an electronegativity, because it doesn't form bonds.

Trends in electronegativity down a group

• As you go down a group, electronegativity decreases. (If it increases up to fluorine, it must decrease as you go down.) The chart shows the patterns of electronegativity in Groups 1 and 7.

Trends in Melting and Boiling Points

Trends in Melting Point (METALS).

Trends in Melting Point and Boiling Point (NON METALS)

Melting and boiling points across a period The chart shows how the melting and boiling points of the elements change as you go across the period. The figures are plotted in Kelvin rather than °C to avoid having negative values.

Table of physical data

ElementProton number

SymbolMelting point(K)

Boiling point(K)

sodium 11 Na 371 1156

magnesium 12 Mg 922 1380

aluminium 13 Al 933 2740

silicon 14 Si 1683 2628

phosphorus 15 P 317 553

sulphur 16 S 392 718

chlorine 17 Cl 172 238

argon 18 Ar 84 87

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Explanation of trend across a PeriodExplanation of trend across a Period1. The Period always begins with metals Example:Sodium, magnesium and aluminium• Sodium, magnesium and aluminium are all metals. They

have metallic bonding, in which positive metal ions are attracted to delocalized electrons. Going from sodium to aluminium:

• the charge on the metal ions increases from +1 to +3 (with magnesium at +2) ...

• the number of delocalized electrons increases ...

• so the strength of the metallic bonding increases and the melting points and boiling points increase. 

2. The Period then comes across the Metalloids.

Example Silicon

• Silicon is a metalloid (an element with some of the properties of metals and some of the properties of non-metals).

• Silicon has giant covalent bonding. It has a giant lattice structure, in which each silicon atom is covalently-bonded to four other silicon atoms in a tetrahedral arrangement. A giant covalent structure

•                                                               • The structure is held together by strong

covalent bonds in all three dimensions.

Now the Period reaches the non-metals.

Example• Phosphorus, sulphur, chlorine and argon

• These  are all non-metals, and they  exist as  small, separate molecules. Phosphorus, sulphur and chlorine  exist as simple molecules, with strong covalent bonds between their atoms.

• Their melting and boiling points are very low because little energy is needed to overcome their bonds.

• Sulphur has a higher melting point and boiling point than the other three because:

phosphorus exists as P4 molecules ... sulphur exists as S8 molecules ... chlorine exists as Cl2 molecules ... argon exists  individual Ar atoms ...

• the strength of the forces decreases as the size of the molecule decreases ...

• so the melting points and boiling points decrease in the order

S8 > P4 > Cl2 > Ar 

IONIC RADIIIONIC RADIUS

• Ions aren't the same size as the atoms they come from. Compare the sizes of sodium and chloride ions with the sizes of sodium and chlorine atoms.

Positive ions• Positive ions are smaller than the atoms they come from.

Sodium is 2,8,1; Na+ is 2,8.

• You've lost a whole layer of electrons, and the remaining 10 electrons are being pulled in by the full force of 11 protons.

Negative ions

Negative ions are bigger than the atoms they come from. Chlorine is 2,8,7; Cl- is 2,8,8. Although the electrons are still all in the 3-level, the extra repulsion produced by the incoming electron causes the atom to expand. There are still only 17 protons, but they are now having to hold 18 electrons.

Trend

• Both cations and anions increase in size down a group.

PHYSICAL AND CHEMICAL PROPERTIES OF THE GROUP I ELEMENTS

• The Alkali Metals are: • Lithium • Sodium • Potassium • Rubidium • Cesium • Francium

ATOMIC AND PHYSICAL PROPERTIES OF THE GROUP 1 ELEMENTS

Trends in Atomic Radius

You can see that the atomic radius increases as you go down the Group.

Trends in First Ionisation Energy

Notice that first ionisation energy falls as you go down the group.

Trends in Electronegativity

Trends in Melting and Boiling Points

Chemical Properties of Group I

• REACTIONS OF THE GROUP 1 ELEMENTS WITH WATER

All of these metals react vigorously or even explosively with cold water. In each case, a solution of the metal hydroxide is produced together with hydrogen gas.                                                                        This equation applies to any of these metals and water -

Details for the individual metals

• Lithium• Lithium's density is only about half that of water so it

floats on the surface, gently fizzing and giving off hydrogen. It gradually reacts and disappears, forming a colourless solution of lithium hydroxide.

• Sodium

• Sodium also floats on the surface, but enough heat is given off to melt the sodium (sodium has a lower melting point than lithium and the reaction produces heat faster) and it melts almost at once to form a small silvery ball that dashes around the surface.

• A white trail of sodium hydroxide is seen in the water under the sodium, but this soon dissolves to give a colourless solution of sodium hydroxide.

• Potassium• Potassium behaves rather

like sodium except that the reaction is faster and enough heat is given off to set light to the hydrogen.

• This time the normal hydrogen flame is contaminated by potassium compounds and so is coloured lilac (a faintly bluish pink).

• Rubidium• Rubidium is denser than water and so sinks. It

reacts violently and immediately, with everything spitting out of the container again. Rubidium hydroxide solution and hydrogen are formed.

• Caesium

• Caesium explodes on contact with water, quite possibly shattering the container. Caesium hydroxide and hydrogen are formed

Reactions with Oxygen

• Group I elements combine with oxygen to form metal oxides which are basic.

GENERAL EQUATION

Metal + Oxygen Metal oxide

However, each of the metals has its own behaviour as you move down the group.

Lithium• Lithium forms the normal oxide

Li+ O2-

SodiumForms an oxide and the peroxide

Sodium Oxide

When there is a lot of oxygen available then the following reaction occurs.

Reactions with the Halogens

• Group I react with the halogens to for metal halides.

• General Equation

Metal + Halogen Metal Halide

The halides formed are usually a white solid

• Metal + Fluorine Metal Fluoride

• Metal + Bromine Metal Bromide

• Metal + Iodine Metal Iodide

Reaction with dilute acids

• Group I metals react violently with dilute acids to produce hydrogen gas and a salt.

For example, sodium will react with dilute hydrochloric acid to give sodium chloride solution and hydrogen gas.                                            

Reaction with Hydrogen

When heated with hydrogen they react to form hydrides (metal hydride)

The oxidation number of hydrogen is -1

Reactions with Hydrogen

• Halogens react with Hydrogen to produce Hydrogen Halides

Examples:

Physical properties• The hydrogen halides are colourless gases at room

temperature, producing steamy fumes in moist air. Hydrogen fluoride has an abnormally high boiling point for the size of the molecule (293 K or 20°C), and could condense to a liquid on a cool day.

• The hydrogen Halides react with water and form solutions of strong acids

• Hydrogen chloride gas is very soluble in water, reacting with it to produce hydrochloric acid.

• The familiar steamy fumes of hydrogen chloride in moist air are caused by the hydrogen chloride reacting with water vapour in the air to produce a fog of concentrated hydrochloric acid.

Hydrobromic acid and hydriodic acid as strong acids

• Hydrogen bromide and hydrogen iodide dissolve in (and react with) water in exactly the same way as hydrogen chloride does. Hydrogen bromide reacts to give hydrobromic acid; hydrogen iodide gives hydriodic acid. Both of these are also strong acids.

Hydrofluoric acid as an exception

• By contrast, although hydrogen fluoride dissolves freely in water, hydrofluoric acid is only a weak acid

Reactions with water

• Fluorine will react with water to produce O2 gas

• The other halogens will behave differently in water forming an acid

Reactions with Metals

• The halogens combine with metals to give a salt containing the halide ion:

Example:

The salt formed is usually white and when dissolved in water gives a colourless solution.

•The only metal halides that are insoluble in water are usually halides of lead and silver

TESTING FOR HALIDE IONS

• Since the salt formed are colourless when aqueous there must be a way to test for the presence of Cl- Br- and I-

• Using silver nitrate solution

By adding a few drops of this solution and looking for the formation of a precipitate.

• Silver nitrate solution is added to give:

:

ion presentobservation

F- no precipitate

Cl- white precipitate

Br- very pale cream precipitate

I- very pale yellow precipitate

The chloride, bromide and iodide precipitates are shown in the photograph:

.

The chemistry of the test

• The precipitates are the insoluble silver halides - silver chloride, silver bromide or silver iodide.

Silver fluoride is soluble, and so you don't get a precipitate

The chemistry of the test

• The precipitates are the insoluble silver halides - silver chloride, silver bromide or silver iodide.

Silver fluoride is soluble, and so you don't get a precipitate

Reactions of halogens with one another

• Since the reactivity of these elements decreases going down the group then a halogen positioned higher in the group will react with a lower halogen halide.

Bromide added to Chlorine

• Bromide ion is added to chlorine water. We see that the reaction has produced bromine.

2Br-(aq ) + Cl2(aq ) --> 2Cl-(aq ) + Br2(aq )

Bromide ion is added to chlorine water

and is shaken

The reaction has produced bromine

Iodide added to Chlorine

• Iodide ion is added to chlorine water. we see that the reaction has produced iodine.

2I-(aq ) + Cl2(aq ) --> 2Cl-(aq ) + I2(aq )

Iodide ion is added to chlorine water

Chloride added to Bromine

• Chloride ion is added to bromine water.

• we see that no reaction occurrs.

2Cl-(aq ) + Br2(aq ) --> No reaction

Chloride ion is added to bromine water

... and is shaken.

No reaction took place.

Iodide added to Bromine

• Iodide ion is added to bromine. we see that the reaction has produced iodine.

• 2I-(aq ) + Br2(aq ) --> 2Br-(aq ) + I2(aq )

Iodide ion is added to bromine water

... and is shaken.

The reaction has produced iodine.

Chloride added to Iodine

2Cl-(aq ) + I2(aq ) --> No chemical reaction took place

Chloride ion is added to iodine water

... and is shaken.

No chemical reaction took place.

Bromide added to Iodine

2Br-(aq ) + I2 --> No chemical reaction took place

Trends Across a Period.

• Reducing strength decreases (ability to donate electrons)

• Oxidizing strength increases (ability to accept electrons)

• as atoms become less able to release electrons they

have a greater tendency to acquire additional electrons.

• Elements on the left of the table form positive ions and those to the right form negative ions.

Properties of the Period 3 Elements

Atomic Properties• Electronic structures

• In Period 3 of the Periodic Table, the 3s and 3p orbitals are filling with electrons. The electronic structures for the eight elements are:

Na [Ne] 3s1

Mg [Ne] 3s2

Al [Ne] 3s2 3px1

Si [Ne] 3s2 3px1 3py

1

P [Ne] 3s2 3px1 3py

1 3pz1

S [Ne] 3s2 3px2 3py

1 3pz1

Cl [Ne] 3s2 3px2 3py

2 3pz1

Ar [Ne] 3s2 3px2 3py

2 3pz2

First ionisation energyThe pattern of first ionisation energies across Period 3

Explaining the pattern• First ionisation energy is governed by:

• the charge on the nucleus;

• the distance of the outer electron from the nucleus;

• the amount of screening by inner electrons;

• whether the electron is alone in an orbital or one of a pair.

The upward trend• In the whole of period 3, the outer electrons are in 3-

level orbitals. These are all the same sort of distances from the nucleus, and are screened by the same electrons in the first and second levels.

• The major difference is the increasing number of protons in the nucleus as you go from sodium across to argon. This causes greater attraction between the nucleus and the electrons and so increases the ionisation energies.

The fall at aluminium• You might expect the aluminium value to be

more than the magnesium value because of the extra proton.

• But the fact that aluminium's outer electron is in a 3p orbital rather than a 3s.

• The 3p electron is slightly more distant from the nucleus than the 3s, and partially screened by the 3s electrons as well as the inner electrons. Both of these factors offset the effect of the extra proton.

The fall at sulphur• The screening is identical in phosphorus and sulphur

(from the inner electrons and, to some extent, from the 3s electrons), and the electron is being removed from an identical orbital.

• The difference is that in the sulphur case the electron being removed is one of the 3p2 pair. The repulsion between the two electrons in the same orbital means that the electron is easier to remove than it would otherwise be.

Atomic radius

• The trendThe diagram shows how the atomic radius changes as you

go across Period 3.

• The figures used to construct this diagram are based on:

• metallic radii for Na, Mg and Al;

• covalent radii for Si, P, S and Cl;

• the van der Waals radius for Ar because it doesn't form any strong bonds.

• The general trend is atomic size decreases across the period (ignoring the noble gases)

ElectronegativityThe trend

The trend across Period 3 looks like this:

Physical Properties

Structures of the elementsStructures of the elements• The structures of the elements change as you go across

the period. The first three are metallic, silicon is giant covalent, and the rest are simple molecules.

Electrical conductivityElectrical conductivity

• Sodium, magnesium and aluminium are all good conductors of electricity.

• Silicon is a semiconductor.

• None of the rest conduct electricity.

Melting and boiling points

• The chart shows how the melting and boiling points of the elements change as you go across the period.

.

The sizes of the melting and boiling points are governed entirely by the sizes of the molecules. Remember the structures of the

molecules:

CHEMICAL REACTIONS OF THE PERIOD 3

ELEMENTS

Reactions with waterReactions with waterSodium

• Sodium has a very exothermic reaction with cold water producing hydrogen and a colourless solution of sodium hydroxide.

.                                                 

Magnesium• Magnesium has a very slight reaction with cold

water, but burns in steam.• A very clean coil of magnesium dropped into cold cold

waterwater eventually gets covered in small bubbles of hydrogen which float it to the surface. Magnesium hydroxide is formed as a very thin layer on the magnesium and this tends to stop the reaction.

.                                            

•Magnesium burns in steamsteam with its typical white flame to produce white magnesium oxide and hydrogen.

Aluminium

• Aluminium powder heated in steam produces hydrogen and aluminium oxide. The reaction is relatively slow because of the existing strong aluminium oxide layer on the metal, and the build-up of even more oxide during the reaction.

Silicon•The common shiny grey lumps of silicon with a rather metal-like appearance are fairly un-reactive. Most sources suggest that this form of silicon will react with steam at red heat to produce silicon dioxide and hydrogen.

Phosphorus and sulfurThese have no reaction with water.

Chlorine

• Chlorine dissolves in water to some extent to give a green solution. A reversible reaction takes place to produce a mixture of hydrochloric acid and chloric(I) acid (hypochlorous acid).

ArgonThere is no reaction between argon and water.

Reactions with oxygen

SodiumSodium burns in oxygen with an orange flame to produce a white solid mixture of sodium oxide and sodium peroxide.For the simple oxide:

For the peroxide:

MagnesiumMagnesium burns in oxygen with an intense white flame to give white solid magnesium oxide.

• Aluminium• Aluminium will burn in oxygen if it is powdered,

otherwise the strong oxide layer on the aluminium tends to inhibit the reaction. If you sprinkle aluminium powder into a Bunsen flame, you get white sparkles. White aluminium oxide is formed.

SiliconSilicon will burn in oxygen if heated strongly enough. Silicon dioxide is produced.

• Phosphorus• White phosphorus catches fire spontaneously in air, burning

with a white flame and producing clouds of white smoke - a mixture of phosphorus(III) oxide and phosphorus(V) oxide.

• The proportions of these depend on the amount of oxygen available. In an excess of oxygen, the product will be almost entirely phosphorus(V) oxide.

For the phosphorus(III) oxide:

                                    For the phosphorus(V) oxide:

• Sulphur• Sulphur burns in air or oxygen on gentle heating with

a pale blue flame. It produces colourless sulphur dioxide gas.

Chlorine and argonDespite having several oxides, chlorine will not react directly not react directly with oxygenwith oxygen. Argon doesn't react either.

With an excess of oxygen produces sulfur trioxide

2S + 3O2 2SO3

PHYSICAL PROPERTIES OF THE

PERIOD 3 OXIDES

A quick summary of the trends

The oxides

• The oxides we'll be looking at are:

Na2O MgO Al2O3 SiO2 P4O10 SO3 Cl2O7

P4O6 SO2 Cl2O

Oxides of the Period 3Oxides of the Period 3

Na2O MgO Al2O3 SiO2 P4O10

and

P4O6

SO2

and

SO3

Cl2O

and

Cl2O7

ionic ionic ionic Covalent Covalent Covalent Covalent

Electrical conductivity

• None of these oxides has any free or mobile electrons. That means that none of them will conduct electricity when they are solid.

• The ionic oxides can, however, undergo electrolysis when they are molten. They can conduct electricity because of the movement of the ions towards the electrodes and the discharge of the ions when they get there.

The trend in acid-base behaviourThe trend in acid-base behaviour

• The trend is from strongly basic oxides on the left-hand side to strongly acidic ones on the right, via an amphoteric oxide (aluminium oxide) in the middle. An amphoteric oxide is one which shows both acidic and basic properties.

Chemistry of the individual oxides

Sodium oxide• Sodium oxide is a simple strongly basic oxide. It is

basic because it contains the oxide ion, O2-, which is a very strong base with a high tendency to combine with hydrogen ions.

Reaction with water• Sodium oxide reacts exothermically with cold water

to produce sodium hydroxide solution. Depending on its concentration, this will have a pH around 14.

Reaction with acids• As a strong base, sodium oxide also reacts with acids.

For example, it would react with dilute hydrochloric acid to produce sodium chloride solution.

Magnesium oxideMagnesium oxide is a simple basic oxide, because it also contains oxide ions. However, it isn't as strongly basic as sodium oxide because the oxide ions aren't so free.

• Reaction with water• If you shake some white magnesium oxide powder

with water, nothing seems to happen - it doesn't look as if it reacts. However, if you test the pH of the liquid, you find that it is somewhere around pH 9 - showing that it is slightly alkaline

• Some magnesium hydroxide is formed in the reaction, but this is almost insoluble - and so not many hydroxide ions actually get into solution.

• Reaction with acids• Magnesium oxide reacts with acids as you would

expect any simple metal oxide to react. For example, it reacts with warm dilute hydrochloric acid to give magnesium chloride solution.

Aluminium oxideAluminium oxide

• Aluminium oxide is amphoteric. It has reactions as both a base and an acid.

• Reaction with water• Aluminium oxide doesn't react in a simple way with

water in the sense that sodium oxide and magnesium oxide do, and doesn't dissolve in it. Although it still contains oxide ions, they are held too strongly in the solid lattice to react with the water.

• Reaction with acidsReaction with acids• Aluminium oxide contains oxide ions and so reacts

with acids in the same way as sodium or magnesium oxides. That means, for example, that aluminium oxide will react with hot dilute hydrochloric acid to give aluminium chloride solution.

• Reaction with basesReaction with bases• Aluminium oxide has also got an acidic side to its nature, and

it shows this by reacting with bases such as sodium hydroxide solution.

• Various aluminates are formed - compounds where the aluminium is found in the negative ion. This is possible because aluminium has the ability to form covalent bonds with oxygen.

• With hot, concentrated sodium hydroxide solution, aluminium oxide reacts to give a colourless solution of sodium tetrahydroxoaluminate.

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Silicon dioxide (silicon(IV) oxide)Silicon dioxide (silicon(IV) oxide)

• Silicon dioxide has no basic properties - it doesn't contain oxide ions and it doesn't react with acids. Instead, it is very weakly acidic, reacting with strong bases.

• Reaction with water• Silicon dioxide doesn't react with water, because of

the difficulty of breaking up the giant covalent structure.

• Reaction with bases

• Silicon dioxide reacts with sodium hydroxide solution, only if it is hot and concentratedonly if it is hot and concentrated. A colourless solution of sodium silicate is formed.

The phosphorus oxidesThe phosphorus oxides

• Phosphorus(III) oxide (P4O6)

• Phosphorus(III) oxide reacts with cold water to give a solution of the weak acid, H3PO3 - known variously as phosphorous acid, orthophosphorous acid or phosphonic acid.

• P4O6 + 6 H2O 4H3PO3

Phosphorus(V) oxidePhosphorus(V) oxide

• Phosphorus(V) oxide reacts violently with water to give a solution containing a mixture of acids, the nature of which depends on the conditions. We usually just consider one of these, phosphoric(V) acid, H3PO4 - also known just as phosphoric acid.

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The sulphur oxidesThe sulphur oxides• We are going to be looking at sulphur dioxide, SO2,

and sulphur trioxide, SO3.• Sulphur dioxide• Sulphur dioxide is fairly soluble in water, reacting

with it to give a solution of sulphurous acid (sulphuric(IV) acid), H2SO3. This only exists in solution, and any attempt to isolate it just causes sulphur dioxide to be given off again.

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• Sulphur trioxideSulphur trioxide

• Sulphur trioxide reacts violently with water to produce a fog of concentrated sulphuric acid droplets.

The chlorine oxideThe chlorine oxide

• Chlorine(VII) oxide• Chlorine(VII) oxide is the highest oxide of chlorine -

the chlorine is in its maximum oxidation state of +7. It continues the trend of the highest oxides of the Period 3 elements towards being stronger acids.

• Chlorine (VII) oxide reacts with water to give the very strong acid, chloric (VII) acid - also known as perchloric acid. The pH of typical solutions will, like sulphuric acid, be around 0.

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PROPERTIES OF THE PERIOD 3 CHLORIDES

• The chlorides

• The chlorides we'll be looking at are:

NaCl MgCl2 AlCl3 SiCl4 PCl5 S2Cl2

PCl3

Reactions with waterReactions with water

• As an approximation, the simple ionic chlorides (sodium and magnesium chloride) just dissolve in water.

• The other chlorides all react with water in a variety of ways.

• The reaction with water is known as hydrolysis.

Sodium chloride, NaCl

• Sodium chloride is a simple ionic compound consisting of a giant array of sodium and chloride ions.

• A small representative bit of a sodium chloride lattice looks like this:

This is normally drawn in an exploded form as:

• The strong attractions between the positive and negative ions need a lot of heat energy to break, and so sodium chloride has high melting and boiling points.

• It doesn't conduct electricity in the solid state because it hasn't any mobile electrons and the ions aren't free to move. However, when it melts it undergoes electrolysis.

• Sodium chloride simply dissolves in water to give a neutral solution.

Magnesium chloride, MgClMagnesium chloride, MgCl22

• Magnesium chloride is also ionic, but with a more complicated arrangement of the ions to allow for having twice as many chloride ions as magnesium ions

• Again, lots of heat energy is needed to overcome the attractions between the ions, and so the melting and boiling points are again high.

• Solid magnesium chloride is a non-conductor of electricity because the ions aren't free to move. However, it undergoes electrolysis when the ions become free on melting.

• Magnesium chloride dissolves in water to give a faintly acidic solution (pH = approximately 6).

Aluminium chloride, AlClAluminium chloride, AlCl33

• At room temperature, solid aluminium chloride has an ionic lattice with a lot of covalent character.

• Solid aluminium chloride doesn't conduct electricity at room temperature because the ions aren't free to move.

• The reaction of aluminium chloride with water is dramatic. If you drop water onto solid aluminium chloride, you get a violent reaction producing clouds of steamy fumes of hydrogen chloride gas.

• The aluminium chloride reacts with the water rather than just dissolving in it.

Silicon tetrachloride, SiCl4Silicon tetrachloride, SiCl4

• Silicon tetrachloride is a colourless liquid at room temperature which fumes in moist air.

• It doesn't conduct electricity because of the lack of ions or mobile electrons.

• It fumes in moist air because it reacts with water in the air to produce hydrogen chloride

• It is a violent reaction to produce silicon dioxide and fumes of hydrogen chloride.

The phosphorus chloridesThe phosphorus chlorides

• There are two phosphorus chlorides - phosphorus(III) chloride, PCl3, and phosphorus(V) chloride, PCl5.

Phosphorus(III) chloride (phosphorus trichloride), PCl3

• This is another simple covalent chloride - again a fuming liquid at room temperature.

• It doesn't conduct electricity because of the lack of ions or mobile electrons.

• Phosphorus(III) chloride reacts violently with water. You get phosphorous acid, H3PO3, and fumes of hydrogen chloride (or a solution containing hydrochloric acid if lots of water is used).

Disulphur dichloride, SDisulphur dichloride, S22ClCl22

• Disulphur dichloride is a simple covalent liquid - orange and smelly!

• Disulphur dichloride reacts slowly with water to produce a complex mixture of things including hydrochloric acid, sulphur, hydrogen sulphide and various sulphur-containing acids and anions (negative ions).

• There is no way that you can write a single equation for this - and one would never be expected in an exam.

PHYSICAL AND CHEMICAL PROPERTIES

OF THE GROUP VII ELEMENTS

The Halogens are: Fluorine

Iodine Bromine

Chlorine

Astatine

• The name Halogen is derived from the Greek word ‘halos’ meaning salt former

• All halogens have the outer shell configuration s2p5

• The halogens all occur as diatomic molecules by forming a single covalent bond to get a stable electron configuration.

Trends in Atomic Radius

Explaining the increase in atomic radius

• The radius of an atom is governed by- the number of layers of electrons around the nucleus- the pull the outer electrons feel from the nucleus.

• Compare fluorine and chlorine:F 2,7 Cl 2,8,7

• In each case, the outer electrons feel a net pull from the nucleus.

Trends in ElectronegativityTrends in Electronegativity

Explaining the decrease in electronegativityExplaining the decrease in electronegativity

• This is easily shown using simple dots-and-crosses diagrams for hydrogen fluoride and hydrogen chloride.

.

Trends in Melting Point and Boiling PointTrends in Melting Point and Boiling Point

• The bonding pair of electrons between the hydrogen and the halogen feels the same net pull of 7+ from both the fluorine and the chlorine.

• (This is exactly the same sort of argument as you have seen in the atomic radius.)

• However, in the chlorine case, the nucleus is further away from that bonding pair. That means that it won't be as strongly attracted as in the fluorine case.

• The larger pull from the closer fluorine nucleus is why fluorine is more electronegative than chlorine is.

• Explaining the trends in melting point and boiling pointExplaining the trends in melting point and boiling point

• All of the halogens exist as diatomic molecules - F2, Cl2, and so on.

• The intermolecular attractions between one molecule and its neighbours are van der Waals dispersion

• As the molecules get bigger there are obviously more electrons which can move around and set up the temporary dipoles which create these attractions.

• The stronger intermolecular attractions as the molecules get bigger means that you have to supply more heat energy to turn them into either a liquid or a gas - and so their melting and boiling points rise.