MODULES 6 and 7 – STRUCTURE AND CHEMICAL BONDING

6.1: ELEMENTS, COMPOUNDS AND MIXTURES

ESSENTIAL DEFINITIONS ATOM: The smallest particle that cannot be broken down by chemical means, composed

of protons, neutrons and electrons ION: an atom or group of atoms that carries an electrical charge MOLECULE: a particle of two or mote atoms joined together by covalent bonds COMPOUND: a substance made of two or more different types of atoms joined together

by chemical bonds. Compounds consist of fixed proportions of the different types of atoms they are composed of. Compounds often have very different properties to the elements from which they are made.

MIXTURE: a mixture contains two or more elements or compounds that are NOT chemically bonded together. A mixture does not have fixed proportions of each element or compound in it.

ALLOY: An alloy is a mixture of metal atoms with other metal atoms OR non metal atoms. The properties of a metal are changed by making it into an alloy. Metals are made into alloys to improve their strength, hardness or resistance to corrosion.

CHEMICAL BONDS: Chemical bonds are attractive forces holding together two or more atoms of the same or different elements together. There are generally two kinds of bonding formed in chemical reactions; ionic and covalent.

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THE UNREACTIVE ELEMENTS

Atoms of most elements form compounds with atoms of other elements. The common exceptions to this general rule are the noble gases, elements of Group 8 or 0, which exist as un-bonded, discrete (separate) atoms - e.g.

Element Symbol Electronarrangement

Helium He 2

Neon Ne 2,8

Argon Ar 2,8,8

All these atoms have complete outer (valence) shells of electrons. Due to their unreactivity, they are most commonly used as inert gases in situations where reactions are unwanted/unneeded, e.g. helium in party balloons or blimps, argon in light bulbs, neon in signs.

Some of the most reactive elements are found in the Groups next to the Group 0 elements, for example

Group Element Symbol Electronarrangement

1 sodium Na 2,8,1

7 chlorine Cl 2,8,7

Full valence shells of electrons are associated with a more stable situation in terms of energy. Atoms which do not have full valence shells try to lose, gain or share electrons to acquire full, valence shells. Atoms react with atoms of their own type or with other atoms in order to achieve a lower energy state, which is provided by gaining full valence electron shells.

Electrons can be either:• transferred: between atoms of elements where one needs to lose electrons and the

other to gain electrons for each to have a full valence shell (metals with non-metals). This is called ionic bonding.

• shared: between atoms of elements that both need to gain electrons for a full valence shell (non-metals with non-metals). This is called covalent bonding.

Only the valence electrons (in the outer shell) are involved in chemical reactions.

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7.1: IONIC BONDING & FORMULAE [Text pgs 34 - 35]

Ions can be:• single atoms that have lost or gained electrons (e.g. Na+, or Cl-)• Groups of atoms joined together with an overall charge (e.g. SO42-).

Two atoms or ions with the same electron structure are called isoelectronic, e.g. the following particles all have the electron structure of 2,8 and they are considered to be isoelectronic.

Oxide ion O2- Fluoride ion F-Neon atom Ne Sodium ion Na+Magnesium ion Mg2+ Aluminium ion Al 3+

Ionic bonding most commonly occurs in compounds where a metal and non-metal element react together.

• Metal atoms usually react by losing electrons. The metal atoms then form positive ions.

e.g. 11Na: 2, 8, 1 (atom) 11Na+: 2, 8 (ion) + e-

• Non-metal atoms usually react by gaining electrons. The non-metal atoms become negative ions.

e.g. 17Cl: 2, 8, 7 (atom) + e- 17Cl-: 2, 8, 8 (ion)

The neighbouring, oppositely charged ions then attract each other (electrostatic attraction).When sodium reacts with chlorine, an electron moves from the sodium atom to the chlorine atom. The Ionic bond is the electrostatic attraction between positive and negative ions.

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YOU SHOULD KNOW THE NAMES AND FORMULAE OF ALL THE IONS ON THIS PAGE

Monoatomic ions form when an atom loses or gains electrons to form ions. The charge on an ion is deduced from the number of electrons gained or lost. (Atoms in brackets () are unlikely to form ions. Group 8 or 0 is not shown since its atoms do not form ions except under extreme conditions).

Common monoatomic ions

1H+

3Li+ 4Be2+ (B) (C) 7N3- 8O2- 9F-

11Na+

12Mg2+

13Al3+ (Si) (P) 16S2- 17Cl-

19K+ 20Ca2+ 35Br-

Many transition metals can have multiple (more than one) ionic charges. Iron is a good common example of this as it forms both +2 and +3 ions. The charge of the iron ion in a compound is shown by writing the charge in Roman numerals in the middle of the name.

FeO = Iron (II) oxide Fe2O3 = Iron (III) oxide.

OTHER USEFUL MONOATOMIC IONS

POLYATOMIC IONS Groups of atoms bonded (covalently) together, but gain or lose an electron to another atom during the process and thus come out with an overall charge.Common polyatomic ions

NH4+ ammonium OH- hydroxideNO3- nitrate PO43- phosphate

CO32- carbonate MnO4- permanganate ormanganate(VII)

SO42- sulphate CH3COO- ethanoate

SO32- sulphite CN- cyanide

HCO3- hydrogen carbonate or bicarbonate CrO42- chromate

Cr2O72- dichromate

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Ag+ Silver Zn2+

Zinc

Pb2+

Lead (II) Sn2+

Tin (II)

Pb4+ Lead (IV) Sn4+ Tin (IV)Fe2+

Iron (II) Cu+ Copper (I)

Fe3+

Iron (III) Cu2+

Copper (II)

Cr3+

Chromium (III) Mn2+

Manganese (II)

Cr6+ Chromium (VI) Mn4+ Manganese (IV)

6.2.2 DETERMINATION OF THE FORMULAE OF IONIC COMPOUNDS

When an atom loses an electron to form a positive ion the electron is gained by another atom to form a negative ion.

Two methods of working out the formula of an ionic compound are shown:

1. THE CROSS OVER METHOD.• Take the number of the charge on the positive ion (not the +) and put it as a subscript

(bottom number) on the negative ion. • Take the number of the charge of the negative ion (not the -) and put it as the subscript

(bottom number) of the positive ion. • Simplify the 2 numbers if possible.

Examples:

1. Na1+ + Cl1- (Na1+)1(Cl1-)1 = NaCl

+1 from the sodium ions

-1 from the chloride ion

2. Al3+ + O2- (Al3+)2(O2-)3 = Al2O3

+6 from two aluminium ions

-6 from the three oxide ions

2. THE BALANCE (SEE-SAW) METHOD.Balance up the charges by imagining the ions are on a see-saw. The number of positive charges must equal (balance) the number of negative charges for the see-saw to be in balance.

NaCl Al2O3

Both of these methods generate the same result (formula).

Note: Brackets ( ) must be used when a formula contains more than one Polyatomic ion.

For example aluminium compounds:Aluminium oxide - contains the monoatomic oxide ion (O2-) Formula Al2O3

Aluminium sulphate contains polyatomic sulphate ions (SO42-) Formula Al2(SO4)3

Remember: The formula of the sulphate ion is placed inside brackets because there is more than one sulphate ion in the simple, ionic formula.

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7.1.7 METALLIC BONDING [TEXT P.44 – 45]Metals consist of 3 dimensional lattices (regular, repetitive 3D arrangements) of metal ions (cations) surrounded by their valence electrons. These valence electrons move freely from metal ion to metal ion (delocalised electrons) throughout the structure.

The electrostatic attraction between the cations and the moving "sea of mobile negative electrons" is known as the metallic bond and is very strong.

Metallic bonding is found in pure metals and alloys (mixtures of metals).

A metallic lattice. (3-D arrangement)

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UNIT 7: COVALENT BONDING & MOLECULES [Text pgs 36 - 39]

Covalent bonds occur when two or more non-metal atoms share electrons.

The number of electrons shared depends on how many electrons are needed by an atom to complete the outer, valence shell. The total number of electrons in a complete valence shell is usually 8. (An exception is hydrogen, which needs only one extra electron to make it isoelectronic with helium with 2 electrons in its complete valence shell.)

This is shown for a selection of atoms in the table below.

0

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Group I, II and III atoms are metallic and do not usually form covalent bonds.

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Dot and Cross Diagrams (also called Lewis structures)Dot and Cross diagrams show the valence (outer shell) electrons only in simple covalent compounds or polyatomic ions (e.g. ammonium NH4+). They are used to work out the shape of the molecule, using the number of electron pairs that surround the central atom.

The method is:1. Find the formula of the compound.2. Find the atomic number (Z) of each element.3. Work out the number of valence electrons of each element.4. Draw the electrons around the central atom and pair them up with electrons from the

others. One of the atom’s electrons is drawn with a cross, the other with dots. Generally the central atom will have up to 4 pairs of electrons around it, this makes it stable with 8 outer shell electrons, e.g.

Example: Methane, CH4

Non-polar vs. Polar Covalent BondsSome covalent bonds involve exactly equal sharing of the bonded pair(s) of electrons (e.g. Cl with Cl, O with O) or close enough to equal sharing (e.g. C with H) that the shared pair of electrons is equidistant from the two atoms and the electronic charge is evenly distributed around both atoms. This is referred to as a non-polar (pure) covalent bond.

Not all covalent bonds demonstrate even sharing of the electron pair or pairs in the covalent bond. When one of the atoms (e.g. F, O or Cl) has a stronger attraction for the shared pair of electrons in the bond than the other atom (e.g. C or H) the shared pair of electrons moves closer to that atom (e.g. F, O or Cl). The area around that atom gains a slightly negative charge (-) and the area around the other atom a slightly positive charge (+). The bond is a dipole or polar.

A polar covalent bond or dipole has oppositely charged ends separated by some distance. It is similar to a magnet with magnetic North and South poles.

Or more simply:

Dot and Cross diagram

HH

HH C

x

xCarbon has 4 valence electrons and needs four more.Each hydrogen has 1 valence electron and needs 1 more. Sharing the available electrons gives each atom a full valence shell.

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In the extreme case of polar bonding, the atom of one element pulls so much more strongly on the electrons than the other element’s atom that it literally removes the electron from the weaker atom. Instead of electron sharing, electron transfer has occurred. A positive ion and a negative ion are formed. This is the basis for ionic bonding.

Non-polar vs Polar Molecules

Whether or not the entire molecule shows polarity depends on the polarity of its bonds in the molecule AND its overall shape

Classifying an entire molecule as polar or non-polar is determined by the following rules:1) If all the bonds in the molecule are non-polar, then the entire molecule is non-polar.

2) If one or more of the bonds are polar, then polarity of the entire molecule depends on its shape:

a) if the shape is symmetrical so all the dipoles cancel out, then the entire molecule will be non-polar (usually tetrahedral and trigonal planar);

b) if the shape is NOT symmetrical so the dipoles do not cancel out, then the entire molecule is polar.

Examples of Polar vs. Non-polar Molecules

H2O is a polar molecule .The oxygen atom has a stronger attraction for the bonded electrons than the hydrogen atoms.

The covalently bonded electrons are closer to the oxygen atom making it slightly negative (-) with respect to the hydrogen atom being slightly positive (+)

Dipoles are created on each bond.

Lone pairs (non-bonding) of electrons on oxygen repel the bonded pairs into a bent shaped molecule.

The molecule is not symmetrical in shape and the dipoles do not cancel out.

Thus water is a polar molecule.Simplified diagram of waterNote how the resulting molecule has opposite ends/sides of opposite charge. This is true for all polar molecules.

CH4 is a non-polar molecule.

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In methane (CH4)The carbon and hydrogen have similar attractions for the bonded electrons.

The covalent bond electrons are shared almost equally creating no dipoles.

Since all bonds are non-polar, the molecule is non-polar.

Note even if all the bonds had been polar (e.g.CCl4), due to symmetrical shape, all the dipoles would cancel out and the molecule would still be non-polar.

If only one, two or three of the bonds had been polar (e.g. CH3Cl, CH2Cl2 or CHCl3), then the dipoles would not cancel and the molecule would be polar.

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