Chemistry IB OUP

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  • 425C H E M I C A L B O N D I N G A N D S T R U C T U R E

    C H E M I C A L B O N D I N G A N D S T R U C T U R EIonic bonding

    FORMULAS OF IONIC COMPOUNDSIt is easy to obtain the correct formula as the overall charge of the compound must be zero.

    lithium uoride L i + F - magnesium chloride M g 2+ Cl-2 aluminium bromide A l 3+ Br-


    sodium oxide Na+2 O 2- calcium sulde C a 2+ S 2- iron(III) oxide Fe3+



    potassium nitride K+3N3- calcium phosphide Ca2+


    2 iron(II) oxide F e 2+ O 2-

    Note: the formulas above have been written to show the charges carried by the ions. Unless asked specically to do this it is

    common practice to omit the charges and simply write LiF, MgC l 2 , etc.

    IONIC BOND When atoms combine they do so by trying to achieve a noble

    gas conguration. Ionic compounds are formed when electrons

    are transferred from one atom to another to form ions with

    complete outer shells of electrons. In an ionic compound the

    positive and negative ions are attracted to each other by strong

    electrostatic forces, and build up into a strong lattice. Ionic

    compounds have high melting points as considerable energy is

    required to overcome these forces of attraction.

    The classic example of an ionic compound is sodium chloride N a + C l - , formed when sodium metal burns in chlorine. Chlorine is

    a covalent molecule, so each atom already has a noble gas conguration. However, the energy given out when the ionic lattice is

    formed is sufcient to break the bond in the chlorine molecule to give atoms of chlorine. Each sodium atom then transfers one

    electron to a chlorine atom to form the ions.

    The charge carried by an ion depends on the number of electrons the atom needed to lose or gain to achieve a full outer shell.

    Cations Anions Group 1 Group 2 Group 3 Group 15 Group 16 Group 17

    +1 +2 +3 -3 -2 -1

    L i + N a + K + M g 2+ C a 2+ A l 3+ N 3- P 3- O 2- S 2- F - C l - B r -

    Thus, in magnesium chloride, two chlorine atoms each gain one electron from a magnesium atom to form M g 2+ C l 2 - . In magnesium

    oxide two electrons are transferred from magnesium to oxygen to give M g 2+ O 2- . Transition metals can form more than one ion. For

    example, iron can form F e 2+ and F e 3+ and copper can form C u + and C u 2+ .

    IONS CONTAINING MORE THAN ONE ELEMENT (POLYATOMIC IONS)In ions formed from more than one element the charge is often spread (delocalized) over the whole ion. An example of a positive

    ion is the ammonium ion NH4+, in which all four NH bonds are identical. Negative ions are sometimes known as acid radicals as

    they are formed when an acid loses one or more H + ions.

    hydroxide O H - carbonate CO32-

    nitrate N O 3 - (from nitric acid, HN O 3 ) hydrogencarbonate HCO3


    sulfate SO42- ethanoate C H

    3 CO O - (from ethanoic acid, C H 3 COOH)

    hydrogensulfate HS O 4 - phosphate PO

    43- (from phosphoric acid H



    The formulas of the ionic compounds are obtained in exactly the same way. Note: brackets are used to show that the subscript

    covers all the elements in the ion e.g. sodium nitrate, NaNO3, ammonium sulfate, (NH



    4 and calcium phosphate, Ca




    {from carbonic acid, H 2 C O 3 }{from sulfuric acid, H 2 S O 4 }

    IONIC BOND AND PROPERTIES OF IONIC COMPOUNDSIonic compounds are formed between metals on the left of the periodic table and non-metals on the

    right of the periodic table; that is, between elements in groups 1, 2, and 3 with a low electronegativity

    (electropositive elements) and elements with a high electronegativity in groups 15, 16, and 17.

    Generally the difference between the electronegativity values needs to be greater than about 1.8 for

    ionic bonding to occur.

    Ions in solid ionic compounds are held in a crystal lattice. The ionic bond is the sum of all the

    electrostatic attractions (and repulsions) within the lattice. A large amount of energy is required to

    break the lattice so ionic compounds tend to have high melting points. Many are soluble in water as

    the hydration energy of the ions provides the energy to overcome the lattice enthalpy. Solid ionic

    compounds cannot conduct electricity as the ions are held in xed positions. When molten the ions

    are free to move and conduct electricity as they are chemically decomposed at the respective electrodes.

    ClNa Cl-Na+


    11 protons11 electrons


    17 protons17 electrons

    [Ne]11 protons10 electrons

    [Ar]17 protons18 electrons




    + +


    Sodium chloride (melting point 801 C) lons held strongly in ionic lattice

    839353_IBSG CHEMISTRY_Ch04.indd 25 7/14/14 1:44 PM

  • 26 C H E M I C A L B O N D I N G A N D S T R U C T U R E

    Covalent bondingSINGLE COVALENT BONDSCovalent bonding involves the sharing of one or more pairs of electrons so that

    each atom in the molecule achieves a noble gas conguration. The simplest covalent

    molecule is hydrogen. Each hydrogen atom has one electron in its outer shell. The two

    electrons are shared and attracted electrostatically by both positive nuclei resulting

    in a directional bond between the two atoms to form a molecule. When one pair of

    electrons is shared the resulting bond is known as a single covalent bond. Another

    example of a diatomic molecule with a single covalent bond is chlorine, C l 2 .

    COORDINATE (DATIVE) BONDSThe electrons in the shared pair may originate from the same

    atom. This is known as a coordinate covalent bond.

    NH H



    C == O


    O H




    coordinate bond

    Carbonmonoxide ammonium ion hydroxonium ion

    Sulfur dioxide and sulfur trioxide are both sometimes shown

    as having a coordinate bond between sulfur and oxygen or

    they are shown as having double bonds between the sulfur

    and the oxygen. Both are acceptable.

    SO O


    8 valence electrons around S

    10 valence electrons around S

    2 coordinate bonds, 8 valence electrons

    around S

    12 valence electrons around S

    SO O S


    O OS


    O O

    LEWIS STRUCTURESIn the Lewis structure (also known as electron dot structure) all the valence electrons are shown. There are various different

    methods of depicting the electrons. The simplest method involves using a line to represent one pair of electrons. It is also acceptable

    to represent single electrons by dots, crosses or a combination of the two. The four methods below are all correct ways of showing

    the Lewis structure of uorine.

    F F xx xx xxx x x x

    x xx xF F F FF F xx x

    x x

    x x

    Sometimes just the shared pairs of electrons are shown, e.g. FF. This gives information about the bonding in the molecule, but it is

    not the Lewis structure as it does not show all the valence electrons.


    CH H


    Hmethane ammonia

    NH H








    F F


    The carbon atom (electron conguration 1 s 2 2 s 2 2 p 2 ) has four

    electrons in its outer shell and requires a share in four more

    electrons. It forms four single bonds with elements that only

    require a share in one more electron, such as hydrogen or

    chlorine. Nitrogen (1 s 2 2 s 2 2 p 3 ) forms three single bonds with

    hydrogen in ammonia leaving one non-bonded pair of electrons

    (also known as a lone pair). In water there are two non-bonded

    pairs and in hydrogen uoride three non-bonded pairs.

    MULTIPLE COVALENT BONDSIn some compounds atoms can share more than one pair of

    electrons to achieve an noble gas conguration.



    C C H





    carbon dioxide





    ClH H Cl

    BOND LENGTH AND BOND STRENGTHThe strength of attraction that the two nuclei have for the shared

    electrons affects both the length and strength of the bond.

    Although there is considerable variation in the bond lengths and

    strengths of single bonds in different compounds, double bonds

    are generally much stronger and shorter than single bonds. The

    strongest covalent bonds are shown by triple bonds.

    Length Strength / nm / kJ mo l 1

    Single bonds ClCl 0.199 242 CC 0.154 346

    Double bonds C=C 0.134 614 O=O 0.121 498

    Triple bonds CC 0.120 839 NN 0.110 945

    e.g. ethanoic acid:



    0.124 nm

    0.143 nm

    double bond between C and O shorter and

    stronger than single bond






    BOND POLARITYIn diatomic molecules containing the same element (e.g. H 2 or

    Cl2) the electron pair will be shared equally, as both atoms exert

    an identical attraction. However, when the atoms are different

    the more electronegative atom exerts a greater attraction for the

    electron pair. One end of the molecule will thus be more electron

    rich than the other end, resulting in a polar bond. This relatively

    small difference in charge is represented by + and -. The bigger

    the difference in electronegativities the more polar the bond.

    H F+ -


    H O+ 2-


    H N+ +3-