5/28/2018 Chapter 4 (Part 1)

1/31

Chemical Bonds

5/28/2018 Chapter 4 (Part 1)

2/31

Learning Outcomes describe ionic (electrovalent) bonding, as in sodium chloride and magnesium oxide,

including the use of dot-and-cross diagrams

describe, including the use of dot-and-cross diagrams,

(i) covalent bonding, as in hydrogen, oxygen, chlorine, hydrogen chloride, carbon

dioxide, methane, ethene

(ii) co-ordinate (dative covalent) bonding, as in the formation of the ammonium

ion and in the Al2Cl 6molecule

explain the shapes of, and bond angles in, molecules by using the qualitative model

of electron-pair repulsion (including lone pairs), using as simple examples: BF3

(trigonal), CO2 (linear), CH4 (tetrahedral), NH3 (pyramidal), H2O (non-linear), SF6

(octahedral)

describe covalent bonding in terms of orbital overlap, giving and bonds

explain the shape of, and bond angles in, the ethane, ethene and

benzene molecules in terms of and bonds

5/28/2018 Chapter 4 (Part 1)

3/31

Learning Outcomes predict the shapes of, and bond angles in, molecules analogous to those specified

above.

describe hydrogen bonding, using ammonia and water as simple examples of

molecules containing N-H and O-H groups

explain the terms bond energy, bond length and bond polarity and use them to

compare the reactivities of covalent bonds

describe intermolecular forces (van der Waals forces), based on permanent and

induced dipoles, as in CHCl 3(l); Br2(l) and the liquid noble gases

describe metallic bonding in terms of a lattice of positive ions surrounded by

mobile electrons describe, interpret and/or predict the effect of different types of bonding (ionic

bonding, covalent bonding, hydrogen bonding, other intermolecular interactions,

metallic bonding) on the physical properties of substances

deduce the type of bonding present from given information

show understanding of chemical reactions in terms of energy transfers associatedwith the breaking and making of chemical bonds

5/28/2018 Chapter 4 (Part 1)

4/31

Chemical Bonds

5/28/2018 Chapter 4 (Part 1)

5/31

Chemical Bonds Chemical bond is the intermolecular forces that hold atoms together in

molecules.

Bonding involves only the valence electrons. There are 2 main types of chemical bond:

a) Ionic bond

- electrons are transferred from one atom to another

b) Covalent bond

- bonding occurs when atoms share electrons

5/28/2018 Chapter 4 (Part 1)

6/31

Ionic Bonding Ionic bonds are formed by one atom transferring electrons to another

atom to form ions. Ions are atoms, or groups of atoms, which have lost or

gained electrons.

Ions of opposite charge will attract one another, thus creating an ionic

bond.

It is the electrostatic force that hold the ions of opposite charge together.

5/28/2018 Chapter 4 (Part 1)

7/31

Ionic BondsExample: NaCl

5/28/2018 Chapter 4 (Part 1)

8/31

Ionic compound

NaCl, have an extended crystal lattice with non-metal anions

electrostatically attracted to adjacent metal cations and metal cations

electrostatically attracted to adjacent non-metal anions

5/28/2018 Chapter 4 (Part 1)

9/31

ExerciseDescribe the ionic bond formation in the following compounds:

a) Magnesium oxide

b) Calcium sulfide

c) Potassium nitride

d) Magnesium chloride

e) Calcium phosphide

5/28/2018 Chapter 4 (Part 1)

10/31

Covalent bond

Covalent bonds are formed as a result of the sharing of one or

more pairs of bonding electrons.

Example : H2

Two hydrogen atoms share their single electrons and form a

covalent bond.

H + H H H H-H

electrons to share a shared pair a covalent bond

of electrons

The Covalent Bond Model

5/28/2018 Chapter 4 (Part 1)

11/31

Covalent BondsExample : H2and H2O molecule

5/28/2018 Chapter 4 (Part 1)

12/31

Bonding electrons

Pairs of valence electrons that are shared

between atoms in a covalent bond

Nonbonding electrons

Pairs of valence electrons on an atom that are

not involved in electron sharing

Bonding and Nonbonding Electrons

5/28/2018 Chapter 4 (Part 1)

13/31

Number of covalent bonds depends on the

number of unpaired valence electrons of the

central atom.

To obey octet rule by sharing electrons from

other atoms in the molecule.

Covalent Compounds

5/28/2018 Chapter 4 (Part 1)

14/31

Single covalent bond

A covalent bond in which 2 atoms share 1 pair of

electron

Double covalent bond

The sharing of two pairs of electrons

Triple covalent bond

Three pairs of electrons are shared

Multiple Covalent Bonds

5/28/2018 Chapter 4 (Part 1)

15/31

Multiple Covalent Bond

Example of covalent compound with a double bond is oxygen gas.

O = O

Example of covalent compound with a triple bond is nitrogen gas.

N N

5/28/2018 Chapter 4 (Part 1)

16/31

Explain the bond formation in the following compounds.

a) Hydrogen fluoride

b) Methane

c) Carbon dioxide

d) Ethene

Exercise

5/28/2018 Chapter 4 (Part 1)

17/31

Covalent Compound

Silicon dioxide, SiO2, is a molecular compound. It is also a mineral called quartz (left).

Quartz is found in nearly every type of rock. Most sand grains (center) are bits of quartz.

Glass is made from sand.

5/28/2018 Chapter 4 (Part 1)

18/31

Both electrons of a shared pair come from one of the

two atoms involved in the bond

One atom has a pair of nonbonding electrons Another atom has two or more empty space in its

valence shell.

The result of overlap of a filled and a vacant orbital.

Co-ordinate Covalent Bond

5/28/2018 Chapter 4 (Part 1)

19/31

Co-ordinate covalent bond is represented

by an arrow.

Once a co-ordinate bond is formed, it

cannot be distinguished from other

covalent bonds. All electrons are identical and all the bonds

are of the same strength.

Co-ordinate Covalent Bond

5/28/2018 Chapter 4 (Part 1)

20/31

Both electrons from the nitrogen are shared with the upperhydrogen

Other example: Carbon monoxide

5/28/2018 Chapter 4 (Part 1)

21/31

Co-ordinate Covalent Bond

At high temperatures aluminium chloride exists as molecules withthe formula AlCl3(electron deficient).

At lower temperatures two molecules of AlCl3 combine to form a

molecule with the formula Al2Cl6 .

The molecules are able to combine because lone pairs of electronson two of the chlorne atoms form co-ordinate bonds with the

aluminium atoms.

AlCl

Cl Cl

Cl Al Cl

Cl

5/28/2018 Chapter 4 (Part 1)

22/31

Geometry of the atoms around a central atom can be

predicted by the electron-pair repulsion theory.

This theory is used to determine the shape whichlooks at the geometry of the electron groups around a

central atom.

Molecular Geometry

5/28/2018 Chapter 4 (Part 1)

23/31

Bond angles

The shape and bond angles of a covalently bonded molecule

depend on:

a) the number of pairs of electrons around each atom

b) whether these pairs are lone pairs or bonding pairs

Lone pairs occupy more space than bonding electron pairs.

Double bonds occupy more space than single bonds.

LP-LP > LP-BP > BP-BP

Lone pairs are more repulsive than bonding pairs

5/28/2018 Chapter 4 (Part 1)

24/31

Shape of Molecules

Sets

(group of bonding

pairs/number of

bonded atoms)

Lone

Pairs

Shape

2 0 Linear 180o

2 2 Bent 104.5o

3 0 Trigonal planar 120o

3 1 Trigonal pyramidal 107.3 o

4 0 Tetrahedral 109.5o

5/28/2018 Chapter 4 (Part 1)

25/31

Shape of Molecules

5/28/2018 Chapter 4 (Part 1)

26/31

5/28/2018 Chapter 4 (Part 1)

27/31

5/28/2018 Chapter 4 (Part 1)

28/31

Shapes of Molecules

Silicon tetrachloride has a tetrahedralstructure.

4 bonding pairs of electrons and no lone

pairs

Equal repulsive forces of each bonding

pair of electrons

All ClSiCl bond angles being 109.5o

5/28/2018 Chapter 4 (Part 1)

29/31

Shapes of Molecules

Lone pairbond pair repulsion is greater than bond pairbond

pair repulsion

Bonding pairs of electrons are pushed closer together.

The HNH bond angle is about 107o

3 bonding pairs of electrons

1 lone pair

5/28/2018 Chapter 4 (Part 1)

30/31

Examples

Arrangement of

electron pairs on

central atom

Number of

bonding

electron pairs

Example

Linear 2 BeCl2

Trigonal planar 3 BCl3

Tetrahedral 4 CH4

5/28/2018 Chapter 4 (Part 1)

31/31

Learning Check

1. Predict the shapes for the following molecules:

a) hydrogen sulfide, H2S

b) phosphine, PH3

c) boron chloride, BCl3 (SPECIAL!! BCl3 does not fulfill octet rule)

d) carbon dioxide, CO2

e) nitrogen trichloride, NCl3