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G 11

Chemical Bonding

Starting points

The force which hold atoms

Ionic Bonding

Isoelectronic Series

Ionic Lattice

Metallic Bonding

PRIOR KNOWLEDGE

working out the electron arrangement.

group number and valence electrons.

chemical reactions and formation of ions.

level of nuclear attraction for the negatively

charged electrons.

https://www.youtube.com/watch?feature=player_detailpage&v=Qf07-8Jhhpc

Some valuable learning material on chemical bonding.

CHEMICAL BONDING

It is the electrostatic attraction between

the oppositely charged ions.

Ionic Bond

11Na

11Na+ + 1e-

2, 8, 1

2, 8

17Cl + 1e-

17Cl-

2, 8, 7

2, 8, 8

Na+ Cl-

Electrostatic attraction

Ionic bond

Na+ ion is smaller than the parent atom sodium.

Cl- ion is larger than the parent atom chlorine.

(argon)

For example:

IONS WHICH HAVE ISOELECTRONIC (SAME)

STRUCTURE AS Ne (Neon)

Na+

Mg2+

Al3+

+ 1e-

+ 2e-

+ 3e-

Na

Mg

Al

;

;

The above elements in groups 1, 2 and 3 have only

1, 2, or 3 electrons in their outer shell.

These elements at the beginning of the period lose electrons

to form positive ions.

C/W

Deduce the ions formed when elements in groups, 5, 6

and 7 gain electrons. 2. Do they have the isoelectronic structure?

3. Which Nobel gas configuration would

these ions acquired?

2, 8

2, 8

2, 8

Explain the chemical bonding?

C/W

Magnesium oxide

Calcium chloride

Sodium nitride

IONIC COMPOUNDS HAVE LATTICE STRUCTURE

An ionic lattice means:

There is regular arrangement of positive

metal ions and negative non-metal ions.

These ions are held together by a overall,

net attractive force, ionic bond.

The strength of an ionic lattice is measured

by its lattice enthalpy.

The lattice enthalpy is the energy required to

decompose ( break down) one mole of ionic lattice into gaseous ions.

A single ionic lattice contains a huge number of

ions in a regular repeating units, called unit cells.

Ionic solids are said to possess a giant structure

Construct comparison charts or posters of the four different crystal types and their properties.

HOME WORK

METALLIC BONDING

Delocalized electrons:

The mobile (free to move)

valence electrons.

Metallic bonding is the electrostatic attraction

between the metal ions and the delocalized

electrons.

Metallic bonding is non-directional:

all of the valence electrons are

attracted to the nuclei of all the

meta ions.

The presence of delocalized electrons

accounts for the physical properties

of metals.

This model is often known as the

electron-sea model, where the

delocalized electrons form the ‘sea’

THE MELTING POINTS OF METALS

The melting Point?

The melting point is an approximate measure of the strength

of the metallic bonding in a metal lattice.

The higher the melting point, the stronger the bonding.

The factors controlling the strength of metallic bond:

The delocalized electrons in the

valence shell of the metal atoms.

The size of the metal ion.

As the number of valence electrons per atom increases and

ionic radius decreases, the stronger the metallic bonding.

The unusual properties of metals

Ductile property of a metal: Metals can be drawn into

long wires ( ductility ) without breaking.

Malleable property of a metal: The metallic bonding in a metal

is strong and flexible and so metals can be hammered into thin

sheets ( malleability ).

In a metal the valence electrons do not belong to

any particular atom.

if sufficient force is applied the metal, one layer

of metal atoms can be slide over another.

Application of force

G11

COVALENT BONDING and

DEDUCING Lewis STRUCTURES

TOPIC:

STARTING POINTS:

Covalent bond; ‘sharing of electrons’

Types of covalent bonds

Covalent molecules

Rules to draw Lewis structures

Deducing Lewis structures for covalent molecules

LESSON 2

COVALENT BONDING (sharing of electrons )

Covalent bond

It is the electrostatic attraction between a pair of electrons and

positively charged nucleus nucleus

For example

Hydrogen molecule

The two hydrogen atoms are held together

because their nuclei are both attracted to

the electron pair which is shared between them.

Oxygen atoms can each form two covalent bonds. Two pairs of electrons are shared in an oxygen molecule (O2) - a double bond.

DOUBLE BOND

TRIPLE BOND

A triple bond is when three pairs of electrons

are shared between two atoms in a molecule

– nitrogen molecule.

C/W: Deduce the number of covalent bonds in

C2H2 ethyene (acetylene) molecule.

Using Lewis structures to deduce the formation

of covalent bonds

Lewis structures ( electron dot diagrams ) only include the

outer or valence electrons since these are the only electrons

are involved in bonding.

Rules to draw Lewis structures (hand-out)

C/W: Deduce the Lewis structures of :

O2, N2, H2O, C2H4, C2H6, HCl, NH3

OH-, CO32-, CN-

Drawing Lewis Structures for

Molecules and Ions

Steps to follow, when writing Lewis Structures for

molecules and ions – Hand-out.

Deduce the Lewis ( electron dot ) structures;

for example, H2O, CH4, CCl4 , HCN, CN-, F-

CO32-, O2, N2, Li+ , Cl2, etc.

For learners please use:

CO-ORDINATE ( dative ) BONDING

For example formula equation for the reaction between

ammonia and hydrochloric acid.

NH3 + HCl NH4Cl

Representing co-ordinate bonds

+Cl-

C/W: Formation of Co-ordinate covalent bonds.

Hand-out

PREDICTING THE TYPE OF BONDING FROM

ELECTRONEGATIVITY VALUE

ELECTRONEGATVITY: It is the ability or power of an atom in a

covalent bond to attract shared pair of electrons to itself.

Example Two atoms A and B; if the atoms are equally

electronegative, both have the same tendency to attract the

shared pair of electrons.

What happens if B is slightly more electronegative than A?

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

(read as "delta") means

"slightly"

( slightly positive and slightly negative;

Polar bond )

The most electronegative

element

The least electronegative

elements

Fluorine ( F )

On the Pauling Scale it is assigned

a value of 4.0

Caesium ( Cs ) and Francium ( Fr )

On the Pauling Scale

assigned values at 7.0

Patterns of electronegativity in the

Periodic Table

Electronegativity generally increases

across a period owing to:

increasing nuclear charge

decreasing atomic radius

Example trend in electronegativity

across period 3 elements

Electronegativity decreases on moving

down a group due to:

Atomic size increases

Greater shielding effect

Example trend in electronegativity

down groups 1 and 7.

Predicting the type of chemical bonding based upon the electronegativity difference.

Hand-out/worked example

Polarity of Bonds and Molecules

Bond Polarity means separation of charge.

Experimentally, bond polarity is measured by its dipole moment.

Higher electron density

( partial negative charge )

Low electron density

( partial positive charge )

The larger difference between the electronegativities of the atoms leading to a more polar covalent bond.

2.1

3.0

Electronegativities values

The polarity in the bond can also be represented by a arrow indicating a dipole , (two charges separated by a distance).

The tip of the arrow points toward the more electronegative atom.

Molecular Polarity

The polarity of the molecule is the sum of all of the bond polarities in the molecule.

Formaldehyde ( CH2O ), Carbon dioxide ( CO2 )

For examples:

polar

non-polar

CO2 is a linear molecule, the dipoles cancel each other.

Water is a bent molecule with polar O-H bonds. The bond dipole moments add to give a resultant dipole (m = 1.85 D) directed toward the more electronegative oxygen.

Polarization

The separation of charge in a polar bond is termed as polarization.

polarization

When two electrical charges of opposite sign are separated

by a small distance, a dipole is established. The size of the

dipole is measured by its dipole moment

Dipole moment ( )

= charge X distance = Q x r

Units = debyes (D)

To establish the polarity of some liquids

Practical Chemistry

C/W:

CCl4

CHCl3

(i) Deduce the polar and non-polar bonds.

(ii) The molecules as a whole, polar or non-polar?

Bond polarity CCl4 molecule

( tetrachloromethane )

individual bonds

are polar

molecule as a whole

is non-polar

INTERMOLECULAR FORCES

Forces between the molecules

Van der Waals’ Forces

Dipole-dipole forces

Hydrogen bonding

Van der Waals’ Forces

Van der Waals’ forces are short range attractive forces

and they originated due to the formation of temporary dipoles.

The origin of Van der Waals’ forces

the electron density is evenly spread

around the nucleus

Example

At a given instant, the more of the electron

density is at one end of the molecule;

giving the molecule a temporary dipole

Instantaneous dipole

1.

2.

As the right hand molecule approaches, its electrons will tend to be attracted by the slightly positive end of the left hand one.

3

This sets up an induced dipole in the approaching molecule,

resulting in the formation of another dipole.

4

The formation of induced dipoles is rapidly transmitted from

one molecule to another through the liquid or solids.

5

The forces of attraction between temporary

or induced dipoles are known as van der Waals’ forces.

Factors which influence van der Waals’ forces

The strength of van der Waals’ forces is influenced by two factors:

molecular size

molecular shape

The strength of van der Waals’ forces increases with increasing

molecular mass.

The molecules increase in size and contain high electron density

(a greater number of electrons).

Electron cloud can distorted increasingly easily (greater polarizability) .

molecular size

Boiling point increases with increasing number of electrons –

( molecular masses )

Example: Boiling points of the Halogens and the Noble gases.

Molecular Shape

Molecules with a large surface area allow a closer contact

between molecules.

This gives rise to greater or more extensive van der Waals’ forces

of attraction than in molecules of similar molecular mass but more

compact shapes due to branching.

EXAMPLE

Hydrocarbons molecules; Butane and 2-methylepropane

both have molecular formula, C4H10 .

Butane has a higher boiling point because the van der waals’ forces

are greater.

The molecules are longer and can lie closer together than the shorter, fatter 2-methylpropane molecules.

Dipole-dipole forces

In a polar molecule, the molecules have permanent dipole moments.

A dipole-dipole force exists between polar molecules because the positive

end of the dipole of one molecule will electrostatically attract the negative

end of the dipole of another molecule.

For example; dipole-dipole force in solid hydrogen chloride

Polar molecule

The strength of dipole-dipole forces:

the size of the dipole moment; the larger the dipole moment,

the more polar the molecules of the substances and greater

the strength of the dipole-dipole force.

Dipole moments and boiling points of propane,

methoxymethane, chloromethane and ethane nitrile

Substance with similar molecular masses:

higher the dipole moment

stronger the dipole-dipole attraction

higher the boiling point

Van der Waals’ Vs Dipole-Dipole forces

When molecules have very different molecular masses,

van der Waals’ forces are more significant than dipole-dipole.

The molecules with the largest relative molecular mass

has the strongest intermolecular attraction.

When the molecules have similar molecular masses,

dipole-dipole forces are more significant. The most polar

molecules has the strongest intermolecular attraction.

Tetrachloromethane ( CCl4 )

Trichloromethane ( CHCl3)

polar molecule

non-polar

molecule

example

Compare the boiling points of

van der Waals’ forces ( weak forces )

are operating between the molecules

strong dipole-dipole attraction between

one molecule and its neighbours

Boiling points

So which has the highest boiling point? CCl4 does, because it is a bigger molecule with more electrons ( greater molecular mass ).

Starting Points

Hydrogen bonding

Solubility and hydrogen bond

Practical chemistry

HYDROGEN BONDING

The origin of hydrogen bonding

A hydrogen atom covalently bonded to

nitrogen, oxygen, and fluorine.

These three atoms are small and

highly electronegative.

A lone pairs of electron on the

electronegative atoms

Hydrogen bonding between

water molecules

The electrostatic attraction that holds the hydrogen atom of one molecule to the oxygen atom of another molecule is the example of Hydrogen bonding.

Representation of

hydrogen bond

Example

Example

Hydrogen bonding in ammonia molecule, The nitrogen atom has one lone pair of electrons. This means that each ammonia molecule can form one hydrogen bond.

Nitrogen is larger and less electronegative than fluorine and hence the resulting hydrogen bonding in ammonia is weaker than hydrogen bond formed by hydrogen fluoride.

Hydrogen bonding in ammonia

C/W: Deduce hydrogen bonding in liquid hydrogen fluoride.

Variation in boiling temperature

down groups and across periods

C/W

Order of Strength of Intermolecular

Forces

In order of decreasing strength

Hydrogen bonding dipole-dipole force van der Waals’ force

>

>

In order of increasing strength

<

Van der Waals’ force dipole-dipole force hydrogen bonding

C/W

C/W

Put the following molecules in order of increasing boiling point and explain your choice.

CH3CHO, CH3CH2OH, CH3CH2CH3

C/W: 1. CONSTRUCT A POSTER AND EXPLAIN

THE FACT THAT ICE FLOATS ON WATER IS

THE EVIDENCE OF THE POWER of

HYDROGEN BOND.

2. ROLE OF HYDROGEN BONDING IN BIOLGICAL

MOLECULES

Bonding (shared) and Lone (unshared)

pairs of electrons

lone pairs of electrons

(unshared)

bonding pair of electron

shared

In water molecule:

2 lone pairs

2 bonding pairs

No. charge centres = 4

double or triple bonds behave a single

charge center

Valence Shell Electron Pair

repulsion theory

(VSEPR)

VSPER is used to predict the:

the shapes of molecules and ions

The VSPER theory states that:

the electron pairs around the central atom

repel each other

bonding pairs and lone pairs of electrons arrange

themselves to be far apart as possible

bonding pairs and lone pairs are also called negative

charge centers

Shapes of Molecules and Bond angles

The order of the repulsion strength of lone pairs

and bond pairs of electron is:

lone pair-lone repulsion lone pair-bond pair repulsion

bond pair-bond pair repulsion

species with two negative charge centres

Example BF2

shape: linear

generic formula: MX2

M = central atom and X is the bonding atoms

Species with three negative

charge centers

Example BF3

shape: trigonal planner

generic formula: MX3

species with four negative

charge centres

shape: tetrahedral

generic formula:

MX4

five charge

centers

six charge

centers