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UNIVERSITY OF TECHNOLOGY, JAMAICA

DIVISION OF CHEMISTRY

GROUP: MT-1/ DN-1/ EH-1/ AS-1

MODULE: GENERAL CHEMISTRY (CHY2021)

_____________________________________________________________________________

UNIT 4: CHEMICAL BONDING

Assigned Reading: Chapt. 8, 9 Chemistry – The Central Science, 11th

edition

Generally: What is a bond? Then what is a chemical bond?

Answer: A chemical bond is the attractive forces that hold atoms or ions together to form

molecules, or crystals.

There are three (3) main types of chemical bonds:

(i) ionic bond (electrovalent bond)

(ii) covalent bond (normal covalent and dative covalent bond)

(iii) metallic bond

The types of bonds that are present in a substance are largely responsible for its physical

and chemical properties.

Lewis Symbols and the Octet Rule

Valence electrons reside in the outer shell and are the electrons which are going to be

involved in chemical interactions and bonding (valence comes from the Latin valere, "to be

strong").

Electron-dot symbols (Lewis symbols):

o convenient representation of valence electrons

o allows you to keep track of valence electrons during bond formation

o consists of the chemical symbol for the element plus a dot for each valence electron

For examples, sulfur has an electron configuration of [Ne]3s23p

4, thus there are six valence

electrons. Its Lewis symbol would therefore be:

Note:

The dots (representing electrons) are placed on the four sides of the atomic symbol (top,

bottom, left, right)

Each side can accommodate up to 2 electrons

The number of valence electrons in the table below is the same as the column number of the

element in the periodic table (for representative elements only)

Atoms often gain, lose, or share electrons to achieve the same number of electrons as the

noble gas closest to them in the periodic table.

Because all noble gasses (except He) have filled s and p valence orbitals (8 electrons), many

atoms undergoing reactions also end up with 8 valence electrons.

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This observation has led to the Octet Rule: Atoms tend to lose, gain, or share electrons

until they are surrounded by 8 valence electrons

Note: there are many exceptions to the octet rule (He and H, for example), but it provides a

useful model for understanding the basis of chemical bonding.

I Ionic Bonds

Ionic bonds electrostatic forces that exist between ions of opposite charge.

These are strong bonds.

Ionic substances usually result from the interaction of metals and non-metals.

Note that the electrons (e-) are completely transferred from the metal to non-metal in

the formation of ionic bonds. This results in the formations of positive metal ions

(cations) and negative non-metal ions (anions).

E.g.

&

Properties

All pure ionic compounds are solids at room temperature. None are liquids or gases.

The boiling point and melting point are high in substances with ionic bond (1000 – 1500

oC)

Many ionic compounds are water soluble (they are polar and therefore will dissolve in

polar solvents such as water). When dissolved, they dissociate into ions and conduct

electricity. Cations go to cathode; and anions go to anode during electrolysis.

Ionic compounds are usually brittle and crystalline.

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II Normal covalent Bond

Unlike ionic bonds, covalent bonds are formed by the sharing of e- between two atoms, to

form a molecule that is more stable than the individual atoms.

Two H atoms with one e- in the outermost shell (valence shell) will share e

- to form a

molecule that is more stable than the individual atoms.

E.g.

The objective for the sharing of electrons is to stabilize the atoms, i.e. to make atoms with

a full valence shell of electrons. This is known as the octet rule

NB: 2, 8 or 18 electrons may complete the octet of electrons required by atoms.

Covalent bonds are formed by the sharing of:

(i) a single pair of electrons, e.g. H―H, Cl—Cl, H—Cl, etc.

(ii) two pairs of electrons, e.g. O=O, O=C=O, O=S=O, etc.

(iii) three pairs of electrons, e.g. C≡N, N≡N, C≡C, etc.

NB: ― one pair of electrons (i.e. 2 e-); = two pairs of electrons; ≡ three pairs of electrons

Dative Covalent Bond

Both atoms that form the bond do not contribute an electron each during sharing. Instead,

ONLY one of the atoms contributes both electrons for bond formation.

E.g. A + B AB normal covalent bond

But A + B AB dative covalent bond

All four (4) bonds are equivalent (another e.g. is H3O+)

NB: There is no real difference between a normal covalent and a dative bond; only how they

are formed.

Properties

Substances with covalent bonds are molecules.

Usually have low m.p. and are non-conductors, e.g. iodine, sucrose or glucose.

Are usually low mp solids, or liquids (e.g. ethanol) or gases (CO2 and CH4) at room

temperature.

They are usually sparingly soluble in water but normally soluble in organic solvents.

III Metallic Bonds

This is the type of bond present between the atoms in metals.

Metallic bonds are different from ionic bonds.

Then H+ + N

HHH N

HHH

H +

.H . H+ .H . Hsharing of e

-

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Here, there is an array of positive ions that remain fixed while their loosely held electrons

move freely throughout the lattice structure. In other words, the metal cations (Mn+

) are

surrounded by a ‘sea’ of mobile electrons, hence the concept of the ‘electron-sea’ model to

describe the type of bonding found in metals. The electrons are communally shared by all

the metal cations. Metals are able to conduct electricity due to the motion of these electrons.

Remember, that ionic compounds do not conduct electricity in solid state, since there are no

‘free and mobile’ electrons or ions.

Properties

This strong bonding generally results in dense, strong materials with high melting and

boiling points.

Metals are good conductors of electricity and heat because these 'free' electrons carry the

charge of an electric current when a potential difference (voltage!) is applied across a piece

of metal.

Typical metals also have a silvery surface (but remember this may be easily tarnished by

corrosive oxidation in air and water), malleable, ductile.

IV Polarity in Covalent Bonds

A nonpolar covalent bond is one in which the electrons are shared equally between two

atoms, e.g. H2, Cl2.

A polar covalent bond is one in which one atom (in a molecule) has a greater attraction

for the electrons than the other atom. (If this relative attraction is great enough, then the

bond is an ionic bond)

In polar covalent bonds, electrons are not shared equally.

Bond polarity is due to differences in electronegativity (EN) between two atoms in a covalent

bond, i.e. the ability of an atom in a molecule to attract the shared electrons in that bond.

Electrons spend more time around the atom that has the higher electron affinity.

The result is a dipole, in which one atom has a partial positive charge and the other has a

partial negative charge.

The H—F molecule is polar, i.e. the centres of +ve and –ve charges do not coincide. In the

same instance, the F—F molecule is non-polar. Polar compounds have a permanent dipole

moment (), while non-polar compounds do not have a permanent dipole moment ().

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Electronegativity and Bond Polarity

Compound

F—F

H—F

LiF

Electronegativity difference

4.0 – 4.0 = 0

4.0 – 2.1 = 1.9

4.0 – 1.0 = 3

Type of bond Nonpolar

covalent

Polar

covalent

Ionic

The degree of polarity of a molecule is described by its dipole moment, = Qr

Where, Q is the charges on the atoms and r is the distance between the charge centres.

V Intra and intermolecular Forces

Remember that particles in gas are relatively far apart and exert no attractive forces.

In liquids and solids, the distance between particles is very small or negligible and there is

very little empty space.

In liquids and in some molecular solids, attractive forces exist between the MOLECULES

and are known as intermolecular forces.

In contrast, intramolecular forces hold the ATOMS in a molecule or covalent compound

together. In ionic compounds, electrostatic forces hold the ions together.

Intermolecular forces are much weaker than intramolecular forces.

As a group, intermolecular forces are called van der Waal forces.

To understand the properties of solids and liquids, we have to look at some types of

intermolecular forces.

Ion-Dipole

Attractive electrostatic forces can also exist between polar and charged (ionic) molecules, and

these are termed ion-dipole forces

Involves an interaction between a charged ion and a polar molecule (i.e. a molecule with a

dipole)

Cations are attracted to the negative end of a dipole

Anions are attracted to the positive end of a dipole

The magnitude of the interaction energy depends upon the dipole moment of the molecule,

the charge of the ion (Q) and the distance (d) from the centre of the ion to the midpoint of

the dipole.

Ion-dipole forces are important in solutions of ionic substances in polar solvents (e.g. a salt

in aqueous solvent).

eg H Cl Cl H

intramolecular intramolecular

inermolecular

stronger stronger

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Dipole-Dipole Forces

These are forces of attraction between polar molecules, i.e. molecules with covalent bonds

having unlike or different atoms. When dipoles come close together, the positive end of one

molecule attracts the negative end of other molecules (like opposite poles of magnets).

In general, dipole-dipole forces are relatively weak (0.1 – 10 kJ/mol) attraction e.g. H-F, H-

Cl, H-Br, ethanol, etc.:

F H

+-

F H

+-

F H

+-

F H

+-

F H

+-

London (Dispersion) Forces

These are the weakest of all intermolecular forces

How do we explain why some non-polar molecules are liquids and even solids when we

expect them to be gases?

This force of attraction is present in atoms and between non-polar molecules, i.e. molecules

with identical atoms, e.g. Cl and Cl; H and H; N and N; etc.

This happens as a result when electrons in a molecule or atom momentarily shift. NB:

electrons are constantly in motion. This results in temporary polarization of the molecule or

atom (i.e. +ve and –ve regions temporarily exist in the same atom or molecule).

When a slight positive charge is developed at one end of one molecule, it induces a slight

negative charge in one end of the molecule next to it (i.e. attracting the electrons of the

neighbouring). This is called an INDUCED DIPOLE.

For a brief moment (an instance), a force of attraction exists between the molecules.

Their strength is influenced by the size and shape of the molecules involved.

The closer the molecules are the stronger will be the attractive forces.

Neutral atoms temporary dipole develops

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Hydrogen Bonding

This is a special type of dipole – dipole force……that involves H.

It is stronger than London and dipole – dipole forces.

It forms when a hydrogen atom that is covalently bonded to a strongly electronegative atom

is at the same time is weakly bonded to another electronegative atom (in another molecule)

that has at least one lone pair of electrons.

The elements that are usually involved in H-bonding are N, O and F

H-bonding is responsible for some special properties exhibited by water and other

molecules, e.g. higher boiling points and greater water solubility of many compounds.

O

H

H

O

H

H

O

H

H

O

H

H

O

H

H

Hydrogen Bond

O-H covalent bond

Hydrogen Bonding in water

VI Drawing Lewis Structures

The general procedure:

1. Sum the valence electrons from all atoms

Use the periodic table for reference

Add an electron for each indicated negative charge, subtract an electron for each

indicated positive charge

2. Write the symbols for the atoms to show which atoms are attached to which, and connect

them with a single bond

You may need some additional evidence to decide bonding interactions

If a central atom has various groups bonded to it, it is usually listed first: CO32-

, SF4

Often atoms are written in the order of their connections: HCN

3. Complete the octets of the atoms bonded to the central atom (H only has two)

4. Place any leftover electrons on the central atom (even if it results in more than an octet)

5. If there are not enough electrons to give the central atom an octet, try multiple bonds (use

one or more of the unshared pairs of electrons on the atoms bonded to the central atom to

form double or triple bonds)

Example 1: Draw the Lewis structure of phosphorous trichloride (PCl3)

This is an example of a central atom, P, surrounded by chlorine atoms

1. We will have 5(P) plus 21 (3*7, for Cl), or 26 total valence electrons

2. The general symbol, starting with only single bonds, would be:

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3. Completing the octets of the Cl atoms bonded to the central P atom:

4. This gives us a total of (18 electrons) plus the 6 in the three single bonds, or 24 electrons total.

Thus we have 2 extra valence electrons which are not accounted for. We will place them on the

central element:

5. The central atom now has an octet, and there is no need to invoke any double or triple bonds to

achieve an octet for the central atom. We are finished.

Example 2: Draw the Lewis structure for the NO+ ion

1. We will have 5 (N) plus 6 (O) minus 1 (1+ ion), or 10 valence electrons

2. The general structure starting only with single bonds would be:

3. Completing the octet of the O bonded to N:

4. This gives us a total of 6 plus 2 for the single bond, or 8 electrons. There are 2 unaccounted for

electrons and we will place them on the N:

5. There are only 4 atoms on the N atom, not enough for an octet, so lets try a double bond

between the N and O:

The oxygen still has an octet, but the N only has 6 valence electrons, so lets try a triple bond:

Each atom now has a valence octet. We are finished.

The brackets with the + symbol are used to indicate that this is an ion with a net charge of 1+

Now, try to derive the Lewis structure of the following: H2CO, ClO3-, NO2

-, PO4

3-.

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MOLECULAR SHAPES: THE VSEPR MODEL

NB: Before using this model, it is expected that students would already have a good grasp on the

drawing of Lewis structures (electron dot-cross diagrams) of molecules/complex ions.

VSEPR MODEL:- Valence-Shell Electron-Pair Repulsion model.

In this model the approximate molecular shapes are determined by the number of valence

electrons around the atoms. However, it’s the TOTAL number of electrons surrounding the

CENTRAL atom of a molecule/compound that is MOST important in the prediction of

molecular shapes.

BASIS OF THE VSEPR MODEL

(i) electrons in bonds and in lone pairs are considered as ‘CHARGE CLOUDS’

(ii) ‘CHARGE CLOUDS’ repel one another and stay apart as far as possible, thus causing

molecules to assume specific shapes.

(iii) Two steps to remember when applying the VSEPR model:

STEP 1

Count the total number of electron-charged clouds surrounding the atom of interest (i.e. the

central atom). This includes no. of bonded electrons + lone pairs/nonbonded electrons.

STEP 2

Predict the arrangement of charge clouds around each atom by assuming that the clouds

are oriented in space to be as far away from one another as possible.

How orientation is achieved – depends on number of charge clouds.

Read up on the VSEPR model, Hybrid Orbitals and Molecular Orbital Theory for tutorial

discussion.

WEAKNESS OF THE VSEPR-MODEL:

Lone pair – Lone pair > Lone Pair – Bonded Pair > Bonded Pair - Bonded Pair

electron-repulsion electron – repulsion electron – repulsion

and not all equal as predicted by the model. This results in varying different number of

shapes for similar number of ‘CHARGE CLOUDS’

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SUMMARY OF VSEPR MODEL

Number

of

bonds

Number of

Lone Pairs

Number

of charge

clouds

VSEPR

Prediction

Molecular

Geometry

Example

2

0

2

Linear

Linear

3 0 3 Trigonal

Planar

Trigonal

Planar

2

1

3

Trigonal

Planar

Bent

4 0 4 Tetrahedral Tetrahedral

3

1

4

Tetrahedral

Trigonal

Pyramidal

2

2

4

Tetrahedral

Bent

5 0 5 Trigonal-

bipyramidal

Trigonal-

bipyramidal

4 1 5 Trigonal-

bipyramidal

Sea-Saw

SF4

3 2 5 Trigonal-

bipyramidal

T-Shape ClF3

2 3 5 Trigonal-

bipyramidal

Linear I3-

B FF

F

S

O

O

H

C

HH

H

N

HH

H

HO

H

Cl P

Cl

Cl

Cl

Cl

Cl Be Cl

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Number

of

bonds

Number of

Lone Pairs

Number

of charge

clouds

VSEPR

Prediction

Molecular

Geometry

Example

6

0

6

Octahedral

Octahedral

5

1

6

Octahedral

Square

pyramidal

4

2

6

Octahedral

Square Planar

Use the ABxUy notation to predict molecular geometry:

where A = central atom, Bx = number of bonded e- clouds, Uy = number of unbonded e

- clouds.

VALENCE BOND THEORY

VSEPR model – Predict shapes, but say nothing about the electronic nature of covalent

bonds.

Valence Bond Theory – quantum mechanical model to describe bonding in molecules.

Covalent bond – 2 atoms approach each other closely enough so that a singly occupied

valence orbital of each atom overlaps. NB: each orbital contain at least one-electron with

opposite spins.

The now – paired electrons in the overlapping orbitals are attracted to the nuclei of both

atoms, and thus bonds the two atoms together.

Hybridization

Using C-atom as an example: EC = 1s2 2s

2 2px

1 2py

1 2pz

0 .

Carbon has 4 valence electrons, 2 of which are paired in a 2s orbital, and 2 of which are

unpaired in different 2p orbitals.

Question? How can carbon form 4 – bonds, if 2 valence electrons are paired (unavailable),

and only 2 available for sharing, i.e. only two available for bonding.

H H+

1s 1s

H H

H2 molecule

SF

F F

F

F

F

Sb

Cl

C Cl

Cl

Cl

XeF

F F

F

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Answer? Electrons must be promoted from the lower 2s orbital to the 2p orbitals, giving an

excited – state configuration with EC = 1s2 2s

1 2px

1 2Py

1 2Pz

1 , which now has all 4 valence

electrons available for bonding.

However, even after excitation, there are 3 2p electrons + 1 2s electrons. We know that

carbon C forms 4 equivalent bonds (e.g. CH4), meaning that it uses 4 equivalent orbitals for

bonding. The ANSWER is that ‘HYBRIDIZATION’ of 3 2p + 1 2s orbitals produce

four equivalent ‘sp3

hybrid orbitals’ as shown below:

NB: for sp3 , this means a combination of one s orbital and three p orbitals, and nothing to

do with the number of electrons involved in bonding.

SCHEMATICALLY (Bonding in methane)

NOTE

Read up on the sp2 hybridization in ethane, and sp hybridization in ethyne then draw the orbitals

for NH3, C2H4, H2O.

Follow: www.newagepublishers.com/samplechapter/000568.pdf

1s

2s

2p

1s

2s

2p

excitation

carbon:

ground state EC

carbon:

excited state EC

1s

4 x 2sp3

carbon:

4 x sp3

hybridization EC

hybridization

(e- promoted)

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Summary of basic relationship between hybridization and bond angles and geometry

Number of electron clouds around central

atom

Type of hybridization and shape

Two e- clouds/groups sp-hybridization; 180

o bond angle; linear

geometry. eg C2H2

Three e- clouds/groups sp

2-hybridization; 120

o bond angle; trigonal

planar; eg C2H4

Four e- cloud/groups sp3-hybridization; 109.5

o bond angle; C2H6,

H2O

Five e- cloud./groups sp3-d hybridization; 90 and 120

o bond angles

MOLECULAR ORBITAL THEORY

MO theory is more complex than the valence bond model, but sometimes gives a more

satisfactory accounting of chemical and physical properties.

Recall

Atomic orbital - , whose square gives the probability of finding the electron within a given

region of space in an atom.

Now: Molecular orbital theory considers the entire molecule rather than concentrating on

individual atoms.

Thus, MO refers to molecules, and are formed by the overlap of two or more molecular

orbitals. AO refers to atoms.

Therefore, molecular orbital – wavefunction (), whose square (2) gives the probability of

finding the electron within a given region of space in a molecule.

Like AOs, MOs have specific energy level and specific shapes. They can be occupied by a

maximum of 2 electrons with opposite spins.

The theory can predict electronic spectra and paramagnetism, when VSEPR and the V-B

Theories don't.

The major draw-back to this theory is that we are limited to talking about diatomic

molecules (molecules that have only two atoms bonded together), or the theory gets very

complex.

The MO theory treats molecular bonds as a sharing of electrons between nuclei. Unlike the

V-B theory, which treats the electrons as localized balloons of electron density, the MO

theory says that the electrons are delocalized. That means that they are spread out over the

entire molecule.

Now, when two atoms come together, their two atomic orbitals react to form two possible

molecular orbitals. One of the molecular orbitals is lower in energy. It is called the bonding

orbital and stabilizes the molecule. The other orbital is called an anti-bonding orbital. It is

higher in energy than the original atomic orbitals and destabilizes the molecule.

Below is a picture of the molecular orbitals of two hydrogen atoms come together to form a

hydrogen molecule:

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Mosby-Year Book Inc.

The MO Theory has five basic rules:

1. The number of molecular orbitals = the number of atomic orbitals combined

2. Of the two MO's, one is a bonding orbital (lower energy) and one is an anti-bonding

orbital (higher energy)

3. Electrons enter the lowest orbital available

4. The maximum # of electrons in an orbital is 2 (Pauli Exclusion Principle)

5. Electrons spread out before pairing up (Hund's Rule)

Below is a molecular orbital energy diagram for the hydrogen molecule. Notice that the two

AO's or atomic orbitals combine to form 2 MO's - the bonding and the anti-bonding molecular

orbitals. Also, notice that the five rules have been followed, the electrons having been placed in

the lowest energy orbital (rule 3) and have paired up(rule 4) and there are only two electrons in

the orbitals(rule 5).

Mosby-Year Book Inc.

If you notice at the very bottom of the above picture, "bond order" is mentioned. If a molecule is

to be stable, it must have a bond order greater that 0. Bond order is calculated as: 1/2 ( # of

electrons in bonding orbitals - # of electrons in anti-bonding orbitals). If the bond order is 0, the

molecule is unstable and won't form. If the bond order is 1 a single bond is formed. If the BO

(bond order) is 2 or 3 a double or triple bond will be formed respectively.

When the 2nd period atoms are bonded to one another, you have both 2s and three 2p orbitals to

contend with. When this happens, you have many more MO's:

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It is possible for the 2s orbital on one atom to interact with the 2pz orbital on the other. This

interaction introduces an element of s-p mixing, or hybridization, into the molecular orbital

theory. The result is a slight change in the relative energies of the molecular orbitals, to give the

diagram shown in the figure below.

Experiments have shown that O2 and F2 are best described by the model in the figure above, but

B2, C2, and N2 are best described by a model that includes hybridization, as shown in the figure

below. In other words, once the total number of electrons in the molecule exceeds 14, we use

the model at the top.

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Exercise: Draw the MO Diagram for the CO and NO+ molecule and calculate their bond order.

Predict the type of magnetism found.

Note: 2s and 2p interactions causes the order of the orbitals to change (see text)