Chemical Bonding 2

Question 1

Define covalent bonding.

Answer 1

Covalent bonding is the electrostatic forces of attraction between the positively charged nucleus of the bonded atoms and their shared pair of electrons.

Question 2

Describe the full process of the formation of a covalent bond.

Answer 2

When the atoms are far apart, no change is observed. As the distance between the atoms decreases, each nucleus starts to attract the other atom’s electron. The attraction continues to draw the atoms closer. However, nucleus-electron attraction, nucleus-nucleus and electron-electron repulsion exist simultaneously. As the distance continues to decrease, the repulsion becomes more significant. At some optimum inter-nucleus distance, the repulsive forces balances the attraction forces. The system has its minimum energy at this point, and a covalent bond is formed. The electron cloud that forms between the bonded atoms is called a molecular orbital.

Question 3

Name the types of covalent bonds.

Answer 3

1) Sigma (σ) bond
2) Pi (π) bond

Question 4

State the maximum number of σ bonds and π bond in a covalent bond.

Answer 4

1 σ bond, 2 π bonds

Question 5

How are Sigma (σ) bonds formed in a covalent bond?

Answer 5

Sigma bonds are formed by the ‘head-on’ overlap of two atomic orbitals.

Question 6

How are Pi (π) bonds formed in a covalent bond?

Answer 6

π bonds are formed by the ‘side-on’ overlap of the two orbitals.

Question 7

Suggest in terms of orbital overlap, why a π bond is weaker than a σ bond.

Answer 7

The side-on overlap for a π bond is less efficient than the head-on overlap for a σ bond. The region of overlap is smaller and the overlap is poorer. Hence, a π bond is weaker than a σ bond.

Question 8

Define bond order.

Answer 8

The number of covalent bonds formed between 2 atoms.

Question 9

State 2 differences between single bonds and multiple bonds.

Answer 9

Multiple bonds are stronger and shorter than single bonds.

Question 10

For single bonds, double bonds and triple bonds, state respectively:
1) the bond order
2) the number of shared electron pairs
3) the type of bond(s) formed

Answer 10

Single bond:
1) 1
2) 1
3) 1σ
Double bond:
1) 2
2) 2
3) 1σ, 1π
Triple bond:
1) 3
2) 3
3) 1σ, 2π

Question 11

Define a bond pair in a covalent compound.

Answer 11

An electron pair shared between the bonded atoms.

Question 12

Define a lone pair.

Answer 12

An electron pair that is part of an atom’s valence shell but not shared with another atom.

Question 13

State the octet rule.

Answer 13

Atoms tend to gain, lose or share electrons so as to have eight electrons in their outer electron shell.

Question 14

Draw the dot-and-cross diagram of a water molecule.

Answer 14

O-H, O-H, with O having 2 lone pairs

Question 15

Draw the dot-and-cross diagram of a nitrogen molecule.

Answer 15

6 electrons shared between the N atoms, with each N atom having 1 lone pair.

Question 16

Draw the dot-and-cross diagram of HCN.

Answer 16

6 electrons shared between the C atom and N atom, with the N atom having 1 lone pair

Question 17

Define electronegativity.

Answer 17

The relative ability of an atom in a molecule to attract the shared electrons in a covalent bond.

Question 18

Describe when will electronegativity increase or decrease in a Periodic Table.

Answer 18

Electronegativity increases across a Period and decreases down the Group.

Question 19

Draw the Lewis Structure of a C2H6 molecule.

Answer 19

All bonds are drawn with one line, no lone pairs

Question 20

Draw the Lewis Structure of a carbon dioxide molecule.

Answer 20

O = C = O

Question 21

Draw the Lewis Structure of a beryllium chloride (BeCl2) molecule.

Answer 21

Cl - Be - Cl

Question 22

Draw the dot-and-cross diagram of PCl4^+.

Answer 22

1 whole square bracket, 1 pair of electrons shared between the P atom and each of the 4 Cl atoms. One electron from P is lost to form the cation

Question 23

Draw the dot-and-cross diagram of hydroxide (OH^-).

Answer 23

1 whole square bracket, O atom gains two electrons: one from H atom, the other gained to form the anion

Question 24

Draw the dot-and-cross diagram of a carbonate ion.

Answer 24

1 whole square bracket,
- 2 lone pairs in each of the 3 O atoms
- 4 electrons shared between the C atom and one of the O atom
- each of the other 2 O atoms with only 2 electrons shared with the C atom gains 1 electron, causing the charge of 2-

Question 25

Draw the dot-and-cross diagram of barium oxide.

Answer 25

square brackets,
- Ba ion has a charge of 2+
- O has a charge of 2-

Question 26

Draw the dot-and-cross diagram of BaO2 (barium peroxide).

Answer 26

square brackets,contains a cation and an anion
- Ba ion has a charge of 2+
- O2 has a charge of 2-
- within oxygen ions, 7 dots and 7 crosses in total, one extra cross and one extra dot form the charge of 2-

Question 27

Define bond length.

Answer 27

The distance between the nuclei of the 2 bonded atoms when their atomic orbitals overlap to form molecular orbitals.

Question 28

Define bond energy.

Answer 28

The average energy required to break a covalent bond in a gaseous substance.

Question 29

The bond energy of the C-C single bond is 350 kJ mol, while the bond energy of a C=C double bond is 610 kJ mol. Why is the bond energy of the double bond less than double of that of the single bond?

Answer 29

A C-C single bond is made up of a σ bond while a C=C bond is made up of a σ bond and a π bond. Since π bonds have more efficient orbital overlap and are stronger than σ bonds, thus the energy of a C=C double bond is not twice that of a C-C single bond.

Question 30

State the relationship between bond length and bond energy.

Answer 30

The shorter the bond length, the higher the bond energy. A larger bond energy means that the bond is stronger, with greater forces of attraction between nuclei and electron clouds.

Question 31

Define non-polar bonds.

Answer 31

Non-polar bonds consist of two atoms with the same electronegativity, and the bonding electrons are shared equally.

Question 32

Define polar bonds.

Answer 32

Polar bonds consist of atoms with different electronegativity, and there is unequal sharing of bonding electrons.

Question 33

In a polar bond, which atom acquire a partial negative charge and why?

Answer 33

The more electronegative atom acquires a partial negative charge as it has a greater tendency to attract the bonding electrons to itself.

Question 34

How to determine the direction of a polar arrow in a polar covalent bond?

Answer 34

The arrow points from the less electronegative atom to the more electronegative atom.

Question 35

Describe the relationship between the difference in electronegativity in the polar bond VS the bond polarity and its dipole moment.

Answer 35

The greater the difference in electronegativity, the greater the bond polarity and the greater its dipole moment.

Question 36

For a covalent bond of same bond order and similar bond length, which type of bond is stronger?

Answer 36

A polar bond is stronger than a non-polar bond.

Question 37

Draw the Lewis structure of a carbon dioxide molecule, describing its shape.

Answer 37

O=C=O
linear shape

Question 38

Draw the Lewis structure of a CH4 molecule, describing its shape.

Answer 38

1C bonding with 4 H, all single bonds
tetrahedral shape (all the bond angles are obtuse)
*all tetrahedral-shaped covalent bonds with only 2 types of atoms are non-polar

Question 39

Draw the Lewis structure of a NH3 molecule, describing its shape.

Answer 39

1N bonding with 3H, all single bonds; N has a lone pair
trigonal pyramidal shape

Question 40

Draw the Lewis structure of a water molecule, describing its shape.

Answer 40

1O bonding with 2H, all single bonds; O has 2 lone pairs
bent / V-shaped

Question 41

Suggest 2 reasons why a non-polar molecule has no overall dipole.

Answer 41

1: There is no polar bond in the molecule.
2: There are polar bonds in the molecule but the dipole moments cancel out such that the overall dipole moment is zero.

Question 42

State the net dipole moment for the following molecules:
1: CO2
2: CCl4

Answer 42

1: zero
2: zero

Question 43

State 3 properties of polar molecules.

Answer 43

1: A polar molecule has an overall permanent dipole.
2: There must be at least one polar bond in the molecule.
3: The dipole moments of the polar bonds do not cancel out such that the net dipole moment is not zero.