Chemical Bonding 2
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Define covalent bonding.
Covalent bonding is the electrostatic forces of attraction between the positively charged nucleus of the bonded atoms and their shared pair of electrons.
Describe the full process of the formation of a covalent bond.
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.
Name the types of covalent bonds.
1) Sigma (σ) bond 2) Pi (π) bond
State the maximum number of σ bonds and π bond in a covalent bond.
1 σ bond, 2 π bonds
How are Sigma (σ) bonds formed in a covalent bond?
Sigma bonds are formed by the ‘head-on’ overlap of two atomic orbitals.
How are Pi (π) bonds formed in a covalent bond?
π bonds are formed by the ‘side-on’ overlap of the two orbitals.
Suggest in terms of orbital overlap, why a π bond is weaker than a σ bond.
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.
Define bond order.
The number of covalent bonds formed between 2 atoms.
State 2 differences between single bonds and multiple bonds.
Multiple bonds are stronger and shorter than single bonds.
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
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
Define a bond pair in a covalent compound.
An electron pair shared between the bonded atoms.
Define a lone pair.
An electron pair that is part of an atom’s valence shell but not shared with another atom.
State the octet rule.
Atoms tend to gain, lose or share electrons so as to have eight electrons in their outer electron shell.
Draw the dot-and-cross diagram of a water molecule.
O-H, O-H, with O having 2 lone pairs
Draw the dot-and-cross diagram of a nitrogen molecule.
6 electrons shared between the N atoms, with each N atom having 1 lone pair.
Draw the dot-and-cross diagram of HCN.
6 electrons shared between the C atom and N atom, with the N atom having 1 lone pair
Define electronegativity.
The relative ability of an atom in a molecule to attract the shared electrons in a covalent bond.
Describe when will electronegativity increase or decrease in a Periodic Table.
Electronegativity increases across a Period and decreases down the Group.
Draw the Lewis Structure of a C2H6 molecule.
All bonds are drawn with one line, no lone pairs
Draw the Lewis Structure of a carbon dioxide molecule.
O = C = O
Draw the Lewis Structure of a beryllium chloride (BeCl2) molecule.
Cl - Be - Cl
Draw the dot-and-cross diagram of PCl4^+.
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
Draw the dot-and-cross diagram of hydroxide (OH^-).
1 whole square bracket, O atom gains two electrons: one from H atom, the other gained to form the anion
Draw the dot-and-cross diagram of a carbonate ion.
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-