2.2 - Electrons, Bonding and Structure

Question 1

How many electrons are in the first 4 quantum energy levels (shells)?

Answer 1

1st = 2
2nd = 8
3rd = 18
4th = 32

Question 2

What is an orbital?

Answer 2

A region of space around the nucleus of an atom that can hold up to two electrons with opposite spins.

Question 3

Describe the shape of s and p-orbitals.

Answer 3

S = circle
P = figure of eight

Question 4

How many orbitals and electrons are in the sub-shells s, p, and d?

Answer 4

S = 1 orbital, 2 electrons
P = 3 orbitals, 6 electrons
D = 5 orbitals, 10 electrons

Question 5

Describe the relative energies of orbitals in sub-shells up to 4p.

Answer 5

From least to most: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p

Note - this is a general trend and can be broken.

Question 6

What are the general rules to remember for electron configurations?

Answer 6

4s fills up before 3d

4s leaves before 3d

Question 7

Define ionic bonding.

Answer 7

The electrostatic attraction between positive and negative ions.

Question 8

What are the ionic charges of groups 1, 2, 3, 4, 5, 6 and 7?

Answer 8

1 = 1+
2 = 2+
3 = 3+
4 =
5 = 3-
6 = 2-
7 = 1-

Question 9

What is a compound ion?

Answer 9

A group of covalently bonded atoms that have a charge.

Question 10

Give the formulae of the following compound ions: nitrate, carbonate, sulphate and ammonium.
What about these: hydroxide, silver and zinc.

Answer 10

Nitrate = NO3 -
Carbonate = CO3 2-
Sulphate = SO4 2-
Ammonium = NH4 +
Hydroxide = OH -
Silver = Ag +
Zinc = Zn 2+

Question 11

Describe the difference between a giant and a simple structure.

Answer 11

Giant - a substance with a huge network of intramolecular forces.
Simple - a substance with a few covalent bonds in it forming a small molecule (e.g. CO2). Held together by intermolecular forces.

Question 12

Describe the structure of ionic crystals.

Answer 12

Ionic crystals are made up of a giant ionic lattice. A lattice is a regular structure. A giant ionic lattice is made up of a giant network of ions that are electrostatically attracted to each other.

Question 13

Describe the bonding and structure of ionic compounds, as well as melting/boiling points and hardness.

Answer 13

Giant structure with ionic bonding (giant ionic lattice).
Very high melting and boiling points - a large amount of energy needs to be put in to overcome the strong electrostatic attractions and separate the ions.
Hard to scratch due to the strong nature of the bonding which requires such a large amount of energy to make any changes to the structure. Very brittle as the ions will only accept being in one particular arrangement (due to their opposite charges).

Question 14

Describe the electrical conductivity and solubility of ionic compounds.

Answer 14

Electrical conductivity in solids - none. The ions are held in a fixed position by the strong ionic bonds. This means they are not free to move. When molten or in solution - conducts. The ions are now free to move as they are no longer held together in the lattice structure.
Many are soluble in water due to the attraction between the ions and the polar water molecules. Insoluble in non-polar solvents.

Question 15

Define covalent bonding.

Answer 15

The strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.

Question 16

What is a lone pair?

Answer 16

Non-bonded electron pairs.

Question 17

What is a dative bond?

Answer 17

The electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms where both electrons come from the same atom.

Question 18

What is meant by the term average bond enthalpy?

Answer 18

A measure of the energy required to break a covalent bond. The higher the average bond enthalpy, the stronger the covalent bond (stronger bonds will require more energy to break).

Question 19

How is the shape of a molecule determined?

Answer 19

Electrons are negatively charged. Electron pairs repel one another equally. Lone pairs repel more than bonded pairs. It is determined by the number of bonded pairs.

Question 20

What is the shape and bond angle of methane? How many pairs does it have?

Answer 20

Tetrahedral - 109.5°

4 bonded pairs, 0 lone pairs.

Question 21

What is the shape and bond angle of ammonia? How many pairs does it have?

Answer 21

Triangular pyramidal - 107°

3 bonded pairs, 1 lone pair.

Question 22

What is the shape and bond angle of water? How many pairs does it have?

Answer 22

Non-linear (bent) - 104.5°

2 bonded pairs, 2 lone pairs.

Question 23

What is the shape and bond angle of carbon dioxide? How many pairs does it have?

Answer 23

Linear - 180°

2 double bonded pairs, 0 lone pairs.

Question 24

What is the shape and bond angle of boron triflouride? How many pairs does it have?

Answer 24

Trigonal planar - 120°

3 bonded pairs, 0 lone pairs.

Question 25

What is the shape and bond angle of sulphur hexaflouride? How many pairs does it have?

Answer 25

Octahedral - 90°

6 bonded pairs, 0 lone pairs.

Question 26

What does electronegativity mean?

Answer 26

The ability of an atom to attract the bonding electrons in a covalent bond.

Question 27

What decides how electronegative an atom is?

Answer 27

A combination of: nuclear charge (number of protons in the nucleus), amount of shielding, distance of the bond from the nuclear charge.

Question 28

What are the most electronegative elements?

Answer 28

F (most), O, N, Cl.

Question 29

Describe how the Pauling electronegativity values can be used to interpret electronegativity.

Answer 29

The higher the value, the mire electronegative an atom is.
The difference in electronegativities between two atoms influences the type of bond they will form: the greater the distance the more ionic the bond will be; a difference of less than 0.4 gives non-polar covalent bonds; a difference between 0.4-2.0 gives polar covalent bonds; a difference of more than 2.0 gives ionic bonds.

Question 30

Give an example of what a polar molecule is, using electronegativity.

Answer 30

HCl - the chlorine is more electronegative than the hydrogen so attracts the bonded electrons more.
This sets up permanent dipoles, with the chlorine slightly positive and the hydrogen slightly negative.
A polar bond has been created.

Question 31

How can you tell if a molecule will be polar or non-polar?

Answer 31

One end of the molecule must be different in charge from the other end of the molecule to be polar - in non-polar molecules, the ends are all the same.

Question 32

Explain the formation of a polar molecule and give an example.

Answer 32

The molecule is asymmetrical (the polar bonds are not arranged symmetrically).
The dipoles do not cancel each other out.
The molecule is polar.
CHCl3 is an example.

Question 33

Explain the formation of a non-polar molecule and give an example.

Answer 33

The molecule is symmetrical (the polar bonds are arranged symmetrically).
The dipoles cancel each other out.
The molecule is non-polar.
CCl4 is an example.

Question 34

What is the difference between intermolecular and intramolecular?

Answer 34

Intermolecular (between) = van der Waals’, permanent dipole-dipole, hydrogen bonds. Forces arise between simple molecular substances.
Intramolecular (within) = covalent and ionic. Within a substance.

Question 35

Describe the different intermolecular forces.

Answer 35

Induced dipole-dipole interactions - Random movement of electrons results in an uneven distribution of electrons. This creates an instantaneous dipole in a molecule/atom. This induces dipoles in neighbouring molecules/atoms.
Permanent dipole-dipole interactions - A difference in electronegativity between 2 atoms means the electrons in the covalent bond are pulled more towards the most electronegative element. This sets up permanent dipoles. The area around the more electronegative element will be slightly negative. The area around the less electronegative element will be slightly positive.
Hydrogen bonds - Formed when hydrogen is bonded to very electronegative elements (O, N, F). Bond is formed between a hydrogen in one molecule and a lone pair on another molecule.

Question 36

Describe and explain the anomalous properties of water.

Answer 36

Relatively high boiling point - H bonds are stronger than other intermolecular forces. More energy is needed to overcome them. This results in a higher boiling point.
Ice is less dense than water - ice is an open lattice (hydrogen bonds are relatively long). In a solid state the H2O molecules are held apart by the H bonds.

Question 37

Describe the structure and bonding of simple covalent substances.

Answer 37

Simple covalent substances form simple molecular lattices.

These are made from covalently bonded molecules attracted to each other by intermolecular forces.

Question 38

Describe the bonding and structure of simple covalent structures and explain some of their properties.

Answer 38

Simple structures with covalent bonding in the molecule and intermolecular forces between the molecules.
Melting and boiling point is low. Molecules are only held together by intermolecular forces which are weak and do not require much energy to break.
Do not conduct. No free electrons or ions.
If they are very polar and capable of forming H bonds with the water molecules then they will be soluble. Non-polar molecules will be insoluble.

Question 39

Describe how to identify the type of molecular force within a substance.

Answer 39

First check that it is a simple molecular/atomic substance (CO2, NH3 or elements like Cl2, Ar).
Is there an electronegativity difference between the atoms? No = Van der Waal forces.
Yes = Is the molecule polar? No = Van der Waal forces.
Yes = Are there any of the following bonds: F-H, O-H, N-H? No = Permanent dipole-dipole attractions.
Yes = Hydrogen bonding.