C7-8. Electronic Structure and Periodicity

Question 1

What is the oxidation number of an element on its own?

Answer 1

This is always zero.

Question 2

How do we assign oxidation numbers in compounds?

Answer 2

• Each atom within the compound is assigned an oxidation number, which is usually identical to its ionic charge.
• The sum of these is always equal to the charge of the complex ion.

Question 3

What are the oxidation numbers of the following exceptional cases of atoms in compounds?
1. H in hydrides
2. O in peroxides
3. O bonded to F

Answer 3

1. This is -1, such as in NaH.
2. This is -1, such as in H2O2.
3. This is +2, such as in F2O.

Question 4

What can we say about the oxidation number of an element bonded to a more electronegative element?

Answer 4

The oxidation number of the less electronegative atom will have a positive oxidation number.

Question 5

What convention must we follow when writing oxidation numbers?

Answer 5

The sign always comes before the number itself, unlike with ionic charges where it comes afterwards.

Question 6

How do we name atoms and compounds which can have different ionic charges? Give an example.

Answer 6

• We present the oxidation number of the atom or compound as a Roman numeral given after the name of an element.
• Nitrate (III) i.e. nitrite, NO2- has a +3 oxidation number, and nitrate (V) i.e. NO3- has a +5 oxidation number.

Question 7

How do we define an oxidation reaction?

Answer 7

This is the loss of electrons, and the gain in oxidation number. Remember the OILRIG mnemonic.

Question 8

How do we define a reduction reaction?

Answer 8

This is the gain in electrons, but the loss in oxidation number. Think of it as, the oxidation number reduces.

Question 9

What is a redox reaction defined as?

Answer 9

A reaction in which species are both oxidised and reduced - i.e. if one of the processes happens, so must the other. This could be for example a metal-acid reaction.

Question 10

What is first ionisation energy defined as?

Answer 10

• The energy required to remove one electron from each atom In one mole of gaseous atoms of an element, in order to form one mole of gaseous 1+ ions.
• Think of it is a measure of how easily an atom loses electrons to form ions.

Question 11

Why must ionisation occur in the gaseous state?

Answer 11

This is because the removal of electrons is only achievable in high-energy states.

Question 12

What is the general formula for 1st ionisation energy?

Answer 12

X(g) = X+(g) + e-

Question 13

Why does first ionisation energy decrease down the group?

Answer 13

• The atomic radius increases due to the addition of new shells
• This also gives rise to more shielding
• Therefore the attraction between valence electrons and the nucleus is reduced

Question 14

Why does first ionisation energy generally increase across the period?

Answer 14

• Atomic radius decreases due to the addition of valence electrons in the same shell, and shielding is the same
• Nuclear charge increases due to the addition of protons
• Increasing attraction between valence electrons and nucleus

Question 15

How can we explain the group 2-3 dip in ionisation energies across the period?

Answer 15

• An electron is removed from a lower energy s-orbital in Group 2, whereas in Group 3 an electron is removed from a higher energy p-orbital.
• This requires less energy.

Question 16

How do we explain the group 5-6 dip in ionisation energies across the period?

Answer 16

• The Group 5 atom’s orbitals are singularly occupied, so the electron added in Group 6 must pair up, giving rise to repulsion as electrons have like charges.
• This means that less energy is required to overcome forces of attraction.

Question 17

What are successive ionisation energies defined as?

Answer 17

The energy required to successively remove electrons from each atom In one mole of gaseous atoms of an element.

Question 18

What do the ‘jumps’ on a graph of successive ionisation energies represent?

Answer 18

These represent the loss of an energy level (shell). It requires more energy to remove an electron from a shell closer to the nucleus, as there is less shielding, and nuclear attraction is greater.

Question 19

Why do successive ionisation energies increase logarithmically?

Answer 19

• As each electron is removed, the valence electrons are drawn closer to the nucleus since there are more protons than electrons.
• This gives rise to an increase in nuclear attraction, due to reduced atomic radius.

Question 20

How do we define periodicity?

Answer 20

A retreating trend of physical, chemical and atomic properties across the period, caused by regular and predictable variations in atomic structure.

Question 21

How are elements classified in to ‘blocks’ of the periodic table?

Answer 21

The ‘block’ an element occupies is defined by which sub-shell its highest-energy electron occupies: s, p or d.

Question 22

Explain the trend in atomic radius across period 3.

Answer 22

• This exhibits a decrease
• Nuclear charge increases, and electrons are added to the same shell; shielding is unchanged
• Greater attraction, electrons drawn closer to nucleus

Question 23

Explain the trend in electronegativity across period 3.

Answer 23

• This exhibits an increase
• Nuclear charge increases and atomic radius decreases
• Giving rise to stronger attraction between the nucleus and the two bonded/shared electrons

Question 24

Describe the trend in boiling point across period 3.

Answer 24

• Increases from sodium to magnesium
• Peaks at silicon
• Steep decrease to sulfur, then slow decrease to chlorine
• Lowest at argon

Question 25

Explain the increase in boiling point between sodium and magnesium.

Answer 25

• These elements undergo metallic bonding
• As we move across the period, the cation gains a higher charge; more delocalised electrons
• The atomic radius is also smaller
• This leads to greater electrostatic attraction

Question 26

Explain why the boiling point of Period 3 metals peaks at silicon.

Answer 26

• Silicon has a giant covalent structure
• This contains a very large (indefinite) number of covalent bonds which are very strong
• Therefore lots of energy is required to overcome the forces of attraction

Question 27

Why are the boiling points of elements from phosphorus to chlorine lower than those before them in the period?

Answer 27

• These elements have a simple molecular structure
• We only have to overcome much weaker London forces between molecules
• So less energy is required to overcome the intermolecular bonds

Question 28

Explain why sulfur has the highest boiling point out of itself, phosphorus and chlorine.

Answer 28

• Consider the molecular formulae of these elements: S8, Cl2, P4
• Sulfur has the most electrons in a relative large molecule
• So London forces are much stronger, requiring more energy to overcome

Question 29

Why does argon have the lowest boiling point of the Period 3 elements?

Answer 29

• Argon is monatomic
• We only have to overcome very weak London forces between atoms
• These require little energy to overcome