Trends In The Periodic Table

Question 1

Atomic Radius (Size of the Atom)

Answer 1

• As you move down a group, from top to bottom, atomic radius increases
• As you move across a period, from left to right, atomic radius decreases

Question 2

As you move down a group, from top to bottom, atomic radius increases

Answer 2

• Number of energy levels increases as you move down a group, number of electrons increases.
• Each subsequent energy level is further from the nucleus than the last.

Question 3

As you move across a period, from left to right, atomic radius decreases

Answer 3

• Electrons are added to the same energy level across the period.
• Protons are being added to the nucleus.
• Concentration of more protons in nucleus creates a “higher effective positive nuclear charge”.
• Stronger force of attraction pulls electrons closer to nucleus.

Question 4

As you move down a group, first ionization energy decreases

Answer 4

• Electrons further from nucleus & higher energy level are easier to remove down a group (Shielding).

Question 5

As you move across a period, first ionisation energy increases

Answer 5

• Atomic radius decreases, atom is smaller across the period.
• Outer electrons are closer and more attracted to nucleus.

Question 6

Nuclear Charge

Answer 6

Greater nuclear charge (atomic number), stronger nucleus pull on electrons in a given energy level (smaller radius, higher IE).

Question 7

The distance of the electron from the nucleus (atomic radius)

Answer 7

• Greater distance between nucleus and it’s electrons in outer energy levels, smaller force of attraction.
• Force of attraction less in energy levels further from nucleus (IE decreases down a group).

Question 8

The number of inner level electrons screening the outer electrons from the force of the nucleus - (screening effect)

Answer 8

• Electrons in higher levels have more electrons between them and nucleus, this ‘Screens’ the attractive force of the nucleus to some extent.
• Inner electrons repel the outer ones to some extent.

Question 9

Force of repulsion by electrons in an orbital

Answer 9

• Electron pair (same orbital) exert a certain amount of repulsion on each other.
• Easier to remove (less energy needed) an electron in pair than one from same level, not paired.

Question 10

Ionisation Energy Notation

Answer 10

X(g) + 1ˢᵗ IE -> X ⁺ (g) + e⁻
X(g) + 2ⁿᵈ IE -> X²⁺ (g) + e⁻
X(g) + 3ʳᵈ IE -> X³⁺ (g) + e⁻

Question 11

Anomalies (IE’s that don’t fit the trend)

Answer 11

• Group 2 & 3
• Group 5 & 6

Question 12

Why is the first ionisation energy of boron less than that of beryllium ?

Answer 12

• Be’s valence e⁻ are in 2s sub-energy level
• B’s valence e⁻ are in 2p > energy of e⁻ in 2s and so could be said to be further from the nucleus
• Less energy is needed to remove e⁻ from B-atom than Be-atom

Question 13

Why is the first ionisation energy of oxygen less than that of nitrogen?

Answer 13

• Both N and Os valence e⁻ are in the 2p sub-energy level
• O has one 2p orbital with an e⁻ pair
• N has one e- in each 2p orbital (3 unpaired e⁻)
• Paired e- in the same orbital repel each other
• Less energy is needed to remove an e- from O-atom than from N-atom

Question 14

Trends in Successive Ionisation Energies (and how this proved the existence of energy levels)

Answer 14

The large jump in ionization energy from one level to the next provides evidence for the existence of energy
levels in the atom, confirming Bohr’s theory

Question 15

Summary of factors affecting Ionisation Energy

Answer 15

• Nuclear Charge
• The distance of the electron from the nucleus (atomic radius)
• The number of inner level electrons screening the outer electrons from the force of nucleus (screening effect)
• Force of repulsion by electrons in an orbital