Periodicity

Question 1

What is ionisation energy?

Answer 1

The energy required to remove completely an electron from the gaseous atom or molecule in its ‘ground state’
First ionisation energy: M(g) → M⁺(g) + e⁻

Question 2

Ionisation energy generally requires?

Answer 2

An input of energy, the ionisation energy will be positive

Question 3

What is the general trend in ionisation energy going from left to right across the periodic table?

Answer 3

Increase in ionisation energy from left to right
Due to the effect of nuclear charge increasing from left to right, pulling electron density in and resulting in more energy needed to remove electrons

Question 4

What is the general trend in ionisation energy going down groups on the periodic table?

Answer 4

Ionisation energy decreases group down groups on the periodic table
Sharp drop occurs when we start a new row
The electron being removed comes from a shell further and further away from the nucleus, and there being more shielding by other shells from nuclear charge

Question 5

In general there is an increase in ionisation energy across a period
But there are deviations from the trend at B. Why?

Answer 5

• From Be to B, the single 2p electron is shielded by the 2s pair and hence is more easily removed

Question 6

In general there is an increase in ionisation energy across a period
But there are deviations from the trend at O. Why?

Answer 6

• Oxygen has a lower ionisation than nitrogen due to electron repulsion
• Oxygen has a paired electron in the p-orbital which will experience repulsion due to the other electron hence will be easier to remove
• It will also result in a more stable half-filled p-orbital when oxygen is ionised

Question 7

There is a general trend of decreasing ionisation energy going down a group as the valence e- is further away from the nucleus
There are deviations of this trend at Ga. Why?

Answer 7

• The reason why Ga is higher than expected, is due to Ga being preceded by the first row of d-block elements
• The corresponding 10 e- in the 3d orbital do not provide very effective shielding for the 4s and 4p e-, as they are much more diffuse
• Hence more energy is required to remove an inital electron from Ga

Question 8

There is a general trend of decreasing ionisation energy going down a group as the valence e- is further away from the nucleus
There are deviations of this trend at Tl. Why?

Answer 8

In: [Kr] 4d10 5s2 5p1
Tl: [Xe] 4f¹⁴ 5d¹⁰ 6s² 6p¹
Tl is preceded by the first set of f electrons
Again the f electrons are weakly shielding - so the out electron experience a higher nuclear charge than expected
Relativistic effects impact on Tl ionisation energy too

Question 9

What is Electron affinity (EA)

Answer 9

Is the energy released when a gaseous atom, molecule or ion in its ‘ground state’ gains an electron
First electron affinity: X(g) + e⁻ → X⁻(g)

Question 10

What is the energy change like for electron affinity?

Answer 10

Since this process is normally favourable and energy will be given out, the electron affinity will be positive

Question 11

What is the general trend in electron affinity going across the periodic table going from left to right

Answer 11

• As we go across from left to right there is a general increase in electron affinity
• i.e. from boron to flourine the number gets more positive meaning more energy is released
• Due to becoming close to having a full valance octet
• But once we have the ocetet (i.e. nobel gases) the electron is being added to a new shell which is much higher in energy, hene the value is negative

Question 12

What is the general trend in electron affinity going down the periodic table?

Answer 12

• Most clearly seen for group 7 (Cl-At) the electron affinity energy is getting smaller going down the group
• Due to the extra electron being added to a shell further away from the nucleus, hence less nuclear charge holding that electron in place and not as stabilised

Question 13

Generally, electron affinity gets smaller down a group
Why is group 1 and group 2 and anomally to this?

Answer 13

• Due to the small size of the atoms in group 1
• Results in electrostatic repulsion between the new electrons being added
• This can destabilise some of the positive stabilisation of a full valance octet

Question 14

Why is nitrogen so negative

Answer 14

• Nitrogen has a symmetrically half filled set of p-orbitals
• Which is lost when an electron is added and it has lost the extra exchange energy
• This trend remains across group 15 but becomes less and less as the valance orbtials get bigger, resulting in less electrostatic repulsions

Question 15

Why are all the electron affinities of group 1 positive?

Answer 15

• When we are adding an extra electron we are filling the s orbital = more stable
• So a positive amount of energy is released

Question 16

Why are all the electron affinities of group 2 negative

Answer 16

• Adding an electron to a group 2 element will result in the p orbtial being filled
• This is less stable than just having the s orbital filled as the p-orbital is higher in energy
• Hence the electron affinity is negative

Question 17

Why are the electron affinities for group 13 not as positive as group 1

Answer 17

• The values are positive for group 13 as it is moving closer towards a symmetrical half filled p-orbital
• But there isnt the same stabilisation of a filled orbital, like there is for group 1

Question 18

Define electronegativity

Answer 18

The ability of an atom to attract electron density towards itself in a molecule
(measured using the Pauling scale)

Question 19

How is electronegativity worked out through Pauling Electronegativity (χ)

Answer 19

• The difference between the measured and expected bond energies is due to the difference in electronegativity between elements
• e.g. (H = 2.2.; Cl = 3.2)
• The covalent bond is polar. This leads to extra electrostatic/coulombic attraction and higher than expected bond strength

Question 20

Use the following equation to work out the electronegativity of Carbon
notes Xa has to be larger than Xb

Answer 20

= 0.102√(〖E_AB-(E_AA×E_BB)〗^(1/2) )
= 0.102√(〖285-(155×348)〗^(1/2) )
=1.58
= 3.98-1.58 = 2.4

Question 21

How does electronegative change going accross the period table from left to right
And down a group

Answer 21

• We increase in electronegativity from left to right and down a group
• Because electrons are further away from the nucleus

Question 22

Electronegativity will vary as a function of….

Answer 22

• What is bonded to the element
• Oxidation state
• Hybridiation

Question 23

How does hybridiation affect electronegativity?

Answer 23

• e.g. sp>sp²>sp³
• the more s charater, the more tightly the electrons are held
• Since electrons penetrate closer to the nucleus in an s orbital

Question 24

The trend in electronegativity mathces the trend in metallic to non-metallic character
How?

Answer 24

• Non-metals have high electronegativity, metals have low electronegativity and metalloids
• Trends in electronegativity will then reflect trends in properties (i.e. heat/charge/bonding)
  (the p-block is the only part of the periodic table with non-metals (s, d and f block all metals)

Question 25

What is the van Ketelaar triangle

Answer 25

For simple s-p compounds the type of bonding can be predicted through electronegativity
With ionic parameter on the y axis: difference in electronegativity of the 2 atoms in the binary compound
And the covalency parameter on the x axis: the average electronegativity of the 2 atoms in the binary compound
This can then be related to properties