Structure 3

Question 1

Define periodicity

Answer 1

The regular
repetition of chemical and
physical properties as you
move across and down the
periodic table due to patterns in electronic configuration.

Question 2

Define a group.

Answer 2

A column of
the periodic table which
contains elements with
similar chemical properties
with the same number of
electrons in their outer or
valence shell.

• 18 groups exist.

Question 3

Define period.

Answer 3

A horizontal
row in the periodic table
which contains elements
with the same number
of shells, and with an
increasing number of
electrons in the outer
or valence shell, as the
period is crossed from left
to right.

Question 4

Going across the periodic table from left to right, the elements gradually change from

Answer 4

metal to non-metal

Question 5

What happens to metallic character down a group and across a period?

Answer 5

decreases across any period and increases down any group. This is due to the electronegativity (increases down group).

Question 6

Define metalloids.

Answer 6

A group
of chemical elements
intermediate in properties
between metals and
non-metals.

Question 7

What are the properties of metals?

Give appearance, physical state, m.p. and b.p. , ductility and malleability, thermal conductivity, and electrical conductivity.

Answer 7

lustrous
solids (except mercury)
typically high (except group 1 and mercury)
ductile and malleable
good thermal and electrical conductors

Question 8

Properties of metalloids?

Answer 8

lustrous
solids
typically high
brittle
average thermal and electrical conductors

Question 9

Properties of non-metals?

Answer 9

dull
solids , liquids , and gases
low (except carbon and silicon)
brittle if solid
poor thermal and electrical conductors

Question 10

which groups make up s-block?

Answer 10

Group I and II + Hydrogen and Helium

Question 11

which groups make up d-block?

Answer 11

Elements which have the form dx ns2.

Question 12

which groups make up f-block?

Answer 12

Lanthanoids and actinoids

Question 13

which groups make up p-block?

Answer 13

Groups 13 to 18. Have the form ns2 to npx

Question 14

Define atomic radius.

Answer 14

The total distance from an atom’s nucleus to the outermost orbital of electron

Question 15

What happens to atomic radius across a group and a period and why?

Answer 15

Across period: decreases. This is due to an increase in the nuclear charge. The electrons are more strongly electrostatically attracted towards the protons. New electrons are added to the same shell.

Across group: increases. Outer electrons enter new shells and are further away. The inner electrons shield the outer electrons, so they cannot feel the full nuclear charge.

Question 16

What happens to ionic radius across a group and period and why?

Answer 16

Down group: Increases. The number of shells increases which create a greater shielding effect for outer electrons. This outweighs nuclear charge.

Across period: if it is increasing negative charge, ionic radii increase as the extra electrons experience repulsion that increase ionic radius. The nuclear charge still remains the same.

If it is increasing positive charge, ionic radii decrease as the nuclear charge remains the same but there are fewer electrons which undergo greater electrostatic force of attraction.

Question 17

Define electronegativity

Answer 17

A measure of its affinity (ability to attract electrons) for the bonding pair of electrons in a
covalent bond

Question 18

Decreasing electronegativity connotes a increase in

Answer 18

metallic character

Question 19

What happens to electronegativity down and across a group and why?

Answer 19

Down group: decreases. This is due to increase in atomic radius, leading to a weaker attraction between nucleus and the shared pair of electrons.

Across period: increases. This is due to an increase in effective nuclear charge, leading to a stronger attraction between the nucleus and the electrons.

Question 20

Define electron affinity

Answer 20

The amount of energy released when one mole of electrons is gained by one mole of atoms of an element in the gaseous state to form one mole of gaseous ions. First EA is exothermic. Successive EA can be endo/exothermic

Question 21

Give the trends of the electron affinity

Answer 21

Down group: the absolute value for first electron affinity generally decreases before going down a lower group. As the atoms become larger the attraction for an additional electron is less.
Electron affinity become less exothermic going down the group.

Across period: absolute value for first electron affinity generally increases. Elements across a period have a higher effective nuclear charge, resulting in stronger attraction between the added electron and the nucleus.

Successive: it becomes less favourable. The repulsion between electrons in the same cloud starts to outweigh the attraction of new electrons to the positively charged nucleus. Therefore, successive electron affinities tend to be less negative or may even become positive.

Question 22

Define ionisation energy.

Answer 22

The ionisation energy (IE) of an element is the amount of energy required to remove one mole of electrons from one mole of atoms of an element in the gaseous state to form one mole of gaseous ions.

Question 23

Trends of ionisation energy

Answer 23

Across a period: increases. There is an increasing effective nuclear charge, so more energy is required in order to remove an electron.

Down a group: first ionisation decreases. Atomic radii increases and electrons in higher energy levels have a weaker attraction to the nucleus. Lesser energy is required to remove an electron.

Question 24

Exceptions to IE trend and why?

Answer 24

Boron has lower IE than Beryllium.

Beryllium has electronic configuration 1s2 2s2
Boron has electronic configuration 1s2 2s2 2p1.

Electrons in p orbitals are of higher energy and further away from the nucleus than electrons in the s-orbitals, therefore requiring lesser energy to remove.

Oxygen has lower IE than Nitrogen.

Nitrogen has electronic configuration 1s2 2s2 2p3
Oxygen has the electronic configuration 1s2 2s2 2p4.

Nitrogen has three singly occupied p orbitals but in oxygen, one orbital is doubly occupied. An electron in a doubly occupied orbital is repelled by the second electron and requires lesser energy to remove.

Question 25

Reaction between alkali metal and water

Answer 25

Example Na

2Na + 2H2O -> 2NaOH + H2.

Highly reactive and stored under oil to prevent reactions.

Question 26

Explain reactivity for group 1 metals with water

Answer 26

Reactivity increases going down the group. Elements have lower ionisation energies so it is easier for those with lower IEs to transfer out valence electrons to water.

Question 27

Reaction between halogens and halide ions

Answer 27

Group 17 elements highly reactive. Reactivity decreases down a group.

Displacement reaction:
Cl2 + 2KBr -> Br2 + 2KCl

Cl is more reactive than Br so it displaces it. It has more electron affinity so gains an extra electron from the bromide ion to form a chloride ion.

Question 28

metals form which types of oxides?

Answer 28

basic

Question 29

Reaction of basic oxide with H2O?

Answer 29

Na2O + H2O -> 2NaOH
MgO + H2O -> Mg(OH)2