Chapter 8 - Periodicity

Question 1

Nickel is a metal with a high melting point.

State the block in the Periodic Table that contains nickel. (1)

Answer 1

D block

Question 2

Give the full electron configuration of the Ni2+ ion (1)

Answer 2

1s2 2s2 2p6 3s2 3p6 3d8 (4s0)

Question 3

Complete the electronic configuration of aluminium.

1s2 …………………………………………………………………………………………. (1)

Answer 3

1s2 2s2 2p6 3s2 3p1

Question 4

State the block in the Periodic Table to which aluminium belongs. (1)

Answer 4

P block

Question 5

The Ne atom and the Mg2+ ion have the same number of electrons.

Give two reasons why the first ionisation energy of neon is lower than the third ionisation energy of magnesium. (2)

Answer 5

• Mg+2 has more protons than Ne

- Mg+2 electrons closer to nucleus

Question 6

By reference to all the atoms involved explain, in terms of electrons, how Na2S is formed from its atoms

Answer 6

• electron transfer between Na to S

- 1 electron from each Na atoms

Question 7

Explain the meaning of the term periodicity as applied to the properties of rows of elements in the Periodic Table.

Describe and explain the trends in atomic radius, in electronegativity and in conductivity for the elements sodium to argon. (13)

Answer 7

Periodicity= trend in the properties of element across a period,
Elements in the same group have similar properties,

Atomic radius= decreases across the row, nuclear charge increases,
More attraction for electrons,

Electronegativity= increases across the row,
Nuclear charge increases,
Atomic radius decreases,
More attraction for shared electrons

Conductivity= decreases across the row,
Na to Al are metals,
Electrons are free to move in metals

Question 8

Write an equation to illustrate the process occurring when the second ionisation energy of magnesium is measured. (1)

Answer 8

Mg+(g) = Mg^+2 (g) + 2e-

Question 9

State and explain the trend in the melting points of the Period 3 metals Na, Mg and Al (3)

Answer 9

• increases,
• strong metallic bond,
• strong attraction between positive ions and delocalized electrons

Question 10

Explain the meaning of the term first ionisation energy (1)

Answer 10

Energy required to remove one mole of 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous +1 ions

Question 11

State and explain the general trend in first ionisation energies for the elements Na to P (3)

Answer 11

• increases,
• nuclear charge increases,
• similar shielding,
• atomic radius decreases

Question 12

State which one of the elements from Na to P has the highest melting point and explain your answer.

Answer 12

• silicon,
• giant covalent,
• strong covalent bonds need to be broken

Question 13

State which one of the elements from Na to P deviates from this general trend and explain why this occurs (3)

Answer 13

• aluminum,
• electron in 3p orbital,
• less energy needed to lose electron

Question 14

Explain why the melting point of sulphur, S8, is higher than that of phosphorus, P4 (1)

Answer 14

-van der waals forces

Question 15

State and explain the trend in melting point of the Group II elements Ca–Ba. (3)

Answer 15

• decreases,
• weak metallic bond,
• increase in size of an atom

Question 16

State the trend in atomic radius from phosphorus to chlorine and explain the trend (3)

Answer 16

• decreases,
• nuclear charge increases,
• similar shielding

Question 17

In terms of structure and bonding, explain why sulfur has a higher melting point than phosphorus (3)

Answer 17

• sulfur molecules larger than phosphorus,
• van der waals forces between molecules are stronger,
• more energy needed to loosen forces between molecules

Question 18

In terms of atomic structure, explain why the van der Waals’ forces in liquid argon are very weak (2)

Answer 18

• argon particles are single atoms with electrons closer to nucleus,
• cannot easily be polarized

Question 19

Explain why nickel is ductile (can be stretched into wires). (1)

Answer 19

Layers can slide over each other

Question 20

Explain, in terms of its structure and bonding, why nickel has a high melting point. (2)

Answer 20

Bonding= strong metallic bonds,
Structure= strong attraction between positive ions and delocalized electrons

Question 21

Describe the bonding in metals. (2)

Answer 21

• lattice of metal surrounded by delocalized electrons,

- strong attraction between positive ions and delocalized electrons

Question 22

Explain why the melting point of magnesium is higher than that of sodium. (3)

Answer 22

• greater nuclear charge,
• stronger attraction between positive ions and delocalized electrons,
• Mg has more delocalized electrons,
• Na is a smaller atom

Question 23

Explain why the first ionisation energy of sulphur is lower than would be predicted from the general trend (2)

Answer 23

• spin pair repulsion,

- repulsion between electrons in the 3p sub shell

Question 24

Explain how the gaseous atoms of rubidium are ionised in a mass spectrometer (2)

Answer 24

• electron gun,

- knocks out electrons

Question 25

State one reason why the first ionisation energy of rubidium is lower than the first ionization energy of sodium (1)

Answer 25

• Rb is a bigger atom,
• electrons are further away from the nucleus,
• more shielding in rubidium

Question 26

By reference to the relevant part of the mass spectrometer, explain how the abundance of an isotope in a sample of rubidium is determined. (2)

Answer 26

• detector,

- current electrical signal related to abundance