Periodicity

Question 1

Which side of the periodic table are metals?

Answer 1

Left

Question 2

What happens to metallic character as you go down a group?

Answer 2

increases

Question 3

Which group are called the alkali metals?

Answer 3

1

Question 4

Which group are called the alkaline earth metals?

Answer 4

2

Question 5

Which group are called the halogens?

Answer 5

7

Question 6

Which group are called the noble gases?

Answer 6

0

Question 7

As you go down a group what happens to the number of electrons in the outer shell?

Answer 7

Nothing, they have the same number of electrons in their outer shell but they are in a higher energy level.

Question 8

Why do elements in the same group have similar chemical properties?

Answer 8

Same number of electrons in their outer shells

Question 9

Why do we call the block containing group 1 and 2 the s block?

Answer 9

The electron(s) in the highest energy level are in s orbitals

Question 10

What is the atomic radius?

Answer 10

The distance from the centre of the nucleus to the outermost electron

Question 11

How do you calculate the atomic radius?

Answer 11

Measure the distance between two covalently bonded atoms nuclei (must be the same element) and divide by 2.

Question 12

What happens to atomic radii as you go across a period?

Answer 12

Decreases

Extra proton in nucleus, increase positive charge, electrons pulled closer.

Question 13

What is the definition of first ionisation energy?

Answer 13

The energy required to remove 1 mole of outermost electrons from one mole of a gaseous element.

Question 14

Describe the pattern of first ionisation energies across a period.

Answer 14

Rise. Then a drop. Then 2 rises. Then a drop. Then 2 rises.

Question 15

Explain the pattern of first ionisation energies across a period.

Answer 15

Rise-increase in nuclear charge. Drop-e- removed from p orbital not s (further away). Risex2-increase nuclear charge. Drop-e- paired in orbital, repel each other so easier to remove. Risex2 increase nuclear charge.

Question 16

What is the definition of electronegativity?

Answer 16

Power of an atom to attract electron density from a covalent bond. The more electronegative, the stronger the pull on electrons.

Question 17

What happens to electronegativity as you go across a period?

Answer 17

Increases (increase in positive nuclear charge)

Question 18

Explain the trend in electrical conductivity across a period

Answer 18

Increases group 1-3 as more delocalised electrons. Decreases at group 4, they form lattices but at high temps can conduct. Group 5-7 don’t at all as they form covalent bonds, no charged mobile particles.

Question 19

Explain the increase in melting/boiling point across group 3 from sodium to aluminium

Answer 19

Held together by metallic bonds, increase boiling points as more delocalised electrons and a higher nuclear charge.

Question 20

Explain the increase in melting point across group 3 from aluminium to silicon

Answer 20

Giant covalent structure in silicon means there’s lots of strong bonds to break as each silicon atom is bonded to 4 others.

Question 21

Explain the decrease in melting/boiling points across group 3 from silicon to argon

Answer 21

Phosphorous-argon are all made of separate molecules which only have LDF holding them together which are easy to break, the smaller the molecules the easier it is to break the LDF.

Question 22

Why is the boiling point of silicon lower than expected (lower than aluminium)?

Answer 22

Most the bonds were broken in the process of melting, and the bonds which are left are lower in energy than in liquid aluminium.

Question 23

What 3 factors affect ionisation energy?

Answer 23

Atomic radius
Number of inner electrons (shielding)
Number of protons (nuclear charge)

Question 24

Why does the first ionisation energy decrease down a group?

Answer 24

More shielding

Larger atomic radius

Question 25

Write the equation for the first ionisation energy of oxygen

Answer 25

O(g) –> O+(g) + e-

Question 26

Which group have the lowest first ionisation energies?

Answer 26

1

Alkali metals

Question 27

What is the bonding/structure of diamond?

Answer 27

Covalent - Each C forms 4 covalent bonds making a giant ionic lattice

Question 28

What is the melting point of diamond?

Answer 28

Very high as lots of strong covalent bonds need to be broken (also makes it hard)

Question 29

What is the electrical conductivity of diamond?

Answer 29

Poor - all electrons are in localised bonds, unable to move

Question 30

What is the solubility of diamond?

Answer 30

Insoluble in all solvents as nothing can provide enough energy to break the strong covalent bonds.

Question 31

What is the structure/bonding of graphite?

Answer 31

Covalent - each C forms 3 covalent bonds in a 2D structure forming layers. LDF hold the layers together forming a giant covalent lattice.

Question 32

What is the melting point of graphite?

Answer 32

Very high as lots of strong covalent bonds to break.

Question 33

How hard is graphite?

Answer 33

Soft as layers can slide over each other

Question 34

What is the electrical conductivity of graphite?

Answer 34

Conducts electricity as each C has a delocalised electron.

Question 35

What is the solubility of graphite?

Answer 35

Insoluble in all solvents as nothing provides enough energy to break the covalent bonds

Question 36

Characteristics of simple molecular bonding

Answer 36

Small separate molecules with LDF holding them together. Bigger molecule, more electrons, stronger forces. Melting/boiling points are low.