7 Periodicity

Question 1

How did Mendeleev arrange the 60 known elements? (2 mark)

Answer 1

By atomic mass, and in groups with similar properties.

Question 2

What 3 things did Mendeleev do that made the periodic table more similar to the one we have today? (3 marks)

Answer 2

If the group properties did not fit, he swapped the elements around.
Left gaps for undiscovered elements.
He predicted properties of the missing elements from the group trends.

Question 3

What are the elements arranged in now? (1 mark)

Answer 3

Atomic number

Question 4

What is periodicity? (1 mark)

Answer 4

When there is a repeating trend in properties of the elements across a period.

Question 5

Where is the s block on the periodic table? (1 mark)

Answer 5

Groups 1 and 2

Question 6

Where is the p block on the periodic table? (1 mark)

Answer 6

Groups 3-12

Question 7

Where is the d block on the periodic table? (1 mark)

Answer 7

Groups 13-18

Question 8

What group did Mendeleev omit entirely, and why he was unaware of it? (2 marks)

Answer 8

Noble gases, group 18, group 0, group 8
They are unreactive

Question 9

What is ionisation energy? (1 mark)

Answer 9

Ionisation energy measures how easily an atom loses electrons to form positive ions.

Question 10

What three factors affect ionisation energy? (3 marks)

Answer 10

Atomic radius
Nuclear charge
Electron shielding

Question 11

What energy does the first electron lost have? (1 mark)

Answer 11

Is on the highest energy level.

Question 12

How does atomic radius affect ionisation energy? (2 marks)

Answer 12

The greater the distance between the nucleus and the outer electrons the less the nuclear attraction.

Question 13

How does nuclear charge affect ionisation energy? (2 marks)

Answer 13

The more protons there are in the nucleus of an atom, the greater attraction between the nucleus and outer electrons.

Question 14

How does electron shielding affect ionisation energy? (2 marks)

Answer 14

Electrons are negatively charged so the inner-shell electrons repel the outer-shell electrons.
This repulsion reduces the attraction between the nucleus and the outer electrons.

Question 15

What is the general trend for ionisation energies in periods? (1 mark)

Answer 15

A general increase.

Question 16

Why does boron have a lower first ionisation energy than beryllium? (4 marks)

Answer 16

The 2p electron in boron
has a higher energy than
one of the 2s electrons in beryllium,
so it requires less energy to be removed (easier to remove).

Question 17

Why does oxygen have a lower first ionisation energy than nitrogen? (4 marks)

Answer 17

In oxygen, all the electrons in the 2p sub-shell are paired, so one of the paired electrons is removed.
In nitrogen, there is a lone electron in the 2p sub-shell, so this electron is removed first.
The paired electrons repel each other (in oxygen), so an electron is easier to remove from an oxygen atom.
Therefore the first ionisation energy of an oxygen atom is lower than the first ionisation energy of a nitrogen atom.

Question 18

Why is there a significant drop in first ionisation energy from neon to sodium? (4 marks)

Answer 18

The highest energy electron in neon is in the 2p subshell.
The highest energy electron in sodium is in the 3s subshell.
The 3s electron has a higher energy than 2p, so it is a lot easier to remove,
therefore the first ionisation energy for sodium is a lot lower than neon’s.

Question 19

Why does neon have a higher first ionisation energy than fluorine? (4 marks)

Answer 19

Neon’s outer shell is full, causing the first ionisation energy to rise significantly.
It also has more electrons, that are attracted to the nucleus, making neon slightly smaller than fluorine, which means more energy is requisites to remove an electron.

Question 20

Why does aluminum have a lower first ionisation than magnesium? (4 marks)

Answer 20

The highest energy electron in magnesium is in the 3s subshell.
The highest energy electron in aluminum is in the 3p subshell.
The 3p subshell has a higher energy than the 3s subshell,
so the first electron is easier to remove in aluminum,
therefore it has a lower first ionisation energy than magnesium.

Question 21

Why does sulphur have a lower first ionisation energy than phosphorous? (4 marks)

Answer 21

None of the electrons in the 3p subshell in phosphorous are paired.
2 of the electrons in the 3p subshell in sulphur are paired,
so they repel each other,
this makes it easier to remove one of them,
therefore the first ionisation energy is lower in sulphur than in phosphorous.

Question 22

Explain the 2, 3, 3 rule.
Hint: Trend in first ionisation energy across groups 2 and 3.

Answer 22

The ionisation energies in the period go up for 2 elements, then down a bit, then up for 3 elements then down a bit, then up for 3 elements (to make a zigzag), then drops back to where the first element was for the new period.

                                  x <-- noblegas
                                /
                            x
                 x       /
                 /\    /
               /    \x
             x
  x       /
  /\    /
/    \x x <-- first element of the period

Question 23

State and explain the trend in first ionisation energies down a group. (4 marks)

Answer 23

Going down a group, the atomic radius increases,
there are more inner shells, so shielding increases,
therefore nuclear attraction to outer electrons decreases,
so the first ionisation energy decreases.

Question 24

What state are all metals at RTP. except one (name the exception)?
(2 marks)

Answer 24

Solid, apart from mercury.

Question 25

What is metallic bonding? (2 marks)

Answer 25

Strong electrostatic attraction between cations and delocalised electrons.

Question 26

Why do simple molecular structures have low melting and boiling points? ( 2 marks)

Answer 26

They have weak induced dipole-dipole interactions (intermolecular forces).
So little energy is needed to overcome these.

Question 27

What 3 elements form giant covalent lattice, in period 2 and 3? (2 marks)

Answer 27

Carbon and silicon and boron

Question 28

How many nitrogen atoms are in a molecule? ( 1mark)

Answer 28

2

Question 29

How many atoms are in a molecule of phosphorous? (1 mark)

Answer 29

4

Question 30

How many atoms are in a sulphur molecule? (1 mark)

Answer 30

8

Question 31

Why are diamond and silicon non-conductors of electricity? (2 marks)

Answer 31

All 4 outer shell electrons are involved in covalent bonding,
so none are available for conducting electricity.

Question 32

What structure and bond angle do diamond and silicon form? (2 marks)

Answer 32

tetrahedral
109.5 degrees

Question 33

Which elements have a simple molecular structure, in periods 2 and 3? (1 mark)

Answer 33

All the elements to the right of carbon.

Question 34

Describe the graph for the trend in melting points of period 2. (3 marks)

Answer 34

Increases across the period until carbon.
Sharp drop (lower than first element in period).
Remaining elements have really low melting points too.

Question 35

Describe the graph for the trend in melting points of period 3. (3 marks)

Answer 35

Increases until silicon.
Sharp drop (lower than first element in period).
Gradual decrease for remaining elements.

Question 36

Why can ionic compounds conduct electricity when molten? (2 marks)

Answer 36

The ionic lattice collapses
So the ions are now able to move and conduct electricity