Periodicity

Question 1

What is the first ionisation energy?

Answer 1

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

Question 2

What is the equation for the first ionisation energy?

Answer 2

X(g)-> X+(g) + e-

Question 3

What is ionisation energy?

Answer 3

A measure of how easily an atom loses electrons to form positive ions.

Question 4

What are 3 factors that affect ionisation energy?

Answer 4

Atomic radius, the greater the atomic radius the weaker the attraction. Nuclear charge, the more protons, the greater the attraction. Electron shielding, as number of shells increase the shielding increases.

Question 5

What is the second ionisation energy?

Answer 5

Energy required to remove one electron from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions

Question 6

What is the equation for second ionisation energy?

Answer 6

X+(g) -> X2+(g) + e-

Question 7

How are elements arranged in a periodic table?

Answer 7

Order of increasing atomic number

Question 8

What is meant by periodicity?

Answer 8

The repeating trends in chemical and physical properties

Question 9

How can the electron configuration be written in short?

Answer 9

The noble gas before the element is used to abbreviate

Question 10

Describe the trend of ionisation energy across period 3.

Answer 10

First ionisation energy increases because
- Nuclear charge increases
- Decreased atomic radius
- Same electron shielding
Which means more energy is needed to remove an electron

Question 11

Why does ionisation energy decrease from Mg to Al?

Answer 11

Al is in the 3p orbital which has a higher energy level than 3s in Mg so its easier to lose an electron

Question 12

Why does ionisation energy decrease from phosphorus to sulphur?

Answer 12

In sulphur, the 2 electrons in the first 3p orbital repel making it easier to lose an electron

Question 13

Why does first ionisation energy decrease between group 2 and 3?

Answer 13

Group 3 outer most electrons are in the p orbitals whereas in group 2 they are in s orbital, so the electrons are easier to be removed

Question 14

Does first ionisation energy increase or decrease at the start of the next period?

Answer 14

Decrease, as there is an increase in atomic radius and increase in shielding

Question 15

Why does first ionisation energy decrease between group 5 and 6?

Answer 15

The group 5 electrons in the p orbitals are single electrons and in group 6 the outermost electrons are spin paired, with some repulsion, making it easier to remove

Question 16

Why does ionisation increase or decrease down a group?

Answer 16

Decreases because
- Shielding increases
- Atomic radius increases
- The increase in protons is outweighed by the increase in distance and shielding

Question 17

What happens in metallic bonding?

Answer 17

Each atom donates its outer shell electrons to a shared pool of electrons (delocalised electrons), its the strongest type of bond between cations and delocalised electrons

Question 18

What keeps atoms in metallic bonds in fixed positions?

Answer 18

The cations are in fixed positions maintaining the structure and shape of the metal

Question 19

What makes the metallic bonds strong?

Answer 19

Strong attraction between positive ions and delocalised electrons

Question 20

Do metals conduct electricity?

Answer 20

They conduct in solid and liquid states, when a voltage is applied the delocalised electrons can move, carrying the charge

Question 21

Are metals soluble?

Answer 21

Metals do not dissolve, any interactions with water would lead to a reaction rather than dissolving.

Question 22

What do non-metallic elements exist as?

Answer 22

Simple covalently bonded molecules, in a solid state they form a simple molecular lattice held together by weak forces

Question 23

How many covalent bonds form between carbon lattices in diamond form and silicon?

Answer 23

They use their 4 outer shells to form 4 covalent bonds

Question 24

What are giant covalent lattices melting and boiling points?

Answer 24

They have high melting and boiling points due to the strong covalent bonds

Question 25

Are giant covalent lattices soluble?

Answer 25

Covalent bonds are far too strong to be broken down by interactions with solvents (insoluble)

Question 26

Are carbon (diamond) and silicon able to conduct electricity?

Answer 26

No, all 4 outer shell electrons are involved in covalent bonding there none are available for conducting electricity

Question 27

Why can graphite conduct electricity?

Answer 27

Only 3 outer electrons are used out of the 4 so the remaining one is released into a pool of delocalised electrons

Question 28

Why does graphite have a lower boiling point than diamond?

Answer 28

The layers in graphite can slide over each other because the forces between them are weak.

Question 29

What does it mean if there is an increase in ionisation energy between 3rd and 4th ionisation energies?

Answer 29

This shows that the fourth electron is being removed from an inner shell, therefore there are 3 electrons in the outer shell and the element must be in group 3 (Al)

Question 30

Explain why xenon has a lower first ionisation energy than neon (Xe is lower down than Ne)

Answer 30

• Xenon has more shells
• Xenon has more shielding
• Therefore nuclear attraction decreases

Question 31

Why does the modern Periodic Table not arrange some elements, such as those in order of increasing atomic mass?

Answer 31

• Arranged in increasing atomic number
• Wouldn’t show properties/trends of the rest of the group

Question 32

Explain the difference in melting point for the elements Na and Mg

Answer 32

• Mg has a greater charge
• Mg has more electrons
• Mg has stronger metallic bonds

Question 33

Sulfur exists as S8 molecules and chlorine as Cl2 molecules. Use this information to explain the difference in their melting points.

Answer 33

• S8 has more electrons than Cl2
• Therefore S8 has stronger intermolecular forces

Question 34

Explain the decrease in the atomic radii across the period from Na to Cl

Answer 34

• Proton charge increases therefore greater pull on electrons
• Shielding remains the same

Question 35

Explain the difference between the first ionisation energies of Li and Na.

Answer 35

• Na has more shielding
• Atomic radius increases
• Less nuclear attraction

Question 36

Predict and explain whether a barium ion is larger, smaller or the same size as a barium atom.

Answer 36

• Smaller due to less shielding

Question 37

Why does ionisation energy decrease from Be to B?

Answer 37

• An electron is being removed from a different sub-shell (p), which has a higher energy level

Question 38

Why does ionisation energy decrease from Be to B?

Answer 38

• An electron is being removed from a different sub-shell (p), which has a higher energy level

Question 39

Explain why a nitrogen atom is larger than an oxygen atom

Answer 39

• N has less protons than O
• Same shielding
• Weaker nuclear attraction in N
• Shell drawn in less for N

Question 40

Explain why a nitrogen atom is larger than an oxygen atom

Answer 40

• N has less protons than O
• Same shielding
• Weaker nuclear attraction in N
• Shell drawn in less for N

Question 41

Suggest why the second ionisation energy of oxygen has a greater value than the first
ionisation energy of oxygen

Answer 41

• O+ is smaller than O
• Shielding is smaller