Chapter 7

Question 1

How are the elements arranged in the periodic table?

Answer 1

• According to increasing atomic number
• In groups (columns) and periods (rows)

Question 2

What do groups in the periodic table show? How does this relate to the elements?

Answer 2

• Groups in the periodic table have the same number of electrons in their outer shell, and similar chemical properties

Question 3

How are groups in the periodic table numbered now?

Answer 3

• Groups 1 and 2 are the same
• The 10 columns of transition metals were named groups 3 to 12 by IUPAC
• Groups 3 to 0 are also therefore known as groups 13 to 18

Question 4

What is periodicity? Required.

Answer 4

• Repeating trends in the physical and chemical properties of elements across a period

Question 5

What are 4 properties that have periodicity?

Answer 5

• Electronic configuration
• Ionisation energy
• Structure
• Melting points

Question 6

What trend do periods have in terms of electronic configuration?

Answer 6

• Each period starts with an electron in the new highest energy level

Question 7

What is first ionisation energy? Required.

Answer 7

• The energy required for the removal of 1 mol of electrons from 1 mol of gaseous atoms of an element

Question 8

Write an equation for the first ionisation energy of chlorine.

Answer 8

• Cl (g) -> Cl+(g) +e- (it is not shown as diatomic as first ionisation energy refers to 1 mol of atoms)

Question 9

What is second ionisation energy?

Answer 9

• The amount of energy needed to turn 1 mol of gaseous 1+ ions into 1 mol of gaseous 2+ ions

Question 10

What 3 factors affect ionisation energy?

Answer 10

• Atomic radius
• Nuclear charge
• Electron shielding

Question 11

How does atomic radius affect ionisation energy?

Answer 11

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

Question 12

How does nuclear charge affect ionisation energy?

Answer 12

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

Question 13

How does electron shielding affect ionisation energy?

Answer 13

• Electrons are all negatively-charged and therefore repel each other, meaning that inner-shell electrons repel outer-shell electrons
• This reduces the attraction between the nucleus and the outer electrons

Question 14

How do successive ionisation energies compare with each other?

Answer 14

• They increase
• There is a big increase in the ionisation energy between different shells of an atom

Question 15

Why do successive ionisation energies increase?

Answer 15

• The atomic radius decreases
• The nuclear attraction between the remaining electrons and nucleus increases (the same number of protons attract fewer electrons)

Question 16

Describe the trends in first ionisation energy in the periodic table.

Answer 16

• They decrease down a group (down across the different periods)
• They generally increase across a period

Question 17

Why does first ionisation energy decrease down a group?

Answer 17

• The atomic radius and shielding increases
• Nuclear attraction on outer electrons decreases

Question 18

Why does first ionisation energy generally increase across a period?

Answer 18

• The nuclear charge increases, which causes the atomic radius to decrease
• There is similar shielding
• The nuclear attraction therefore increases

Question 19

In what 2 places does the first ionisation energy not increase across a period?

Answer 19

• Between groups 2 and 3
• Between groups 5 and 6

Question 20

Why does the first ionisation energy decrease between groups 2 and 3?

Answer 20

• In group 3, the p sub-shell starts to be filled
• As it has more energy than the s sub-shell, it is easier to remove the electron in the p sub-shell
• Group 3 therefore has a lower first ionisation energy than group 2

Question 21

Why does the first ionisation energy decrease between groups 5 and 6?

Answer 21

• In group 6, the p sub-shell electrons start to pair up
• The paired electrons in one orbital repeal each other, making it easier for one of them to be removed
• Therefore, group 6 has a lower first ionisation energy than group 5

Question 22

What does the graph of first ionisation energy against atomic number look like? Give 5 details.

Answer 22

• The highest points are the noble gases
• The lowest points are the group 1 metals
• First ionisation energy generally increases across each period
• There are dips between groups 2 and 3 and 5 and 6
• Each period has a lower first ionisation energy than the previous period (the entire line is shifted down)

Question 23

What is metallic bonding? Required.

Answer 23

• The strong electrostatic attraction between cations and delocalised electrons

Question 24

What structure do metals have?

Answer 24

• They all have a giant metallic lattice structure

Question 25

What are the 2 main properties of metals?

Answer 25

• High electrical conductivity
• High melting and boiling points

Question 26

When can metals conduct electricity, and how?

Answer 26

• When in the solid or liquid states
• This is because the delocalised electrons are free to move through the structure and carry charge

Question 27

Why do most metals have high melting and boiling points?

Answer 27

• Large amounts of energy are needed to overcome the strong electrostatic attraction between the cations and electrons

Question 28

Are metals soluble? Why?

Answer 28

• No
• They react with solvents instead, such as group 1 metals with water

Question 29

What are giant covalent lattices? Required.

Answer 29

• Networks of atoms bonded by strong covalent bonds

Question 30

List 2 elements that form giant covalent lattices.

Answer 30

• Carbon
• Silicon

Question 31

What are the melting and boiling points of giant covalent lattices like, and why?

Answer 31

• They are high because a lot of energy is needed to break the strong covalent bonds

Question 32

Are giant covalent lattices soluble? Why?

Answer 32

• No
• The covalent bonds are too strong to be broken due to interactions with the solvent

Question 33

Can giant covalent lattices conduct electricity? Why?

Answer 33

• Most aren’t as they don’t have any electrons that are free to conduct electricity
• Graphite and graphene are exceptions to this as they only use 3 of their 4 electrons in their outer shell to form covalent bonds

Question 34

What is the difference between graphite and graphene?

Answer 34

• Graphene is a single layer of graphite (in graphite the separate layers are bonded together by London forces)

Question 35

How do melting points vary across periods 2/3?

Answer 35

• The melting points increase from periods 1 to 4
• There is a sharp decrease between groups 4 and 5
• From group 5 to 0, the melting points are low

Question 36

Why do melting points decrease between groups 4 and 5?

Answer 36

• The elements change from giant molecular structures to simple molecular substances

Question 37

Why do melting points increase between groups 1 and 3?

Answer 37

• The larger the charge of the cation (and therefore also the relative amount of delocalised electrons) the stronger the metallic bonds