Chapter 7 - Periodicity

Question 1

What order are the elements arranged in?

Answer 1

Increasing atomic number

Question 2

What do the groups in the periodic table tell you?

Answer 2

• Number of electrons in the outer shell
• Elements that have similar chemical properties

Question 3

What do the periods in the periodic table tell you?

Answer 3

Number of the highest energy electron shell

Question 4

What is periodicity?

Answer 4

Repeating trend in properties of the elements across a period

Question 5

Which groups are the s-block elements?

Answer 5

• Group 1 and 2
• Helium

Question 6

Which groups are the d-block elements?

Answer 6

Transition metals

Question 7

Which groups are the p-block elements?

Answer 7

Groups 3, 4, 5, 6, 7, 8

Question 8

What is the first ionisation energy?

Answer 8

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

Question 9

What is the first ionisation energy of Na_(g) ?

Answer 9

Na_(g) -> Na⁺ _(g) + e^-

Question 10

First electron lost will be the _________ energy level and will experience the _________ attraction from the nucleus

Answer 10

• Highest
• Least

Question 11

What happens to first ionisation energy when atomic radius increases? Explain why

Answer 11

Greater distance between nucleus and outer electrons
Less nuclear attraction
Lower ionisation energy

Question 12

What happens to first ionisation energy when nuclear charge increases? Explain why

Answer 12

More protons
Greater attraction between nucleus and outer electron
Higher ionisation energy

Question 13

What happens to first ionisation energy when there is more electron shielding Explain why

Answer 13

Shielding effect increases when number of inner shells increase
Less attraction between nucleus and outer electrons
Lower ionisation energy

Question 14

Give the general relationship between nuclear attraction and first ionisation energy?

Answer 14

Less attraction between nucleus and outer electrons
Lower ionisation energy

Greater attraction between nucleus and outer electrons
Higher ionisation energy

Question 15

What is the shielding effect?

Answer 15

Inner shell electrons repel outer shell electrons

Question 16

Element has as many ionisation energies as there are ___________

Answer 16

Electrons

Question 17

What is second ionisation energy?

Answer 17

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

Question 18

Why are second ionisation energies greater than the first?

Answer 18

After first electron is lost, remaining electron is pulled closer to nucleus, stronger nuclear attraction = more energy needed to remove electron = higher ionisation energy

Question 19

Electrons with _________ ionisation energies are ___________ to the nucleus

Answer 19

Higher
Closer

Question 20

What does the large difference/jump in ionisation energies represent?

Answer 20

Change from one shell to another

Question 21

What is the trend in first ionisation energies down a group?

Answer 21

• Atomic radius increases
• More inner shells so shielding increases
• Nuclear attraction decreases
• First ionisation energy decreases

s.n. nuclear charge increasing is outweighed by the effects of increased radius and (to a lesser extent) increased shielding

Question 22

What is the trend in first ionisation energies across a period?

Answer 22

• Nuclear charge increases
• Same shell numbers so similar shielding
• Nuclear attraction increases
• Atomic radius decreases
• First ionisation energy increases

Question 23

Why does boron have a lower first ionisation energy than beryllium? (hint: sub-shell is the different)

Answer 23

• Start of filling of 2p sub shell
• 2p sub shell is higher energy than 2s sub shell
• 2p electron easier to remove

Question 24

Why does oxygen have a lower first ionisation energy than nitrogen? (hint: sub-shell is the same)

Answer 24

• Start of electron pairing in the p-orbitals of 2p sub shell
• In O the paired electrons in one of the 2p orbitals repel each other
• Easier to remove one of the paired electrons from O

Question 25

Why does **aluminium** have a lower first ionisation energy than **magnesium**? (hint: sub-shell is the different)

Answer 25

- 3p sub-shell in Al has a higher energy level than 3s sub-shell in Mg - Easier to remove 3p electron from Al

Question 26

Why does **sulphur** have a lower first ionisation energy than **phosphorus**? (hint: sub-shell is the same)

Answer 26

- P has 3 electrons in 3p sub shell which are all unpaired - S has 4 electron in 3p sub shell which has one pair in an orbital - Paired electrons repel so easier to remove an electron from S

Question 27

Describe metallic bonding?

Answer 27

Strong electrostatic attraction between cations and delocalised electrons

Question 28

Describe the electrical conductivity of metallic structures?

Answer 28

Conduct electricity in solid and liquid states since the delocalised electrons are able to carry charge

Question 29

Describe the melting and boiling points of metals?

Answer 29

High M.P. and B.P. as lots of energy needed to break strong electrostatic attraction

Question 30

Describe the solubility of metals?

Answer 30

Do not dissolve in water

Question 31

Describe giant covalent structures?

Answer 31

Boron, carbon and silicon atoms held together by covalent bonds to form giant covalent structures

Question 32

What is the shape of the structures that carbon and silicon atoms form?

Answer 32

Tetrahedral

Question 33

Describe the electrical conductivity of giant covalent structures?

Answer 33

Non conductors of electricity (except graphene and graphite)

Question 34

Why can diamond not conduct electricity?

Answer 34

All four outer shell electrons involved in bonding so no delocalised electrons to carry charge

Question 35

Why can graphene (single layer of graphite) conduct electricity?

Answer 35

Only three of the carbon outer shell electrons are involved in bonding so there are delocalised electrons to carry charge

Question 36

Describe the melting and boiling points of giant covalent structures?

Answer 36

High M.P. and B.P. as lots of energy needed to break strong covalent bonds

Question 37

Describe the solubility of giant covalent structures?

Answer 37

Insoluble in almost all solvents as covalent bonds are too strong to be broken

Question 38

What is the general trend in melting points across period 2 and 3?

Answer 38

- Melting point increases from group 1-4 - Sharp decrease in melting point from group 4-5 - Melting point lower from group 5-0

Question 39

Why is there a sharp decrease in melting point from group 4-5?

Answer 39

Marks change from giant to simple molecular structures

Question 40

What structure to metals form?

Answer 40

Giant metallic lattice structures