C7- Periodicity

Question 1

How did Mendeleev arrange the periodic table?

Answer 1

• in order of atomic mass.
• lined up the elements in groups with similar properties.
• if properties didn’t fit, he swapped elements around, assuming atomic mass measurements were incorrect or new elements were to be discovered.
• predicted properties of missing elements from group trends.

Question 2

How is the periodic table arranged now?

Answer 2

• 114 elements arranged in 7 periods and 18 groups.
• in order of increasing atomic number, from left to right.
• each element in a group has the same number of outer shell electrons and similar properties.
• the number of a period indicates number of the highest energy electron shell.

Question 3

What is periodicity?

Answer 3

The repeating trend in properties of the elements across each period. the most obvious periodicity is the trend from metals to non-metals.
Properties:
- electron configuration.
- ionisation energy.
- structure.
- melting points.

Question 4

What are the blocks of the periodic table?

Answer 4

s-block = group 1+2
d-block= group 3-12 (transition metals)
p-block= group 13-18
f-block= bottom box of elements.

Question 5

What is the trend in electron configuration across a period?

Answer 5

• across period 2: 2s fills then 2p.
• across period 3: same pattern of filling repeated for 3s and 3p sub-shells.
• across period 4: From n=4, only 4s and 4p sub-shells are occupied.
  For each period the s- and p- sub-shells are filled in the same way= a period pattern.

Question 6

What is the trend of electron configuration down a group?

Answer 6

Elements in each group have atoms with the same number of electrons in each sub-shell.
-this similarity in electron configuration gives elements in the same group their similar chemistry.

Question 7

What is ionisation energy?

Answer 7

Measures how easily an atom loses electrons to form positive ions.
Electrons with the largest ionisation energies are from the shell closest to the nucleus.

Question 8

What is first ionisation energy?

Answer 8

The energy required to remove one electron from each atom in one mole of gaseous atoms in an element to form one mole of gaseous 1+ ions.
-Electrons lost will be in the highest energy level and will experience the least attraction from the nucleus.

Question 9

What factors affect ionisation energy and how?

Answer 9

1. Atomic radius- greater the distance between nucleus and outer electrons, less the nuclear attraction. Force of attraction falls with increasing distance.
2. Nuclear charge- the more protons in nucleus, the greater the attraction with outer electrons.
3. Electron shielding- inner shell electrons repel outer shell ones due to negative charge. Called the shielding effect. Reduces attraction between nucleus and outer shell electrons.

Question 10

What are successive ionisation energies?

Answer 10

An element has many ionisation energies.

• after the first electron is lost, the nuclear attraction on remaining electrons increases as they are pulled closer to the nucleus. Therefore more ionisation energy is needed to remove another electron.
• successive ionisation energies are the same as the first ionisation energy, however, the number of ionisation energy is the same as the charge on the ion produced.

Question 11

What can you predict using successive ionisation energies?

Answer 11

-number of electrons in outer shell.
for example: if there is big difference between 3rd and 4th ionisation energies. Shows fourth electron is being removed from inner shell, therefore 3 electrons on outer shell.
- group of the element on periodic table.
-identity of an element.

Question 12

Outline the general trend in first ionisation energies across periods?

Answer 12

• a general increase in first ionisation energy across each period.
• sharp decrease in first ionisation energy from end of one period and the start of the next.

Question 13

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

Answer 13

First ionisation energies decrease down a group.
- although nuclear charge increases, its effect is outweighed by the increased radius and the increased shielding (more inner shells).

Question 14

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

Answer 14

General increase in first ionisation energy across the first 3 periods as:

• nuclear charge increases.
• same shell so similar shielding.
• nuclear attraction increases.
• atomic radius decreases.

Question 15

What are the anomalies in the first ionisation energy across period 2+3?

Answer 15

The general increase falls in two places in each period.

• These drops occur at the same positions of both periods suggesting there might be s periodic cause. (In period 2: Be-B and N-O)
• also linked to existence of sub-shells, their energies and the filling of orbitals.

Question 16

Explain the fall in ionisation energy from Be to B.

Answer 16

• the drop marks the start of the filling of the 2p sub-shell.
• 2p subshell in B has a higher energy than the 2s sub shell in Be.
• therefore the 2p electron in B is easier to remove then the 2s electron is Be and as a result B has a lower first ionisation energy.

Question 17

Explain the fall in ionisation energy from N and O.

Answer 17

• the drop marks the start of electron pairing in the p orbitals of the 2s subshell.
  -in N and O the highest energy electrons are in a 2p subshell.
• in O, the paired electrons in one of the 2p orbitals repel one another, making it easier to remove an electron.
  Therefore oxygen has a lower first ionisation energy than nitrogen.

Question 18

Explain metallic bonding?

Answer 18

Metallic bonding is the strong electrostatic attraction between cations and delocalised electrons. Occurs in metals.

• in a solid metal structure, each atom has donated its - outer shell electron to a shared pool of delocalised electrons which are free to move.
• the cations are in a fixed position, maintaining the structure and shape.
• billions of metal atoms are held together by metallic bonding in a giant metallic lattice.

Question 19

Explain the electrical conductivity in metals.

Answer 19

Metals conduct electricity in solid and liquid state.

- the delocalised electrons are mobile so can move through the structure carrying a charge.

Question 20

Explain the melting and boiling points of metals.

Answer 20

Most metals have high melting and boiling points.
- the mp depends on the strength of the metallic bonds holding the atom in the giant metallic lattice:
For most, high temperatures are necessary to provide the large amount of energy needed to overcome the strong electrostatic attraction.
(Mercury= exception mp: -39)

Question 21

Explain the solubility of metals.

Answer 21

• they do not dissolve.

- any interactions with polar solvents like water leads to a reaction rather than dissolving.

Question 22

Explain giant covalent lattices.

Answer 22

They consist if billions of atoms held together by a network of strong covalent bonds. The non-metals boron, carbon and silicon have this lattice structure.

• C and Si are in group 4 so can use these four electrons in the outer shells to form covalent bonds to other C/Si atoms.
• this results in a tetrahedral structure. Bond angles= 109.5*

Question 23

Explain the MP/BP and solubility of giant covalent lattices.

Answer 23

1. Have high melting and boiling points due to the strong covalent bonds which require a lot of energy to overcome.
2. Insoluble in almost all solvents- the covalent bonds holding atoms in the lattice are way too strong to be broken by interaction with solvents.

Question 24

Explain the electrical conductivity of giant covalent lattices.

Answer 24

• they are non-conductors.
• exceptions are graphene and graphite which are forms of carbons.
• In C and Si the outer shell is full so there are no delocalised electrons.
• however carbon is able to form structures (graphene/graphite) where one of the electrons is available for conductivity.

Question 25

Explain the periodic trend in melting points and why?

Answer 25

Across period 2 and 3:

• melting point increases from group 1 to 14. These elements have giant metallic structures.
• sharp decrease from group 14 to 15. They have giant covalent structures.
• melting points remain relatively low from group 15 to 18. Elements have simple molecular structures with weak london forces.