Chapter 7 - Periodicity

Question 1

What is periodicity?

Answer 1

The term used to describe the repeating patterns seen within groups and periods on the periodic table

Question 2

How did John Dalton arrange the elements in his periodic table?

Answer 2

In order of atomic weight, used pictorial symbols to represent both elements and compounds

Question 3

How did John Newlands arrange his periodic table?

Answer 3

He arranged known elements in order of mass, grouping every 8 elements together in what he called the ‘Law of Octaves’.

Question 4

What were the flaws in Newland’s table?

Answer 4

He left no gaps (assumed all elements had been found), elements in the same group were therefore not similar

Question 5

How did Dmitri Mendeleev arrange his periodic table?

Answer 5

He put the elements in order of atomic mass, and then arranged them so periodic pattern in properties could be seen, left gaps for undiscovered elements using table to predict their properties

Question 6

Which group was missing from Mendeleev’s periodic table and why did he omit it?

Answer 6

Group 8/0

It contains very unreactive (inert) noble gases, none of which had been discovered at the time

Question 7

What is the trend across a period for electron configuration?

Answer 7

Each period starts with an electron in a new highest energy shell
For each period, the s and p sub-shells are filled in the same way

Question 8

What is the trend down a group for electron configuration?

Answer 8

Elements in the same group have atoms with the same number of electrons in their outer shell, and the same number of electrons in each sub shell

Question 9

What is ionisation energy?

Answer 9

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

Question 10

What is first ionisation energy?

Answer 10

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

Question 11

What is the general equation for first ionisation energy?

Answer 11

Element (g) → Element+(g) + e-

Question 12

What is atomic radius?

Answer 12

The distance between the nucleus and the outer electrons

Question 13

What is nuclear charge?

Answer 13

The number of protons there are in the nucleus of an atom

Question 14

What is electron shielding?

Answer 14

The repulsion caused by inner shell electrons that reduces the attraction between the nucleus and outer shell electrons

Question 15

What is the trend in first ionisation energy down a group (and why)?

Answer 15

First ionisation energy decreases down a group
This is because the atomic radius increases, and there are more inner shells so shielding also increases. Therefore the nuclear attraction on outer electrons decreases and so does the 1st IE

Question 16

What is the trend in first ionisation energy across a period (and why)?

Answer 16

It increases across a period
Nuclear charge increases across a period, and there is similar shielding as elements have the same shell, nuclear attraction therefore increases and atomic radius decreases - 1st IE increases

Question 17

Why does 1st IE fall in two places across period 2 and 3?

Answer 17

Paired electrons are easier to remove (repulsion) so it is easier to remove electron from O than N
2p subshell has higher energy than 2s subshell, so 1 electron in 2p is easier to remove, e.g. easier for B than Be

Question 18

What is metallic bonding?

Answer 18

The strong electrostatic attraction between cations (positive ions) and delocalised electrons

Question 19

What is the structure in metals?

Answer 19

Giant metallic lattice

Question 20

What are the properties of metals?

Answer 20

• strong metallic bonds
• high electrical conductivity
• high melting and boiling points

Question 21

Why can metals conduct electricity?

Answer 21

Due to the delocalised electrons which can move through the structure and carry charge

Question 22

Why do metals have high melting and boiling points?

Answer 22

High temperatures are necessary to provide the large amount of energy to overcome the strong electrostatic attraction between cations and electrons

Question 23

What is the solubility of metals?

Answer 23

Metals do not dissolve

Question 24

What are giant covalent structures?

Answer 24

When many billions of atoms are held together by a network of strong covalent bonds to form a giant covalent lattice

Question 25

What are the typical properties of substances with a giant covalent lattice structure?

Answer 25

• high melting and boiling point
• insoluble in almost all solvents
• non conductors of electricity (graphene and graphite are exceptions)

Question 26

What is the structure of diamond?

Answer 26

A tetrahedral arrangement of carbon atoms (bond angles 109.5 by repulsion)

Question 27

Why do giant covalent lattices have high melting and boiling points?

Answer 27

High temperatures necessary to break strong covalent bonds

Question 28

Why are giant covalent lattices insoluble in almost all solvents?

Answer 28

Covalent bonds holding atoms in lattice are far too strong to be broken by interaction with solvents

Question 29

What is the electrical conductivity of diamond, graphite and graphene (and why)?

Answer 29

Diamond - does not conduct as all 4 outer shell electrons are involved in covalent bonding
Graphene and Graphite - conducts as only 3 of outer shell electrons are used in covalent bonding, the remaining is delocalised (can carry charge)

Question 30

What is the structure of graphene?

Answer 30

It is a single layer of graphite, composed of hexagonally arranged carbon atoms linked by strong covalent bonds - thinnest and strongest material ever made

Question 31

What is the structure of graphite?

Answer 31

Composed of parallel layers of hexagonally arranged carbon atoms, layers bonded by weak London forces. Bonding in layers only uses 3 of carbon’s 4 outer shell electrons, so the spare electron is delocalised between layers

Question 32

What is the periodic trend in melting point for Periods 2 and 3?

Answer 32

Melting point increases across from Group 1 to Group 4 (giant structures), then there is a sharp decrease in melting point between Group 4 and 5 as they are simple molecules. Giant structures require more energy to overcome their strong forces, whereas simple molecular structures only have weak forces to overcome