Unit 1 Periodicity

Question 1

What is a periodic pattern?

Answer 1

A pattern which repeats across a period in the periodic table

Question 2

What did Mendeleev use to arrange the then known elements?

Answer 2

Increasing masses and similar chemical properties

Question 3

How is the modern periodic table arranged?

Answer 3

Increasing atomic number

Question 4

In the modern periodic table, what are trends related to?

Answer 4

Bonding and structure

Question 5

What is meant by the covalent radius of an atom?

Answer 5

The covalent radius is a measure of the size of an atom. It is half the bond length.

Question 6

Explain the trend in covalent radius across a period

Answer 6

Across a period from left to right, the covalent radius decreases.
This is due to
Increasing nuclear charge or
Increasing number of protons

Question 7

Explain the trends in covalent radius down a group.

Answer 7

As you move down a group in the periodic table, the covalent radius increases. Atoms increase in size.

The screening effect means the outer electrons are further away from the nucleus and so are not as attracted to the positive charge.

Question 8

What is the density of a substance?

Answer 8

It’s mass per unit volume

Question 9

What happens to density as you move across a period?

Answer 9

The general trend is that in any period, density first increases from group 1 to a maximum in the centre, then decreases again towards group 0

Question 10

What happens to density as you move down a group?

Answer 10

Density increases as the mass of the atom is increasing faster than the volume

Question 11

What do melting and boiling points indicate?

Answer 11

The strength of the forces between particles

Question 12

What happens to mps and bps as you move across a period?

Answer 12

Increase from group 1 to group 4 (group 4 elements are covalent networks), then decrease to group 0

Question 13

What happens to mps and bps as you move down group 1?

Answer 13

Melting and boiling points decrease as there is a decrease in attraction between atoms

Question 14

What happens to mps and bps as you move down group 7?

Answer 14

Melting and boiling points increase as LDF attractions between molecules increase

Question 15

What is meant by the ionisation energy of an atom?

Answer 15

The ionisation energy is the energy involved in removing one mole of gaseous electrons from one mole of atoms in the gaseous state.

Question 16

Magnesium has a 1st ionisation energy of 744 kJmol-1

2nd ionisation energy of 1460 kJmol-1

3rd ionisation energy of 7750 kJmol-1

Explain this trend.

Answer 16

The third ionisation energy shows a massive increase in ionisation energy. This is because the 3rd electron has to be removed from magnesium’s second energy level. This is closer to the nucleus and also a full and complete energy level.

Question 17

Explain the trends in ionisation energy across a period.

Answer 17

Across a period from left to right, the ionisation energy increases. This is due to increased nuclear charge or increasing number of protons.

Question 18

Which group of the periodic table will element X be found in? (Please note that the first element in the graph is not speciified)

Answer 18

Group 4 (largest peaks will be group 0)

Question 19

What is meant by an atom’s electronegativity?

Answer 19

Electronegativity is a measure of an atom’s attraction for the bonding electrons in a covalent bond.

Question 20

Explain the trend in electronegativity across a period.

Answer 20

Across a period from left to right the electronegativity of atoms increases.

As you move from left to right across the periodic table, atoms have a greater charge in their nucleus and a smaller covalent radius. This allows the nucleus to attract the bonding electrons more strongly.

Question 21

Explain the trend in electronegativity down a group.

Answer 21

Going down a group, the electronegativity of atoms decreases.

As you move down a group, atoms increase in size, with more energy levels.

The extra energy levels inner electrons and increased covalent radius keep the bonding electrons further away from the nucleus.

This screening effect on the attractive force of the nuclues means less attraction for the bonding electrons so lower electronegativity.

Question 22

The first 20 elements in the Periodic Table are categorised according to which four bonding and structures (hydrogen to calcium)

Answer 22

metallic (Li, Be, Na, Mg, Al, K, Ca)

covalent molecular (H₂, N₂, O₂, F₂, Cl₂, P₄, S₈ and fullerenes (eg C₆₀))

covalent network (B, C (diamond, graphite), Si)

monatomic (noble gases)

Question 23

What are the properties and structure of elements related to?

Answer 23

The types of bonding present

Question 24

Describe metallic bonding and structure in metallic elements.

Answer 24

Metallic bonding occurs between the atoms of metal elements. The outer electrons are delocalised (free to move).

This produces an electrostatic force of attraction between the positive metal ions and the negative delocalised sea of electrons.

It has no intermolecular forces present.

Question 25

Describe bonding and structure in covalent network elements.

Answer 25

Covalent networks are large, rigid three-dimensional arrangements of atoms held together by strong covalent bonds.

It has no intermolecular forces present.

Question 26

Describe bonding and structure in discrete covalent molecular elements.

Answer 26

Discrete covalent molecules are small groups of fixed number of atoms held together by strong covalent bonds inside the molecule and weak intermolecular forces between the molecules.

Examples (H₂, N₂, O₂, F₂, Cl₂, P₄, S₈ and fullerenes (eg C₆₀))

Sulfur, phosphorous and fullerencs are covalent molecular solids.

Question 27

Describe bonding and structure in monatomic elements.

Answer 27

Group 0 atoms are single, unattached particles. They are stable atoms. Weak London dispersion forces.

Question 28

Why is ionic bonding not present in elements?

Answer 28

It requires at least two elements (a metal transfers electrons non-metal) to form compound with ions.

Question 29

What is metallic bond strength?

Answer 29

A measure of the no. of outer electrons. The greater the no. of electrons, the stronger the metallic bond

Question 30

What are metal boiling points dependant on?

Answer 30

• How many electrons are in the outer shell
• No. of electron shells

Question 31

Why are the noble gases stable?

Answer 31

They have full outer electron shells, meaning they are chemically stable and monatomic

Question 32

Why do noble gases form liquids and solids at low temperatures, despite being monatomic?

Answer 32

Weak intermolecular forces - LDF

Question 33

What is LDF?

Answer 33

The attraction between the temporary dipoles caused by uneven distribution of moving electrons

Question 34

What is a dipole?

Answer 34

An unequal distribution of charge

Question 35

Give examples of discrete covalent molecular solids

Answer 35

Phosphorous (P₄), sulphur (S₈), and carbon (fullerenes C₆₀)

Question 36

Why are phosphorous and sulphur solids at room temperature, whereas chlorine is a gas?

Answer 36

P and S are larger with more atoms. Phosphorous has 4 atoms so forms P₄ molecule and Sulphur has 8 atoms so forms S_{<span>8</span>} molecule.

So they both have more atoms so more electrons so more LDF between their molecules than Chlorine. This requires more energy to break so hence why they have higher melting and boiling points than Chlorine.

Chlorine forms diatomic molecules so has two atoms Cl₂ in a molecule which are gaseous at room temperature as they have less atoms so less electrons so less LDF between their molecules.

Question 37

What two network forms does carbon exist as?

Answer 37

Graphite and Diamond

Question 38

What are some properties of graphite?

Answer 38

• Layers of atoms in hexagonal plates
• 3/4 outer electrons covalently bonded, 4th outer electron moves within layer
• Conducts electricity due to delocalised electron
• Soft, slippery, used as a lubricant

Question 39

What are some properties of diamond?

Answer 39

• Regular tetrahedral structure
• All 4 outer electrons covalently bonded
• Does not conduct electricity
• Hardest natural substance, used in cutting tools