Periodicity

Question 1

How does the periodic table arrange elements?

Answer 1

In order of increasing proton number

Question 2

Elements in the same group have similar what?

Answer 2

Chemical properties

Question 3

Define first ionisation energy

Answer 3

Minimum energy required to remove 1 mole of electrons from 1 mole of gaseous atoms.
It is measured in KJmol^-1

Question 4

Why does atomic radius decrease along a period?

Answer 4

This is due to an increased nuclear charge for the same number of electron shells. The outer electrons are pulled in closer to the nucleus as the increased charge produces a greater electrostatic attraction. As a result the atomic radius for that element is reduced.

Question 5

Why does atomic radius increase down a group?

Answer 5

With each element down a group an electron shells is added. This increases the distance between the outer electrons and the nucleus, reducing the electrostatic force of attraction. More shells also increases the electron shielding. Therefore, the nuclear attraction is reduced further and atomic radius decreases.

Question 6

Why are successive ionisation energies larger than the first ionisation energy?

Answer 6

As electrons are removed, the electrostatic force of attraction between the positive nucleus and the negative outer electron increases. More energy is therefore needed to overcome come this attraction, so ionisation energy increases.

Question 7

What’s the trend in first ionisation energy along a period?

Answer 7

First ionisation energy increases due to a decreasing atomic radius, meaning there is a greater electrostatic force of attraction.

Question 8

What’s the trend in first ionisation energy down a group?

Answer 8

First ionisation energy decreases due to an increasing atomic radius and electron shielding which reduces the effect of the electrostatic force of attraction.

Question 9

When there is a large jump in successive ionisation energies, what does that indicate?

Answer 9

Indicates how many electrons are in the outer shell.

Question 10

Why is there a drop in ionisation energy in group 3, when going along a period?
(Trend is that first ionisation energy follows a general increasing trend)

Answer 10

Outer electron in group 3 elects is in a p orbital rather than an s orbital.
as it is in a higher energy level, which is further away from the nucleus so the electron is held less strongly. There is more shielding provided by the s electrons.

Question 11

Why is there a drop in ionisation energy in group 6 when going along a period?
(Trend is that first ionisation energy follows a general increasing trend along a period)

Answer 11

In group 5 the electron is being removed from a singly occupied orbital.
In group 6 the electron is being removed from an orbital containing 2 electrons.
The repulsion between the 2 electrons in an orbital means that it is easer to remove from shared orbitals.

Question 12

Define metallic bonding

Answer 12

Strong electrostatic attraction between cations and delocalised electrons.

Question 13

What are metallic properties?

Answer 13

Good conductors - sea of delocalised electrons is able to move and carry a flow of charge
Malleable - as uniform layers of positive ions are ale to slide over one another.
The electrostatic forces of attraction between positive ions and delocalised electrons are very strong and therefore require a lot of energy to overcome. meaning metallic structure s have high melting points and are nearly always solid at room temperature.

Question 14

What structure do all metals have?

Answer 14

A giant metallic lattice structure

Question 15

What is a giant covalent lattice?

Answer 15

A network of atoms bonded by many strong covalent bonds. Substances that have a macromoleular structure are covalently bonded into a giant lattice structure. Each atom has multiple covalent bonds which are very strong. Giving the substance a very high melting point.

Question 16

Explain the solid covalent lattice structure of diamond.

Answer 16

Daimond - Macromolecular structure made up of carbon atoms, each of which are boned to four further carbon atoms. Provides rigid structure

Question 17

Explain the solid covalent lattice structure of graphite.

Answer 17

Graphite - macro molecular structure made up of carbon atoms.
Each carbon atoms bonded to three other carbon atoms in flat sheets. Meaning delocalised electrons per carbon atom which can move between layers, meaning graphite can conduct electricity.
There are strong covalent bonds between carbon atoms but weak intermolecular forces between layers.
This means the layers can slide over each other - good lubricant

Question 18

Explain the solid covalent structures or graphene

Answer 18

Consists of 2D sheets of graphite that are just one atom thick. These sheets are formed of hexagonal carbon rings that create a very strong rigid material that is extremely light weight. Delocalised electrons move through each layer allowing it to conduct electricity.

Question 19

There is a variation in melting points across period 2. Lithium and beryllium have what kind of bonding and explain their melting points.

Answer 19

metallic bonding
Melting points increase due to a greater positive charge of the ions. (Li = +1, Be = +2).

Question 20

There is a variation in melting points across period 2. Boron and carbon have what kind of bonding and explain their melting points.

Answer 20

Giant covalent lattices
very strong covalent bonds
Covalent bonds require a lot of energy to break, giving them very high melting points.

Question 21

There is a variation in melting points across period 2. Nitrogen, oxygen, fluorine and neon have what kind of bonding and explain their melting points.

Answer 21

Simple covalent molecules.
Weak van der waals intermolecular forces.
These forces don’t require that much energy to overcome so these molecules have relatively low, similar melting points.

Question 22

There is a variation in melting points across period 3. Sodium, magnesium and aluminium have what kind of bonding and explain their melting points.

Answer 22

Metallic bonding
Their melting points increase due to greater positive charged ions. (Na = +1, Mg = +2, Al = +3).
This means that more electrons are released as free electrons, so the attractive electrostatic forces increase from Na to Al.

Question 23

There is a variation in melting points across period 3. Silicon has what kind of bonding and explain their melting points.

Answer 23

Very strong covalent structure (macromolecular)
These covalent bonds require a lot of energy to break, giving it a very high melting point.

Question 24

There is a variation in melting points across period 3. Phosphorus, sulfur and chlorine have what kind of bonding and explain their melting points.

Answer 24

simple covalent molecules
Weak van der waals forces
These intermolecular forces don’t require much energy to overcome so these molecules have a relatively low, similar melting points.

Question 25

There is a variation in melting points across period 3. Argon has what kind of bonding and explain their melting points.

Answer 25

Noble gas that exists as individual atoms with a full outer shell of electrons. This makes the atom very stable and the van der waals forces between them are very weak. As a result, the melting points increase of argon is very low and it exists as a gas at room temperature.

Question 26

Does atomic radius in group 2 deacrease or increase down the group? Explain why

Answer 26

Increases due to additional electron shells.

Question 27

Why does the reactivity of group 2 metals increase down the group?

Answer 27

Increased electron shielding and increased atomic radius down the group makes the outer electrons easier to lose.

Question 28

Does first and second ionisation energy decrease or increase down group 2?

Answer 28

Due to greater atomic radius and increased amounts of shielding as you go down the group, it makes it easier to remove an electron.

Question 29

Group 2 metals react with water in a redox reaction. What are the products when this reaction occurs and write down the equation

Answer 29

Metal hydroxide and hydrogen

Question 30

Group 2 metal react with oxygen. What are the products when this reaction occurs and write down the equation

Answer 30

Forms metal oxides Can form metal peroxides when reacted with excess oxygen.

Question 31

Group 2 metals react with dilute acids. What is produced in all reactions?

Answer 31

Hydrogen gas.

Question 32

What does a group 2 hydroxide and a dilute acid form?

Answer 32

A salt and water. This is a type of neutralisation reaction. The salt formed depends on the acid used.

Question 33

What is Ca(OH)2 used in?

Answer 33

Agriculture to neutralise acid soils

Question 34

What is Mg(OH)2 and CaCO3 used in?

Answer 34

Treating indigestion. As it is alkaline and can neutralise acids.

Question 35

Does solubility in group 2 hydroxides and sulfates, decrease or increase down the group?

Answer 35

Decrease

Question 36

What do halogens exists as?

Answer 36

Diatomic molecules with single covalent bonds.

Question 37

Does atomic radius of group 17 increase or decrease down the group?

Answer 37

Increases down the group due to additional electron shells.

Question 38

Why does electronegativity decrease down group 17?

Answer 38

Atomic radius and electron shielding increases. This means electrons in the outer shells are less strongly attracted to the nucleus, and so are more easily removed. Therefore, electronegativity decreases down group 17.

Question 39

Why does melting and boiling point increase down group 17?

Answer 39

Group 17 elements are simple covalent structures with weak induced dipole-dipole forces. Therefore strength of these intermolecular forces increases as the relative atomic mass of the molecule increases, so the strength of the induced dipole-dipole interactions increases down the group. Means more energy is required to overcome them, resulting in higher melting and boiling points. This is why fluorine is a gas at room temperature, whereas iodine is a solid.

Question 40

Why does reactivity decrease down group 17?

Answer 40

As atomic radius increases, it becomes harder to gain an electron as the positive attraction of the nucleus is weakened by additional shielding. Therefore, down group 17 it is harder to attract an electron so reactivity decreases.

Question 41

What is a disproportionatation reaction?

Answer 41

A reaction in which a species is both oxidised and reduced.

Question 42

What is the reaction between chlorine and water used in water treatment?

Answer 42

Chlorine reacts with cold water to produce chlorate ions (ClO-) and chloride ions via a disproportionation reaction where chlorine is both oxidised and reduced. Cl2 + H2O = ClO- + Cl- + 2H+ 0 +1 -1

Question 43

What is the reaction between chlorine with cold NaOH used in bleach production?

Answer 43

Sodium chlorate produced through a disproportionation reaction. 2NaOH + Cl2 = NaClO + NaCl + H2O

Question 44

What are the risks and benefits of using chlorine in water treatment?

Answer 44

Risks - toxic chlorine gas, possible risks from formation of chlorinated hydrocarbons Benefits - kills bacteria