OCR A Chemistry - Periodicity

Properties about The Periodic Table

Properties about The Periodic Table

Elements are arranged in increasing atomic number
Groups are vertical columns with elements with the same outermost shell electrons
Periods are vertical columns with elements with the same number of shells

What is periodicity?

Repeating trends in the properties of elements

What is the First Ionisation Energy?

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

Equation for the First Ionisation Energy

E(g) --> E+ + e-

Ionisation Energy: Atomic Radius

The greater the atomic radius, the greater the distance between the nucleus and the outermost shell electrons, therefore less nuclear attraction and less energy is required to remove one electron

Ionisation Energy: Nuclear Charge

The greater the atomic number, the greater the proton number and nuclear charge. Therefore there is greater attraction between the nucleus and the outer electrons if the nuclear charge is greater

Ionisation Energy: Electron Shielding

Electrons are negatively charges particles and the repel each other. This causes inner and outermost shells to repel each other, known as the shielding effect which reduces the attraction between the nucleus and outer shell electrons and therefore less energy is required to remove the outermost electron

Successive Ionisation Energies

As more electrons are removed, more energy is required to remove the electron as the nuclear charge becomes greater than the electrons on the shells

Why are there jumps in Ionisation Energy Graphs?

Jumps are where the electron being removed is from a different shell to the previous one, which is closer to the nucleus and with less shielding

Trends in First Ionisation Energy Across a Period

Nuclear Charge increases
Same Shell - similar shielding
Nuclear Attraction Increases
Atomic Radius decreases
Therefore First Ionisation Energy increases

Trends in First Ionisation Energy Down a Group

Atomic Radius Increases
Shielding Increases - more inner shells
Nuclear Attraction decreases
Therefore First Ionisation Energy decreases

Comparing Beryllium and Boron

The 2p subshell of boron has more energy than the beryllium 2s subshell, the electron is easier to remove

Comparing Nitrogen and Oxygen

The fall marks the pairing of electrons in the 2p subshell.
In oxygen the paired electrons repel one another making it easier to remove the electron therefore the ionisation energy is less than nitrogen

Metallic Bonding

Strong electrostatic attraction between cations (metal ions, these are fixed giving structure) and anions (delocalised electrons, are mobile )
Held together in a Giant Metallic Lattice

Properties of Metals: Electrical Conductivity

In solid and liquid states delocalised electrons are mobile, so they can move through the structure and carry the charge

Properties of Metals: Melting and Boiling Points

Melting point of metals depends on the strength of the metallic bonds. High temperatures are needed to overcome the strong electrostatic attraction, therefore have high melting and boiling points

Solubility of Metals

Metals are not soluble however, there may be some interaction between polar solvents and the charges in the lattice - this would lead to a reaction rather than it dissolving

The structure of atoms bonded by strong covalents bonds are known as

Giant Covalent Lattice

Properties of Non-Metals: Melting and Boiling Points

Have high melting and boiling points because covalent bonds require lots of energy to break them

Properties of Non-Metals: Solubility

Insoluble in all solvents
The covalent bonds are far too strong to be broken by interacting with solvents

Properties of Non-Metals: Electrical Conductivity

All are non-conductors apart from Graphene and Graphite. These can conduct electricity because each carbon atom has a free delocalised electron that is mobile and can carry the charge through the structure

Trends of Melting Points in The Periodic Table

Melting point increases from Group 1 to 14
Sharp decrease between Group 14 and 15 (due to the change from giant to simple molecular structures)
Low melting points for Groups 15 to 18

Why are Group 2 Metals Reducing Agents?

In redox reactions, each metal atom is oxidised losing two electrons. Another species will gain these two electrons and that species will be reduced. Therefore Group 2 Metals are reducing agents

Group 2 Metals Reacting with Oxygen

Form a Metal Oxide
Made of M2+ and O2+ ions
The metal is oxidised and the oxygen is reduced

Group 2 Metals Reacting with Water

Form and Alkaline Hydroxide and Hydrogen Gas
GF: M(OH)2
Down the group the reactivity increases and becomes more vigorous

Group 2 Metals Reacting with Dilute Acids

Form a Salt and Hydrogen Gas
The metal is oxidised and the hydrogen is reduced

Group 2 Trends in Reactivity and IE

Reactivity increases down the group because the atom requires the input of two ionisation energies

Group 2 Oxides and Water

Release OH- ions and form the alkaline solutions of metal hydroxides
Slightly soluble in water
The above further reacts to form the metal hydroxide precipitate

Solubility of Group 2 Hydroxides

Solubility increases down the group
Solutions have more OH- ions and are more alkaline/pH increases

Group 2 Compounds in Agriculture

Calcium Hydroxide is added to increase the pH of acidic soils - it neutralises the acid in the soil forming water

Group 2 Compounds in Medicine

Used as Antacids
Neutralisation reaction takes place between the acid and base

Trends in Boiling Points of Halogens

BP increases down the group
More electrons
Stronger London Forces
More energy required to break the intermolecular forces
Chlorine is a gas and Iodine is a solid at RTP

Redox Reactions of Halogens

Each halogen atom is reduced
Cl2 + 2e- --> 2Cl-
Halogens are Oxidising Agents

Halogen-Halide Displacement Reactions

Reactivity of Halogens decrease down the group
The tested halogen solution is added to the aqueous solutions containing the halide ions of the other two and if it displaces it there is a colour change

Chlorine in Water changes colour to

Pale Green

Bromine in Water changes colour to

Orange

Iodine in Water changes colour to

Brown

To distinguish the colours of the halogens in water, what is added?

An organic non-polar solvent is used eg. cyclohexane

Bromine Solution in Cyclohexane changes colour to

Orange

Iodine Solution in Cyclohexane changes colour to

Violet

Chlorine and Bromine Ions

Chlorine is reduced and Bromine is oxidised

Trends in Reactivity of Halogens

Atom Radius increases
Shielding increases because there are more inner shells
Less Nuclear Attraction
Therfore Reactivity Decreases

What is a disproportionation reaction?

In a redox reaction the same element is both oxidised and reduced

Reaction of Chlorine and Water

Products are Chloric Acid (HClO)and Hydrochloric Acid (HCl)
Chlorine is used in water purification as a disinfectant
Chloric Acid is also a weak bleach

Reaction of Chlorine and Cold Dilute Aqueous NaOH

NaClO is formed as well as NaCl and H2O

Benefits and Risks of Using Chlorine

Chlorine is used to kill bacteria, however in large quantities it can be toxic as its a respiratory irritant

Testing for Halide Ions

Aqueous Halides react with Aqeuous Silver Ions to Form Precipitates of Silver Halides

Ag+ + X- --> AgX

Chlorine - White
Bromine - Cream
Iodine - Yellow

Add aqueous ammonia to test solubility since colours can be difficult to tell apart

Chlorine - Soluble in dilute NH3
Bromine - Soluble in conc. NH3
Iodine - Insoluble in conc. NH3

Carbonate Test

Add nitric acid to the solution
If you see effervescence it could be a carbonate
Bubble the gas through limewater
If it turns cloudy/milky as a white ppt has formed a carbonate ion is present

Sulfate Test

Barium Ions are added to the solution
Barium Sulphate is formed
Its a white ppt and is insoluble in water

What is the sequence of testing?

Carbonate - The other two don't produce bubbles with nitric acid
Sulphate - Barium carbonate is also white and insoluble, so you should rule out the possibility of a carbonate earlier on
Halides

Mixture of Ions?

Carry out test in the same order

Carbonate - Continue adding nitric acid till effervescence stops

Sulphate - Add excess of Barium Nitrate, any ions left will form a ppt of barium sulphate then filter solution to remove the barium sulphate

Halide - Add NH3 to confirm

Ammonium Ions Test

Aqueous ammonium ions and Aqueous hydroxide ions react to form ammonia gas. The mixture is then warmed and tested with damp red litmus paper and should turn blue since its an alkaline