AS - Inorganc | ChemistryStudent

AS-Level Redox

Redox

Atoms are more stable when they have a full outer shell of electrons.
To get a full outer shell of electrons atoms can lose or gain electrons to form charged ions.
If an atom loses electrons it is oxidised.
If an atom gains electrons it is reduced.
Reducing agents get oxidised themselves, causing another species to be reduced.
Oxidising agents get reduced themselves, causing another species to be oxidised.
Half equations show how one species is being oxidised or reduced in a redox reaction.

The Periodic Table

The periodic table is an arrangement of elements based on increasing atomic number.
Each column is called a group.
- Groups show the number of electrons in the outer shell of an element.
Each row is called a period.
- All elements in one period have same number of electron sub-shells.
The s, p and d blocks refer to the type of electron orbital that the outermost electrons are in for an element.

Atomic Radii

Atomic radii decrease across a period in the periodic table.
Atomic radii increase down a group in the periodic table.
Electrons in inner sub-shells block or shield the positive charge of the nucleus from the electrons in the outer sub-shells; this is called inner electron shielding.
The greater the degree of electron shielding, the less tightly outer electrons are pulled into the nucleus so the larger the atomic radii.
Across a period, the number of protons in the nucleus of an element increases but the level of shielding stays the same; this pulls the outer electrons in tighter and decreases the atomic radii.

First Ionisation Energies

The first ionisation energy of an element is the energy required to remove one mole’s worth of electrons from one mole's worth of gaseous atoms.
As a trend, across a period in the periodic table, first ionisation energy increases.
As a trend, down a group in the periodic table, first ionisation energy decreases.
The stability of the ion formed when an atom loses an electron determines the energy required to remove the electron.
Elements in groups 3 and 6 do not follow the general trend for first ionisation energy. It is easier to remove an electron from a group 3 element compared to a group 2 element and from a group 6 element compared to a group 5 element in the same period.

Boiling Points

The melting points and boiling points of a substance are linked to the type of structure the substance has and how much energy is required to break apart that structure.
Simple molecular structures are held together with weak intermolecular forces that require low amounts of energy to break – they have low melting and boiling points.
Giant structures (ionic and covalent) are held together by strong atomic bonding that requires high amounts of energy to break – they have high melting and boiling points.
Across a period in the periodic table, melting points of elements increase from groups 1 to 4 (giant structures), then decrease from groups 5 to 8 (simple molecular substances).

Group 2 Metals

Group 2 Metal Reactions

Reactivity of group 2 metals increases going down the group.
Down the group, outer electrons are held less tightly due to increased inner electron shielding, so they are easier to lose, making the metals more reactive.
Group 2 metals form metal hydroxides when they react with water.
- Beryllium is not reactive enough to form a hydroxide.
- Magnesium does react with water but only slowly.
The solubility of group 2 metal hydroxides increases down the group.
- The solutions are more alkaline as more hydroxide ions are released into solutions (pH of metal hydroxide solutions increases down the group).
The solubility of group 2 metal sulfates decreases down the group, with barium sulfate being completely insoluble.

The Halogens and Electronegativity

Halogens are elements found in group 7 of the periodic table.
Electronegativity is an atom’s tendency to attract a pair of electrons in a covalent bond towards itself. The halogens are highly electronegative elements.
Down group 7, electronegativity decreases.
Fluorine is the most electronegative halogen and forms highly polar bonds when bonded to atoms with low electronegativities. Polarity of such bonds decreases down the group.

Boiling Points

The melting and boiling points of the halogens increase going down group 7.
Halogens exist in elemental form as simple, diatomic molecules held together by weak intermolecular forces (temporary induced dipole-dipole).
The larger the molecules, the larger the intermolecular forces that can arise between them, which require more energy to overcome – giving the substance a high melting point.
The size of the halogen molecules increases down the group, meaning their melting and boiling points also increase.

Oxidising Power

Halogen atoms gain one electron to have a full outer shell – they are reduced.
For a halogen atom to gain an electron, another species must lose an electron (be oxidised).
Halogens are oxidising agents as they can force other species to be oxidised.
Oxidising agents are reduced.
The ability of halogens to act as oxidising agents decreases down the group.
Fluorine is the strongest oxidising agent of the halogens.
Halogens can oxidise less reactive halide ions in displacement reactions.

Reducing Power

Halide ions can give away an electron to become halogen atoms – they can be oxidised.
For a halide ion to lose an electron, another species must gain an electron (be reduced).
Halide ions are reducing agents, as they can force other species to be reduced.
Reducing agents are oxidised!
The ability of halide ions to act as reducing agents increases down the group.
The reactions of halide ion salts with concentrated sulfuric acid show the different strengths of the halide ions as reducing agents.
- Chloride ions do not reduce the sulfur in sulfuric acid.
- Bromide ions can reduce the sulfur from 6+ to 4+ oxidation state.
- Iodide ions can reduce the sulfur from 6+ to 4+ to 2- oxidation state.

Halide Ion Tests

Silver nitrate (AgNO ) can be used to identify halide ions in a solution.
- Silver ions combine with halide ions to produce silver halide precipitates with different colours.
  - Fluoride ions do not produce a precipitate.
  - Chloride ions produce a white precipitate.
  - Bromide ions produce a cream precipitate.
  - Iodide ions produce a yellow precipitate.
- The precipitates have different solubilities in ammonia (NH ), which can be used to help further identify the halides.
  - AgCl(s) dissolves in dilute ammonia.
  - AgBr(s) dissolves in concentrated ammonia only.
  - AgI(s) does not dissolve in concentrated ammonia.
Dilute nitric acid is added before silver nitrate to remove impurities that may also form precipitates with the silver ions.