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.

AS-Level The Periodic Table

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).

AS-Level Group 2 Metals

Group 2 Metals

  • Metals in group 2 of the periodic table are called the alkaline earth metals.

  • Down group 2, atomic radius increases.

  • Down group 2, first ionisation energy decreases.

  • As a trend, melting point decreases down the group (magnesium is an exception).

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.

AS-Level Group 7 (Halogens)

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.

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