AS-Level Redox
Redox
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Atoms are more stable when they have a full outer shell of electrons.
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To get a full outer shell of electrons atoms can lose or gain electrons to form charged ions.
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If an atom loses electrons it is oxidised.
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If an atom gains electrons it is reduced.
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Reducing agents get oxidised themselves, causing another species to be reduced.
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Oxidising agents get reduced themselves, causing another species to be oxidised.
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Half equations show how one species is being oxidised or reduced in a redox reaction.
AS-Level The Periodic Table
The Periodic Table
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The periodic table is an arrangement of elements based on increasing atomic number.
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Each column is called a group.
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Groups show the number of electrons in the outer shell of an element.
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Each row is called a period.
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All elements in one period have same number of electron sub-shells.
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The s, p and d blocks refer to the type of electron orbital that the outermost electrons are in for an element.
Atomic Radii
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Atomic radii decrease across a period in the periodic table.
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Atomic radii increase down a group in the periodic table.
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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.
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The greater the degree of electron shielding, the less tightly outer electrons are pulled into the nucleus so the larger the atomic radii.
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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
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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.
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As a trend, across a period in the periodic table, first ionisation energy increases.
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As a trend, down a group in the periodic table, first ionisation energy decreases.
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The stability of the ion formed when an atom loses an electron determines the energy required to remove the electron.
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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
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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.
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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.
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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.
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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
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Metals in group 2 of the periodic table are called the alkaline earth metals.
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Down group 2, atomic radius increases.
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Down group 2, first ionisation energy decreases.
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As a trend, melting point decreases down the group (magnesium is an exception).
Group 2 Metal Reactions
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Reactivity of group 2 metals increases going down the group.
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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.
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Group 2 metals form metal hydroxides when they react with water.
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Beryllium is not reactive enough to form a hydroxide.
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Magnesium does react with water but only slowly.
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The solubility of group 2 metal hydroxides increases down the group.
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The solutions are more alkaline as more hydroxide ions are released into solutions (pH of metal hydroxide solutions increases down the group).
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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
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Halogens are elements found in group 7 of the periodic table.
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Electronegativity is an atom’s tendency to attract a pair of electrons in a covalent bond towards itself. The halogens are highly electronegative elements.
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Down group 7, electronegativity decreases.
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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
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The melting and boiling points of the halogens increase going down group 7.
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Halogens exist in elemental form as simple, diatomic molecules held together by weak intermolecular forces (temporary induced dipole-dipole).
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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.
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The size of the halogen molecules increases down the group, meaning their melting and boiling points also increase.
Oxidising Power
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Halogen atoms gain one electron to have a full outer shell – they are reduced.
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For a halogen atom to gain an electron, another species must lose an electron (be oxidised).
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Halogens are oxidising agents as they can force other species to be oxidised.
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Oxidising agents are reduced.
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The ability of halogens to act as oxidising agents decreases down the group.
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Fluorine is the strongest oxidising agent of the halogens.
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Halogens can oxidise less reactive halide ions in displacement reactions.
Reducing Power
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Halide ions can give away an electron to become halogen atoms – they can be oxidised.
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For a halide ion to lose an electron, another species must gain an electron (be reduced).
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Halide ions are reducing agents, as they can force other species to be reduced.
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Reducing agents are oxidised!
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The ability of halide ions to act as reducing agents increases down the group.
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The reactions of halide ion salts with concentrated sulfuric acid show the different strengths of the halide ions as reducing agents.
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Chloride ions do not reduce the sulfur in sulfuric acid.
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Bromide ions can reduce the sulfur from 6+ to 4+ oxidation state.
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Iodide ions can reduce the sulfur from 6+ to 4+ to 2- oxidation state.
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Halide Ion Tests
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Silver nitrate (AgNO ) can be used to identify halide ions in a solution.
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Silver ions combine with halide ions to produce silver halide precipitates with different colours.
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Fluoride ions do not produce a precipitate.
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Chloride ions produce a white precipitate.
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Bromide ions produce a cream precipitate.
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Iodide ions produce a yellow precipitate.
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The precipitates have different solubilities in ammonia (NH ), which can be used to help further identify the halides.
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AgCl(s) dissolves in dilute ammonia.
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AgBr(s) dissolves in concentrated ammonia only.
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AgI(s) does not dissolve in concentrated ammonia.
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Dilute nitric acid is added before silver nitrate to remove impurities that may also form precipitates with the silver ions.
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