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A process where an element ends up in two different compounds, one with a higher oxidation state and one with a lower.

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Like diamond,  cubic boron nitride is extremely hard with a high melting point as covalent bonds must be broken to break or melt the solid this leads to its use as a wear- resistant coating on industrial abrasive

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Can form several hexagonal sheets similar to graphite, but the atoms in different layers lie directly above one another, with each boron having a nitrogen atom directly above & below it. This differs from graphite as the layers in graphite are arranged so that atoms on adjoining layers or not directly above one another. Before this the force is  atoms on adjoining layers are not directly above one another. The forces between layers are weak so boron nitride shares the ability for layers to slip over one another with graphite, so it’s used as a lubricant.
The electrical properties of boron nitride are very different from graphite as there are no delocalised electrons present, with electrons localised as lone pairs on nitrogen atoms. The B-N bonds are polar due to the different electronegativities of the 2 atoms. This makes BN an insulator & leads to its use in electronics as a substrate for seminconductors, microwave-transparent windows, & a structural material for seals, electrodes & catalyst carriers in fuel cells & batteries.

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Lead(II) chloride is the most stable chloride of lead. It is a white ionic solid made of Pb²⁺ & Cl⁻ ions. Because it's an ionic compound it doesn't react with water & doesn't dissolve in cold water, it can be dissolved in hot water. This is common with most lead(II) compounds, which are insoluble in cold water.

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As chlorine and chlorate ions have great oxidising power, they are often used in bleaches.
Bleaching is an oxidation reaction with the oxidised form of the coloured material or dye being colourless.
ClO- is able to kill bacteria as the microbe cells are oxidised. - This is the basis of chlorination of water.

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Reactions of chlorine with sodium hydroxide

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Reactions of chlorine with sodium hydroxide

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When chlorine is bubbled through water, a reversible reaction occurs:
Cl2 (g) + H2O (l) ⇌ HCl (aq) + HOCl (aq)
The chlorine is being both oxidised and reduced.

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Addition of sulfuric acid to sodium bromide produces HBr gas.
NaBr(s) + H₂SO₄(conc) --> NaHSO₄(s) + HBr(g)
The sulfuric acid oxidises some of the HBr to form brown fumes of Br₂, & SO₂ gas. The hydrobromic acid is slightly easier to oxidise(E° - +1.09V) & so the sulfuric acid cases the redox reaction:
2HBr(s) + H₂SO₄(conc) --> SO₂(g) + Br₂(g) + 2H₂O(l)
Sulfur is reduced from +6 in H₂SO₄ to +4 in SO₂.
Bromine is oxidised

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Addition of sulfuric acid to sodium chloride produces HCL gas. The hydrochloric acid is difficult to oxidise(E° - +1.36V) & so the sulfuric acid doesn't cause any redox reaction.
NaCl(s) + H₂SO₄(conc) --> NaHSO₄(s) + HCl(g)
Obs- steamy fumes of HCl

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Reactions of concentrated sulphuric acid with sodium halides

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Reactions of concentrated sulphuric acid with sodium halides

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Conc. sulphuric acid is a strong acid & oxidising agent. The different ease of oxidising the halide ions means that the reactions undertaken by the different sodium halides are very different. When any sodium halide is added to concentrated sulphuric acid, the hydrogen halide is formed as a steamy gas.
The sulfuric acid or products formed from it may oxidise the halide in the hydrogen halide to form the halogen if the halide is relatively easy to oxidise. This process becomes easier lower down the group.

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Halogens have different strengths as oxidising agents, their oxidising power decreasing down the group. This ability to remove electrons from other species can be measured using standard electrode potentials for the halogens.

Chlorine is most positive, readily gains electrons to form chloride ions. This also shows it's difficult to oxidise chloride ions.

Iodine is least positive, gains electrons less readily than bromine and chlorine. Also showing that it's easier to to oxidise iodide ions

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These chemicals are sets of diatomic molecules, containing elements in all 3 physical states. Chemical properties of all elements are similar, all being non-metals although the range of stability of oxidation states shows clear patterns on descending the group.

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Reactions of solutions of lead(II) compounds, Pb2+(aq)

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Reactions of solutions of lead(II) compounds, Pb2+(aq)

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Ionic compounds & practically all of them are insoluble in water. The only 2 common compounds which do dissolve are lead nitrate, Pb(NO₃)₂, and lead ethanoate, Pb(CH₃COO)₂. They produce various precipitates:

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CCl4 does not react with water, it forms a separate layer under the water. The carbon atoms cannot combine easily with the water molecules. The lack of reactivity is due to the absence of available d-orbitals in the valence shell meaning the octet can't expand to allow water molecules to combine with carbon atoms.
SiCl₄, reacts very quickly with water in a hydrolysis reaction. The reaction produces fumes of hydrogen chloride gas & silicon dioxide, SiO2, as a solid precipitate. The reaction becomes more vigorous down the group as the bonds in the compound become weaker.
SiCl₄(l) + 2H₂O(l) --> SiO₂(s) + 4HCl(g)
The reasons for the increased reactivity is that silicon has available 3d-orbitals as well as 3s & 3p orbitals involved in boding to the chlorine atoms. The lone pairs of H₂O can form coordinate bonds with these empty d-orbitals, giving a complex molecule that can then eliminate 2 HCl molecules. It can then eliminate 2 more HCl molecules, leaving SiO₂.
Can also see the product written as Si(OH)₄, with the overall reaction as:
SiCL₄(l) + 4H₂O(l) --> Si(OH)₄(s) + 4HCl(g)
Both of these are acceptable,. Hydrated silicon dioxide(SiO₂.2H₂O) has the same overall composition as Silicon hydroxide(Si(OH)₄)& spectroscopic analysis does not distinguish between these.

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Stable chlorides of carbon & silicon are the tetrachloride's, CCL₄ & SiCL₄. They are both colourless liquids containing individual covalent molecules. Molecules found in both are tetrahedral due to the 8 electrons in the valence shell.

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Lead(II) oxide, PbO is an orange solid which contains bonding which is mainly ionic. PbO is an amphoteric oxide, reacting with acids & bases:
PbO(s) + 2HNO₃(aq) --> Pb(NO₃)₂(aq) + H₂O(l)
PbO(s) + 2NaOH(aq) + 2H₂O(l) --> Na₂(Pb(OH)₄)(aq)

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CO₂ is a colourless gas made up of small covalent molecules. This is an acidic oxide as the oxide is soluble in water to give the very weak acid, carbonic acid:
CO₂(g) + H₂O(l) ⇌ H⁺(aq) + HCO₃⁻(aq)
CO₂ will react with alkalis to form a salt. Salts produced in this way car carbonates or hydrogencarbonates:
CO₂(g) + 2NaOH(aq) --> Na₂CO₃(aq) + H₂O(l)
CO₂(g) +NaOH(aq) -->NaHCO₃(aq)

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Classify metal oxides as basic oxides & non-metal oxides as acidic oxides. Some metals are amphoteric oxides. The change from non-metal at the top of group 4 to metal at the bottom is reflected in the acid-base properties.

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Lead(II) oxide, PbO is the most stable oxide of lead. Lead(IV) oxide, PbO2 will act as an oxidising agent as it easily becomes reduced from +4 to +2. All lead(IV) compounds are oxidising agents, so PbO2 may be used for this purpose:
PbO₂(s) + 4HCl(conc) --> PbCl₂(s) + Cl₂(g) + 2H₂O(l)

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CO₂ is the most stable oxide of carbon. CO is the only stable compound to contain carbon in the +2 state. CO will act as a reducing agent as it easily becomes oxidised from +2 to +4. It is used when extraction metals form their oxides. E.g
Iron - Fe₂O₃(s) + 3CO(g) --> 2Fe(s) + 3CO₂(g)
Copper - CuO(s) + CO(g) --> Cu(s) + CO₂(g)
This method can only be used for the oxides of less reactive metals. The oxides of the more reactive metals(anything above zinc) are too stable and will not react.

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Oxidation states shown in group 4 are +2 & +4 with the stability of the +2 oxidation state increasing down the group as the inert pair effect becomes more significant. The most stable oxidation state for all elements in group 4 is +4 apart from lead. The stability of the oxidation states governs redox properties of the compounds, & oxides of carbon and lead are an example of this.

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A range of metals & non-metals. The changes from top to bottom of the group are significant. The first element, carbon, is a non-metal which forms a huge range of covalent compounds with carbon having a +4 oxidation state, whilst the heaviest stable atom is lead, a metal which generally forms ionic compounds with the +2 oxidation state being most stable.

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Boron forms a lot of compound with nitrogen. There are a total of 12 electrons on the 2 atoms. The atomic radii of all the atoms are similar with carbon being almost exactly the acrrage of the radii of boron & nitrogen, with. Similar relationship in their electronegativities. This leads to boron nitride, BN, having several forms which are similar to the several different forms of carbon.

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An atom that doesn't have a full outer shell. When group 3 elements form compounds, they usually have 3 covalent bonds e.g AlCl₃. To fill their outer shell, they will often form coordinate bonds with other compounds or form dimers, as in the case of AlCl₃, in the gas phase, which forms Al₂Cl₆.
Each aluminium atom uses a lone pair on a chlorine atom to form a coordinate bond. The dimer, is no longer electron-deficient. Other atoms can do this too, the compounds are classified as donor-acceptor compounds.
A typical example is the compound formed between electron-deficient BF₃ and the lobe pair on NH₃, the compound formed is no longer deficient.

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Boron & aluminium are the 2 most common elements. They have very different physical properties as one is metal & the other a non-metal.

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Many P-block elements form amphoteric oxides, typically metals close to the line separating the 2 characteristics. They must react with both acids & bases.
For aluminium oxide or hydroxide:
Al₂O₃ + 6HCl --> 2AlCl₃ + 3H₂O or Al(OH)₃ + 3H⁺ --> Al³⁺ + 3H₂O
Al₂O₃ + 2NaOH + 3H₂O -> 2Na[Al(OH)₄] or Al(OH)₃ + OH⁻ --> [Al(OH)₄]⁻
For lead(II) oxide or hydroxide:
PbO + 2HNO ₃ --> Pb(NO₃)₂ + 2H₂O or Pb(OH)₂ + 2H⁺ --> Pb²⁺ + 2H₂O
PbO + 2NaOH + H₂O --> Na₂[Pb(OH)₄] or Pb(OH)₂ + 2OH⁻ --> [Pb(OH₄]2-
Solutions containing amphoteric metal compounds form precipitates when sodium hydroxide added. Precipitates are metal hydroxides. Hydroxides can react with more sodium hydroxide, these precipitates redissolve:
For aluminium:
Al³⁺(aq) + 3OH⁻ --> Al(OH)₃(s) + OH⁻(aq) --> [Al(OH)₄]⁻(aq)
For lead:
Pb²⁺(aq) + 2OH⁻(aq) --> Pb(OH)₂(s) then Pb(OH)₂(s) +2OH⁻(aq) --> [Pb(OH)₄]²⁻(aq)

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In P-block elements at the top of the groups are non-metals, where as elements at the bottom of the groups are metals. This change in properties leads to the characteristic zig-zag line between metals and non-metals.
This change has a significant effect on the bonding and properties of the p-block compounds.

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Elements in the second period can only have 8 electrons surrounding them - making 4 pairs. This limits the number of bonds that can be formed in the first row elements:

Carbon: 4 covalent bonds

Nitrogen: 3 covalent bonds & 1 lone pair

Elements in period 3 and below are able to expand their octet. Meaning every electron in the outer shell can be used to form a covalent bond as there is no longer a limit of 8 electrons in the outer shell. This affects the number of bonds that can be formed for elements in groups 5, 6 & 7:

Phosphorus: 5 covalent bonds, PCL₅

Chlorine: up to 7 covalent bonds, ClO₄⁻

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Stability of the lower oxidation state becomes greater down the group. This tendency for heavier elements to form lower oxidation states is called the inert pair effect.

when an element has an oxidation state of 4, it involves 4 electrons

when it has an oxidation state of 2, the inner 2 do not become involved, this ns² pair is called the inert pair.

The trend occurs in groups 3, 4 and 5. It's only the lower members of this group that show the tendency to lower oxidation states:
+1 for group 3, +2 for group 4 and +3 for group 5.
In group 3 and 4, the lower oxidation state is not observed in the first 2 of each group(exception of +2 in CO), but in group 5 the upper members(nitrogen & phosphorus) exhibit lower oxidation state more frequently.

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Elements in P-block typically show 2 oxidation states which equals the group number and a lower oxidation state which is 2 less.

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Their outer electrons are located in the P-block. Every P-block element has a full s sub-shell in their outer shell, with between 1 and 6 further electrons in their p sub-shell.