Na, Mg, Al, Si, P, S, Cl
why is there increasing variation in oxidation number across period 3?
- in their oxides, the metals Na, Mg & Al display only 1 possible oxidation number each, corresponding to the loss of all their valence electrons
- P, S & Cl have a wider range of oxidation numbers as they can expand their octets -> oxidation number depends on the number of electrons used for bonding
note also: maximum positive oxidation number of an element often corresponds to the number of valence electrons in an atom of the element
why does bonding of the oxides change from ionic to covalent across period 3?
- difference in electronegativity between each element and oxygen decreases across the period, so bonding becomes increasingly covalent.
- covalent bonds usually form between atoms of similar electronegativity while ionic bonds usually form between atoms with a great difference in electronegativity
why is the trend in melting points of period 3 oxides as follows:
Na₂O - 1130C
MgO - 2850C
Al₂O₃ - 2070C
SiO₂ - 1700C
P₄O₁₀ - sublimes at 360C
SO₃ - 17C
- Na₂O, MgO and Al₂O₃ are all ionic oxides -> giant ionic lattice with strong electrostatic FOA between oppositely charged ions -> a large amt of energy required to overcome strong ionic bonds -> higher mp
- Na₂O has lower mp than MgO because Mg²⁺ has smaller ionic radius & higher charge than Na⁺ -> lattice energy of MgO is greater in magnitude than that of Na₂O -> more energy required to overcome stronger electrostatic FOA between Mg²⁺ & O²⁻ than in the Na₂O lattice -> MgO has higher mp
- Al₂O₃ has lower mp than MgO (despite theoretical larger magnitude L.E.) as Al³⁺ has very high charge density which can polarise the O²⁻ ion -> Al₂O₃ has covalent character -> magnitude of L.E. differs from expected
- SiO₂ has high mp as it has giant molecular structure -> strong covalent bonds between Si & O atoms (each Si bonded to 4 O, each O bonded to 2 Si) -> large amt of energy required to break strong covalent bonds -> high mp
- P₄O₁₀ & SO₃ are simple discrete covalent molecules with relatively weak intermolecular dispersion forces between molecules -> comparatively less energy required to overcome these weaker dispersion forces -> lower mp
describe the behaviour of basic period 3 oxides (Na₂O & MgO) with water & acids
Na₂O
- Na₂O dissolves completely in WATER in a vigorous & exothermic reaction, forming a strongly alkaline colourless solution of around pH13
- Na₂O (s) + H₂O (l) → 2Na⁺ (aq) + 2OH⁻ (aq)
- Na₂O dissolves in ACIDS in an exothermic reaction, producing a salt solution
- Na₂O (s) + 2H⁺ (aq) → 2Na⁺ (aq) + H₂O (l)
MgO
- MgO reacts in a very slow reaction with limited solubility in WATER. some Mg(OH)₂ is formed but it is onlysparingly soluble, so not all hydroxide ions are formed and released into the solution, forming a solution of pH9, lower than that of Na₂O
- MgO (s) + H₂O (l) → Mg(OH)₂ (s) [followed by]
- Mg(OH)₂ (s) ⇌ Mg²⁺ (aq) + 2OH⁻ (aq)
- MgO dissolves in ACIDS to produce a salt solution
- MgO (s) + 2H⁺ (aq) → Mg²⁺ (aq) + H₂O (l)
the basic hydroxides, NaOH and Mg(OH)₂ react with acids to form salt solutions in a classic neutralisation reaction
describe the behaviour of basic period 3 hydroxides with acids
- NaOH is a white pellet when solid, which dissolves exothermically to give a colourless solution. when reacting with an acid, there is no visible change, but the resultant mixture becomes warmer
- NaOH (aq) + HCl (aq) → NaCl (aq) + H₂O (l)
- Mg(OH)₂ is a white solid/ppt that is only sparingly soluble in water. it dissolves in acids to give a colourless solution
- Mg(OH)₂ (s) + 2HCl (aq) → MgCl₂ (aq) + 2H₂O (l)
why is Al₂O₃ amphoteric, and describe its behaviour with acids and bases
Al₂O₃ is an ionic oxide containing Al³⁺ and O²⁻ ions, but the high charge density of the Al³⁺ ion gives it great polarizing power, allowing the O²⁻ ion to be slightly polarized, conferring some covalent character to the oxide -> Al₂O₃ displays both basic & acidic properties
- Al₂O₃ dissolves in ACIDS to form a salt solution, with Al³⁺ as the cation
- Al₂O₃ (s) + 6H⁺ (aq) → 2Al³⁺ (aq) + 3H₂O (l)
- Al₂O₃ dissolves in excess BASE to form a salt solution with the aluminate complex ion: [Al(OH)₄]⁻
- Al₂O₃ (s) + 2OH⁻ (aq) + 3H₂O (l) → 2[Al(OH)₄]⁻ (aq)
describe the behaviour of the amphoteric Al(OH)₃ with acids & bases
- Al(OH)₃ dissolves readily in ACIDS to form a colourless salt solution, with Al³⁺ as the cation
- Al(OH)₃ (s) + 3HCl (aq) → AlCl₃ (aq) + 3H₂O (l)
- Al(OH)₃ dissolves in excess BASE to form the colourless aluminate complex ion, [Al(OH)₄]⁻
- Al(OH)₃ (s) + NaOH (aq) → Na[Al(OH)₄] (aq)
describe the behaviour of acidic period 3 oxides (SiO₂, P₄O₁₀, SO₃) with water and bases
SiO₂
- SiO₂, being a giant molecular oxide, requires more vigorous conditions to break up the macromolecular lattice
- SiO₂ does not dissolve in/react with water
- SiO₂ does not react with hot aqueous bases, and requires hot & concentrated strong bases before the reaction proceeds
- SiO₂ (s) + 2OH⁻ (conc) → SiO₃²⁻ (aq) + H₂O (l)
P₄O₁₀
- P₄O₁₀ reacts violently with WATER to give an acidic solution of H₃PO₄, forming a fairly strongly acidic solution of around pH2
- P₄O₁₀ (s) + 6H₂O (l) → 4H₃PO₄ (aq)
- P₄O₁₀ dissolves in strong BASES to form a salt solution. if excess OH⁻ is added:
- P₄O₁₀ (s) + 12OH⁻ (aq) → 4PO₄³⁻ (aq) + 6H₂O (l)
SO₃
- SO₃ reacts in a violent and very exothermic reaction in WATER, producing a very acidic mist of H₂SO₄ droplets, forming a strongly acidic solution of around pH1
- SO₃ (l) + H₂O (l) → H₂SO₄ (aq)
- SO₃ reacts directly with BASES to form salt solutions
- SO₃ (l) + 2OH⁻ (aq) → SO₄²⁻ (aq) + H₂O (l)
why is the trend of melting point of period 3 chlorides as follows:
NaCl - 801C
MgCl₂ - 714C
AlCl₃ - sublimes at 180C
SiCl₄ - -70C
PCl₅ - sublimes at 160C
- NaCl & MgCl₂ have giant ionic lattice structure with strong electrostatic FOA between oppositely charged ions -> a large amt of energy is required to overcome the strong ionic bonds -> high mp. NaCl & MgCl₂ are white crystalline solids at rt
- AlCl₃ is simple molecular in structure in the vapour phase, forming gaseous Al₂Cl₆ dimers. in the vapour phase, the following equilibrium exists: Al₂Cl₆ (g) ⇌ 2AlCl₃ (g). as the temperature increases, POE shifts right, favouring the monomers. AlCl₃ is a pale yellow powder at rt
- SiCl₄ exist as simple discrete covalent molecules with relatively weak intermolecular forces between molecules -> comparatively less energy is required to overcome these weaker forces -> lower mp. SiCl₄ is a colourless fuming liquid at rt
- PCl₅ is considered to have a simple molecular structure and is an off-white to pale yellow powder at rt
describe the behaviour of ionic chlorides (NaCl & MgCl₂) with water
NaCl
- NaCl dissolves in water to form a colourless solution of neutral pH.
- Na⁺ has low charge density and does not react with water molecules (no hydrolysis), so the solution remains neutral at pH7
MgCl₂
- MgCl₂ dissolves in water to form a colourless solution of slightly acidic pH
- MgCl₂ (s) + 6H₂O (l) → [Mg(H₂O)₆]²⁺ (aq) + 2Cl⁻ (aq)
- [Mg(H₂O)₆]²⁺ (aq) + H₂O (l) ⇌ [Mg(H₂O)₅OH]⁺ (aq) + H₃O⁺ (aq)
- Mg²⁺ has higher charge density than Na⁺, so the hydrated magnesium ion, undergoes slight hydrolysis to form a very weakly acidic solution of around pH6.5
- POE of 2nd eq lies on the left, but enough H₃O⁺ is produced to lower pH away from neutral
describe the behaviour of AlCl₃ with water & Al (III) ions with sodium carbonate
rxn with a limited amt of water
- when a few drops of water are added to solid AlCl₃, steamy white fumes of HCl are evolved and a white solid that is insoluble in water remains
- AlCl₃ (s) + 3H₂O → Al(OH)₃ (s) + 3HCl (g)
rxn with a large amt of water
- in an excess of water, AlCl₃ dissolves to form a colourless solution of acidic pH, with Al³⁺ forming the complex [Al(H₂O)₆]³⁺
- AlCl₃ (s) + 6H₂O → [Al(H₂O)₆]³⁺ (aq) + 3Cl⁻ (aq)
- [Al(H₂O)₆]³⁺ (aq) + H₂O (l) ⇌ [Al(H₂O)₅OH]²⁺ (aq) + H₃O⁺ (aq)
- the very high charge density of Al³⁺ sufficiently polarises & weakens the O-H bonds in the H₂O ligands such that they break readily to donate a proton. [Al(H₂O)₆]³⁺ hydrolyses to a larger extent than [Mg(H₂O)₆]²⁺, resulting in a distinctly acidic solution of around pH3
rxn with OH⁻
- adding a controlled amount of OH⁻ ions can further deprotonate the water ligands, until a white ppt of aluminium hydroxide is formed
- [Al(H₂O)₅OH]²⁺ (aq) + OH⁻ (aq) ⇌ [Al(H₂O)₄(OH)₂]⁺ (aq) + H₂O (l)
- [Al(H₂O)₄(OH)₂]⁺ (aq) + OH⁻ (aq) ⇌ [Al(H₂O)₃(OH)₃] (s) + H₂O (l)
- if an excess amt of OH⁻ is added, the white ppt dissolves to form a colourless soluion, forming the aluminate ion
- [Al(H₂O)₃(OH)₃] (s) + OH⁻ (aq) ⇌ [Al(OH)₄]⁻ (aq) + 3H₂O (l)
rxn with sodium carbonate
- a solution of Al³⁺ is usually acidic enough to react with sodium carbonate and produce effervescence of CO₂ gas & white ppt of aluminium hydroxide
- 2[Al(H₂O)₆]³⁺ (aq) + 3CO₃²⁻ (aq) → 2[Al(H₂O)₃(OH)₃] (s) + 3H₂O (l) + 3CO₂ (g)
describe the behaviour of SiCl₄ & PCl₅ with water
- both SiCl₄ & PCl₅ react violently with water to produce fumes of HCl
- if water is added in large excess, HCl fumes may not be observed as they mostly dissolve in water, forming a strongly acidic solution of pH1 containing hydrochloric acid
- SiCl₄ & PCl₅ undergoes completely hydrolysis:
- SiCl₄ (l) + 4H₂O (l) → SiO₂.2H₂O (s) + 4HCl (aq)
- PCl₅ (s) + 4H₂O (l) → H₃PO₄ (aq) + 5HCl (aq)
- upon dropwise addition of water, PCl₅ produces POCl₃:
- PCl₅ (s) + H₂O (l) → POCl₃ (l) + 2HCl (g)
why does reducing power increase down group 2 (i.e. tendency to be oxidised increases)?
reactivity of group 2 metals increases down the group, so as atomic radii increases, the metal atoms lose their electrons more readily so they form cations more readily -> reducing power/tendency to be oxidised of group 2 metals increases
why does thermal stability of group 2 carbonates increase down the group?
- down group 2, cationic radius increases down the group while the charge remains the same -> charge density decreases down the group -> polarising power of the cation decreases down the group
- down the group, the cation is less able to distort the electron cloud of the carbonate, weakening the C-O bonds within the carbonate anion to a smaller extent -> covalent bonds within the carbonate anion are less likely to be broken down the group
- hence ease of decomposition decreases/thermal stability increases, so higher temperatures are required to decompose the compound
larger anions are more susceptible to polarisation of the electron cloud. only polyatomic anions are susceptible to decomposition as monoatomic anions cannot be broken down further
why does the volatility of group 17 elements decrease down the group?
volatility: how easily the substance in the liquid phase may be converted to the gas phase
- group 17 elements exist as simple, non-polar diatomic molecules with dispersion forces
- down the group, the size of electron cloud and hence polarisability of the halogen molecule increases -> more energy is required to overcome the strong dispersion forces between the molecules -> volatility decreases down group 17 & mp & bp increases
why does oxidising power of group 17 elements decrease down the group (i.e. tendency to be reduced decreases)?
- the outermost electron shell of halogens contains 7 electrons, so their chemistry is dominated by a tendency to gain a completely filled outermost electron shell -> halogens tend to be reduced in a redox reaction
- oxidising power can be measured by the E° value -> E° becomes less positive down the group -> oxidising power decreases down the group as they become less likely to be reduced in a redox reaction
how do group 17 elements displace each other?
- a halogen higher in the group can oxidise a halogen below it -> a more reactive halogen (higher in the group) displaces a less reactive one (lower in the group) from its compound
how do group 17 elements react with aqueous solutions containing iron (II) ions?
- in the presence of a halogen oxidising agent, an aqueous solution of Fe²⁺ is oxidised to Fe³⁺
- only Cl₂ & Br₂ are able to oxidise Fe²⁺ to Fe³⁺ (calculate E°)
why can chlorine and bromine oxidise S₂O₃ to SO₄²⁻, but iodine only oxidises S₂O₃ to S₄O₆²⁻?
oxidation state of S in
S₂O₃: +2
SO₄²⁻: +6
S₄O₆²⁻: +2.5
- iodine is a weaker oxidising agent than chlorine and bromine (since it is lower in group 17) -> oxidises sulfur to a smaller extent than chlorine & bromine
hydrogen halides: HF, HCl, HBr, HI
why does thermal stability of hydrogen halides decrease down the group?
- down the group, atomic radius increases from F to I -> H-X bond length increases -> bond strength decreases -> less energy is required to break the H-X bond -> thermal stability decreases
what are the conditions required for hydrogen halides to decompose?
- HCl does not decompose even on strong heating
- HBr decomposes on strong heating & brown fumes of bromine are observed
- HI decomposes when a red hot rod is plunged into it & violet fumes of I₂ are observed
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atomic structure18
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chem bonding46
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gaseous state16
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energetics24
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kinetics19
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chem equilibria14
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intro to organic chem16
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isomerism15
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alkanes21
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alkenes16
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arenes25
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solubility equilibria15
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acid base equilibria18
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halogen derivatives29
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hydroxy compounds39
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carbonyl compounds18
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carboxylic acid & derivatives32
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nitrogen compounds23
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electrochemistry13
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periodicity21
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transition elements18
