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3.2.1-3 Inorganic Chemistry Y1 (Group 7 Halogens (Testing for halide ions:…
3.2.1-3 Inorganic Chemistry Y1
Classification
elements with highest energy electrons in s-orbitals are in the s block
elements with highest energy electrons in p-orbitals are in the p block
elements with highest energy electrons in d-orbitals are in the d block
Period 3 Elements
atomic radius:
decreases across the period
all outer electrons are in third shell, and as more electrons are added, the force of attraction between the positive nucleus and the negative electrons increases, pulling them in closer, and atomic radius decreases
first ionisation energy:
increases across the period
number of protons and electrons increases across a period but outer electrons are still in third energy shell; increased charge on nucleus means it is harder to remove an electrons, as it is more strongly attracted to the nucleus (which has a greater positive charge)
drop in ionisation energy from group 2 to group 3: aluminium has outer most electron in p shell which is at a higher energy level than the s shell
drop in ionisation energy from group 5 to 6: in sulfur, electrons start to pair up in orbitals; electron repulsion
melting points:
giant structures on the left have a higher melting and boiling point
molecules and atomic structures have lower melting and boiling points
melting point increases across the period from sodium to aluminium; strength of metallic bond increases; increased charge of metal ion and so increased number of delocalised electrons holding metallic lattice
silicon has giant covalent structure so had a much higher melting point
melting point decreases from phosphorus to argon due to decreased strength of van der wals intermolecular forces (except sulfur, which is a larger molecule so has stronger van der wals forces)
Group 1, group 2, and group 3 elements (sodium, magnesium and aluminium) are metals (silicon in group 4 is a semi metal)
Group 5, group 6 and group 7 elements (phosphorus, sulfur, chlorine) are non-metals (argon in group 0 is a noble gas)
Group 2 Alkaline Earth Metals
melting points:
decrease slightly down the group; sea of electrons is further away from the nucleus so weaker force of attraction and weaker metallic bonds
atomic radius:
increases down the group; each element has an extra filled electron shell than the one before
first ionisation energy:
decreases down the group; atomic radii increases, and outermost electron is further away from the nucleus; weaker force of attraction due to larger distance, and electron shielding; electron is easier to remove
reaction with water:
increased reactivity down the group; due to decrease ionisation energy down the group; lower activation energy needed and so faster reaction
Uses:
magnesium used to extract titanium from TiCl₄
Mg(OH)₂ in medicine for indigestion; neutralise excess stomach acid
Ca(OH)₂ in agriculture; treat acidic soil
removing SO₂ from flue gas
acidified BaCl₂ solution tests for sulfate ions; produce white precipitate is sulfate ions present; must be acidified as otherwise it would absorb carbon dioxide from the atmosphere and make the solution cloudy; hydrogen carbonate ions and hydrogen sulfite ions are formed. and these ions form soluble compounds of barium
BaSO₄ in medicine; outline the gut in medical X-rays
Relative solubilities of hydroxides
Mg(OH)₂ is sparingly soluble
more soluble down the group
white solids
Relative solubilities of sulphates
BaSO₄ is insoluble
less soluble down the group
Relative solubility of carbonates
decrease down the group
Group 7 Halogens
electronegativity:
decreases down the group
shared electrons in a covalent bond get further away from the nucleus as the atoms get larger; increased shielding by more electron shells
boiling point:
increase down the group
larger atoms have more electrons so van der Waals forces are stronger
the lower the boiling point, the more volatile the element
Oxidising ability:
usually react by gaining electrons tp become negative ions with a charge of -1
decreases down the group
Reducing ability:
halide ions lose electrons to become halogen molecules
increase down the group; larger the ion, the further away the outer electron is from the nucleus so it is easier to lose
Testing for halide ions:
metal halides react with silver ions in aqueous solution to form a precipitate of the insoluble silver halide; not silver fluoride as it is soluble in water
add dilute nitric acid to the halide solution; get rid of any soluble carbonate of hydroxide impurities; if still present would form insoluble silver carbonate or silver hydroxide
add a few drops of silver nitrate solution and the halide precipitate is formed
AgCl: white precipitate; dissolves in dilute ammonia
AgBr: cream precipitate; dissolves in concentrated ammonia
AgI: yellow precipitate; insoluble in concentrated ammonia
Chlorine and chlorate ions
reaction of chlorine with water
Cl₂(g) + H₂O(l) ⇌ HClO(aq) + HCl(aq)
forms chloric(I) acid and hydrochloric acid; disproportionation reaction: when oxidation state of some atoms of the same element increase and others decrease
reaction takes place when chlorine is used to purify water for drinking, and in swimming pools; chloric acid is a bleach and also an oxidising agent
2Cl₂(g) + H₂O(l) --> 4HCl(aq) + O₂(g)
green --> colourless
NaCIO(s) + H₂0 ⇌ Na⁺(aq) + OH⁻(aq) + HClO(aq)
equilibrium moves to the left in alkaline solution, so swimming pools must be kept slightly acidic
use of chlorine in water treatments
benefits to health of water treatment by chlorine outweighs toxic effects
reaction of chlorine with cold dilute aqueous NaOH
NaClO is an oxidising agent and the active ingredient in household bleach
Cl₂(g) + 2NaOH(aq) --> NaCIO(aq) + NaCl(aq) + H₂0(l)
disproportionation reaction
atomic radii
increase down the group due to increasing number of electron shells
displacement reactions (oxidation):
halogens react with metal halides; halide in compound will be displaced by a more reactive halogen but not by a less reactive one
halogen will (always) displace the ion of a halogen below it in the periodic table
can't investigate fluorine in aqueous solution because it react with water
sodium halides and concentrated sulfuric acid
sodium chloride
NaCl(s) + H₂SO₄(l) --> NaHSO₄(s) + HCl(g)
acid base reaction; no oxidation state has changed
steamy fumes of hydrogen chloride are seen
sodium bromide
NaBr(s) + H₂SO₄(l) --> NaHSO₄(s) + HBr(g)
acid base reaction; similar to sodium chloride reaction
2H⁺ + 2Br⁻ + H₂SO₄(l) --> SO₂(g) + 2H₂O(l) + Br₂(l)
redox reaction; exothermic reaction; some of the bromine vapourises
steamy fumes of hydrogen bromide and brown fumes of bromine are seen
sodium iodide
steamy fumes of hydrogen iodide, black solid of iodine, and yellow solid can be seen; can also smell bad egg smell of hydrogen sulfide gas
acid base reaction
Nal(s) + H₂SO₄(l) --> NaHSO₄ (s) + Hl(g)
iodine ions reduce the sulfur even further; better reducing agents than bromide ions
8H⁺ + 8I⁻ + H₂SO₄(l) --> H₂S(g) + 4H₂O(l) + 4I₂(s)
Sammer Sheikh