Inorganic Chemistry: Paper 1

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79 Terms

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periodicity
repeating pattern/trends (of physical or chemical properties)
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period 3: atomic radius
decreases along the period

increased nuclear charge for same number of electron shells and shielding, produces greater attraction to outer electron
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period 3: ionisation energy
increases along the period

decreased atomic radius and increased nuclear charge, more energy required to remove electrons
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period 3: melting point (Na, Mg, Al)
metallic bonding

high melting point as there is a strong electrostatic attraction between delocalised electrons and positive metal ions

Al the highest because greater positive charge on ions and more free electrons, greater electrostatic attraction
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period 3: melting point (Si)
macromolecular

highest melting point because must overcome strong covalent bonds, requiring lots of energy.
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period 3: melting point (P, S, Cl)
simple molecule

lower melting point as only held together by weak van der Waal intermolecular forces, requiring little energy

S > P > Cl because S8, P4, Cl2 and greater the Mr, the greater the surface area for van der Waals, more van der Waals, more energy to overcome
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period 3: melting point (Ar)
simple molecule

exist as single atoms so weak van der Waals and small Mr
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group 2: atomic radius
increase down the group

add new shell, increasing distance to outer electron, increased shielding, reducing nuclear attraction
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group 2: ionisation energy
decreases down the group

new shells, atomic radius increases, increased shielding, reduces nuclear attraction, less energy to remove an electron
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group 2: melting point
decreases down the group

metallic bonding: larger the ion within the lattice weakens the attractive force as it acts over a larger distance between positive nucleus and delocalised electrons
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group 2: reacting with water
M(s) + 2H2O(l) → M(OH)2(aq) + H2(g)

increase in reactivity down the group

Mg very slow: forms layer of Mg(OH)2 which is sparingly soluble so stops the reaction

weak alkaline solution
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Mg with steam
Mg(s) + H2O(g) → MgO(s) + H2(g)
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Mg in extraction of titanium
TiCl4 + 2Mg → 2MgCl2 + Ti

displacement reaction
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RP4: test for group 2 cations (NaOH)

1. initially add 10 drops NaOH

MgCl2, CaCl2, SrCl2 = slight white ppt

BaCl2 = colourless solution


2. add excess NaOH

MgCl2 = white ppt

CaCl2, SrCl2 = slight white ppt

BaCl2 = colourless

because Mg(OH)2 is sparingly soluble
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RP4: test for group 2 cations (H2SO4)

1. initially add 10 drops H2SO4

BaCl2, SrCl2 = white ppt

MgCl2, CaCl2 = slight white ppt


2. add excess H2SO4

BaCl2, SrCl2 = white ppt

CaCl2 = slight white ppt

MgCl2 = colourless solution
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group 2: solubility of hydroxides
increases in solubility down the group

Mg(OH)2 is sparingly soluble
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use of Mg(OH)2
milk of magnesia is used as an antacid as it neutralises stomach acid, for indigestion
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use of Ca(OH)2
slaked lime is used in agriculture to raise the pH of a field (is basic)
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use of CaO/CaCO3
to remove SO2 from flue gases, prevent the release into atmosphere

CaO + 2H2O + SO2 → CaSO3 + 2H2O

CaCO3 + 2SO2 + H2O → Ca(HSO3)2 + CO2
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group 2: solubility of sulfates
decrease in solubility down the group

BaSO4 is insoluble
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use of BaSO4
barium meal is digested and pass through digestive system, allowing the outlining of the gut by medical x-rays

it is completely insoluble so will not dissolve into blood, despite being toxic
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RP4: test for sulfate ions

1. acidify solution with HNO3/HCl

(to remove CO2, which if present forms a false positive)


2. add BaCl2

white precipitate if present (as BaSO4 is formed)
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RP4: test for carbonate ions

1. add HCl
2. collect gas formed
3. bubble through limewater (Ca(OH)2)

cloudy, white solution if present
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RP4: test for ammonium ions

1. add equal amount of NaOH
2. heat the sample
3. place moist, red litmus over the mouth of tube as gas give off is ammonia

will turn blue

NH4+ + OH- → NH3 + H2O
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RP4: test for hydroxide ions
turns red litmus paper blue

either by dipping it in NaOH solution

or

place moist, red litmus paper in petri dish with ammonia solution (on filter paper) on the other side, ammonia vapour will turn paper blue
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group 7: boiling point
increases down the group

atomic radius increases increasing the strength of the van der Waal forces, greater Mr has greater surface for them to act on

I = solid, Br = liquid, Cl = gas
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group 7: reducing agent
must be oxidised itself, lose its own electron

halide ions (I-) are best, most likely to lose electrons

increase down the group

atomic radius and shielding increases, reduces nuclear attraction
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group 7: oxidising agent
must be reduced itself, gain electrons

halogen (F2) molecules are best, most likely to gain electrons

decrease down the group

atomic radius and shielding increases, reduces nuclear attraction less attraction to electrons
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group 7: electronegativity
decrease down the group

atomic radius and shielding increases, reduces nuclear attraction less attraction to electrons
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group 7: displacement reactions
Cl displaces Br = orange/red colour

Cl/Br displaces I = black precipitate
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RP4: test for halide ion

1. add HNO3

prevents false positive from carbonate ions


2. add AgNO3

Ag+(aq) + X- → AgX(s)

AgCl = white precipitate

AgBr = cream precipitate

AgI = yellow precipitate
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RP4: test for silver halide ions solubility

1. add dilute ammonia

AgCl will dissolve, precipitate disappears

AgCl + NH3 → \[Ag(NH3)2\]+ + Cl-


2. add concentrated ammonia

AgBr will dissolve

AgI is insoluble
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sodium **chloride** + sulfuric acid
H2SO4(l) + NaCl(s) → HCl(g) + NaHSO4(s)

HCl = white misty fumes
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sodium **bromide** + sulfuric acid
H2SO4(l) + NaBr(s) → HBr(g) + NaHSO4(s)

2HBr(g) + H2SO4(l) → Br2(g) + SO2(g) + 2H2O(l)

HBr = white misty fumes

Br2 = reddish-brown gas

SO2 = choking fumes?

Br- acting as reducing agent
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sodium **iodide** + sulfuric acid
H2SO4(l) + NaI(s) → HI(g) + NaHSO4(s)

2HI(g) + H2SO4(l) → I2(s) + **SO2(g)** + 2H2O(l)

6HI(g) + H2SO4(l) → 3I2(s) + **S(s)** + 4H2O(l)

8HI(g) + H2SO4(l) → 4I2(s) + **H2S(g)** + 4H2O(l)

HI = white misty fumes

I2 = violet/purple vapour

SO2 = choking fumes?

S = yellow solid

H2S = toxic, bad smelling gas
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chlorine with cold, dilute NaOH(aq)
Cl2(aq) + 2NaOH(aq) → NaCl(aq) + NaClO(aq) + H2O(l)

NaClO =household bleach that can kill bacteria, water treatment, and bleach paper and textiles
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chlorine and water (**no sunlight**)
Cl2 + H2O → ClO- + Cl- + 2H+

Cl2 + H2O → HClO + HCl

disproportionation reaction
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chlorine and water (**sunlight**)
2Cl2 + 2H2O → 4HCl + O2

UV light is present
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chlorine: water treatment
(+) kills disease-causing microorganisms

prevents algae growth

(-) chlorine is harmful, respiratory and carcinogenic

organic chlorine compounds harmful to environment

(=) however, benefits outweigh the negatives
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sodium with water
reacts vigorously with cold water, fizzing H2 gas and forms a ball

producing strong alkaline solution

2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)

sodium is more reactive than magnesium as it requires less energy to remove 1 electron than 2, more energy need to form Mg2+ ions
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period 3: oxides
Na2O, MgO, Al2O3, SiO2, P4O10, SO2 and SO3
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sodium with oxygen

magnesium with oxygen

aluminium with oxygen
2Na + 1/2O2 → Na2O(s)

Mg + 1/2O2 → MgO(s)

2Al + 1 1/2O2 → Al2O3(s)
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silicon with oxygen

phosphorus and oxygen

sulfur and oxygen
Si + O2 → SiO2(s)

P4 + 5O2 → P4O10(s)

S + O2 → SO2(g)
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sodium, magnesium, and aluminium oxide structure and bonding
giant ionic lattice

strong attractive forces between ions

MgO > Na2O (mp) Mg2+ attract the O2- more greatly

MgO > Al2O3 (mp) Al3+ distorts electron cloud of O2- and some covalent character lowers E
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silicon oxide structure and bonding
macromolecular

many strong covalent bonds to overcome, requiring lots of energy
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phosphorus and sulfur oxide structure and bonding
simple molecular

weak intermolecular bonds, require little E to overcome
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ionic oxides with water (not aluminium)
Na2O + H2O → 2NaOH

MgO + H2O → Mg(OH)2

alkaline solutions pH 12-14 and 9-10 respectively
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simple covalent oxides with water
eg. P4O10 + 6H2O → 4H3PO4(aq)

acidic solutions, pH 0-2
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silicon and aluminium oxide with water
SiO2 = insoluble

Al2O3 = amphoteric, insoluble
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amphoteric
will react with acid and base to form a salt

(acting as either acid/base)
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basic oxides
eg. 2HCl + MgO → MgCl2 + H2O

react with acid to form salt

MgO and Na2O
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acidic oxides
eg. 12NaOH + P4O10 → 4Na3PO4 + 6H2O

react with base to form salt

SiO2, P4O10, SO2 and SO3
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amphoteric oxides
acting as acid with base

2NaOH + Al2O3 + 3H2O → **2NaAl(OH)4**

sodium tetrahydroxoaluminate

acting as base with acid

Al2O3 + 3H2SO4 → **Al2(SO4)3** + 3H2O
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transition metal
elements with incomplete d-subshell that can form at least one stable ion with an incomplete d-subshell
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complex
central metal atom/ion surrounded by ligands
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ligand
molecule/ion that forms a co-ordinate bond with a transition metal by donating a pair of electrons
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co-ordination number
number of co-ordinate bonds to the central metal atom/ion
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chelate effect
bidentate/multidentate more energetically favourable than monodentate

because entropy change is always positive as there is a net increase in the number of particles

so ΔG = negative
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cis-trans isomerism
occurs in square planar and octahedral complexes

cis = 90°

trans = 180°

could cause change in properties due to polarity
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optical isomerism
occurs in tetrahedral and octahedral complexes

a mirror image of the 2 complexes that are not superimposable
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cis-platin
Pt bound to 2 Cl- ions and 2 NH3 molecules, ligands 90° from the same molecule, square planar structure

binds to DNA, prevents cell division, and tumour growth
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Tollens’ reagent
Ag+ forms linear complex

\[Ag(NH3)2\]+

reduced to metallic silver to distinguish aldehydes and ketones
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formation of coloured ions
colour arises when some wavelengths of visible light are absorbed and the remaining wavelengths are transmitted

d-electrons move from ground state to an excited state when light is absorbed
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energy of visible light equation
**E = hf = hc/λ**

E = energy of visible light (J)

h = Planck’s constant (6.63 x10-34)

f = frequency of light (Hz or s-1)

c = speed of light (3 x8 ms-1)

λ = wavelength of light (m)
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factors affecting change in energy, electron promotion
type of ligand

co-ordination number

oxidation state (change in)
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colour: type of ligand
different ligands will split d-orbital by different energies

depends on the repulsion of each ligand
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colour: co-ordination number
influences strength of metal ion-ligand interaction
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colour: oxidation state
strength of metal-ion ligand interactions varies due to nuclear charge

eg. Mn(II) = Fe(III) electron configuration

however, Fe(III) has greater nuclear charge so its change in energy is greater, absorbing visible light of greater energy.
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heterogenous catalyst
catalyst in a different state/phase from the reactants

reactions occur at the active site on the surface
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homogenous catalyst
catalyst in the same state/phase as the reactants

reactions proceed through an intermediate species
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contact process
2SO2 + O2 → 2SO3

catalyst = V2O5

**2SO2 + 2V2O5 → 2V2O4 + 2SO3**

**O2 + 2V2O4 → 2V2O5**
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Haber process
**N2 + 3H2 → 2NH3**

catalyst = Fe

N2 and H2 diffuse to Fe surface and adsorb to it, weakening their covalent bonds so reaction can take place

product molecule desorbs from surface
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I- and S2O82-
overall = S2O82- + 2I- → I2 + 2SO42-

catalyst = Fe2+ ions


1. **S2O82- + 2Fe2+ → 2SO42- + 2Fe3+**
2. **2I- + 2Fe3+ → I2 + 2Fe2+**

Fe acts as reducing and oxidising agent

both negative ions so require catalyst as they repel and activation energy is high
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autocatalysis example
overall:

**2MnO4- + 5C2O42- + 16H+ → 2Mn2+ + 10CO2 + 8H2O**

catalyst = Mn2+
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how Mn2+ acts as catalyst
**4Mn2+** + MnO4- + 8H+ → **5Mn3+** + 4H2O

**2Mn3+** + C2O42- → **2Mn2+** + 2CO2
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Lewis acid and base
Lewis acid = accept lone pair

Lewis base = donate lone pair

metal complex = central metal ion = acid, ligand = base
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acidity of M3+ and M2+
acidity of M3+ is greater than M2+

M3+ is smaller with a higher charge density than M2+, more strongly polarising

M3+ pulls the O’s electrons of H2O more greatly, weakening the O-H bond, H+ is more easily released
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acidity equations
eg. \[Fe(H2O)6\]3+ ⇌ \[Fe(H2O)5OH\]2+ + H+
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amphoteric character in complexes
aluminium hydroxides dissolve in acids and base

in acid = \[Al(H2O)6\]3+

in neutral = Al(H2O)3(OH)3

in base = **[Al(OH)4]-**