Inorganic Chemistry: Paper 1

periodicity

repeating pattern/trends (of physical or chemical properties)

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

period 3: ionisation energy

increases along the period

decreased atomic radius and increased nuclear charge, more energy required to remove electrons

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

period 3: melting point (Si)

macromolecular

highest melting point because must overcome strong covalent bonds, requiring lots of energy.

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

period 3: melting point (Ar)

simple molecule

exist as single atoms so weak van der Waals and small Mr

group 2: atomic radius

increase down the group

add new shell, increasing distance to outer electron, increased shielding, reducing nuclear attraction

group 2: ionisation energy

decreases down the group

new shells, atomic radius increases, increased shielding, reduces nuclear attraction, less energy to remove an electron

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

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

Mg with steam

Mg(s) + H2O(g) → MgO(s) + H2(g)

Mg in extraction of titanium

TiCl4 + 2Mg → 2MgCl2 + Ti

displacement reaction

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

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

group 2: solubility of hydroxides

increases in solubility down the group

Mg(OH)2 is sparingly soluble

use of Mg(OH)2

milk of magnesia is used as an antacid as it neutralises stomach acid, for indigestion

use of Ca(OH)2

slaked lime is used in agriculture to raise the pH of a field (is basic)

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

group 2: solubility of sulfates

decrease in solubility down the group

BaSO4 is insoluble

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

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)

RP4: test for carbonate ions

1. add HCl

2. collect gas formed

3. bubble through limewater (Ca(OH)2)

cloudy, white solution if present

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

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

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

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

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

group 7: electronegativity

decrease down the group

atomic radius and shielding increases, reduces nuclear attraction less attraction to electrons

group 7: displacement reactions

Cl displaces Br = orange/red colour

Cl/Br displaces I = black precipitate

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

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

sodium **chloride** + sulfuric acid

H2SO4(l) + NaCl(s) → HCl(g) + NaHSO4(s)

HCl = white misty fumes

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

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

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

chlorine and water (**no sunlight**)

Cl2 + H2O → ClO- + Cl- + 2H+

Cl2 + H2O → HClO + HCl

disproportionation reaction

chlorine and water (**sunlight**)

2Cl2 + 2H2O → 4HCl + O2

UV light is present

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

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

period 3: oxides

Na2O, MgO, Al2O3, SiO2, P4O10, SO2 and SO3

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)

silicon with oxygen

phosphorus and oxygen

sulfur and oxygen

Si + O2 → SiO2(s)

P4 + 5O2 → P4O10(s)

S + O2 → SO2(g)

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

silicon oxide structure and bonding

macromolecular

many strong covalent bonds to overcome, requiring lots of energy

phosphorus and sulfur oxide structure and bonding

simple molecular

weak intermolecular bonds, require little E to overcome

ionic oxides with water (not aluminium)

Na2O + H2O → 2NaOH

MgO + H2O → Mg(OH)2

alkaline solutions pH 12-14 and 9-10 respectively

simple covalent oxides with water

eg. P4O10 + 6H2O → 4H3PO4(aq)

acidic solutions, pH 0-2

silicon and aluminium oxide with water

SiO2 = insoluble

Al2O3 = amphoteric, insoluble

amphoteric

will react with acid and base to form a salt

(acting as either acid/base)

basic oxides

eg. 2HCl + MgO → MgCl2 + H2O

react with acid to form salt

MgO and Na2O

acidic oxides

eg. 12NaOH + P4O10 → 4Na3PO4 + 6H2O

react with base to form salt

SiO2, P4O10, SO2 and SO3

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

transition metal

elements with incomplete d-subshell that can form at least one stable ion with an incomplete d-subshell

complex

central metal atom/ion surrounded by ligands

ligand

molecule/ion that forms a co-ordinate bond with a transition metal by donating a pair of electrons

co-ordination number

number of co-ordinate bonds to the central metal atom/ion

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

cis-trans isomerism

occurs in square planar and octahedral complexes

cis = 90°

trans = 180°

could cause change in properties due to polarity

optical isomerism

occurs in tetrahedral and octahedral complexes

a mirror image of the 2 complexes that are not superimposable

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

Tollens’ reagent

Ag+ forms linear complex

\[Ag(NH3)2\]+

reduced to metallic silver to distinguish aldehydes and ketones

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

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)

factors affecting change in energy, electron promotion

type of ligand

co-ordination number

oxidation state (change in)

colour: type of ligand

different ligands will split d-orbital by different energies

depends on the repulsion of each ligand

colour: co-ordination number

influences strength of metal ion-ligand interaction

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.

heterogenous catalyst

catalyst in a different state/phase from the reactants

reactions occur at the active site on the surface

homogenous catalyst

catalyst in the same state/phase as the reactants

reactions proceed through an intermediate species

contact process

2SO2 + O2 → 2SO3

catalyst = V2O5

**2SO2 + 2V2O5 → 2V2O4 + 2SO3**

**O2 + 2V2O4 → 2V2O5**

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

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

autocatalysis example

overall:

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

catalyst = Mn2+

how Mn2+ acts as catalyst

**4Mn2+** + MnO4- + 8H+ → **5Mn3+** + 4H2O

**2Mn3+** + C2O42- → **2Mn2+** + 2CO2

Lewis acid and base

Lewis acid = accept lone pair

Lewis base = donate lone pair

metal complex = central metal ion = acid, ligand = base

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

acidity equations

eg. \[Fe(H2O)6\]3+ ⇌ \[Fe(H2O)5OH\]2+ + H+

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]-**