Module 3 Flashcards
Trend in First Ionisation Energy down Group
- atomic radius increases
- more electron shielding
- nuclear attraction decreases
- outweighs increasing nuclear charge
= first ionisation energy decreases
Trend in First Ionisation Energy across Period
- nuclear charge increases
- similar electron shielding
- nuclear attraction increases
- atomic radius decreases
= first ionisation energy increases
Exceptions:
Filling from 2s to 2p sub shell - new highest sub shell energy (easier to remove)
Start of electron pairing in p orbitals of 2p sub shell - increases repulsion (easier to remove)
Trend in Melting Point across a Period
Giant Metallic Structure -> Giant Covalent Structure -> Simple Molecular Structure
Increases G1-14(4)
Decreases from G5-18(0)
GCS= B, C, Si
Group 2 Reactions
- Reducing Agents
2 electrons on the outer s-subshell are lost and oxidised to form a 2+ ion
Another species will gain 2 electrons to be reduced
Group 2 Reaction with oxygen
Metals burn with a brilliant white light and form white metal oxide
2M(s) + O2(g) -> 2MO(s)
- 2M is oxidised 0 -> +2 (+4)
- 2O is reduced 0 -> -2 (-4)
Group 2 Reaction with water
Water and Mg react slowly - more vigorous reaction as you go down the group
- Bubbles of hydrogen gas produced and white precipitate of M(OH)2 formed
- Precipitate is more soluble in water down the group - less precipitate obtained
M(s) + 2H2O -> M(OH)2 + H2(g)
- 1M is oxidised 0 -> +2 (+2)
- 2H is reduced 0 -> -1 (-2)
- 2H are unchanged +1 -> +1
Group 2 Reaction with acids
Metals react with dilute acids to form a salt and hydrogen gas
M(s) + Acid e.g. 2HCl(aq) -> MCl2(aq) + H2(g)
-1M is oxidised 0 -> +2 (+2)
- 2H is reduced +1 -> 0 (-2)
Trend in reactivity down Group 2
Reactivity increases down Group 2
Ionisation energy decreases
Electronegativity decreases
- atomic radius increases
- electron shielding increases
- less attraction between nucleus and outer shell electrons
- outweighs the increasing nuclear charge
It is easier to an electron from the outer shell
It requires a lower total energy input
It loses tendency to attract a bonding pair of electrons
The metals are stronger reducing agents as you go down Group 2
Group 2 Hydroxides
Group 2 oxides react with water to produce OH- ions and M(OH)2 in an alkaline solution
MO(s) + H2O(l) -> M+2(aq) + OH-(aq)
- Forms precipitate once saturated
M+2(aq) + 2OH-(aq) -> M(OH)2(s)
Solubility of Group 2 hydroxides in water increases down the group
Resulting solutions contain more OH-(aq) ions and are more alkaline
-pH increases down the group
-Mg(OH)2(s) very slightly soluble
Low OH- concentration and pH around 10
-Ba(OH)2 soluble in water
Greater OH- concentration and pH around 13
Uses of group 2 compounds
Used to neutralise acids
- Mg(OH)2(s)+ CaCO3(aq) -> MgCl2(aq) + 2H2O(l)
As antacids to treat indigestion
- Ca(OH)2(s) + 2H+(aq) -> Ca+(aq) + 2H2O(l)
As lime to neutralise acidic soil and form water to increase the pH
Physical Properties of Halogens
Exist as diatomic molecules at RTP
Form simple molecular lattice structure in solid state
Trend in Boiling Point down Group 7
Boiling point increases down the group
- More electrons
- Stronger London forces
- More energy is required to overcome the intermolecular forces
State at RTP changes from gas -> liquid -> solid F2 (g)- pale yellow Cl2 (g) - pale green Br2 (l) - red/brown I2 (s) - shiny grey/black At2 - ?
Trend in Reactivity down Group 7
Reactivity decreases down group 7
- Atomic radius increases
- Electron shielding increases
- Less attraction between the nucleus and outer shell electrons
- The tendency to gain electrons decreases as it is harder to attract an electron from another species
Weaker oxidising agents down the group
Halogen Reactions
- Oxidising agent
7 outer shell electrons - 2 in the outer s-subshell and 5 in the outer p-subshell
1 electron on the outer p-subshell is gained in each halogen atom and reduced to form a 1- ion
Another species will lose 2 electrons to be oxidised
Cl2 + 2e- -> 2Cl-
Cl2 Reaction with Water
Disproportionation reaction
Cl2 (aq) + H2O (l) -> HClO(aq) + HCl (aq)
- 1Cl is oxidised 0 in Cl2 -> +1 in HClO
- 1Cl is reduced 0 in Cl2 -> -1 in HCl
Forms solution with chloric acid (HClO) and chlorate ions (ClO-) - kills bacteria in water treatment
Indicator turns red then bleaches in Cl2 + Water
Cl2 Reaction with cold dilute NaOH(aq)
Disproportionation reaction
Cl2 (aq) + 2NaOH(aq) -> NaClO(aq) + NaCl(aq) + H2O(l)
- 1Cl is oxidised 0 in Cl2 -> +1 in NaClO
- 1Cl is reduced 0 in Cl2 -> -1 in NaCl
Forms solution with large conc. of chlorate ions (ClO-) - bleach
Chlorine uses and risks
Uses
- water treatment
Kills bacteria - protects against diseases e.g. typhoid, cholera
Makes water potable
Risks
- Can react with organic hydrocarbons in drinking water to produce chlorinated hydrocarbons
May cause cancer
- extremely toxic gas
- respiratory irritant in small concentrations - can be fatal in large concentrations
Carbonate test
E.g. Na2CO3(aq) + 2HNO3(aq) -> 2NaNO3(aq) + CO2(g) + H2O(l)
- Add dilute nitric acid to sample (sulfuric acid and hydrochloric acid interfere with results)
- Effervescence - carbonate ions may be present
- Test gas product is CO2 by bubbling through limewater (turns cloudy/milky)
CO2(g) + Ca(OH)2 -> CaCO3(s) + H2O(l)
Sulphate test
Sulphate ions react with barium ions to form a insoluble white BaSO4 precipitate
Ba+2(aq) + SO4-2(aq) -> BaSO4(s)
- Acidify the test solution by adding nitric acid (Sulfuric acid cannot be used because it contains sulfate ions - interferes with test)
- Add Ba(NO3)2(aq) / BaCl2(aq) (do not use BaCl2 if testing for halide ions)
- Observe if white precipitate is formed
Reaction used in Barium Meal
Halide test
Halide ions react with silver ions to form silver halide precipitates
Ag+(aq) + X-(aq) -> AgX(s)
- Acidify the halide solution (aq) using a few drops of dilute nitric acid
- Add AgNO3(aq) to the sample
- Different coloured silver halide precipitates form, depending on the halide ions present
Ag+(aq) + Cl-(aq) -> AgCl(s) - white
Ag+(aq) + Br-(aq) -> AgBr(s) - cream
Ag+(aq) + I-(aq) -> AgI(s) - yellow - Add dilute NH3 (aq) If dissolves - AgCl
Add conc. NH3 (aq) If dissolves - AgBr If insoluble - AgI
Sequence of Anion tests
- Carbonate test
Neither sulphate / halide produce bubbles with dilute acid - no possibility of incorrect conclusion
If no bubbles produced then proceed - Sulphate test
Adding Ba+ to a solution containing carbonate ions produces BaCO3
BaCO3 is also a white precipitate - false positive result - Halide test
Adding Ag+ to a solution containing carbonate/sulphate ions produces Ag2CO3/Ag2SO4
Ag2CO3/Ag2SO4 are also white precipitates - false positive result
Ammonium ion test
Ammonium ions react with hydroxide ions to form ammonia(g) when heated
NH4+(aq) + OH-(aq) -> NH3(g) + H2O(l)
- Add NaOH(aq) to ammonium ion solution
- Warm mixture and ammonia gas is produced
- Test gas using damp red litmus paper - turns blue if positive (alkaline)
Enthalpy Change
Enthalpy: heat energy stored in a chemical system
ΔH = H(products) - H(reactants)
Exothermic Reaction
Energy transferred from the chemical system to the surroundings
- releases heat energy (surrounding temp ↑)
- ΔH is negative
Endothermic Reaction
Energy transferred to the system from the surroundings
- absorbs heat energy (surrounding temp ↓)
- ΔH is positive
Standard enthalpy change of formation
The enthalpy change that occurs when 1 mole of a compound is formed from its constituent elements under standard conditions and states (298K, 100kPa)
Standard enthalpy change of combustion
The enthalpy change that takes place when one mole of a substance reacts undergoes complete combustion under standard conditions and states (298K, 100kPa)
Standard enthalpy change of neutralisation
The enthalpy change that accompanies the reaction of an acid by a base to form one mole of H2O(l) under standard conditions and states (298K, 100kPa)
Average bond enthalpy
Energy required to break one mole of a specified type of bond in a gaseous molecule
- endothermic (+ΔH)
Limitations:
- only average value
- actual energy value varies in different chemical environments
- energy for change in state not accounted for
ΔrH = Σ(reactants bond enthalpy) - Σ(products bond enthalpy)
Hess’ Law (Indirect Enthalpy Change)
If a reaction can take place by more than one route and the initial + final concentrations are the same - total enthalpy change same for each route
ΔfH = C - B ΔcH = B - C
Homogenous catalysts
Catalyst in the same physical state as reactants
- reacts to form an intermediate (breaks down to produce product + regenerate catalyst)
E.g. Cl• in O3 depletion
Heterogenous catalysts
Catalyst with a different physical state as reactants
- reactant molecules are adsorbed onto catalyst surface then leave by desorption
E.g. Fe(s) in Haber process
Uses of catalysts
Reduces temperature needed for the process + energy requirements
- less electricity / fossil fuels used (fewer pollutants)
- cut costs
- increase sustainability + profitability
le Chatelier’s Principle
When a system in dynamic equilibrium is subjected to an external change, the system readjusts itself to minimise the effect of the change and to restore equilibrium
- rate of forward reaction = rate of reverse reaction
- concentrations of reactants/products remain constant
aA + bB ⇌ cC + dD
Kc = [C]^c [D]^d / [A]^a [B]^b
Halogen-Halide Displacement Reaction
A more reactive halogen can displace a less reactive halide ion from its solution - causes a colour change
Cl2 reacts with Br- and I-
Br2 reacts with I
I2 does not react
Cl2 = pale green Br2 = orange I2 = brown (water) violet (cyclohexane)
E.g. Cl2 + 2Br- -> 2Cl- + Br2
- 2Br is oxidised -1 -> 0 (+2)
- 2Cl is reduced 0 -> -1 (-2)
