Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

Chemistry > Module 3 > Flashcards

Module 3 Flashcards

You may prefer our related Brainscape-certified flashcards:
Periodic Table flashcards Chemistry 101 flashcards AP® Chemistry flashcards Organic Chemistry 101 flashcards
1
Q

Trend in First Ionisation Energy down Group

A
  • atomic radius increases
  • more electron shielding
  • nuclear attraction decreases
  • outweighs increasing nuclear charge
    = first ionisation energy decreases
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

Trend in First Ionisation Energy across Period

A
  • 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)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

Trend in Melting Point across a Period

A

Giant Metallic Structure -> Giant Covalent Structure -> Simple Molecular Structure
Increases G1-14(4)
Decreases from G5-18(0)
GCS= B, C, Si

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

Group 2 Reactions

A
  • 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
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

Group 2 Reaction with oxygen

A

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)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

Group 2 Reaction with water

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

Group 2 Reaction with acids

A

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)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

Trend in reactivity down Group 2

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

Group 2 Hydroxides

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

Uses of group 2 compounds

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

Physical Properties of Halogens

A

Exist as diatomic molecules at RTP

Form simple molecular lattice structure in solid state

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

Trend in Boiling Point down Group 7

A

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 - ?
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

Trend in Reactivity down Group 7

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

Halogen Reactions

A
  • 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-
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

Cl2 Reaction with Water

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

Cl2 Reaction with cold dilute NaOH(aq)

A

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

17
Q

Chlorine uses and risks

A

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

18
Q

Carbonate test

A

E.g. Na2CO3(aq) + 2HNO3(aq) -> 2NaNO3(aq) + CO2(g) + H2O(l)

  1. Add dilute nitric acid to sample (sulfuric acid and hydrochloric acid interfere with results)
  2. Effervescence - carbonate ions may be present
  3. Test gas product is CO2 by bubbling through limewater (turns cloudy/milky)
    CO2(g) + Ca(OH)2 -> CaCO3(s) + H2O(l)
19
Q

Sulphate test

A

Sulphate ions react with barium ions to form a insoluble white BaSO4 precipitate
Ba+2(aq) + SO4-2(aq) -> BaSO4(s)

  1. Acidify the test solution by adding nitric acid (Sulfuric acid cannot be used because it contains sulfate ions - interferes with test)
  2. Add Ba(NO3)2(aq) / BaCl2(aq) (do not use BaCl2 if testing for halide ions)
  3. Observe if white precipitate is formed

Reaction used in Barium Meal

20
Q

Halide test

A

Halide ions react with silver ions to form silver halide precipitates
Ag+(aq) + X-(aq) -> AgX(s)

  1. Acidify the halide solution (aq) using a few drops of dilute nitric acid
  2. Add AgNO3(aq) to the sample
  3. 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
  4. Add dilute NH3 (aq) If dissolves - AgCl
    Add conc. NH3 (aq) If dissolves - AgBr If insoluble - AgI
21
Q

Sequence of Anion tests

A
  1. Carbonate test
    Neither sulphate / halide produce bubbles with dilute acid - no possibility of incorrect conclusion
    If no bubbles produced then proceed
  2. Sulphate test
    Adding Ba+ to a solution containing carbonate ions produces BaCO3
    BaCO3 is also a white precipitate - false positive result
  3. Halide test
    Adding Ag+ to a solution containing carbonate/sulphate ions produces Ag2CO3/Ag2SO4
    Ag2CO3/Ag2SO4 are also white precipitates - false positive result
22
Q

Ammonium ion test

A

Ammonium ions react with hydroxide ions to form ammonia(g) when heated
NH4+(aq) + OH-(aq) -> NH3(g) + H2O(l)

  1. Add NaOH(aq) to ammonium ion solution
  2. Warm mixture and ammonia gas is produced
  3. Test gas using damp red litmus paper - turns blue if positive (alkaline)
23
Q

Enthalpy Change

A

Enthalpy: heat energy stored in a chemical system

ΔH = H(products) - H(reactants)

24
Q

Exothermic Reaction

A

Energy transferred from the chemical system to the surroundings

  • releases heat energy (surrounding temp ↑)
  • ΔH is negative
25
Q

Endothermic Reaction

A

Energy transferred to the system from the surroundings

  • absorbs heat energy (surrounding temp ↓)
  • ΔH is positive
26
Q

Standard enthalpy change of formation

A

The enthalpy change that occurs when 1 mole of a compound is formed from its constituent elements under standard conditions and states (298K, 100kPa)

27
Q

Standard enthalpy change of combustion

A

The enthalpy change that takes place when one mole of a substance reacts undergoes complete combustion under standard conditions and states (298K, 100kPa)

28
Q

Standard enthalpy change of neutralisation

A

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)

29
Q

Average bond enthalpy

A

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)

30
Q

Hess’ Law (Indirect Enthalpy Change)

A

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
31
Q

Homogenous catalysts

A

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

32
Q

Heterogenous catalysts

A

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

33
Q

Uses of catalysts

A

Reduces temperature needed for the process + energy requirements

  • less electricity / fossil fuels used (fewer pollutants)
  • cut costs
  • increase sustainability + profitability
34
Q

le Chatelier’s Principle

A

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

35
Q

Halogen-Halide Displacement Reaction

A

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)

Chemistry (16 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions