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Unit 3 Questions Flashcards

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1
Q

Give the equation for the oxidation of iron with permanganate

A

5Fe(2+) + MnO4(-) + 8H(+) –> 5Fe(3+) + Mn(2+) + 4H2O

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2
Q

What is the usual oxidation state of NO3?

A

+1 overall

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3
Q

What happens at the cathode of an electrochemical cell?

A

Positive cations gain electrons - reduction

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4
Q

What happens at the anode of an electrochemical cell?

A

Negative anions lose electrons - oxidation

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5
Q

Draw a diagram of an electrochemical cell with Mg(2+) and Cu(2+)

A

See diagram A

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6
Q

Write out the half cell diagram thing for Mg and Cu

A

Mg(s) | Mg(2+)(aq) :: Cu(2+)(aq) | Cu(s)

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7
Q

What is the purpose of the salt bridge in an electrochemical cell?

A

To allow ions to flow from one solution to the other, completing the circuit, without the solutions mixing

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8
Q

Why is the voltmeter in an electrochemical cell high resistance?

A

To minimise the current lost as heat

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9
Q

Which half cell will be the positive electrode?

A

The one with the most positive E(theta) value

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10
Q

Which half cell goes on the left?

A

The most negative E(theta), oxidation

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11
Q

Give the definition of the standard electrode potential (E(theta))

A

The potential difference when any half-cell is connected to the standard hydrogen electrode under standard conditions.

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12
Q

Draw the standard hydrogen electrode

A

See diagram B

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13
Q

Draw a metal-solution half cell with Zinc

A

See diagram C

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14
Q

Draw a mixed ion half cell with Fe

A

See diagram D

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15
Q

Why is the standard electrode potential of hydrogen 0?

A

Because it is taken as the standard and all the other E(theta) values are measured with respect to it.

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16
Q

Give the chemical definitions of EMF

A

The potential difference across a cell when it takes no current, and as such the maximum amount of energy which can be given by the cell

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17
Q

Give the colours of Cu(2+)(aq) + 2e(-) <–> Cu(s)

A

Blue <–> copper colour

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18
Q

Give the colours of Fe(3+)(aq) + e(-) <–> Fe(2+)(aq)

A

Yellow <–> pale green

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19
Q

Give the colours of Br2(aq) + 2e(-) <–> 2Br(-)(aq)

A

Orange <–> colourless

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20
Q

Give the colours of Cr2O7(2-)(aq) + 14H(+)(aq) +6e(-) <–> 2Cr(3+)(aq) + 7H2)(l)

A

Orange <–> dark green

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21
Q

MnO4(-)(aq) + 8H(+)(aq) + 5e(-) <–> Mn(2+)(aq) + 4H2O(l)

A

Purple <–> pale pink/colourless apparantely

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22
Q

You know what?

A

A w—–, etc :)

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23
Q

How do fuel cells generate power?

A

Use electrochemical methods to get energy from fuels, typically hydrogen

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24
Q

What happens at the anode of a fuel cell?

A

Hydrogen is oxidised to H+ ions

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25
Q

What happens at the cathode of a fuel cell?

A

Oxygen gas is reduced to water

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26
Q

What catalyses the reactions in a fuel cell?

A

A platinum electrode

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27
Q

What is the purpose of the polymer/electrolyte membrane in the hydrogen fuel cell?

A

It allows only positive ions into the cathode, electrons have to go via an external circuit

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28
Q

State two different ways to connect hydrogen fuel cells

A

In series for higher voltage or parallel for higher current density

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29
Q

Give two advantages of hydrogen fuel cells

A
  1. Clean, only waste product is water
  2. High efficiency
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30
Q

Give two disadvantages of hydrogen fuel cells

A
  1. Expensive (platinum)
  2. Storage difficulties - hydrogen is flammable and explosive
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31
Q

I have something for you!

A

~ # ~
It’s a waffle! You’ve earned it :)

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32
Q

What is an amphoteric substance?

A

A substance that acts as an acid in basic conditions and vice-versa

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33
Q

I am looking for some p-block elements. Where on the periodic table might I find them?

A

You must seek the crossover from metal to non-metal and ionic to covalent. There you will find what you seek.

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34
Q

Is magnesium amphoteric? Give its equation(s) for the reaction with hydroxide ions and the observations

A

It is not amphoteric.
Dropwise and excess: Mg(2+)(aq) + 2OH(-)(aq) –> Mg(OH)2(s)
This forms a white precipitate

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35
Q

Is zinc amphoteric? Give its equations(s) for the reaction with hydroxide ions and the observations

A

It is amphoteric.
Dropwise: Zn(2+)(aq) + 2OH(-)(aq) –> Zn(OH)2(s)
Excess: Zn(OH)2(s) + 2OH(-)(aq) –> Zn(OH)4(aq)
Dropwise forms white precipitate, excess forms colourless solution

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36
Q

What is the highest oxidation state p-block elements can reach?

A

Their group number

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37
Q

What is an inert pair?

A

An ns2 pair of electrons not involved in bonding

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38
Q

What effect does the inert pair effect have on the oxidation state and how does this change down the group?

A

Elements that don’t have access to the inert pair show a state 2 lower than usual.
This lower oxidation state becomes more stable as you descend the group

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39
Q

Explain octet expansion and its effect on the oxidation state

A

Compounds in groups V and VI need 5 or 6 covalent bonds to reach their maximum oxidation state. This is not problem for period 3 onwards as they have access to their d orbitals, but those before are limited: N to 2 and O to 2 bonds.

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40
Q

What is a dimer?

A

A species created when two molecules join together

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41
Q

What kinds of compound are boron and aluminium often found in?

A

Electron deficiency compounds (BF3 or AlCl3) due to each having 2 valence electrons and therefore three covalent bonds (incomplete octet)

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42
Q

What kind of species do B and Al often react with?

A

Lone pair species (removes electron deficiency with dative bond)

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43
Q

Draw the bonding of two AlCl3 molecules

A

See diagram E

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44
Q

At what temperature does the dimer Al2Cl6 dissolve back into AlCl3?

A

Dative bonds break above 200 degrees C

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45
Q

Name one use of the dimer Al2Cl6

A

As a Friedel-Crafts catalyst for chlorination or alkylation of benzene

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46
Q

Draw the mechanism for the reaction of CH3Cl and AlCl3. What kind of reaction is this?

A

See diagram F
It is an electrophilic substitution reaction

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47
Q

Why does BF3 form donor-acceptor compounds?

A

Because the B is electron deficient and will try to react with compounds with lone pairs, accepting a donated electron

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48
Q

Draw an example of a donor-acceptor reaction

A

See diagram G

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49
Q

What are ionic liquids

A

Organic salts with melting points below 100 degrees. They have a wide temperature range at which they are liquid - unusually so for a salt

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50
Q

How are ionic liquids formed?

A

Large organic chlorocompounds, R-Cl, react with AlCl3. Forms dative bond to AlCl3 followed by ionisation to R(+)AlCl4(-)

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51
Q

Give two uses and two advantages of ionic liquids

A
  1. Solvents 2. Catalysts
  2. They are recyclable 2. Organic products are imiscible in liquids, which means one can separate them easily
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52
Q

What is BN isoelectric to and what does this mean?

A

Carbon - has the same electronic configuration and therefore also similar properties

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53
Q

Draw hexagonal boron nitride

A

See diagram H

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54
Q

Describe the structure of hexagonal boron nitride

A

Same as graphite basically but atoms lie directly above one another with no delocalised electrons

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55
Q

What are the three main useful properties of hexagonal boron nitride?

A
  1. Excellent lubricating properties - weak VdWs between layers
  2. Insulator because of the lack of delocalised electrons (N is more electronegative than B, so the bond is polar)
  3. Can be bent so the edges overlap to form a nanotube. Pack with carbon fullerene and expose to intense electron beam for a semiconductor
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56
Q

Describe the structure of tetrahedral boron nitride

A

Tetrahedral - basically like diamond

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57
Q

Explain two uses of cubic boron nitride.
Also two further properties if you feel like it

A
  1. Cutting tools/wear resistant coating: second hardest known material, behind diamond.
  2. Support for catalysts because of its large surface area in powder form

It also has high thermal conductivity and is chemically inert

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58
Q

Name two trends in group IV

A

Increasing metallic character and stability of the +2 state down the group

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59
Q

Describe how the bonding and structure changes as we descend group IV

A

Top elements (C and Si) are non-metals with giant covalent structures
Bottom elements (Sn and Pb) have metallic bonding (lattice of metal ions in sea of delocalised electrons, etc)

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60
Q

Describe the oxidation states, appearance, bonding, redox character, and use of carbon monoxide, giving equations where needed for the latter

A
  • Carbon is stable as +VI and only exists as +II in CO
  • Colourless gas
  • Covalent
  • Acts as a reducing agent as it tries to reach the more stable +IV
  • Used to extract metals from oxides
    Fe2O3(s) + 3CO(g) –> 2Fe(l) + 3CO2(g) (could be any oxide)
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61
Q

Hey!

A

You’re doing great!

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62
Q

Describe the bonding and appearance of carbon dioxide

A

Simple covalent colourless gas :)

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63
Q

Give the equation for the reaction of carbon dioxide with hydroxide ions

A

CO2 + 2OH(-) –> CO3(2-) + H2O
As it is an acidic oxide, it forms a carbonate ion and water

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64
Q

What does carbon dioxide form when it reacts with water? Give the equation

A

Carbonic acid
CO2 + H2O –> H2CO3

65
Q

Describe the appearance and bonding of lead (II) oxide

A

Yellow solid with ionic bonding - most stable state (inert pair effect increase down group)

66
Q

Describe the appearance and stabilityof lead (IV) oxide

A

Brown solid, unstable and decomposes readily into lead (II) oxide (when heated, actually) and releases oxygen

67
Q

Lead (II) oxide is amphoteric. Show this using equations

A

PbO + 2HNO3 –> Pb(NO3)2 + H2O (acidic conditions)
PbO + 2NaOH –> Na2PbO2 + H2O (basic conditions)

68
Q

Explain the redox properties of lead (IV) oxide

A

Good oxidising agent as the +II state is more stable

69
Q

Which two elements form a yellow precipitate with I- ions?

A

Silver and lead

70
Q

Meine Liebe!

A

Your work ethic is admirable :)

71
Q

Give a two similarities and a difference between CCl4 and SiCl4

A

Both colourless covalent liquids, but CCl4 is insoluble in water (cannot be hydrolysed) whereas SiCl4 can be hydrolysed in water to form an oxide and HCl gas

72
Q

Does CCl4 have access to a d-orbital?

A

No; there are none on the outer level (second) and the higher ones are too high energy to reach

73
Q

Give the reaction for the hydrolysis of SiCl4 as well as any observations.
How does the reaction start?

A

SiCl4(l) + 2H2O(l) –> SiO2(s) + 4HCl(g)
Makes a white solid and bubbles.
The reaction starts when a lone pair from the oxygen bonds to the silicon (d-orbital)

74
Q

Describe the solubility of Pb(2+) compounds

A

All insoluble save nitrate and ethanoate

75
Q

Describe the appearance and ability to be hydrolysed of PbCl2 and PbCl4

A

PbCl2 - colourless liquid, will hydrolyse in water to form lead (IV) oxide and HCl gas
PbCl4 - white solid with ionic bonds. Insoluble in cold water - only hot

76
Q

Reactions of Pb(2+) with group VII and other lovely things
Anion…………..Dropwise…………..Excess
Cl(-)
I(-)
SO4(2-)
CO3(2-)
OH(-)

A

Anion………Dropwise………………Excess
Cl(-)…………….White ppt…………………..No change
I(-)………………Canary yellow ppt………No change
SO4(2-)……….White ppt…………………..No change
CO3(2-)……….White ppt…………………..No change
OH(-)…………..White ppt…………………..Colourless solution

77
Q

Write the general formula for the precipitation of Pb(2+)

A

Pb(2+)(aq) + 2X(-)(aq) –> MX2(s)

78
Q

Write the formula for lead with excess NaOH solution

A

Pb(OH)2 + 2OH(-) –> Pb(OH)4 ions in solution

79
Q

Which of iodine and iodide is the oxidising and which the reducing agent?

A

Iodine = oxidising agent
Iodide = reducing agent

80
Q

How does reactivity behave in group VII (the halogens)?

A

Decreases down the group

81
Q

……….Cl2………………………Br2……………………………I2
Cl-………………………………………………………………………….
Br-…………………………………………………………………………
I-……………………………………………………………………………

A

……….Cl2………………………Br2……………………………I2
Cl-……./……………………………No reaction………………No reaction
Br-…..Orange……………………/………………………………No reaction
I-……..Reddish brown……….Reddish brown…………../

82
Q

What do halogens often produce?

A

Poisonous fumes

83
Q

Does chlorine oxidise bromine? Why?

A

Yes, because of chlorine’s higher E(theta) value

84
Q

Does iodine oxidise bromine? Why?

A

No, because of iodine’s lower E(theta) value

85
Q

Why is the -I oxidation state more stable in chlorine than in bromine or iodine?

A

Because E(theta) is a good measure of oxidising power and chlorine has the highest E(theta), making it the most oxidising

86
Q

Write the equations for the reactions of Cl2 with NaOH in warm and cold conditions. What is the oxidation state of the chlorine? What kind of reaction is it?

A

When cold: 2NaOH(aq) + Cl2(g) –> NaClO(aq) + NaCl(aq) + H2O (chlorate I)
When warm: 6NaOH(aq) + 3Cl2(g) –> NaClO2(aq) + 5NaCl(aq) + 3H2O (chlorate V)
They are both disproportionation reactions

87
Q

How can we make chlorates I and V and what are their respective uses?

A

Chlorate I: react with NaOH when cold - used as a steriliser
Chlorate V: react with NaOH when warm - used as weed killer

88
Q

Write the reaction of iodine with NaOH

A

6NaOH(aq) + 3I2(g) –> NaIO3(aq) + 5NaI(aq) + 3H2O(l)

89
Q

What happens to oxidation states as we descend group VII?

A

Electronegativity decreases, making the element less oxidising and the higher oxidation states more stable

90
Q

What do we observe with all sodium halides and concentrated H2SO4?

A

Effervescence (hydrogen halide), and turns damp litmus paper pink (acidic gas)

91
Q

Sodium halides with concentrated H2SO4
Sodium Halide…….Observations
NaCl………………………………………………..
NaBr……………………………………………….
NaI………………………………………………….

A

Sodium Halide…….Observations
NaCl……………………….White misty fumes (HCl), colourless solution
NaBr……………………….Orange/brown fumes (Br2), colourless solution (NaHSO4)
NaI………………………….Purple fumes (HI), black solid (iodine)

92
Q

What is the general equation for the reaction with sodium halides and concentrated H2SO4? Which halogens does this apply to? What happens to the others?

A

NaX(s) + H2SO4 –> NaHSO4 + HX(g)
Only for F and Cl - strong oxidising agents
NaBr and NaI will reduce H2SO4, being strong reducing agents

93
Q

Hey!

A

You are amazing!! app, etc :)

94
Q

What are transition elements?

A

Elements with partially-filled d-orbitals in their ion form

95
Q

Give the four main characteristics of transition metals

A
  • Variable valency (different oxidation states)
  • Form complex ions by dative bonding
  • Form coloured ions
  • Both metals and compounds are used as catalysts
96
Q

Why are many oxidation states possible for transition metals?

A

Because the ionisation energies for the d-orbital electrons are all very similar

97
Q

Give the possible oxidation states for Chromium, Manganese, and Iron, highlighting the stable ones

A

Cr: +II, +III, +VI
Mn: +II, +III, +IV, +VI, +VII
Fe: +II, +III, +VI

98
Q

What is a ligand?

A

A small molecule with a lone pair that can bond to a transition metal

99
Q

Define a complex

A

Ligands joined to a transition metal by dative bonds

100
Q

What is a monodentate?

A

A ligand that has one atom that can bond to a metal ion

101
Q

What is a bidentate?

A

A ligand that has two atoms that can bond to a metal ion

102
Q

Describe briefly the structure of a complex

A

A central cation (usually a transition metal with vacant d-orbitals) surrounded by outer groups called ligands which have a lone pair of electrons

103
Q

Name the shapes complexes can hold

A

Octahedral (six ligands), tetrahedral or occasionally square planar (four ligands), linear (two ligands)

104
Q

Give three examples of ligands

A

H2O, NH3, Cl(-)

105
Q

Draw the transition metal cycle thing (you know the one)

A

See diagram I

106
Q

Explain why transition metals display such vivid colour

A

When a ligands approaches a transition metal, its 5 3d orbitals differ slightly in energy, become non-degenerate, and split into two groups. This splitting will allow an electron to be transferred from a lower to a higher orbital by absorbing visible light. All frequencies are absorbed, with the exception of the one which is emitted.

107
Q

What does the colour of a transition metal depend on?

A

The way orbitals split as well as the ligand

108
Q

What is the colour of Cr(3+)?

A

Dark green

109
Q

What is the colour of Cr2O7(2-)?

A

Orange

110
Q

What is the colour of Co(2+)?

A

Pink!

111
Q

What is the colour of Fe(3+)?

A

Red-brown

112
Q

What is the colour of CrO4(2-)?

A

Yellow

113
Q

What is the colour of MnO4(-)?

A

Purple

114
Q

What is the colour of Fe(2+)?

A

Pale green

115
Q

What is the colour of Cu(2+)?

A

Pale blue

116
Q

Come on!

A

Keep going!

117
Q

………..Colour of Solution……Dropwise OH(-)…..Excess OH(-)
Cr(3+)……………………………………………………………………………………………….
Cu(2+)………………………………………………………………………………………………
Fe(2+)………………………………………………………………………………………………
Fe(3+)………………………………………………………………………………………………
Zn(2+)………………………………………………………………………………………………
General dropwise equations and excess for those needed

A

…………….Colour of Solution….Dropwise OH(-)…..Excess OH(-)
Cr(3+)…….Indigo/violet……………….Grey/green ppt………Green solution
Cu(2+)……Blue…………………………..Blue ppt…………………..No change
Fe(2+)…….Pale yellow…………………Dirty green ppt……….No change
Fe(3+)…….Orange………………………Rusty brown ppt……..No change
Zn(2+)…….Colourless………………….White ppt……………….Colourless solution
M(2+)(aq) + 2OH(-)(aq) –> M(OH)2(s) or M(3+)(aq) + 3OH(-)(aq) –> M(OH)3(s)
Cr(OH)3(s) + 3OH(-)(aq) –> Cr(OH)6(l)
Zn(OH)2(s) + 2OH(-)(aq) –> Cr(OH)4(aq)

118
Q

How do heterogeneous catalysts work?

A

Solid catalysts for gas pahse or solution reactions. Reactants adsorbed to surface of solid, bringing them together

119
Q

Why are d-block elements good homogeneous catalysts?

A

Because their partially filled d-orbitals and variable oxidation states allow them to bond to relevant molecules and oxidise/reduce them toe make them more reactive

120
Q

Name three transition metal catalysts and the reactions they catalyse

A
  1. Iron: haber process for making ammonia
  2. Nickel: catalytic hydrogenation of alkenes
  3. Vanadium (V) oxide: contact process for making H2SO4
121
Q

Nearly there!

A

It gets easier now :)

122
Q

Name 5 ways to measure rate

A
  1. Gas volume at constant pressure with gas syringe
  2. Gas pressure at constant volume when reactants and products are both gaseous
  3. Change in mass when a dense gas is released
  4. Colorimetry when colour changes
  5. Sampling and quenching when none of the above are viable
123
Q

Name two advantages and two disadvantages of sampling and quenching

A
  1. Large range of reactions
  2. No quenching needed for heterogeneous catalysts
  3. Labour and time intensive
  4. Does not work for homogeneous catalysts apparantely?
124
Q

What is the rate equation for reaction aA + bB –> cC + dD

A

rate = k[A]^y[B]^z

125
Q

Give the units of k when the order of reaction is 0, 1, or 2

A

0: mol dm^(-3) s^(-1)
1: s^(-1)
2: mol^(-1) dm^(3) s^(-1)

126
Q

What do rate equation show?

A

The relationship between the rate and concentration of reactants

127
Q

Draw the graphs for [X] and rate over time for order 0-2 reactions

A

See diagram J

128
Q

Which step is the rate determining step?

A

The slowest one

129
Q

What happens to k when the temperature increases?

A

It also increases

130
Q

Give the fancy equation for k

A

k=Ae^(-Ea/RT) where A=frequency factor (same units as k) and R=gas constant on data sheet

131
Q

What is the graphical way to find Ea and A?

A

Plot graph of ln(k) agains 1/T. Grad = -Ea/R. Y intercept = ln(A)

132
Q

Well done!

A

We’re getting there!

133
Q

Give the definition and an example of enthalpy change of atomisation

A

The enthalpy change to form one mole of atoms in the gas phase
Na(s) –> Na(g)

134
Q

Give the definition and an example of enthalpy change of lattice formation

A

The enthalpy change when one mole of ionic compound forms from ions of elements in the gas phase
Na(+)(g) + Cl(-)(g) –> NaCl(s)

135
Q

Give the definition and an example of enthalpy change of lattice breaking

A

The enthalpy change when one mole of an ionic compound is broken into ions of the elements in the gas phase
NaCl(s) –> Na(+)(g) + Cl(-)(g)

136
Q

Give the definition and an example of enthalpy change of hydration

A

The enthalpy change when one mole of gaseous ions is surrounded by water molecules to make a solution
Na(+)(g) + aq –> Na(+)(aq)

137
Q

Give the definition and one example of enthalpy change of solution

A

The enthalpy change when one mole of an ionic compound dissolves in water to form a solution
NaCl(s) + aq –> NaCl(aq)

138
Q

Which are the largest endothermic and exothermic enthalpy changes?

A

The larges endothermic term is the ionisation energy and the largest exothermic the lattice energy

139
Q

How can we tell the stability of a compound from the enthalpy changes?

A

The more negative /\Hf, the more stable the compound

140
Q

How do we use enthalpy changes to measure the covalent character of a compound?

A

Difficult to measure lattice energy directly - get experimental values using Bonn-Haber cycles and calculate theoretical value using ionic model. If the two values are very different, then the compound is ionic with a substantial covalent character

141
Q

How does the lattice energy vary with charge and size of ions?

A

Small ions have more concentrated charge => stronger attraction => more exothermic lattice energies
Greater charge on ion => greater attraction to others => more exothermic lattice energies

142
Q

How can we tell the solubility of a compound from the enthalpy changes?

A

More exothermic /\Hsoln => more soluble

143
Q

Give the equation for /\Hsoln from other enthalpies

A

/\Hsoln = /\HLB + /\Hhyd where generally /\Hhyd > /\HLB

144
Q

What does /\Hat equal for diatomic gases?

A

Half the bond energy

145
Q

Describe the solubility of hydroxide ions as you descend the group

A

Bigger ions down group => less energy to break lattice => more soluble

146
Q

Describe the solubility of sulphate ions as you descend the group

A

Bigger ions down group => lower attraction from larger ion to polar water molecule => hydration energy less exothermic => less soluble

147
Q

Nearly there now!

A

You’re doing great :)

148
Q

What does entropy measure?

A

The disorder of a system

149
Q

/\S surrounding =

A

/\S surrounding = /\H/T where /\H is in J not KJ

150
Q

/\G =

A

/\G = /\H - T/\S (/\H and /\S must have same units)

151
Q

Give the equation for Kc. Which states can it include?

A

Kc = ([C]^c [D]^d)/([A]a [B]b)
For solutions and gases

152
Q

Give the equation for Kp. Which states can it include?

A

Kp = (PC^c PD^d)/(PA^a PB^b)
For gases only

153
Q

What is the partial pressure?

A

The share of the total pressure proportional to the mole fraction

154
Q

What does Kc<1 mean for the spontaneity of a reaction?

A

Kc<1 means more reactants than products at equilibrium => /\G is positive and the reaction is not spontaneous

155
Q

Do catalysts, pressure, and concentration effect the equilibrium constants?

A

Only temperature can

156
Q

What does equilibrium tell us about a reaction?

A

The relative stability of reactants and products and the energy changes, but nothing about the mechanism of the reaction

157
Q

What tells us about the mechanism of a reaction if not equilibrium?

A

The reaction rates

158
Q

Yes! You’re nearly there!

A

Not far to go now! :D

Chemistry (9 decks)

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