Unit 3 Questions Flashcards
(158 cards)
1
Q
Give the equation for the oxidation of iron with permanganate
A
5Fe(2+) + MnO4(-) + 8H(+) –> 5Fe(3+) + Mn(2+) + 4H2O
2
Q
What is the usual oxidation state of NO3?
A
+1 overall
3
Q
What happens at the cathode of an electrochemical cell?
A
Positive cations gain electrons - reduction
4
Q
What happens at the anode of an electrochemical cell?
A
Negative anions lose electrons - oxidation
5
Q
Draw a diagram of an electrochemical cell with Mg(2+) and Cu(2+)
A
See diagram A
6
Q
Write out the half cell diagram thing for Mg and Cu
A
Mg(s) | Mg(2+)(aq) :: Cu(2+)(aq) | Cu(s)
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
8
Q
Why is the voltmeter in an electrochemical cell high resistance?
A
To minimise the current lost as heat
9
Q
Which half cell will be the positive electrode?
A
The one with the most positive E(theta) value
10
Q
Which half cell goes on the left?
A
The most negative E(theta), oxidation
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.
12
Q
Draw the standard hydrogen electrode
A
See diagram B
13
Q
Draw a metal-solution half cell with Zinc
A
See diagram C
14
Q
Draw a mixed ion half cell with Fe
A
See diagram D
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.
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
17
Q
Give the colours of Cu(2+)(aq) + 2e(-) <–> Cu(s)
A
Blue <–> copper colour
18
Q
Give the colours of Fe(3+)(aq) + e(-) <–> Fe(2+)(aq)
A
Yellow <–> pale green
19
Q
Give the colours of Br2(aq) + 2e(-) <–> 2Br(-)(aq)
A
Orange <–> colourless
20
Q
Give the colours of Cr2O7(2-)(aq) + 14H(+)(aq) +6e(-) <–> 2Cr(3+)(aq) + 7H2)(l)
A
Orange <–> dark green
21
Q
MnO4(-)(aq) + 8H(+)(aq) + 5e(-) <–> Mn(2+)(aq) + 4H2O(l)
A
Purple <–> pale pink/colourless apparantely
22
Q
You know what?
A
A w—–, etc :)
23
Q
How do fuel cells generate power?
A
Use electrochemical methods to get energy from fuels, typically hydrogen
24
Q
What happens at the anode of a fuel cell?
A
Hydrogen is oxidised to H+ ions
25
What happens at the cathode of a fuel cell?
Oxygen gas is reduced to water
26
What catalyses the reactions in a fuel cell?
A platinum electrode
27
What is the purpose of the polymer/electrolyte membrane in the hydrogen fuel cell?
It allows only positive ions into the cathode, electrons have to go via an external circuit
28
State two different ways to connect hydrogen fuel cells
In series for higher voltage or parallel for higher current density
29
Give two advantages of hydrogen fuel cells
1. Clean, only waste product is water
2. High efficiency
30
Give two disadvantages of hydrogen fuel cells
1. Expensive (platinum)
2. Storage difficulties - hydrogen is flammable and explosive
31
I have something for you!
~ # ~
It's a waffle! You've earned it :)
32
What is an amphoteric substance?
A substance that acts as an acid in basic conditions and vice-versa
33
I am looking for some p-block elements. Where on the periodic table might I find them?
You must seek the crossover from metal to non-metal and ionic to covalent. There you will find what you seek.
34
Is magnesium amphoteric? Give its equation(s) for the reaction with hydroxide ions and the observations
It is not amphoteric.
Dropwise and excess: Mg(2+)(aq) + 2OH(-)(aq) --> Mg(OH)2(s)
This forms a white precipitate
35
Is zinc amphoteric? Give its equations(s) for the reaction with hydroxide ions and the observations
It is amphoteric.
Dropwise: Zn(2+)(aq) + 2OH(-)(aq) --> Zn(OH)2(s)
Excess: Zn(OH)2(s) + 2OH(-)(aq) --> [Zn(OH)4](2-)(aq)
Dropwise forms white precipitate, excess forms colourless solution
36
What is the highest oxidation state p-block elements can reach?
Their group number
37
What is an inert pair?
An ns2 pair of electrons not involved in bonding
38
What effect does the inert pair effect have on the oxidation state and how does this change down the group?
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
39
Explain octet expansion and its effect on the oxidation state
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.
40
What is a dimer?
A species created when two molecules join together
41
What kinds of compound are boron and aluminium often found in?
Electron deficiency compounds (BF3 or AlCl3) due to each having 2 valence electrons and therefore three covalent bonds (incomplete octet)
42
What kind of species do B and Al often react with?
Lone pair species (removes electron deficiency with dative bond)
43
Draw the bonding of two AlCl3 molecules
See diagram E
44
At what temperature does the dimer Al2Cl6 dissolve back into AlCl3?
Dative bonds break above 200 degrees C
45
Name one use of the dimer Al2Cl6
As a Friedel-Crafts catalyst for chlorination or alkylation of benzene
46
Draw the mechanism for the reaction of CH3Cl and AlCl3. What kind of reaction is this?
See diagram F
It is an electrophilic substitution reaction
47
Why does BF3 form donor-acceptor compounds?
Because the B is electron deficient and will try to react with compounds with lone pairs, accepting a donated electron
48
Draw an example of a donor-acceptor reaction
See diagram G
49
What are ionic liquids
Organic salts with melting points below 100 degrees. They have a wide temperature range at which they are liquid - unusually so for a salt
50
How are ionic liquids formed?
Large organic chlorocompounds, R-Cl, react with AlCl3. Forms dative bond to AlCl3 followed by ionisation to R(+)AlCl4(-)
51
Give two uses and two advantages of ionic liquids
1. Solvents 2. Catalysts
1. They are recyclable 2. Organic products are imiscible in liquids, which means one can separate them easily
52
What is BN isoelectric to and what does this mean?
Carbon - has the same electronic configuration and therefore also similar properties
53
Draw hexagonal boron nitride
See diagram H
54
Describe the structure of hexagonal boron nitride
Same as graphite basically but atoms lie directly above one another with no delocalised electrons
55
What are the three main useful properties of hexagonal boron nitride?
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
56
Describe the structure of tetrahedral boron nitride
Tetrahedral - basically like diamond
57
Explain two uses of cubic boron nitride.
Also two further properties if you feel like it
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
58
Name two trends in group IV
Increasing metallic character and stability of the +2 state down the group
59
Describe how the bonding and structure changes as we descend group IV
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)
60
Describe the oxidation states, appearance, bonding, redox character, and use of carbon monoxide, giving equations where needed for the latter
- 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)
61
Hey!
You're doing great!
62
Describe the bonding and appearance of carbon dioxide
Simple covalent colourless gas :)
63
Give the equation for the reaction of carbon dioxide with hydroxide ions
CO2 + 2OH(-) --> CO3(2-) + H2O
As it is an acidic oxide, it forms a carbonate ion and water
64
What does carbon dioxide form when it reacts with water? Give the equation
Carbonic acid
CO2 + H2O --> H2CO3
65
Describe the appearance and bonding of lead (II) oxide
Yellow solid with ionic bonding - most stable state (inert pair effect increase down group)
66
Describe the appearance and stabilityof lead (IV) oxide
Brown solid, unstable and decomposes readily into lead (II) oxide (when heated, actually) and releases oxygen
67
Lead (II) oxide is amphoteric. Show this using equations
PbO + 2HNO3 --> Pb(NO3)2 + H2O (acidic conditions)
PbO + 2NaOH --> Na2PbO2 + H2O (basic conditions)
68
Explain the redox properties of lead (IV) oxide
Good oxidising agent as the +II state is more stable
69
Which two elements form a yellow precipitate with I- ions?
Silver and lead
70
Meine Liebe!
Your work ethic is admirable :)
71
Give a two similarities and a difference between CCl4 and SiCl4
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
Does CCl4 have access to a d-orbital?
No; there are none on the outer level (second) and the higher ones are too high energy to reach
73
Give the reaction for the hydrolysis of SiCl4 as well as any observations.
How does the reaction start?
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
Describe the solubility of Pb(2+) compounds
All insoluble save nitrate and ethanoate
75
Describe the appearance and ability to be hydrolysed of PbCl2 and PbCl4
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
Reactions of Pb(2+) with group VII and other lovely things
**Anion..............Dropwise..............Excess**
Cl(-)
I(-)
SO4(2-)
CO3(2-)
OH(-)
**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
Write the general formula for the precipitation of Pb(2+)
Pb(2+)(aq) + 2X(-)(aq) --> MX2(s)
78
Write the formula for lead with excess NaOH solution
Pb(OH)2 + 2OH(-) --> [Pb(OH)4](2-) ions in solution
79
Which of iodine and iodide is the oxidising and which the reducing agent?
Iodine = oxidising agent
Iodide = reducing agent
80
How does reactivity behave in group VII (the halogens)?
Decreases down the group
81
**..........Cl2...........................Br2.................................I2**
**Cl-**.....................................................................................
**Br-**....................................................................................
**I-**.......................................................................................
**..........Cl2...........................Br2.................................I2**
**Cl-**......./.................................No reaction..................No reaction
**Br-**.....Orange......................../....................................No reaction
**I-**........Reddish brown..........Reddish brown............../
82
What do halogens often produce?
Poisonous fumes
83
Does chlorine oxidise bromine? Why?
Yes, because of chlorine's higher E(theta) value
84
Does iodine oxidise bromine? Why?
No, because of iodine's lower E(theta) value
85
Why is the -I oxidation state more stable in chlorine than in bromine or iodine?
Because E(theta) is a good measure of oxidising power and chlorine has the highest E(theta), making it the most oxidising
86
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?
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
How can we make chlorates I and V and what are their respective uses?
Chlorate I: react with NaOH when cold - used as a steriliser
Chlorate V: react with NaOH when warm - used as weed killer
88
Write the reaction of iodine with NaOH
6NaOH(aq) + 3I2(g) --> NaIO3(aq) + 5NaI(aq) + 3H2O(l)
89
What happens to oxidation states as we descend group VII?
Electronegativity decreases, making the element less oxidising and the higher oxidation states more stable
90
What do we observe with all sodium halides and concentrated H2SO4?
Effervescence (hydrogen halide), and turns damp litmus paper pink (acidic gas)
91
Sodium halides with concentrated H2SO4
**Sodium Halide.......Observations**
**NaCl**........................................................
**NaBr**.......................................................
**NaI**..........................................................
**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
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?
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
Hey!
You are amazing!! app, etc :)
94
What are transition elements?
Elements with partially-filled d-orbitals in their ion form
95
Give the four main characteristics of transition metals
* Variable valency (different oxidation states)
* Form complex ions by dative bonding
* Form coloured ions
* Both metals and compounds are used as catalysts
96
Why are many oxidation states possible for transition metals?
Because the ionisation energies for the d-orbital electrons are all very similar
97
Give the possible oxidation states for Chromium, Manganese, and Iron, highlighting the stable ones
Cr: +II, **+III**, **+VI**
Mn: **+II**, +III, +IV, +VI, **+VII**
Fe: **+II**, **+III**, +VI
98
What is a ligand?
A small molecule with a lone pair that can bond to a transition metal
99
Define a complex
Ligands joined to a transition metal by dative bonds
100
What is a monodentate?
A ligand that has one atom that can bond to a metal ion
101
What is a bidentate?
A ligand that has two atoms that can bond to a metal ion
102
Describe briefly the structure of a complex
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
Name the shapes complexes can hold
Octahedral (six ligands), tetrahedral or occasionally square planar (four ligands), linear (two ligands)
104
Give three examples of ligands
H2O, NH3, Cl(-)
105
Draw the transition metal cycle thing (you know the one)
See diagram I
106
Explain why transition metals display such vivid colour
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
What does the colour of a transition metal depend on?
The way orbitals split as well as the ligand
108
What is the colour of Cr(3+)?
Dark green
109
What is the colour of Cr2O7(2-)?
Orange
110
What is the colour of Co(2+)?
Pink!
111
What is the colour of Fe(3+)?
Red-brown
112
What is the colour of CrO4(2-)?
Yellow
113
What is the colour of MnO4(-)?
Purple
114
What is the colour of Fe(2+)?
Pale green
115
What is the colour of Cu(2+)?
Pale blue
116
Come on!
Keep going!
117
**...........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
**................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](3-)(l)
Zn(OH)2(s) + 2OH(-)(aq) --> [Cr(OH)4](2-)(aq)
118
How do heterogeneous catalysts work?
Solid catalysts for gas pahse or solution reactions. Reactants adsorbed to surface of solid, bringing them together
119
Why are d-block elements good homogeneous catalysts?
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
Name three transition metal catalysts and the reactions they catalyse
1. Iron: haber process for making ammonia
2. Nickel: catalytic hydrogenation of alkenes
3. Vanadium (V) oxide: contact process for making H2SO4
121
Nearly there!
It gets easier now :)
122
Name 5 ways to measure rate
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
Name two advantages and two disadvantages of sampling and quenching
1. Large range of reactions
2. No quenching needed for heterogeneous catalysts
1. Labour and time intensive
2. Does not work for homogeneous catalysts apparantely?
124
What is the rate equation for reaction aA + bB --> cC + dD
rate = k[A]^y[B]^z
125
Give the units of k when the order of reaction is 0, 1, or 2
0: mol dm^(-3) s^(-1)
1: s^(-1)
2: mol^(-1) dm^(3) s^(-1)
126
What do rate equation show?
The relationship between the rate and concentration of reactants
127
Draw the graphs for [X] and rate over time for order 0-2 reactions
See diagram J
128
Which step is the rate determining step?
The slowest one
129
What happens to k when the temperature increases?
It also increases
130
Give the fancy equation for k
k=Ae^(-Ea/RT) where A=frequency factor (same units as k) and R=gas constant on data sheet
131
What is the graphical way to find Ea and A?
Plot graph of ln(k) agains 1/T. Grad = -Ea/R. Y intercept = ln(A)
132
Well done!
We're getting there!
133
Give the definition and an example of enthalpy change of atomisation
The enthalpy change to form one mole of atoms in the gas phase
Na(s) --> Na(g)
134
Give the definition and an example of enthalpy change of lattice formation
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
Give the definition and an example of enthalpy change of lattice breaking
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
Give the definition and an example of enthalpy change of hydration
The enthalpy change when one mole of gaseous ions is surrounded by water molecules to make a solution
Na(+)(g) + aq --> Na(+)(aq)
137
Give the definition and one example of enthalpy change of solution
The enthalpy change when one mole of an ionic compound dissolves in water to form a solution
NaCl(s) + aq --> NaCl(aq)
138
Which are the largest endothermic and exothermic enthalpy changes?
The larges endothermic term is the ionisation energy and the largest exothermic the lattice energy
139
How can we tell the stability of a compound from the enthalpy changes?
The more negative /\Hf, the more stable the compound
140
How do we use enthalpy changes to measure the covalent character of a compound?
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
How does the lattice energy vary with charge and size of ions?
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
How can we tell the solubility of a compound from the enthalpy changes?
More exothermic /\Hsoln => more soluble
143
Give the equation for /\Hsoln from other enthalpies
/\Hsoln = /\HLB + /\Hhyd where generally /\Hhyd > /\HLB
144
What does /\Hat equal for diatomic gases?
Half the bond energy
145
Describe the solubility of hydroxide ions as you descend the group
Bigger ions down group => less energy to break lattice => more soluble
146
Describe the solubility of sulphate ions as you descend the group
Bigger ions down group => lower attraction from larger ion to polar water molecule => hydration energy less exothermic => less soluble
147
Nearly there now!
You're doing great :)
148
What does entropy measure?
The disorder of a system
149
/\S surrounding =
/\S surrounding = /\H/T where /\H is in J not KJ
150
/\G =
/\G = /\H - T/\S (/\H and /\S must have same units)
151
Give the equation for Kc. Which states can it include?
Kc = ([C]^c [D]^d)/([A]a [B]b)
For solutions and gases
152
Give the equation for Kp. Which states can it include?
Kp = (PC^c PD^d)/(PA^a PB^b)
For gases only
153
What is the partial pressure?
The share of the total pressure proportional to the mole fraction
154
What does Kc<1 mean for the spontaneity of a reaction?
Kc<1 means more reactants than products at equilibrium => /\G is positive and the reaction is not spontaneous
155
Do catalysts, pressure, and concentration effect the equilibrium constants?
Only temperature can
156
What does equilibrium tell us about a reaction?
The relative stability of reactants and products and the energy changes, but nothing about the mechanism of the reaction
157
What tells us about the mechanism of a reaction if not equilibrium?
The reaction rates
158
Yes! You're nearly there!
Not far to go now! :D
