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Unit 3 Flashcards

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1
Q

What colour is [Co(H₂O)]²⁺?

A

Pink

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2
Q

What colour is [CoCl₄]²⁻?

A

Blue

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3
Q

Which Groups in the periodic table is the P-Block?

A

Groups 3 - 0 (Inclusive)

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4
Q

What is Octet expansion?

A

The ability of elements to use available d-Orbitals so that there are more than 8 electrons in the outer shell.

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5
Q

What is a Hypervalent molecule?

A

One formed by octet expansion.

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6
Q

Why can Nitrogen not expand its octet?

A

There is no available D-Orbital for it to promote an electron to.

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7
Q

What is an electron deficient atom?

A

One that does not have a full outer shell of electrons (<8).

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8
Q

What is the inert pair effect?

A

The reluctance of the S² pair of electrons in the bonding level (outer shell) to be involved in bonding.

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9
Q

What is the trend of the inert pair effect?

A

The S² pair are less likely to be involved in bonding as you go down the group.
The inert pair effect increases down the group.

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10
Q

Define ‘Amphoteric Character’.

A

A substance that can act as an acid or a base.

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11
Q

What is the Equation for Aluminium acting as a base?

A

Al₂O₃ +6HCl –> 2AlCl₃ + 3H₂O

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12
Q

What is the Equation for Aluminium acting as an acid?

A

Al₂O₃ + 2NaOH + 3H₂O –> 2Na[Al(OH)₄]

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13
Q

What is the Equation for Lead acting as a base?

A

PbO + 2HNO₃ –>Pb(NO₃)₂ + H₂O

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14
Q

What is the Equation for Lead acting as an acid?

A

PbO + 2NaOH + H₂O –> Na₂[Pb(OH)₄]

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15
Q

Explain why aluminium chloride forms dimers and the bonding in the dimer.

A

The aluminium in AlCl₃ only has 6 electrons in its outer shell so accepts electrons in a co-ordinate bond from the chlorine of another AlCl₃ molecule forming a dimer.

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16
Q

What is a doner acceptor compound?

A

Where one compound donates a pair of electrons to an electron deficient compound forming a co-ordinate bond.

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17
Q

Give an example of a Donor acceptor compound and explain the bonding?

A

NH₃BF₃

Boron trifluoride accepts the lone pair from the Nitrogen in ammonia.

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18
Q

Which Carbon allotrope is comparable to Cubic Boron Nitride?

A

Diamond

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19
Q

Which Carbon allotrope is comparable to Hexagonal Boron Nitride?

A

Graphite

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20
Q

What are the 3 differences between Hexagonal BN and Graphite?

A

RIP (Register, Insulator, Polar)
1. Atoms in adjacent layers are in register (Aligned) in Hexagonal BN, they are not in Graphite.
2. Hexagonal BN is an insulator because the free electrons in graphite are the lone pair on the Nitrogen in BN.
3. The electronegativity difference between Boron and Nitrogen gives the BN a polar bond, this is not present in graphite.

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21
Q

What are the uses of Hexagonal Boron Nitride? (5)

A

Semi-conductors - electronics
Ceramics
Microwaves
Catalyst carrier - fuel cells and batteries
Nano tubes out of single layers

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22
Q

What are some properties of Diamond and Cubic BN? (3)

A

Hard, Strong, Extremely high melting point.

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23
Q

Explain some differences between Diamond and Cubic BN. (2)

A

Diamond is harder due to lack of electronegativity difference.
Cubic BN is more stable so is preferred to diamond for grinding certain materials.

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24
Q

Explain the REDOX properties of the oxides of Carbon.

A

CO₂ is more stable than CO because of the inert pair effect being less at the top of the group. Therefore, CO can act as a reducing agent as the carbon is easily oxidised from +2 to +4.

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25
Give an example of CO being used as a reducing agent. (Iron blast furnace reaction)
Extraction of iron in the blast furnace: Fe₂O₃(s) + 3CO(g) --> 2Fe(s) + 3CO₂(g)
26
Explain the REDOX properties of the oxides of Lead.
PbO is more stable than PbO₂ because of the inert pair effect increasing down the group. Therefore PbO₂ will act as an oxidising agent as the lead is easily reduced from +4 to +2.
27
Give an example of PbO₂ being used as an oxidising agent. (Reacting with HCl(conc.))
PbO₂(s) + 4HCl(conc.) --> PbCl₂(s) + Cl₂(g) + 2H₂O(l)
28
Explain the acidic nature of CO₂ with two equations.
Dissolves in water to form carbonic acid (weak). CO₂(g) + H₂O(l) <--> H⁺(aq) + HCO₃⁻(aq) Reaction with alkalis: CO₂(g) + NaOH(aq) --> NaHCO₃(aq)
29
Using CCl₄, SiCl₄ and PbCl₂'s reaction with water, explain the changes in bonding down group 4.
CCl₄ does not dissolve in water because it does not have available D-orbitals in the valence shell so can not expand it's octet to form a co-ordinate bond with the lone pair of electrons from the water molecule. SiCl₄ reacts rapidly with water because silicon has vacant valence shell D-orbitals with which water can form co-ordinate bonds. PbCl₂ contains ionic bonding so does not react with water and will only dissolve in hot water.
30
Give the equation and observations of the reaction between SiCl₄ and Water.
SiCl₄(l) +H₂O(l) --> SiO₂(s) + 4HCl(g) Steamy, white fumes of HCl(g) White ppt of SiO₂(s)
31
Explain the reaction between NaCl and H₂SO₄. (Electrode potentials)
NaCl(s) + H₂SO₄(aq) --> NaHSO₄(s) + HCl(g) Steamy fumes of HCl No further reaction takes place because H₂SO₄ has a more negative electrode potential than the Cl₂/2Cl⁻.
32
Explain the reaction between NaBr and H₂SO₄. (Electrode Potentials)
NaBr(s) + H₂SO₄(aq) --> NaHSO₄(s) + HBr(g) Steamy fumes of HBr Br can be oxidised by H₂SO₄ so a REDOX reaction takes place. 2HBr(s) + H₂SO₄(aq) --> SO₂(g) + Br₂(g) + 2H₂O(g) Colourless gas Orange fumes of bromine
33
Explain the reaction between NaI and H₂SO₄. (Electrode Potentials)
NaI(s) + H₂SO₄(aq) --> NaHSO₄(s) + HI(g) Steamy fumes of HI I can be oxidised by H₂SO₄ so a REDOX reaction takes place. 2HI(g) + H₂SO₄(aq) --> SO₂(g) + I₂(g) + 2H₂O(g) Purple I₂ fumes which can sublime into a black/brown solid or solution. Iodine is an even stronger reducing agent so further reduction of H₂SO₄ can also occur. 6HI(g) + H₂SO₄(aq) --> 3I₂(g) + S(s) + 4H₂O(g) 8HI(g) + H₂SO₄(aq) --> 4I₂(g) + H₂S(s) + 4H₂O(g) Yellow solid - Sulfur Rotten eggs smell - H₂S
34
What is a disproportionation reaction?
A REDOX reaction where atoms of the same element are both oxidised and reduced to form different products.
35
Give the equation for the reaction of chlorine with cold, dilute Sodium hydroxide. Explain the REDOX nature of this equation.
Cl₂(g) + 2NaOH(aq) --> NaCl(aq) + NaOCl(aq) + H₂O(l) Disproportionation reaction - chlorine is both oxidised and reduced.
36
Give the equation for the reaction of chlorine with heated, concentrated Sodium hydroxide. Explain the REDOX nature of this equation.
3Cl₂(g) + 6NaOH(aq) --> 5NaCl(aq) + NaClO₃(aq) + 3H₂O(l) Disproportionation reaction - chlorine is both oxidised and reduced.
37
Give the uses of chorine and chlorate ions.
Their oxidising power is the reason they are used as bleach. Used in clothing and cleaning. Chorine is also used as disinfectant in water supplies as the oxidising power allows them to kill bacteria - The microbe cells are oxidised.
38
Define 'Transition Metal'
A Metal which forms stable ions and has partially filled D-orbitals.
39
What is the electron structure of Fe²⁺? (1s²....)
1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶
40
What is the electron structure of Fe³⁺? (1s²....)
1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵
41
Why is it easy for transition metals to convert between oxidation states?
There is a very small difference in energy between the different ions because the ionisation energy of electrons in the same D-orbital is very similar.
42
What is a transition metal complex?
A central transition metal ion surrounded by molecules or ions that form co-ordinate bonds with the metal ion centre.
43
Define 'Ligand'
A species that donates a lone pair of electrons to the central metal ion in a complex to form a co-ordinate bond.
44
What is the difference between a monodentate and a bidentate ligand?
Monodentate - Donates one lone pair to the central metal ion in a complex. Bidentate - Donates two lone pairs to the central metal ion in a complex.
45
What is the coordination number of a transition metal complex?
Number of atoms donating a lone pair to the central metal ion.
46
What colour is [Cr(H₂O)₆]³⁺?
Dark Green
47
What colour is CrO₄²⁻?
Yellow
48
What colour is Cr₂O₇²⁻?
Orange
49
What colour is MnO₄⁻?
Purple
50
What colour is [Co(H₂O)₆]²⁺?
Pink
51
What colour is [CoCl₄]²⁻?
Blue
52
What colour is [Fe(H₂O)₆]²⁺?
Pale Green
53
What colour is [Fe(H₂O)₆]³⁺?
Yellow
54
What colour is [Cu(H₂O)₆]²⁺?
Pale Blue
55
What colour is [CuCl₄]²⁻?
Yellow or Green
56
Why do transition metal complexes have colour?
When ligands approach the metal ion, they cause the energy of three d orbitals to become different to the other two. (three lower, two higher). This splits the d orbitals to give two sets of orbitals close together in energy. Electrons in the lower 3d orbitals absorb energy from visible light, which promotes them to the higher 3d orbitals. The wavelengths of light that are not absorbed are reflected, giving the colour of the complex.
57
What equation links Planck constant, wavelength, energy, speed of light?
E = hc/λ
58
How to convert energy to KJmol⁻¹?
E x Nₐ / 1000
59
Why do Copper(1) Complexes appear colourless?
There are no available D-orbitals in the valence shell to accept a co-ordinate bond or to allow electrons to move between orbitals so they appear colourless.
60
What colour is [Cu(H₂O)₂(NH₃)₄]²⁺
Royal Blue
61
Why can transition metals act as catalysts? (3)
They have partially filled 3d energy levels so can form complexes with reacting molecules. They have variable Oxidation states. Products are released and they return to their original oxidation state.
62
What are the three steps for writing half equations?
AHC! (Sneeze?) Atoms, Hydrogens, Charge
63
Describe the sampling and quenching method for determining rates of reactions.
small samples of the solution are taken at regular time intervals (sampling) and immediately put in ice water (quenching). This cools and dilutes the reaction mixture effectively stopping the reaction. The sample is then analysed to determine the rate.
64
What are the advantages and disadvantages of the sampling and quenching method?
HRT (lol) (Homogenous, Range, Time intensive) 1. Only appropriate when reaction mixture is homogenous - if not, the sample is not representative of the overall mixture. 2. Method can be used for a large range of different reactions. 3. Method is labour and time intensive, as each sample must be analysed individually.
65
Why is quenching not appropriate when a heterogenous catalyst is being used.
Taking the sample away from the catalyst will immediately slow the reaction so quenching is not nessecary.
66
What is the equation used to calculate rate of a reaction.
Rate = ΔConc. / ΔTime
67
Give the rate equation for a general reaction.
A + B --> Products Rate = k[A]ᵐ[B]ⁿ m = order of A n = order of B
68
What is the order of reaction with respect to a particular reactant?
The power which that reactant is raised in the rate equation.
69
What is the overall order of a reaction?
The sum of all orders in the rate equation.
70
What is the rate determining step?
The slowest step in a reaction.
71
What does the rate equation show about the rate determining step?
The reactants.
72
Give the Arrhenius equation.
k = Ae⁽⁻ᴱᵃ/ᴿᵀ⁾
73
what does each letter represent in the Arrhenius equation?
k - Rate Constant A - Frequency Factor e - Mathematical Constant 'e' Ea - Activation Energy R - Gas Constant (8.31) T - Temperature in Kelvin
74
What is the expression for Kw?
Kw = [H⁺][OH⁻]
75
What colour is PbI₂?
Bright Yellow
76
What is a homogenous catalyst?
A catalyst in the same physical state as the reaction.
77
Give the half equation for the reduction of manganate (VII) ions.
MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O
78
MASH
Metal + Acid = Salt + Hydrogen
79
AASW
Alkali + Acid = Salt + Water
80
BASW
BASW = Base + Acid = Salt + Water
81
CASCW
Carbonate + Acid = Salt + Carbon Dioxide + Water
82
Give the four neutralisation reaction anagrams.
MASH AASW CASCW BASW
83
What colour is Mg(OH)₂ ppt? What happens when you add excess NaOH and excess NH₃?
White ppt White ppt remains. White ppt remains.
84
What colour is Cu(OH)₂ ppt? What happens when you add excess NaOH and excess NH₃?
Blue ppt Blue ppt remains. Blue precipitate dissolves and a deep blue solution forms.
85
What colour is Fe(OH)₂ ppt? What happens when you add excess NaOH and excess NH₃?
Green ppt Green precipitate remains. Green precipitate remains.
86
What colour is Zn(OH)₂ ppt? What happens when you add excess NaOH and excess NH₃?
White ppt White precipitate dissolves and a colourless solution forms. White precipitate dissolves and a colourless solution forms.
87
What colour is Pb(OH)₂ ppt? What happens when you add excess NaOH and excess NH₃?
White ppt White precipitate dissolves and a colourless solution forms. White ppt remains.
88
What colour is Al(OH)₃ ppt? What happens when you add excess NaOH and excess NH₃?
White ppt White precipitate dissolves and a colourless solution forms. White ppt remains.
89
What colour is Cr(OH)₃ ppt? What happens when you add excess NaOH and excess NH₃?
Green ppt Green ppt dissolves giving a dark green solution. Green ppt dissolves giving a dark green solution.
90
What colour is Fe(OH)₃ ppt? What happens when you add excess NaOH and excess NH₃?
Brown ppt Red-Brown ppt remains. Red-Brown ppt remains.
91
Describe the three tests for metal ions.
Add sodium hydroxide or ammonia solution to aqueous metal ions. Add excess NaOH Add Excess NH₃
92
What is the first law of thermodynamics?
Total energy in a closed system is never lost or gained, only transformed.
93
What is the second law of thermodynamics?
Entropy will always tend to increase in a closed system not in equilibrium.
94
Define 'Entropy' (ΔS)
The disorder of a system High entropy = less order
95
From highest to lowest, order the entropy of solids, liquids and gases.
gas, liquid, solid
96
What is the Gibbs free energy equation?
ΔG = ΔH - TΔS
97
What does each letter represent in the Gibbs free energy equation?
G - Gibbs free energy H - Enthalpy T - Temperature in Kelvin S - Entropy
98
What do negative and positive Gibbs free energy values mean?
Negative - reaction will occur spontaneously (Feasible). Positive - reaction will not occur spontaneously (not feasible).
99
Give the equation for entropy change.
ΔS = S(products) - S(reactants)
100
State Le Chatelier's principle.
If a system in dynamic equilibrium experiences a change then the position of equilibrium will shift to oppose that change.
101
Give the general expression for all equilibrium constants.
K = products/reactants
102
Give the general equation for Kc.
Kc = [products]ᵇᵃˡᵃⁿᶜᶦⁿᵍ ⁿᵒ/[reactants]ᵇᵃˡᵃⁿᶜᶦⁿᵍ ⁿᵒ Where [] = Conc.
103
Give the general equation for Kp.
Kp = P(Products)/P(Reactants) Where P = Partial Pressures
104
What is the only thing that changes equilibrium constants?
Temperature
105
Give the Ka equation and what its used for.
Ka = [H⁺][A⁻] / [HA] Calculating [H⁺] of weak acids to find pH.
106
Give the pKa expression.
pKa = -log(Ka)
107
What is the value of Kw.
10⁻¹⁴ mol² dm⁻⁶
108
Give the Kw expression.
Kw = [H⁺][OH⁻]
109
Give the buffer equation for acids and their salts.
[H⁺] = Ka [acid] / [salt]
110
If [H⁺] = [A⁻] what can the Ka equation become?
Ka x [HA] = [

Chemistry (3 decks)

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