deck_8929427 Flashcards
1
Q
Define acid
A
dissociates in water and releases H+ ions
2
Q
Define alkali
A
dissociates in water to release OH- ions
3
Q
Neutralisation equation
A
H+ + OH-1 H2O
4
Q
Bronsted Lowry Acid:
A
Proton donor
5
Q
Bronsted Lowry Base:
A
Proton acceptor
6
Q
pH = 1 (x 10 the concentration of H+ ions) pH =
A
2
7
Q
What makes an acid strong?
A
It fully dissociates
8
Q
As a Strong acid fully dissociates, pH of a strong acid can be calculated from
A
concentration of acid
9
Q
pH =
A
-log10[H+]
10
Q
Sulphuric acid is dibasic, what does this mean?
A
H2SO4 2H+ + SO42-
11
Q
[H+] =
A
10-pH
12
Q
Weak acids form an acid dissociation constant called
A
Ka
13
Q
Ka =
A
[products] / [reactants]
14
Q
[H+] = for weak acids
A
root([HA] Ka)
15
Q
[HA] is what?
A
concentration of weak acid
16
Q
HCl + H2O –>
A
H3O+ + Cl-
17
Q
What is H3O+
A
A hydronium ion, also known as an oxonium ion
18
Q
What is water acting as in the equation:HCl + H2O H3O+ + Cl-
A
a base
19
Q
What is a bronsted-lowry base?
A
A Bronsted Lowry base is a substance which accepts protons in solution
20
Q
What is a bronsted-lowry acid?
A
A Bronsted Lowry acid is a substance which releases or donates protons in solution
21
Q
What does a bronsted-lowry acid-base reaction involve?
A
the transfer of a proton from one base to another
22
Q
What does ammonia do in water? include equation
A
Ammonia, in water, accepts a protonNH3 + H2O NH4+ + OH-
23
Q
In the equation,NH3 + H2O NH4+ + OH-, what is water acting as?
A
An acid
24
Q
For a bronsted-lowry base, what is the pH of an acid?
A
So for the Bronsted Lowry theory, a base does not need to have a pH>7
25
What are substances that can act as either a base or an acid called?
Substances that can act as either an acid or a base are amphoteric
26
What are some old defnitions of acids?
.The definition of acids has developed since the time of the ancient Greeks.Simpler ideas involved substances that had a sour taste, contained hydrogen, hydrogen ions or had a pH lower than 7
27
.When a chemical reacts with an acid it is the __________ from the acid which is driving the reaction forward
hydrogen ion, H+,
28
HCl (aq) + NaOH (aq)
NaCl (aq) + H2O (l)
29
How did the Bronsted-Lowry name come about?
Johannes Nicolaus Bronsted and Martin Thomas Lowry did not work together but both chemists formulated the idea that acids are proton donors and bases are protons acceptors
30
The Brønsted–Lowry acid–base theory is a development of these earlier ideas and states that:
.An acid is a proton donor.A base is a proton acceptor
31
.How many protons, neutrons and electrons does a Hydrogen ion, H+ have?
.A hydrogen ion is just a proton
32
We know that HCl forms a covalent bond between the hydrogen and chlorine so what happens when it dissolves in water to become hydrochloric acid?
.HCl(g) + aq H+(aq) + Cl-(aq)
33
HCl(g) + aq H+(aq) + Cl-(aq).On closer inspection is?
HCl(g) + H2O(l) H3O+(aq) + Cl-(aq)
34
.H3O+(aq) is known as the hydronium ion, also known as _______ ion
hydroxonium
35
.In this equation:HCl(g) + H2O(l) H3O+(aq) + Cl-(aq)what is HCl and H2O acting as?
The HCl is a proton donor (a Brønsted–Lowry acid).The H2O is a proton acceptor (a B-L base)
36
In the reverse equation of HCl(g) + H2O(l) H3O+(aq) + Cl-(aq) what happens to the roles?
.In the reverse equation, the roles reverse:H3O+(aq) + Cl-(aq) HCl(g) + H2O(l)
37
once an acid has ‘donated’ a proton it would become able to ‘accept’ a proton back and hence act as a base, what is this called?
We call these pairs of chemicals conjugate acid-base pairs
38
Give an example of a conjugate acid-base pair?
HCl and Cl- = Acid and conjugate base
39
What are the acid base pairs in HCl + H2O H3O+ + Cl-
HCl = acid 1Cl- = base 1H2O = base 2H3O+ = acid 2
40
What are the acid base pairs in CH3COOH + H2O CH3COO- + H3O+?
CH3COOH = acid 1CH3COO- = base 1H2O = base 2H3O+ = acid 2
41
What are the acid base pairs in NH3 + H2O NH4+ + OH-?
Base 2 Acid 1 Acid2 Base 1
42
What are the acid base pairs in HCO3- + HCl H2CO3 + Cl-?
Base 2 Acid 1 Acid 2 Base 1
43
Complete the equation for the conjugate acid-base pair: CH3CH(OH)COOH + CH3CH2CH2COOH
CH3CH(OH)COO- + CH3CH2CH2COOH2+
44
List the roles of H+ in reactions
spectator ions:Acid + metal salt + hydrogenSolid carbonates and soluble carbonates:Acid + carbonate Water + carbon dioxidebase (metal oxides): Acid + base salt + waterAlkali: H+ + OH- H2O
45
What did Soren Sorenson do?
introduced simple numbers to represent the colours of indicators using an electrochemical cell to measure the hydrogen content
46
What did Soren Sorenson find?
.He found the ion concentration had a very large range of values of powers of 10 (10-1 to 10-14)
47
What is a strong acid? With general equation
.A strong acid is one which completely dissociates into ions in a solutionHA H+ + A-
48
So for a strong acid, [H+] =
concentration of the acid ([HA])
49
pH =
-log10[H+]
50
What does the equation pH = -log10[H+] tell us?
.It tells us the relative hydrogen ion concentration of a given solution
51
What must we remember with the equation pH = -log10[H+]?
.THIS ONY WORKS FOR STRONG ACIDS
52
.An increase in H+ x10, it reduces the pH by how much?
1
53
What is pH of 1M HCl?
0
54
What is Sorenson's pH scale?
.The logarithmic scale means that a shift of one pH unit means a 10x change in the acidity and alkalinity of the solution.Theoretically there is no limit to the pH scale
55
What type of acid is HCl, what does this mean?
monobasic = [H+]
56
What type of acid is H2SO4, what does this mean?
dibasic = 2[H+]
57
What type of acid is H3PO4, what does this mean?
tribasic = 3[H+]
58
[H+] =
10-pH
59
What is the [HCl] of a solution with pH 1.8?
0.016 moldm-3
60
What is the [H2SO4] of a solution with pH 1.8?
0.008 moldm-3 as dibasic acid
61
50cm^3 of 0.1M HCl is diluted to 100cm^3 with water, what is the change in pH?
Before dilution, pH = 10n dilution, HCl conc is halved to 0.05MAfter dilution, pH = 1.30Change = 0.3
62
What is a strong acid?
A strong acid is one which completely dissociates into ions in a solution
63
WHat does pH = -log10[H+] only work for?
ONLY WORKS FOR STRONG ACIDS
64
[H^+] equation for changing concentraation
[H+] = [H+]old x (old volume / new volume)
65
How are weak acids different to strong ones?
.Weak acids do not dissociate fully like strong acids
66
What happens when a weak acid dissociates?
An equilibrium is established between the ions and the acid
67
As an equibrlium is established when weak acids dissociate, what can we calculate?
we can calculate an equilibrium constant Ka (Acids dissociation cons tat)
68
What does Ka equal for, HA H+ + A-?
Ka = ( [H+] [A-] ) / [HA]
69
For the Ka equation what two assumptions need to be made?
.A pure acids will split evenly into equal quantities of [h+] and [A-].The concentration of the acid doesn’t change as the equilibrium lies so far to the left it can be considered equal
70
does [H+] = [H+ aq]?
yes
71
.A sample of ethanoic acid of concentration 0.04 moldm-3.What is the value of Ka?
2.5x10-9 moldm-3
72
Steps to calculating pH of weak acid
.Step 1: Calculate [H+] from ka and [HA].Step 2: Calculate pH
73
Are the approximations for weak acids justified?
.The first approximation assumes that the dissociation from water is negligible.[H+(aq)]eqm = [A-(aq)]eqm - if pH>6 then water dissociates and [H+] from water will be more significant than from dissociation of acid.This approximation breaks down for very weak acids or very dilute solutions.Second approximation assumes that the concentration of the acid [HA] is much greater than the [H+] concentration at equilibrium.[HA]start >> [H+]eqm [HA]eqm = [HA]start – [H+]eqm so [HA]eqm = [HA]start.Not valid as acid gets stronger as [H+] becomes more significant and real difference between [HA]eqm = [HA]start – [H+]eqm.Not justified for stronger weak acids / very dilute solutions.
74
.The value of Ka for weak acids is almost always very small and difficult to compare numbers with negative indices, how do we get around this?
create a scale of more useable numbers we often use the value of pKa, when talking about the aciditing of weak acids
75
What is pKa often used for?
to compare acids in biological systems
76
pKa equation
pKa = -log10Ka
77
What do the values of pKa mean?
The higher the value of pKa, the weaker the acidThe lower the value of pKa, the stronger the acid
78
How does wine show behaviour of typical dibasic and tribasic acids?
Wines often contain traces of sulphurous acid, H2SO3, added as a preservative.Sulphurous acid is dibasic and its dissociation is shown as:H2SO3 H+ + HSO3 - (pKa = 1.92)HSO3 - H+ + SO3 2- (pKa = 7.18).For the first dissociation, H2SO3, acts as a weak acid.From the pKa values, HSO3, is a far weaker acid than H2SO3.The behaviour is typical of dibasic and tribasic acids
79
Where strong acids completely dissociate to release all H+ ions into solution, weak acids only partially dissociate, explain this in terms of equations
Strong: HA → H+ + A-Weak: HA ⇌ H+ + A-
80
As weak acids form an equilibrium their dissociation can be represented by the acid dissociation constant, Ka, whats the equation for this?
Ka = "[H+(aq)][A-(aq)]" /"[HA(aq)]"
81
Explain why a larger value for Ka means a lower pH
.The larger the value for Ka the more the equilibria lies to the right and so the lower the pH i.e. higher [H+]
82
What alters Ka?
Temperature
83
.As the values for Ka are very small it is easier to give their negative logarithm, pKa, instead, what is the equation for this?
pKa = - logKa
84
pKa = - logKa, what is the inverse equation for this?
The inverse for this equation is: Ka = 10-pKa
85
This means that the weaker an acid the:._____ Ka._____ pKa
.Smaller Ka.Larger pKa
86
.The [H+] of a weak acid depends on the value of what?
Ka and [HA]
87
.The equilibrium concentrations, [ ]eqm, can be used to determine a value for Ka, what is the equation for this?
Ka = "[H+]eqm [A-]eqm" /"[HA]eqm"
88
There are two approximations that need to be made when calculating Ka, what are they?
.HA dissociation forms equal [H+] and [A-].The change in [HA] is negligible so [HA]eqm = [HA]start
89
Using the approximations, Ka can be determined by what equation?
Ka = "[H+]eqm 2" /"[HA]start"
90
How can pH be calculated for a weak acid?
[H+] = √("Ka x [HA]" ) → pH = -log[H+]
91
.A value for Ka can be determined experimentally, how?
by using a pH meter to get the pH of a standard solution
92
.The larger the value for Ka, the greater the _________
dissociation
93
When calculating values for Ka there are issues with the approximations made, what are these issues?
• At pH values >6 water dissociation is significantTherefore doesn’t work for very weak acids or very dilute solutions• If [H+] concentration is significant there will be a difference between [HA]eqm and [HA]startTherefore doesn’t work stronger weak acids with Ka > 10-2 mol dm-3 or very dilute solutions
94
All aqueous solutions contain which ions?
H+ and OH- ions
95
H2O
H+ + OH-
96
In ____ [H+] > [OH-]In ____ [OH-] > [H+]In ____ [H+] = [OH-]
In acids [H+] > [OH-]In alkalis [OH-] > [H+]Neutral [H+] = [OH-]
97
For every 500,000,000 H2O molecules, only _ dissociates
1
98
For every 500,000,000 H2O molecules, only 1 dissociates, what does this mean for the equilbrium?
the equilibrium is on the left hand side
99
Do the Kc equation for water equilibrium
Kc = ([H+][OH-]) / [H2O]
100
[H2O] x Kc =
[H+] x [OH-]
101
[H2O] is such a large excess it can be classed as a ____
constant
102
What is Kw?
ionic product of water
103
Kw = [H2O] x Kc, so replace [H2O] x Kc with Kw in the correct equation
Kw = [H+] x [OH-]
104
At 298K, Kw =
1x10-14 mol2dm-6
105
Why is the pH of pure water at 298K, 7?
Kw = [H+] x [OH-]1x10-14 = [H+] x [OH-]The concentrations are the same and so it can be written as:1x10-14 = [H+]21x10-7 = [H+]pH = 7
106
Effect of temperature on Kw
The dissociation of water is endothermicTemperature increase will move the equilibrium to the right, and so Kw will increaseTemperature decrease will move the equilibrium to the left, and so Kw will decreaseWater will always remain neutral at all temperatures, pH may change with temperature, but [H+] = [OH-] all the time and so it is neutral
107
.Water ionises very slightly, acting as both an acid and as a base – setting up an ________
equilibrium
108
.Water dissociates a very, very small amount according to the equation – it must do, otherwise it would not _____________
conduct electricity
109
.1 dm3 (1000g) of water is mainly ________ H2O
undissociated
110
. [H2O(l)] = 1000/18 = 55.6 mol dm-3 (a constant), why?
.1 dm3 (1000g) of water is mainly undissociated H2O
111
.If we know the concentration of hydroxide ions we can rearrange Kw to give us the value of what?
[H+]
112
What is the pH of 0.4 mol dm-3 NaOH?
Kw = [H+][OH-] = 1.00 × 10–14 mol2dm–6 (1.00 × 10_14 )/([0.4])=[H+] = 2.5 x 10-14 mol dm-3 pH = -log10[H+] = -log10[2.5 x 10-14 ] = 13.60
113
What is the pH of a solution with [OH-] = 2 x 10-2 mol dm-3 at 25oC ?
Step 1 : calculate [H+] from Kw and [OH-] Kw = [H+][OH-] = 1.00 x 10-14 [H+] = Kw = 1.00 x 10-14 = 5.00 x 10-13 mol dm-3 [OH-] 2.00 x 10-2Step 2 : Use calculator to find pH pH = - log [H+] = -log (5.00 x 10-13) = 12.30
114
What are the concentrations of H+ (aq) and OH- (aq) in a solution of pH 3.25 at 25oC ?
Step 1 : Use calculator to find [H+(aq)][H+] = 10-pH = 10-3.25 = 5.62 x 10-4 mol dm-3Step 2 : Calculate [OH-] from Kw and [H+]Kw = [H+][OH-] = 1.00 x 10-14[OH-] = Kw = 1.00 x 10-14 = 1.78 x 10-11 mol dm-3 [H+] 5.62 x 10-4
115
For pH values that are whole numbers, it is easy to work out the [H+] and [OH-] concentrations as the indices add up to what?
-14
116
.The pH of weak bases can be calculated via a similar method to that used for ______
weak bases
117
.In an aqueous solution, there will always be both H+ (aq) and OH- (aq) ions present such that ________
[H+ (aq)][OH- (aq)] = Kw
118
.A solution is ______ when [H+ (aq)] > [OH- (aq)] .A solution is ______ when [H+ (aq)] = [OH- (aq)].A solution is ______ when [OH- (aq)] > [H+ (aq)]
.A solution is acidic when [H+ (aq)] > [OH- (aq)] .A solution is neutral when [H+ (aq)] = [OH- (aq)].A solution is alkaline when [OH- (aq)] > [H+ (aq)]
119
.So a solution that is acidic will still contain __ ions, it is just that there are more __ions (and vice versa in an ______ solution)
OH-H+Alkaline
120
.The value of Kw controls the ______ of each ion
Concentrations
121
• Kw can also tell us the pH of ______
pure water
122
• As the water splits into equal concentrations of OH- and H+ ions we can make what assumption?
Kw = [H+][OH-] = 1.00 × 10–14 mol2dm–6 = [H+]2 = 1.00 × 10–14 = [H+] = 1.00 × 10–7 pH = -log10[H+] = -log10[1.00 x 10-7 ] = 7
123
When is neutral not neutral?
.A neutral solution is defined by a equal number of moles of H+ ions and OH- NOT by a pH of 7 as you may have been taught.Whilst this value is 7 at approximately 298K, the value of Kw increases with temperature
124
.The pOH scale measures what?
the concentration of hydroxide ions
125
.Therefore you can solve fpr pH, pOH and [OH-] given just ___
[H+]
126
pH + pOH =
14
127
.Kw = [H+][OH-] =
1x10-14
128
Define Enthalpy of reaction ΔrH
enthalpy change that accompanies a reaction in the molar quantities shown in a chemical equation under standard conditions, with all reactants and products in their standard states
129
Define Enthalpy of formation ΔfH
enthalpy change when one mole of compound is formed from its constituent elements under standard conditions
130
Define Enthalpy of combustion ΔcH
enthalpy change when one mole of substance is burnt completely in excess oxygen under standard conditions
131
Water slightly ionises, what equation will show this?
H2O(l) H+(aq) + -OH(aq)
132
Ka of :H2O(l) H+(aq) + -OH(aq)=
[H+][-OH] / [H2O]`
133
Ka x [H2O] =
Kw
134
Kw =
[H+][-OH]
135
What is Kw?
Kw is the ionic product of water, at 25’C it equals 1x10-14 mol2dm-6
136
How does ionic bonding work?
1. Loss of an electron(s) by an element2. Gain electrons by a second element3. Attraction between positive and negative ions
137
Na + Cl, ionisation electron addinity and latice equations
Ionisation Energy – Na e- + Na+ +496kJmol-1Electron Affinity – e- + Cl Cl- -349kJmol-1Lattice Enthalpy – Cl- + Na+ NaCl -766kJmol-1
138
What is electron affinity?
.The enthalpy change when one mole of electron is added to one mole of atoms in the gaseous phase to form one mole of -1 ions
139
Is repulsion between two electrons exothermic or endothermic, why?
.Repulsion between two negatively charged things requires energy so exothermic
140
What is lattice enthalpy?
.ΔHlatt is the enthalpy change when 1 mole of ionic substance is formed from its gaseous ions under standard conditions
141
.The strength of the ionic bond is related to the lattice enthalpy, how?
more exothermic the greater the ionic bonding
142
Lattice enthalpy can't be measured, why?
Cannot be measured directly as cannot form one mole of ionic lattice from gaseous ions
143
What are the factors affecting lattice enthalpy? How?
Charge – .The greater the charge on the ions, the stronger the attraction – therefore, more exothermic lattice enthalpySize – .Smaller ions can pack together more tightly, therefore there is greater attraction and more exothermic lattice enthalpy
144
Why, when moving down group two, do the temperatures required to break the lattice increase?
.They have decreasing charge densities.Mg2+ is a smaller ion than Ba2+ , so the +2 charge occupies a smaller volume – this means Mg2+ has a higher charge density than Ba2+.Mg2+ can distort the electron clouds within the CO32- ion (called polarisation), this weakens the covalent bonding in the ion and reduces its decomposition temperature
145
What is a buffer?
a solution that minimises pH change when a small amount of acid or alkali is added
146
What two types of buffers are they?
Weak Acid and Salt of Weak AcidExcess Weak Acid and Strong Base
147
Suggest and explain a Weak Acid and Salt of Weak Acid buffer
Weak Acid – Ethanoic Acid – CH3COOH(aq) CH3COO-(aq) + H+(aq)Salt of Weak Acid – Sodium Ethanoate – CH3COONa(s) + (aq) CH3COO-(aq) + Na+(aq)¬Buffer Contains – CH3COOH(aq) CH3COO-(aq) + H+(aq)When add H+, H+ reacts with conjugate base, equilibrium will move to the left to reduce the amount of H+, pH is constantAdding –OH (alkali) –-OH + H+ H2OConc of H+ decreases, equilibrium moves to the right to increase the conc of H+, pH is constant
148
Suggest and explain an excess Weak Acid and strong base buffer
CH3COOH(aq) + NaOH(aq) --> CH3COONa(aq) + H2O(l)CH3COONa(aq) CH3COO-(aq) + Na+(aq)CH3COOH(aq) CH3COO-(aq) + H+(aq)When add H+, H+ reacts with conjugate base, equilibrium will move to the left to reduce the amount of H+, pH is constantAdding –OH (alkali) –-OH + H+ H2OConc of H+ decreases, equilibrium moves to the right to increase the conc of H+, pH is constant
149
Excess methanoic acid is reacted with potassium hydroxide, explain how a buffer solution is produced and how pH is controlled when 5cm3 of HCl is added.
The solution now contains HCOOK(aq) HCOO- (aq) + K+(aq) and HCOOH(aq) HCOO-(aq) + H+ equilibriums in solution – the buffer. When 5cm3 of HCl is added, it dissociates into H+ and Cl-, the H+ ions increase the concentration of the H+ already in the solution, so the HCOOH(aq) HCOO-(aq) + H+ point of equilibrium moves to the right – decreasing the conc of H+ in solution, and so keeping the pH constant.
150
Ka x ( [HA]/[A-] ) = ? what does each part mean?
[H+] = Ka x ( [HA]/[A-] )[HA] = concentration of weak acid[A-] = concentration of conjugate base
151
50cm3 of 1.2M NaOH reacts with 250cm3 of ethanoic acid (1M), Ka = 1.74x10-5, what is the pH?
CH3COOH CH3OO- + H+CH3COOH + NaOH CH3COO- + Na+ + H2O1. Moles weak acid = 250x1 / 1000 = 0.25mol2. Moles NaOH = 50x1.2 / 1000 = 0.06mol3. Moles A- = 0.06mol4. Moles HA = 0.19mol5. [H+] = Ka ([HA] / [A-]) = 1.74x10-5 x (0.19/0.06) = 5.51x10-56. pH = 4.26
152
Equation to work out lattice enthalpy from born haber cycle
L.E = FORMATION – sum(ATOM + I.E + E.A)orL.E = FORMATION – sum(REST)
153
What are Born Haber Cycles used for?
Born Haber cycles can be used to calculate a measure of ionic bond strength based on experimental data
154
Draw the born haber cycle for NaCl
check notes or google
155
Draw the born haber cycle for MgCl2
check notes or google
156
Draw the born haber cycle for CuO
check notes or google
157
Blood must contain a pH of what?
Blood must contain a pH of 7.40 +- 0.05
158
What is the most important buffer in blood?
the carbonic acid – hydrogencarbonate
159
Carbonic acid dissociation equation
H2CO3 H+ + HCO3-
160
If blood becomes too acidic (acidosis), what presents?
fatigue, shortness of breath, shock, death
161
If blood becomes too alkaline (alkalosis), what presents?
spasms, light-headed, nausea
162
What happens if the body produces more acidic products?
.The H+ concentration in the blood will increase, the equilibrium of the hydrogencarbonate dissociation will shift to the left, the H+ ions will be used up to form more hydrogencarbonate, the H+ concentration will return to normal, the pH will have had minimal change.If the acid level continued to rise, fatigue, shortness of breath, shock, or death would eventually set in.This would occur when we have ran out of HCO3- ions
163
Carbon dioxide dissolves in water to form what? What is the equation for this?
carbonic acid:CO2 + H2O H2CO3
164
Inhalation of high levels of CO2 means what?
Inhalation of high levels of CO2 mean that the equilibrium in the blood must shift to the right to form more H2CO3, this means that the level of H2CO3 in the blood increase, so in the equilibrium H2CO3 H+ + HCO3- the equilibrium shifts to the right to counter this, this forms more H+ ions in the blood, making it more acidic, and resulting in acidosis which would lead to fatigue, shortness of breath, shock, and eventual death. Increased [H+] means enzymes would also be denatured.
165
Define enthalpy of solution, with an example
DeltasolutionH is the enthalpy change when one mole of ionic compound is completely dissolved in water under standard conditions.e.g. NaCl(s) Na+(aq) + Cl-(aq)
166
What is enthalpy of hydration with example
- Delta Hhydration¬ is the enthalpy which takes place when one mole of gaseous ions is dissolved in water forming one mole of aqueous ions under standard conditions- Na+(g) + aq Na+(aq)
167
Why is enthalpy of hydration always exothermic?
These are exothermic as bonds are formed between the ions and water molecules
168
Factors Affecting Enthalpy of Hydration
Charge – the higher charge on the ion, the greater the attraction for the H2O molecules, therefore a more exothermic hydrationSize – smaller ions have a greater charge density compared to the larger ions, this creates a greater attraction for H2O molecules, therefore a more exothermic hydration
169
Draw the general cycle of enthalpy of solution, hydration and lattice enthalpy
check notes
170
Equation that links lattice enthalpy, enthalpy of solution, and enthalpy of hydration
lattice enthalpy + enthalpy of solution = sum of the enthalpy of hydrations
171
Draw the titration curve for when HCl is added to NaOH
check notes
172
Draw the titration curve for when NaOH is added to HCl
check notes
173
Draw strong acid strong base titration curve
check notes
174
Draw strong acid weak base titration curve
check notes
175
Draw weak acid strong base titration curve
check notes
176
Draw weak acid weak base titration curve
check notes
177
The shape of the titration curve depends on what?
- Substance being titrated i.e. acid or base- Type of acid or base i.e. strong or weak
178
Describe and explain parts of titration curve
1. Excess of base: pH decreases as acid added2. Vertical section: acid base concentration similar, pH alters rapidly3. Excess of acid: pH decreases slightly as acid is added4. The equivalence point: halfway up the vertical section
179
What is the equivalence point?
The point halfway between the two horizontal lines on the titration curveThe equivalence point is the volume required to have completely reacted the acid and base toetehr, stoichiometric volumesThe pH at the equivalence point depend son the type of acid and base
180
Where is the equivalence point on each titration curve?
Equivalence point = 7 on SA-SBEquivalence point < 7 on SA-WBEquivalence point > 7 on WA-SBNo equivalence point on WA-WB
181
Why is the equivalence point 0 on SASB titration curves?
- H+ ions completely neutralised by OH- ions, so only water and salt present, neutral
182
Why is the equivalence point < 7 on SAWB titration curves?
- A weak base, like NH3, will have a strong conjugate acid, NH4+ , which will react with water to produce H3O+ so the pH is less than 7
183
Why is the equivalence point > 7 on WASB titration curves?
- A weak acid, HA, will have a strong conjugate base, A-, which can react with water to produce OH- ions, so the pH is more than 7
184
What are acid-base indicators?
Acid-Base indicators are weak acids that have different coloured conjugate basesIn acidic conditions the indicator equilibrium is shifted towards the weak acid (HIn)As the system becomes more basic the equilibrium shifts towards the conjugate base (In-), altering the colour
185
What is the end point?
The end point is when equal [HIn] and [In-] are present, the colour will therefore lie between both extremesEach indicator will have a different pH value for the end point as they each have different ka values
186
HIn
H+ + In-
187
How is an acid-base indicator chosen?
The indicator required for a titration must have a colour change that lies within the vertical section of a titration curve
188
A chemical reaction will proceed when?
A chemical reaction will proceed if the products are energetically more stable than the reactants
189
What is entropy?
Entropy is a measure of the dispersal of energy in a system, the more disordered a system the greater the dispersal of energy = higher entropywe define entropy as a measure of disorder
190
What happens to entropy over time?
Entropy must increase over time
191
When is a system in a state of high entorpy?
When its degree of disorder is high
192
As order within a system increases, its entropy decreases, why?
This can be explained in terms of probability: disordered states are simply more likely to exist (or emerge) than ordered states.The spontaneous direction of change is from a less probable to a more probable state
193
What does entropy always do?
The total entropy always increases, and the process is irreversible
194
What is the unit for entropy
Sin JK^-1mol^-1
195
Why is S (entropy) alwasy potisive?
All substances process some degree of disorder because particles are always in constant motion
196
How does the entropy of each state differ?
Solid has lowest entropy, and gas has the highest
197
Draw a graph of waters change in enthalpy as temperaurer increases
check notes
198
Systems that are more chaotic have a _____ entropy value
higher
199
(s) (s) + (g) , what is delta S
+
200
(g) 2(g) , what is delta S
+
201
What is delta S?C2H5OH(l) C2H5OH(g)
+
202
What is delta S?C2H2(g) + 2H¬2(g) C2H6(g)
-
203
What is delta S?NH4Cl(s) + aq NH4Cl(aq)
+
204
What is delta S?4Na(s) + O2(g) 2Na2O(s) deltaS
-
205
Define the standard entropy change
The standard entropy change is the entropy change that accompanies a reaction in the molar quantities expressed in the equation, under standard conditions
206
Sum for delta S^theta
delta S^theta = sum(S^theta products) – sum(S^theta reactants)
207
Why are exothermic reactions more preferable in nature?
as the products are more stable than reactantsthe key is not the decrease in energy but the associated increase in entropy of the surroundings.
208
deltaSsurroundings is proportional to what?
#NAME?
209
deltaSsurroundings =
(-deltaHsystem)/T
210
deltaStotal =
deltaSsystem + deltaSsurroundings
211
deltaStotal >
0
212
What can we learn about the entropy of the sun by the photosynthesis equation?
6CO2(g) + 6H2O(l) UV light C6H12O6(s) + 6O2(g)Negative entropy changeEntropy change in sun must be so positive it outweighs every plant on earth
213
What is gibbs free energy equation?
deltaG = deltaH – T deltaS < 0
214
How do you get to gibbs free energy equation?
deltaStotal¬ = deltaSsystem + deltaSsurroundings > 0deltaStotal¬ = deltaSsystem – (deltaHsystem)/T > 0T deltaStotal = T deltaSsystem – deltaHsystem > 0-T deltaStotal = -T deltaSsytem + deltaHsystem < 0deltaG = deltaH – T deltaS < 0
215
Using deltaG explain when a reaction is and is not feasible
deltaG must be negative (<0) for a reaction to be feasible i.e. proceedIf deltaG is positive (>0) then a reaction is not feasible
216
What is deltaG at the points of feasibility?
At the point of feasibility, we can say deltaG = 0 (assume that deltaH and deltaS don’t vary with temperature)
217
At low temperatures what does deltaGsystem equal? why?
- At low temperatures, deltaGsystem = deltaH (-T deltaS becomes negligible) so for a reaction to occur it needs to be exothermic
218
At high tempertaures deltaGsystem ewuals what? why?
- At high temperatures, deltaGsystem = -T deltaS (deltaH becomes negligible) so for a reaction to occur it needs to have a positive deltaS as – T deltaS needs to be less than 0
219
Limitations to deltaG equation
just because the value is negative and so feasible it doesn’t mean it occurs, the reaction rate might be incredibly slow or the activation energy too high
220
What is an oxidation number?
An oxidation number shows the charge of an atom if all of its bonds were considered totally ionic
221
What are the oxidation number rules?
1. Elements in their natural state = 02. The total oxidation states in a molecule = 03. The total oxidation states in an ion = the chargeGroups 1, 2, 3 - State +1, +2, +3Fluorine - State -1Hydrogen -State +1-1 in metal hydridesOxygen - State -2-1 in peroxidesChlorine -State -1+ (varies) in chlorates
222
What is an oxidising agent?
takes electrons from what is oxidised
223
What is a reducing agent?
gives electrons to what is reduced
224
Write the half equations and full equation for 4Na + O2 2Na2O
4Na 4Na+ + 4e-O2 + 4e- 2O2-4Na + O2 4Na+ + 2O2-
225
Write the half equations for Cu2O + H2SO4 Cu + CuSO4 + H2O
Cu+ Cu2+ + e-2Cu+ Cu + Cu2+
226
Steps of a redox titration using iron sulphate and sulphuric acid and potassium manganate (VII)
1. Make up a 150cm3 standard solution using 7g of iron sulphate by weighing by difference in a volumetric flask, record your mass measurements in a table2. a. Pipette 25cm3 of the standard solution into a conical flaskb. Add 10cm3 of 1M H2SO4 to the conical flaskc. Fill the burette with a standard solution of 0.02M potassium manganate (VII)d. Carry out the titrations until the ed point is reached, the first permanent pink colour, record your results in a tablee. Repeat until you have two concordant resultsf. Calculate the mean titre
227
Reduction: MnO4- + 8H+ + 5e- Mn2+ + 4H2OOxidation: Fe2+ Fe3+ + e- or 5Fe2+ 5Fe3+ + 5e-MNO4- + 8H+ + 5Fe2+ Mn2+ + 4H2O + 5Fe3+What is the oxidising and reducing agent?
Fe2+ is the reducing agentMnO4- is the oxidizing agent
228
Mass of impure hydrated iron II sulphate = 6.97g25cm3 of hydrated iron (II) sulphate used taken from a 250cm3 volumetric flask full of the dissolved iron tabletsMean titre of 0.02M potassium manganate (VII) = 23.4cm3Calculate the percentage purity of the impure sample of FeSO4.7H2O
93.30%
229
What are the redox equations for the iodine/thiosulphate titrations?
I2 + 2e- 2I-2S2O32- S4O62- + 2e-I2 + 2S2O32- 2I- + S4O62-
230
What colour is iodine?
Iodine is a yellow brown solution
231
What colour is iodide?
Iodide is a straw colour
232
How do you know when an iodine/thiosulphate titration is complete?
Iodine is a yellow brown solutionIodide is a straw colourThe above colour change is the indication for the redox reaction to be complete
233
Is an anode positive or negative?
Positive
234
Is a cathode positive or negative?
Negative
235
Draw a general cell
check notes
236
What is a cell? What is it made of? What does it do?
- A cell has two half cells- The two half cells must be connected with a salt bridge- Simple half cells will consist of a metal (acts as an electrode) and a solution of a compound containing that metalo For example, Cu and CuSO4- These two half cells will produce a small voltage if connected into a circuit (become a battery or cell)
237
Draw the standard hydrogen half-cell
Check notes
238
What is the standard hydrogen half-cell?
- H+/H2 half-cell is chosen to produce standard potentials- H+ + e- ½H2 - 1/2H2 H+ + e-- Conditionso 298Ko 1atmo 1M H+- Both equations have a voltage of 0.0 volts
239
Why is platinum used in the standard hydrogen half-cell?
Platinum is inert and so does not take part in the reaction
240
How can you find the value of any electron potential?
If the standard hydrogen half-cell is connected to another cell, you can find the value of any electron potential (as the standard is0.0 volts).
241
Why a salt bridge? Using potassium chloride and copper as an example
- The salt bridge is used to connect up the circuit- The free moving ions conduct the charge- A salt bridge is usually made from a piece of filter paper (or material) soaked in a salt solution, usually potassium nitrate- The salt should be unreactive with the electrodes and electrode solutionso For example, potassium chloride would not be suitable for copper systems as chloride ions can form complexes with copper ions, a wire is not used because the metal wire would set up its own electrode system with the solutions
242
What is standard electrode potential?
The emf of a half cell compared with a standard hydrogen electrode at standard conditions of 298K, concentration 1.0M and 1 atm of pressure.
243
What is emf?
The emf is the difference in the positive and negative sides of the cell.
244
Describe a two ion cell system?
Both of the ions in the equation will be at a concentration of 1M, platinum electrode is used as it is INERT.2 aqueous ions in a half equation, means they both have to be 1.0M solutions in the same half cell
245
Draw a manganese/lead system
check notes
246
General properties of d-block elements
The d-block elements have high melting and boiling points.The d-block elements are good conductors of both electricity and heat.
247
What elements have been used in coinage for many years?
copper, silver, nickel, and zinc
248
What element is used extensively in construction and production of tools?
Iron
249
What element is used for electrical cables and water pipes?
Copper
250
Titanium has great strength, what are the elements applications?
It has many aerospace and medical applications (for example joint replacement).
251
Sc orbital configuration
[Ar] 4s2 3d1
252
Fe orbital configuration
Fe – [Ar] 4s2 3d6
253
Ni orbital configuration
Ni – [Ar] 4s2 3d8
254
Zn orbital configuration
Zn – [Ar] 4s2 3d10
255
Cr orbital configuration
Cr – [Ar] 4s1 3d5
256
Cu orbital configuration
Cu – [Ar] 4s1 3d10
257
How do copper and chromium minimize repulsions?
Copper and chromium minimize repulsions by being half full or full, chromium has 4s and 3d orbitals half full, copper has 3d full
258
Fe2+ orbital configuration
Fe2+ - [Ar] 4s0 3d6
259
Fe3+ orbital configuration
Fe3+ - [Ar] 4s0 3d5
260
Cu2+ orbital configuration
Cu2+ - [Ar] 4s0 3d9
261
Cr3+ orbital configuration
Cr3+ - [Ar] 4s0 3d3
262
Mn2+ orbital configuration
Mn2+ - [Ar] 4s0 3d5
263
Mn4+ orbital configuration
Mn4+ - [Ar] 4s0 3d3
264
Sc3+ orbital configuration
Sc3+ - [Ar] 4s0 3d0
265
Zn2+ orbital configuration
Zn2+ - [Ar] 4s0 3d10
266
Define a transition element
A transition element is a d-block element that forms at least one ion with an incomplete d sub-shell.
267
Which 'd' block elements do not fit the transition element definition?
- Scandium and zinc
268
What are some characteristic properties of transition metals and there compounds, with examples. (not conduction, melting, boiling etc.)
- They form compounds in which the transition element has different oxidative stateso Fe2+ = +2o Fe3+ = +3- They form colored compounds- The elements and their compounds can act as catalystso Fe in Haber processo Ni in hydrogenation of alkenes
269
A species containing a transition element in its highest oxidation state is often a what?
strong oxidizing agent
270
What does the observed colour of a solution depend on?
The observed colour of a solution depends on the wavelengths absorbed
271
Why does copper sulphate solution appear blue?
Copper sulphate solution appears blue because the energy absorbed corresponds to red and yellow wavelengths, wavelengths corresponding to blue light aren’t absorbed.
272
Draw the exam colour chart
check notes
273
If the colour we observe is red, what colour has been absorbed?
Cyan
274
What colour is cu2+
white
275
What colour is [Cu(H2O)6]2+
blue
276
Transition metals form complex ions or ?
coordination compounds
277
WHat do transition metals form when they make complex ions?
ligand forms bonds with the central transition metal ion
278
Examples of complex ions
[Cr(H2O)6]3+ [CuCl4]2-
279
What is a ligand?
a molecule or ion that can donate a pair of electrons with the transition metal ion to form a coordinate bond
280
What does monodentate mean?
Monodentate ‘one tooth’ means each ligand donates just one pair of electrons.
281
Examples of ligands and how many loan pairs they have
H2O (oxygen has two loan pairs) NH3 (nitrogen has a loan pair) Cl- (chloride has a loan pair)CN- (carbon has a loan pair) OH- (oxygen ha a loan pair)
282
Draw [Cu(H2O)6]2+
check notes
283
Describe the structure and bonding in [Cu(H2O)6]2+
- The central ion is Cu2+- The ligands are water moleculeso Each molecule donates a pair of electrons from the O atom to the Cu2+ to form a co-ordinate bond- The co-ordination number is 6o This indicates the number of coordinate bonds to the central metal ion
284
What is the oxidation of Co in [Co(H2O)5Cl]+
2
285
How to name complex ions?
- The name gives the metal ions and its oxidation state last, and the name/number of ligands before- Pre-fixes di, tri, tetra, penta, hexa used- Ligands are listed alphabetically, with prefixes not allowed to alter this order
286
[Cr(H2O)4Cl2]+ name
¬¬Tetraaquadichlorochromium (III) ion
287
[Co(H2O)5Cl]+ name
Pentaaquamonochlorocobalt (II) ion
288
If the complex ion is an anion, what do you do?
the suffix ‘-ate’ follows the metal
289
Name of [Fe(CN)6]4-
Hexacyanoferrate (II) ion
290
What does cobalt become in an anion complex ion?
cobaltate
291
What does aluminium become in an anion complex ion?
aluminate
292
What does chromium become in an anion complex ion?
chromate
293
What does vanadium become in an anion complex ion?
vanadate
294
What does copper become in an anion complex ion?
cuprate
295
What does iron become in an anion complex ion?
ferrate
296
What does nickel become in an anion complex ion?
nickelate
297
What does the value of pKa show?
The higher the value of pKa, the weaker the acidThe lower the value of pKa, the stronger the acid
298
Equation for finding hydrogen ion concentration in buffer
[H+] = ka x ([HA]/[A-])
299
In an electrochemical cell, where does oxidation happen?
the half-cell with the most negative standard electrode potential
300
What is the relationship between electrode potential and metal reactivity?
More negative electrode potential = more reactive the metal
301
Which equation links E (cell), E(positive electrode), and E(negative electrode)
E (cell) = Eo (positive electrode) – Eo (negative electrode)
302
The more reactive a metal is, the more likely it is to what?
to lose electrons and form a positive ion
303
More reactive metals have more negative electrode potentials and vice versa for what?
Non-metals
304
in general if the electrode potential for the reaction is positive it is regarded as being what?
feasible
305
What is the problem with predicting the feasibility of electrode reactions?
- The value of E says something about the feasibility of the reaction under standard conditions only- The value of E says something about the feasibility of the reaction, but does not say anything about the rate of the reaction
306
What is the anticlockwise rule?
Arrange the redox half equations so that:- The electrons are on the left- The largest negative electrode potentials are at the top- The reaction will then take place anticlockwise around the half equations- E (cell) = E (bottom) – E(top)
307
When does predicting feasibility of electrode reactions go wrong?
When the conditions are not standard, this includes:- A change in concentration- If the half equations are in equilibriumIf the kinetics are not favorable:- The rate of reaction may be slow so reaction does not appear to happen- If the reaction has a high activation energy
308
Define catalyst
a substance that increases the rate of a chemical reaction by providing an alternative reaction pathway of lower activation energy (Ea)
309
What happens when Al3+, Ca2+, Mg2+ Add NaOH or KOH or NH4OH
white precipitate
310
Description of solution, observation with NaOH, and equation for Copper Cu2+
Transparent blue Pale blue precipitateInsoluble in excess Cu2+(aq) + 2OH-(aq) Cu(OH)2(s)
311
Description of solution, observation with NaOH, and equation for Iron(II) Fe2+
Pale green Dark green precipitateTurns brown on contact with airInsoluble in excess Fe2+¬(aq) + 2OH-(aq) Fe(OH)2(s)
312
Description of solution, observation with NaOH, and equation for Iron(III) Fe3+
Orange/brown Orange/brown precipitateInsoluble in excess Fe3+(aq) + 3OH-(aq) Fe(OH)3(s)
313
Description of solution, observation with NaOH, and equation for Chromium(III) Cr3+
Violet Grey-green precipitateSoluble in excess giving dark green solution Cr3+ + 3OH-(aq) Cr(OH)3(s)Cr(OH)3(s) + 3OH-(aq) [Cr(OH6]3-(aq)
314
Description of solution, observation with NaOH, and equation for Manganese(II) Mn2+
Pale pink Off white precipitateRapidly turning brown on contact with airInsoluble in excess Mn2+(aq) + 2OH-(aq) Mn(OH)2(s)
315
- Cu2+(aq) + 2e- Cu(s)If we increase the concentration of Cu2+ ions then:
- Equilibrium moves to oppose the charge- Electrons are removed from the system- The electrode potential becomes more positive
316
What shape are four coordinate complex ions usualy? with examples
- Tetrahedral is the most common shape- E.g. [CuCl4]2- and [CoCl4]2-
317
When not tertrahedral, what shape are 4 coordinate complexes, with examples
- Some 4 co-ordinate complex ions are square planar in shape, with the ligands arranged at the corners of a square- E.g. [Ni(NH3)2Cl2] (cis and trans)
318
When does a square planar shape take place in complex ions? example
- These occur in complexes with 8-d electrons in the d subshell.- E.g. Pt(II), Pd(II), Au(III)
319
What is cis platin used in?
- Cancer treatment in testicular cancer and useful for ovarian, head and neck, and lung cancer- Extremely toxic
320
Why did they change from cisplatin to carboplatin?
- Improved chemical stability relative to cisplatin due to chelation by cyclobutane dicarboxylic acid- Essentially equivalent antitumour activity to cisplatin
321
WHat is oaplatin used in?
treatment of colorectal cancer
322
Types of cancer therapy
- Surgery- Radiotherapy- Chemotherapyo Cytotoxico Targeted Anti-endocrine Novel targeted agents- Immuno-therapy- Gene therapy
323
What is mustard Gas? What came about because of mustard gas?
- Potent vesicant agent that burns eyes, skin and respiratory tractMustard Gas = war gasNitrogen Mustard = anticancer drug
324
How does cisplatin treat cancer?
- The cisplatin binds to DNA and causes a critical structural change n the DNA – a bend of 45 degrees- This stops cell replication and leads to apoptosis (cell death)
325
Key points of lactic acid
- 2-hydroxy propanoic acid- Also known as lactic acid- The second carbon is a chiral center- The mirror image is non-superimposableo One is found in sour milko The other is found in anaerobic respiration
326
What is a fuel cell?
a fuel cell is a device that converts chemical energy into electrical energy, water, and heat through electrochemical reactions.
327
How ddoes a fuel cell work? Are there many or one usually and why?
• Fuel and air react when they come into contact through a porous membrane (electrolyte) which separates them• This reaction results in a transfer of electrons and ions across the electrolyte from the anode to the cathode• If an external load is attached to this arrangement, a complete circuit is formed and a voltage is generated from the flow of electrical currentThe voltage generated by a single cell is typically rather small (< 1 volt), so many cells are connected in series to create a useful voltage.
328
Differences between a fuel cell and a battery
Hydrogen Fuel Cell – • Open system• Anode and cathode are gases in contact with a platinum catalyst• Reactants are externally supplied, no recharging requiredGalvanic Cell (Battery) – • Closed system• Anode and cathode are metals• Reactants are internally consumed, need periodic recharging
329
Differences between a fuel cell and an internal combustion engine
Fuel Cell – • Output is electrical work• Fuel and oxidant react electrochemically• Little to no pollution producedInternal Combustion Engine – • Output is mechanical work• Fuel and oxidant react combustively• Use of fossil fuels can produce significant pollution
330
Similarities between a fuel cell and in internal combustion engine
• Both use hydrogen-rich fuel• Both use compressed air as the oxidant• Both require cooling
331
Draw an alkaline fuel cell
check notes
332
What are the half equations, volatages, cell potential, and overall equation for an alkaline fuel cell?
Half Equations – 2H20 (l) + 2e- H2 (g) + 2 OH- (aq) E = -0.83V½ O2 (g) + 2e- 2 OH- (aq) E = +0.40VCell Potential – = 0.4- -0.83 = 1.23VOverall Equation – H2 + ½ O2 H2O
333
Why Methanol not Hydrogen?
Some new fuel cells use methanol rather than hydrogen as the fuel because- Liquid methanol is easier to store then hydrogen gas- Methanol can be generated from biomass
334
Define electrochemical reaction
A reaction involving the transfer of electrons from one chemical substance to another
335
Define electrode
An electrical terminal that conducts an electric current into or out of a fuel cell (where the electrochemical reaction occurs).
336
Define electrolyte
A chemical compound that conducts ions from one electrode to the other
337
What is an electrochemical cell consisted of?
An electrochemical cell consists of 2 electrodes + 1 electrolyte
338
What is rhe equation for the number of optical isomers
- The number of isomers follows the equation: 2n, where n is the number of chiral centres
339
Why are optical isomers called such?
The isomers are called optical isomers as they can rotate plane-polarized light (light which only travels in one plane)
340
WHat makes a 6 co-ordinate compound cis or trans?
The bond angle between the two ligands which are different, e.g. 90' cis, 180' trans
341
What is a bidentate ligand? Example
Bidentate – ‘two tooth’ ligandsMost common is ethane-1, 2-diamine
342
Draw [Ni(NH2CH2CH2NH2)3]2+
Check notes
343
Describe and explain EDTA4-
A Hexadentate Ligand- EDTA4- has 6 lone pairs, each of which can form a co-ordinate bond- 1 EDTA ion reacts with 1 metal ion- Ethylenediaminetetraacetic acid (EDTA)
344
What is ligand substitution?
The addition of another ligand to a solution containing the aqua transition metal ion results in a substitution reaction
345
What occurs for ligand substitution?
- One or more ligands is exchanged for another- A change in colour of the solution is observed- Sometimes the complex ion changes shape/coordinate number
346
Colour of [Cu(H2O)6]2+, and its colour with dropwise and excess, ammonia and hydrochloric acid
Complex Ion [Cu(H2O)6]2+Pale blue solution Addition of Ammonia Dropwise: Pale blue precipitation of copper(II) hydroxideExcess: Blue precipitate redissolves, forming a deep blue solutionAddition of Concentrated HClDropwise: Begins to turn greenExcess: Begins to turn yellow
347
Colour of [Cr(H2O)6]3+, and its colour with dropwise and excess, ammonia
Complex ion [Cr(H2O)6]3+Violet solutionAmmonia Dropwise: Grey/green precipitateExcess: Precipitate redissolves to produce a purple solution
348
Equation for addition of NaOH to [Cu(H2O)6]2+
[Cu(H2O)6]2+ + 2OH- [Cu(H2O)4(OH)2]2+ + 2H2O
349
Equation for addition of NH3 to [Cu(H2O)6]2+ in dropwise and excess
[Cu(H2O)6]2+ + 2NH3 [Cu(H2O)4(OH)2]2+ + 2NH4+[Cu(H2O)6]2+ + 4NH3 [Cu(NH3)4(H2O)2]2+ + 4H2O
350
Equation for addition of HCl to [Cu(H2O)6]2+
[Cu(H2O)6]2+ + 4Cl- [CuCl4]2- + 6H2O
351
Equation for addition of ammonia to [Cr(H2O)6]3+
[Cr(H2O)6]3+ + 6NH3 [Cr(NH3)6]3+ + 6H2O
352
What is the stability constant?
Kstab – the equilibrium constant existing between a transition metal ion surrounded by water ligands and the complex formed when the same ion has undergone a ligand substitution reactionLike KC but for equilibrias including complex ions
353
Colour of Sc3+
Colourless
354
Colour of Ti2+
colourless
355
Colour of Ti3+
lilac
356
Colour of Ti+4
colourless
357
Colour of Ti5+
colourless
358
Colour of V2+
lilac
359
Colour of V3+
green
360
Colour of V4+
blue
361
Colour of V5+
yellow
362
Colour of Cr2+
blue
363
Colour of Cr3+
green
364
Colour of Cr4+
colourless
365
Colour of Cr5+
colourless
366
Colour of Cr6+
orange
367
Colour of Mn2+
pale pink
368
Colour of Mn+3
colourless
369
Colour of Mn4+
dark purple
370
Colour of Mn5+
colourless
371
Colour of Mn6+
green
372
Colour of Mn7+
lilac
373
Colour of Fe2+
pale green
374
Colour of Fe3+
pale yellow
375
Colour of Fe4+
colourless
376
Colour of Fe5+
colourless
377
Colour of Fe6+
colourless
378
Colour of Co2+
Pink
379
Colour of Co3+
green
380
Colour of Co4+
colourless
381
Colour of Co5+
colourless
382
Colour of Ni2+
green
383
Colour of Ni3+
colorless
384
Colour of Ni4+
colourless
385
Colour of Cu1+
colourless
386
Colour of Cu2+
blue
387
Colour of Cu3+
colourless
388
Colour of Zn2+
colourless
389
[Cu(H2O)6]2+ colour
blue solution
390
Cu(OH)2 colour
Blue precipitate
391
[Cu(NH3)4(H2O)2]2+ colour
Deep blue solution
392
[CuCl4]2- colour
Yellow solution
393
HCl + Cu2+(aq) reaction colour and why
the complex can look green as the reaction is reversible so both blue and yellow species present
394
[Fe(H2O)6]2+ colour
pale green solution
395
Fe(OH)2 colour
Green precipitate, if left in air a reddy brown colour appears (Fe2+ oxidizes to Fe3+)
396
[Fe(H2O)6]3+ colour
Yellow solution
397
Fe(OH)3 colour
Reddy brown precipitate
398
[Mn(H2O)6]2+ colour
Very pale solution
399
Mn(OH)2 colour
Light brown precipitate which darkens in air
400
[Cr(H2O)6]3+ colour
Violet solution
401
Cr(OH)3 colour
Grey/green precipitate
402
[Cr(OH)6]3- colour
Green solution
403
[Cr(NH3)6]3+
Purple solution
404
[Cu(H2O)6]2+ + 4NH3
[Cu(H2O)6]2+ + 4NH3 [Cu(NH3)4(H2O)2]2+ + 4H2OBlue solution Deep blue solution
405
[Co(H¬2O)6]2+ + 6NH3
[Co(H¬2O)6]2+ + 6NH3 [Co(NH3)6]2+ + 6H2O
406
[Cu(H2O)6]2+ + 4Cl-
[Cu(H2O)6]2+ + 4Cl- [CuCl4]2- + 6H2OBlue solution Yellow/green solution
407
[Co(H2O)6]2+ + 4Cl-
[Co(H2O)6]2+ + 4Cl- [CoCl4]2- + 6H2OPink solution Blue solution
408
Cu2+ + 2OH- Cu(OH)2 colours
Cu2+ + 2OH- Cu(OH)2Blue solution blue precipitate
409
Mn2+ + 2OH- Mn(OH)2 colours
Mn2+ + 2OH- Mn(OH)2Very pale pink solution pale brown precipitate
410
Fe2+ + 2OH- Fe(OH)2 COLOURs
Fe2+ + 2OH- Fe(OH)2Green solution green precipitate
411
Fe3+ + 3OH- Fe(OH)3 colours
Fe3+ + 3OH- Fe(OH)3Yellow/brown solution brown precipitate
412
Cr3+ + 3OH- Cr(OH)3 colours
Cr3+ + 3OH- Cr(OH)3Green solution green precipitate
413
[Mn(H2O)6]2+ + 2NH3
[Mn(H2O)6]2+ + 2NH3 Mn(H2O)4(OH)2 + 2NH4+
414
[Fe(H2O)6]3+ + 3NH3
[Fe(H2O)6]3+ + 3NH3 Mn(H2O)3(OH)3 + 3NH4+
415
Cr(H2O)3(OH)3 + 3OH-
Cr(H2O)3(OH)3 + 3OH- [Cr(OH)6]3- + 3H2OGreen precipitate green solution
416
Cr(H2O)3(OH)3 + 3H+
Cr(H2O)3(OH)3 + 3H+ [Cr(H2O)6]3+Green precipitate green solution
417
Cr(OH)3(H2O)3 + 6NH3
Cr(OH)3(H2O)3 + 6NH3 [Cr(NH3)6]3+ + 3H2O + 3OH-Green precipitate purple solution
418
Cu(OH)2(H2O)4 + 4NH3
Cu(OH)2(H2O)4 + 4NH3 [Cu(NH3)4(H2O)2]2+ + 2H2O + 2OH-Blue precipitate deep blue solution
