Random Stuff Flashcards
(135 cards)
1
Q
Enthalpy of formation
Write equation for formation of Na2O
A
Enthalpy when one mole of a substance is formed from its constituent elements with all substances in their standard states
2Na(s) + 1/2O2(g) -> Na2O (s)
Mostly exothermic
2
Q
Enthalpy of combustion
Combustion of H2
A
Enthalpy change when 1 mole of a substance undergoes complete combustion in oxygen with all substances in their standard states
H2(g) + 1/2O2(g) -> H2O (g)
3
Q
First Ionisation enthalpy
A
Enthalpy change when 1 mole of gaseous atoms loses 1 electron per atom to produce 1 mole of gaseous 1+ ions
Endothermic
4
Q
Second ionisation energy is
A
Endothermic
5
Q
First electron affinity
A
Enthalpy change when 1 mole of gaseous atoms gains 2 electron per atom to produce 1 mole of gaseous 1- ions
Exothermic
6
Q
Secon electron affinity is
A
Endothermic cus adding -ve to -ve
7
Q
Enthalpy of atomisation
A
Enthalpy change when 1 mole of gaseous atom si produces from an element in its normal state
Endothermic
1/2I2 (s) -> I(g)
8
Q
A
9
Q
Hydration enthalpy
Do for Mg2+ (g)
A
Enthalpy change when 1 mole of gaseous ions becomes hydrated/dissolved in water
Exothermic
Mg2+ (g) + aq -> Mg2+ (aq)
10
Q
Enthalpy of solution
Equation for MgCl2 (s)
A
Enthalpy change when one mole of an ionic solid dissolves in an amount of water large enough so that the dissolved ions are well separated and do not interact with eachother
MgCl2(s) + aq -> Mg2+ 9aq) + 2Cl- (aq)
11
Q
Bond dissociation enthalpy
A
Enthalpy change when one mole of covalent bonds i broken in the gaseous state
I2(g) -> 2I(g)
Endothermic
12
Q
Lattice enthalpy of formation
Equation for MgCl2
A
Enthalpy change when 1 mole of a solid ionic compound is formed from It’s constituent ions in the gas phase
Mg2+ (g) + 2Cl- (g) -> MgCL2 (s)
13
Q
Lattice enthalpy of dissociation
Equation for MgCL2
A
Enthalpy change when 1 mole of a solid ionic compound is broken up into its constituent ions in the gas phase
MGCl2 (s) ->Mg2+ (g) + 2Cl- (g)
14
Q
Angle in non linear
A
104.5
15
Q
Angle in pyramidal
A
107
16
Q
Angle in tetrahedral
A
109.5
17
Q
Angle in trigonal bipyramidal
A
90 and 120
18
Q
Angle in octahedral
A
90
19
Q
Grou 2 hydroxides solubility down the group
A
More soluble
20
Q
Mg(OH)2 solubility
A
Sparingly soluble
21
Q
Ba(OH)2 solubility
A
Soluble
22
Q
Mg with water
A
Slowly in cold
Mg + 2H2O -> Mg(OH)2 + H2
Heated with steam reacts vigorously
Mg +H2O -> MgO + H2
23
Q
What is used to extract titanium from TiCl4
A
Magnesium
24
Q
Equation for extraction Ti from TiCl4
A
TiCL4(g) + 2MG (l) -> Ti (s) + 2MgCl2 (l)
25
Other use of magnesium
Milk of magnesia
To neutralise excess stomach acid
26
Use of Ca(OH)2
Slaked lime to raise pH of the soil
27
What needs to be removed from flute gases
SO2
28
What is used to remove SO2 from flue gases
CaO or CaCO3
29
How to remove SO2 using CaO
CaO (s) + 2H2O (l) + SO2(g) + 1/2 O2 (g) -> CaSO4.2H2O (s)
30
How to remove SO2 using CaCO3
CaCO3 + 1/2O2 (g) + SO2(g) -> CaSO4 (s) + CO2
31
How to test for sulfate ions
Acidify with HNO3
If there's effervescence, CO2 gas, so CO3 2- ions were present
Add BaCl2
This forms white ppt
32
BaSO4 use irl
Barium meal to outline intestines for X ray
33
Chemical used for testing of halides
Nitric acid to remove CO3 2-
Then AgNO3
Then add dil and conc ammonia
34
NaCl and H2SO4
What is sulfuric acid here
H2So4 + NaCl -> HCl + NaHSO4
HCl white fumes
Proton acceptor
35
NaBr + H2SO4
H2SO4 + NaBr -> HBr + NaHSO4
2HBr + H2SO4 -> Br2 + SO2 + 2H2O
Or
2H2SO4 + 2NaBr -> Na2SO4 + SO2 + Br2 + 2H2O
Orange fumes
Chocking gas
White fumes
36
NaI + H2SO4
H2SO4 (l) + NaI (s) → HI (g) + NaHSO4 (s)
H2SO4 (l) + 2HI (g) → I2 (s) + SO2 (g) + 2H2O (l)
H2SO4 (l) + 6HI (g) → 3I2 (s) + S (s) + 4H2O (l)
H2SO4 (l) + 8HI (g) → 4I2 (s) + H2S (g) + 4H2O (l)
White fumes
Chocking gas (SO2)
Purple vapour (I2)
Yellow solid of
37
SO3 + KOH
-> KHSO4
38
MgO + H3PO4
3MgO + 2H3PO4 -? Mg3(PO4)2 +3H2O
39
Chlorine + water
Cl2 + H2O -> HCl + HClO
40
Chlorine and water in sunlight
2CL2 + 2H2O => 4HCl + O2
41
Chlorine in water treatment
Makes HClO which kills bacteria
42
Chlorine + NaOH
Cold, dilute aqueous NaOH
CL2 + 2NaOH -> NaCl + NaClo + H2O
43
What's NaClO used for
NaClO is active ingredient in bleach
44
How to identify NH4+
(Ammonium ions)
Mix with warm dilute NaOH
Damp red litmus paper will turn blue cus ammonia is formed
45
How to test for group 2
NaOH
Or
H2SO4 and see if it dissolves
#solubility
46
How alkaline are group 2 hydroxides
Slightly
47
Test for OH-
Should turn red litmus paper blue
48
Test for CO3 2-
Add HNO3
Which produces CO2
Which can be passed through lime water which turns cloudy
49
Test for SO4 2- ions
Acidify with HNO3 to remove CO3 2-
Add BaCl2
This will form white ppt
50
Na and water
Vigorous exorthermic
Sodium floats on surface and rapidly melts
2Na + 2H2O + 2NaOH + H2
51
Mg + cold water
Mg + 2H2O -> Mg(OH)2 + H2
Slow
Small number of bubbles on Mg ribbon
52
Mg with steam
Mg + H2O -> MgO + H2
White bright flame
53
Na + O2
4Na + O2 -> 2Na2O
Bright yellow flame
White solid
54
Mg + oxygen
2Mg + O2 -> 2MgO
Bright white flame
White solid
55
Al + O2
4Al + 3O2 -> 2Al2O3
Bright white flame
White powder
56
Si + O2
Si + O2 -> SiO2
Bright white sparkles
White powder
57
What colour does P burn with
Yellow or white flame
Makes white clouds
58
S + O2
SO2
Blue flame
Toxic fumes
59
Which of the period 3 oxides are ionic
Sodium oxide
Magnesium oxide
Aluminium oxide
60
Boiling points of these
Na2O < Al2O3
61
Which period 3 oxides are giant covalent
SiO2
62
Which period 3 oxides are simple covalent
The rest
63
How does pH of the period 3 oxides change across the group
Decreases
13 -> 1
64
Which one is amphoteric
Al2O3
65
P4O10 with water
P4O10 + 6H2O -> 4H3PO4
Vigorous
PH 1
66
Structure of H3PO4
3 OH and one O with = bond
67
SO2 + H2O
H2SO3
Vigorous
68
SO3 + water
H2SO4
69
sodium oxide + HCl
Na2O (s) + 2HCl (aq) → 2NaCl (aq) + H2O (l)
70
Magnesium oxide + HCl
MgO (s) + 2HCl (aq) → MgCl2 (aq) + H2O (l)
71
Aluminium oxide + H2SO4
Al2O3 (s) + 3H2SO4 (aq) → Al2(SO4)3 (aq) + 3H2O (l)
72
Aluminiun oxide + NaOH
Al2O3 (s) + 2NaOH (aq) + 3H2O (l) → 2NaAl(OH)4 (aq
73
SiO2 + NaOH
SiO2 (s) + 2NaOH (aq) → Na2SiO3 (aq) + H2O (l)
74
P4O10 + NaOH
P4O10 (s) + 12NaOH → 4Na3PO4 + 6H2O (l)
75
SO2 + NaOH
SO2 (g) + 2NaOH (aq) → Na2SO3 (aq) + H2O (l)
76
SO3 + NaOH
SO3 (g) + 2NaOH (aq) → Na2SO4 (aq) + H2O (l)
77
Al2O3 + H2O
No reaction
Is insoluble in water
78
SiO2 + H2O
No reaction
79
Explain the behaviour of period 3 oxides with water
Metal oxides (sodium Mg) contain O 2- which is a strong base and will readily produce OH- through a reaction with water
Al2O3 doesnt react cus O 2- ions held too strongly in ionic lattice, so ions cant be separated
SiO2 giant covalent so doesnt react
The rest will react to make acidic solutions
80
Is Na2O or MgO more alkaline
Na2O cus its more soluble
81
Ligand
Molecule or ion that forms a coordinate bond with a transition metal by donating a pair of electrons
82
Colour of [CuCl4]2-
Green
83
What can a half cell be composed of
A metal dipped in its ions
Or
A Pt electrode with 2 aqueous ions
84
Which way do electrons flow in terms of reactivity
From a more reactive to a less reactive
85
What does the salt bridge have
KNO3
86
Which one gets reduced
Most positive
So most negative half cell will be oxidised
NO PRoblem
86
87
88
Standard hydrogen electrode
1 Moldm-3 H+ on left with Pt electrode
H2 gas going in at 298 K,
On the right: 1moldm-3 Cu2+
With copper electrode
89
What if u wanted to use a diprotic acid for the H+
Half the conc
90
Another way to calculate E cell
Reduced - oxidised
91
Cell notation
Most negative goes on the left
Reduced/oxidised form|| oxidised form | reduced form
92
3 types of cells
Non-rechargeable, rechargeable or fuel cells
93
Components of lithium ion battery
Electrode A - LiCoO2
Electrode B - graphite (C)
Electrolyte - lithium salt dissolved in an organic solvent
94
Electrolyte
Part of a battery that acts as a conductive pathway for ions to move from one electrode to another
95
Half equation at positive electode
Li+ + CoO2 + e- <-> Li+[CoO2]-
96
Half equation at negative electdoe
Li <--> Li+ + e-
97
98
Overall equation
Li + CoO2 <-> Li+[CoO2]-
99
How do rechargeable batteries work
Current is applied over cell
Electrons move in opposite direction
Potential difference is reversed
electrode reactions reverse by applying reverse potential
100
Fuel cell
1. Hydrogen gas fed in. This reacts with OH-
2H+ + 4OH- -> 4H2O + 4e-
Water is released on the left
2. Electrons produce in this travel through platinum electrode.
3. These electrons are used in the component to power smth like a engine
4. O2 is fed in on the right and reacts with water
O2+ H2O + 4e- -> 4OH-
The electrons come from the ones from before. Here they are on the cathode
These oh- ions are used to react with the new initial H+ on the left
5. Electrolyte carries OH- from the cathode to the anode (made of KOH) (right to left)
101
What are between the different thing
Ion exchange membranes
These line the Pt electrode and allow OH- ions to pass through but NOT H2 and H2O
102
Overall equation
2H2 + O2 -> 2H2O
103
Pros of fuel cells
- more effienct than internal combustion engines
- more energy is converted into kinetic energy
- combustion engines waste a lot as thermal energy
- many begin less are battery powered, but these dont need to be recharged, you only need a ready supply of O2 and H2
- only waste product is water, no CO2 emitted so better for environment
104
Cons of fuel cells
- H2 is highly flammable so it must be stored and transported correctly
- expensive to store and transport. Need to do it in pressurised containers
- energy is required to make the H2 and O2. Fossil fuels are used to pass water through an electrolysis process.
-
105
Definition of an electrochemical series
List of electrode potentials in numerical order
106
Why may different complex ions have different electrode potential even if the central metal ion si the same
Diff ligands
107
Why is an aqueous electrolyte not used for a lithium cell
Lithium would react with the electrolyte
E of Li+ more negative than E of water
108
Why may the value of a real cell be diff than calculated
Non standard conditions
109
It is difficult to ensure consistency with the setup of a standard hydrogen
electrode. A Cu2+(aq) / Cu(s) electrode (Eo = +0.34 V) can be used as a
secondary standard.
A student does an experiment to measure the standard electrode potential for
the TiO2+(aq) / Ti(s) electrode using the Cu2+(aq) / Cu(s) electrode as a
secondary standard.
A suitable solution containing the acidified TiO2+(aq) ion is formed when
titanium(IV) oxysulfate (TiOSO4) is dissolved in 0.50 mol dm−3 sulfuric acid to
make 50 cm3 of solution.
Page 4 of 16
AQA Chemistry A-Level - Electrode Potentials and Cells QP PhysicsAndMathsTutor.com
(b) Describe an experiment the student does to show that the standard
electrode potential for the TiO2+(aq) / Ti(s) electrode is −0.88 V
The student is provided with:
• the Cu2+(aq) / Cu(s) electrode set up ready to use
• solid titanium(IV) oxysulfate (Mr = 159.9)
• 0.50 mol dm−3 sulfuric acid
• a strip of titanium
• laboratory apparatus and chemicals.
Your answer should include details of:
• how to prepare the solution of acidified TiO2+(aq)
• how to connect the electrodes
• measurements taken
• how the measurements should be used to calculate the s
a) Weigh 7.995 / 8.00 g TiOSO4
(1b) Dissolve in / add (allow react with) (0.50 mol dm-3) sulfuric acid
(1c) transfer to volumetric flask and make up to the mark
Stage 2: Set up cell
Content can be shown in a labelled diagram
(2a) piece of Ti immersed in (1 mol dm−3 acidified) TiO2+(aq) / the solution
(2b) (connect solutions with) salt bridge or description
(2c) (connect metals through high R) voltmeter
Stage 3: Measurements and calculation
(3a) record voltage/potential difference/emf of the cell
(3b) Ecell = ERHS – ELHS
Ecell = Ecopper – Etitanium
(3c) ELHS = ERHS – Ecell OR Ecell should be +1.22 V if Cu on RHS (or −1.22 if
Cu electrode on LHS)
110
PKa in terms of Ka
Ka = 10 ^-pKa
111
What do u use to calculate strong acid pH or if its in excess
PH = log [H+]
112
What do u use if weak acid is in excess
Ka = [H+][A-]/[HA]
113
What do u use for strong base excess
Kw = [H+][OH-]
114
If ur reacting weak acid and strong base, and acid is in excess
What's A- equal to in the end
Moles of OH- added
115
C to K
+ 273
116
Units of Pressure in PV = nRT
Pascals
117
Enthalpy change
Heat energy change at constant pressure
118
Why is the enthalpy of hydration of Ca2+ less exothermic than that of Mg2+
Ca2+ less charge dense
So weaker attraction to O delta -
119
Define mass number
Number of protons and neutrons
120
Bonding in P4O10
VdW
121
Catalyst for Haber process
Iron
Heterogenous
122
How does the catalyst for Haber process catalyse
H2 and N2 reactants adsorb onto active sites of iron
Bonds weaken and reactions takes place
Products desirable from the surface
123
Why may a metal not catalyse a reaction
only 1 oxidation state
124
Why is a reaction slow before the catalyst is added
Two negative ions repeal eachother
So high activation energy is high
125
Why is [Fe(H2O)6]3+ have a lower pH than [Fe(H2O)6]2+
Fe3+ more charge dense
Fe3+ more polarising
So more OH bonds break and H+ ions released
126
Why is [H2O] not shown in Kw
[H2O] is almost constant
127
Why does Kw increase as temp increases
Forward reaction is endothermic
So equilibrium shifts to the RHS to oppose the temp increase
128
Why is the PH probe washed
Diff solutions must not contaminate eachother
To wash off any residual solution which could interfere with the reading
129
Why is vol of titre added more slowly near end of titration
To avoid missing en point
130
Why is something have hgiher entropy than other
It is more DISORDERED
131
Best oxidising agent in group 7
Cl2
132
Worst oxidising agent
I2
133
Best reducing agent
I-
134
Worst reducing agent
3Cl-
