Paper 1 Flashcards
1
Q
Relative mass of an electron
A
0.00055
2
Q
Relative atomic mass
A
Weighted average mass of an atom of an element divided by 1/12 of the mass of carbon 12
3
Q
Relative isotopic mass
A
The mass of an atom of that isotope divided by 1/12 of the mass of carbon twelve
4
Q
First ionisation energy
A
The energy required to remove one mole of electrons from one mile of atoms in its gaseous form
5
Q
What factors affect ionisation energy
A
- the more protons the more attracted the element is to the nucleus
- shielding - the electrons shielding the nucleus repel the valance electrons reducing the affective nuclear charge
- the greater the distance - lower the IE (which subshell it is being removed from)
6
Q
Reasons for the trend in ionisation energy across a period
A
Atomic radius decreases because number of protons increases holding the electrons in more closely
The shielding is constant
7
Q
Reasons for the trend in ionisation energy down a period
A
Proton no is increasing, however so is the shielding, and the atomic radius
Therefore IE increases
8
Q
Evidence of electronic configuration
A
1) emission spectra - electrons are promoted from a ground state to a higher energy state. The electron drops back down and emits a photon with the energy of the band gap. Evidence of discrete quantum shells
2) successive ionisation energies provide evidence for quantum shells and also the group which they belong - big jump in successive ionisation energies occur after the number of electrons in outer shell is removed
3) discontinuities in IE across a group provide evidence of subshells
9
Q
What is an orbital
A
Region within an atom that can hold up to two electrons with opposite spins
10
Q
Shapes S and P
A
S- is spherical
P- infinity sign
Draw on axis
11
Q
Conditions for electrons to fill subhsells
A
Fill singly before pairing up
Two electrons in the same orbital must have opposite spin
12
Q
Order of electron orbitals being filled up
A
1s(2) 2s 2p(6) 3s 3p 4s 3d (10) 4p 5s
(Electronic configuration determines the chemical properties of an element)
13
Q
Where are the s p and d block
A
D transition metal area
P right
S left and hydrogen
14
Q
What is periodicity
A
Repeating pattern across different periods
Recurring trends that are shown in the properties of an element
15
Q
Trends in melting and boiling points of elements in periods 2 and 3
A
Group 2 - increase steadily from Li to C, dramatic fall between carbon and nitrogen then they decrease
Group 3 - increase from Na to Si (Mh and Al) are similar. Then decrease
Li to Carbon Silicon and Boron are higher because they are giant structures therefore the covalent bonds are very strong
Nitrogen to Florine are diatomic therefore only have weak IMFs. Neon is monatomic
16
Q
Ionic bonding
A
Strong electrostatic attraction between oppositely charged ions
17
Q
What affects the strength of ionic bonds
A
Force varies inversely with sum of ionic radii
Force of attraction is proportional to product of the charge
The geometry
18
Q
Trends in ionic radius down a group and for a set of isoelectric ions
A
Down a group: ions have more electronshells. Ions get larger. Radii increase
Isoelectronic: The greater the atomic number the smaller the radius
Therefore greater nuclear charge
19
Q
Covalent bond
A
Strong electrostatic attraction between two nuclei and the shared pair of electrons between them
20
Q
Relationship between bond strength and bond length
A
The longer the bond length the weaker the bond
21
Q
What determine the shape of a simple molecule or ion
A
The repulsion between the election pairs that surround a central atom
They repel to take up position of maximum separation- to minimise repulsion
Lone pairs repel more than bonding pairs
22
Q
Evidence for the existence of ions
A
Conduction of electricity- lattice breaks down when melted or dissolved, ions can move and carry charge. No electricity conducted in the solid
23
Q
Two bonds
No lone pair
A
Linear
180
Cl-Be-Cl
24
Q
Three bond pairs
A
Trigonal planar
120
BCl3 Ethene
Sulphur trioxide
25
Two bond pairs
| One lone pair
Bent 120
Sulphur dioxide
26
Four bond pairs
Tetrahedral
109.5
Wedges and slashed lines
Methane
(NH4)+
27
Three bond pairs
| One lone pair
Trigonal pyramidal
107
Ammonia
28
Two bond pairs two lone pairs
Bent
104.5
Water
29
Five bond pairs
Trigonal bi-pyramidal
120,90
PCl5
30
6 bond pairs
Octahedral
90
SF6
31
Electronegativity
Ability of an atom to attract the bonding electrons in covalent bonding
Ionic and convent bonding are the extremes of a continuum of bonding type and that electronegativity differences lead to a bond polarity in the bonds and molecules
More than 1.5 1 ionic
Less than - polar covalent
Zero- non polar covalent
32
Polar molecules
Even if bonds are polar
Symmetry in the molecule would mean that electron density is equally pulled in each direction - forces cancel eg BCl3 or linear, anything symmetrical
33
London forces
Instantaneous dipole/ induced dipole
The electron density in covalent molecules oscillate within their atomic obtains or covalent bond. This causes a temporary dipole. These attractions repel the electrons in neighbouring molecules and induce a dipole. These are constantly fluctuating
These are usually stronger than permeant dipole dipole attractions
34
Permanent dipole - dipole
Electronegativity differences lead to polarity. The charges attract neighbouring polar molecules
35
Hydrogen bonding
A type of IMF formed between a δ+ hydrogen and lone pair in δ- oxygen /fluorine /nitrogen in different molecules (180 bond angle)
Hydrogen is the smallest atom so there is no shielding electrons therefore it is a partially exposed nucleus
FON are small
High electronegativity difference leads to a strong bond
36
How many H bonds in
H20
NH3
HF
Water has two hydrogens and two lone pairs -> can form 4 hydrogen bonds
HF - two 2H bond
NH3 - one with N Lp and one with one hydrogens. The other two are wasted
FHF BOND ANGLE IS 180 HFH bond angle is 109.5
37
Melting and boiling temperature of water and alcohols
Higher due to strong Hydrogen bonds
Alcohols have a lower volatility than alkanes of a similar number of e
38
Density of ice
Ice contains interlocking rings of 6 molecules held together by hydrogen bonds, since the distance across a ring and a larger than approaching molecules, water is more dense than ice
Average distance the molecules are apart is larger in ice than water
39
Trends in boiling temperature of alkanes
Chain length increases, more electrons, more London forces. Higher bp
More branching, molecules not stacked, less London forces, lower bp
40
Trends in boiling temperatures of hydrogen halides
Bp increases from HCl-HI no of electrons increases more London forces. (Stronger then decrease in polarity)
HF is an anomaly as it is much higher. This is because it can hydrogen bond
41
Choice of solvents
* water can dissolve some ionic compounds, energy released hydrating the ions out-ways energy required to separate ions
* water dissolved simple alcohols as they can hydrogen bond
* water is a poor solvent for some compounds (including polar halogenoalknes) as they cannot hydrogen bond -> use non aqueous solvents which have similar intermolecular forces to those in the solvent)
42
Metallic bonding
Strong electrostatic attraction between metal ions and delocalised electrons
43
Where are giant lattices present
* ionic solids (giant ionic lattices)
* covalently bonded solids (diamond, graphite, silicon(IV) oxide) (giant covalent lattices)
* solid metals (giant metallic lattices)
(Other covalent bonds such as I2, H20, is simple molecular)
44
Graphite
Layered hexagonal rings
Three sigma bonds to other carbons per carbon
Forth electron delocalised above and below the plane of the rings
Which are bonded by London forces
GRAPHENE IS ONE ATOM THICK GRAPHITE
45
Diamond structure
Four sigma bonds to carbons per carbon
3D tetrahedral
Bond angle 109.5
Silicon has a similar structure
46
Trend in reactivity down group 2
Reactivity increase
IE decreases -> group two reacts by loosing electrons
47
Group two plus oxygen
2Ca + O2 -> 2CaO
Burn in air
48
Group two plus chlorine
Ca + Cl2 -> CaCl2
When heated
These ionic chlorides can dissolve in water producing solutions that contain the hydrated cations
49
Magnesium and water
COLD
Mg +2H20 -> Mg(OH)2 + H2
HOT (steam)
Mg + H20 (g) -> MgO + H2
50
Group two plus water
Magnesium different in heat
Ca + 2H20 -> Ca(OH)2 + H2
Reacts rapidly in cold water
51
Group 2 oxides with water
MgO + H20 -> Mg(OH)2 (s) (slow)
React in same way
But fast
Ca(OH)2 ⇌ Ca^2+ + 2OH^-
Lime water
SrO and Ba are soluable
Ba^2+ and 2OH-
52
Group 2 oxides with dilute acid
MO + 2H+ -> M^2+ + H20
With a sulphuric acid can form MSO4
Or nitrate with nitric acid
53
Hydroxides with dilute acid
M(OH)2 + 2H+ -> M^2+ + 2H20
Forms the salts
54
Solubility of group 2 hydroxides
Increases down a group
55
Solubility of group two sulphate
Decreases down the group
56
Trend in thermal stability of nitrates and carbonates of groups one and two
Depends on polarising power of the cation. And polarisability of anion
Compounds containing highly polarising cations decompose more easily. Charge same down a group , radius increases therefore decomposes less easily down a group. Group 2 is more charge dense, larger and smaller radius, decomposes more easily
Nitrates all decompose in group two, only Li in group one
Carbonated all decompose
57
Decomposition of group one nitrates
Only lithium
4LiNO3 -> 2Li2O +4NO2 + O2
all Decompose on strong heating in a different way
They melt
2NaNO3 -> 2NaNO2 + O2
58
Decomposition of group two nitrates
2Mg(NO3)2 -> 2MgO + 4NO2 + O2
All of them decompose this way
On heating
59
Decomposition of group one carbonates
Li2CO3 -> Li2O + CO2
Only lithium
60
Decomposition of group 2 carbonates
CaCO3 -> CaO + CO2
All on heating
61
Formation of flame test colours
Electrons absorb energy from the flame and are promoted, this is not stable so they drop down, releasing energy as light. The frequency depends on the energy gap E=Hf therefore the colour is different for each element
62
Li flame colour
Crimson
63
Na flame colour
Yellow
64
K flame colour
Lilac
65
Mg flame colour
No colour
66
Ca flame colour
Yellow red
67
Sr flame colour
Red
68
Ba flame colour
Apple green
69
Experimental procedure for thermal decomposition
* Place the same amount (no of moles) of each carbonate/nitrate in test tubes and fix a delivery tube
* light a Bunsen burner at a set distance away from the tube
* measure the time taken for a certain amount of gas to evolve
70
Experimental procedure for a flame test
Dip a nichrome wire in HCl on a watch glass. Place in hottest part of flame until no colour is observed.
dip again in HCl and then into the solid to be tested
Place in the hottest part of flame and observe the colour
* electron is promoted to excited state
* unstable so drops back down and emitted a photon of light
* band gaps are quantised -> specific colour seen for each element
71
F state at room temperature
Gas
Pale yellow
72
Cl state at room temperature
Gas
Green
73
Br state at room temperature
Volatile liquid
74
I state at room temperature
Solid which sublimes when heated
Grey black solid sublimed into violet vapour
75
Trend in electronegativity down group 7
The atom is larger, therefore attraction to bonding electrons decreases and therefore the electronegativity decreases down a group
76
Trend in reactivity of group 7
React by gaining electrons
There is a decrease in the exothermic electron affinity as the electron is not brought so close to the nucleus
Fluorine is an exception because the atom is so small that the repulsion between the incoming electron and the seven electrons reduced energy liberated energy
AE decreases down the group therefore reactivity decreases down a group
77
Cl Aq colour
Organic
Pale green
Pale yellow/green
78
Br aq colour and in organic colour
Orange
Orange
(Bromine liquid is brown)
79
I aq colour and organic colour
Pale brown but only slightly soluble
Forms I3^-1 Deep red brown in the presence of I2 and I^-
Violet
80
Redox reactions of halogens and halides in aq followed by organic
Chlorine gas would oxidise bromide and iodine to form Br2 + 2Cl^- therefore would turn orange or violet to green
Cl2(g) + 2Br- (aq) -> Br2 (aq) + 2Cl- (aq)
If iodine grey ppt of iodine forms
Bromine displaces iodine in the same way
Violet to orange
Grey ppt of iodine forms
Reactivity as oxidising agents decrease down a group
81
Halogens plus group 1 and2 metals
Heat
Na+1/2Cl2 -> NaCl
Mg + I2 -> MgI2
Or
2Fe + 3Cl2 -> 2FeCl3
Fe + I2 -> FeI2 (less poweful oxidising agent)
82
Disproportionation
A single species is simultaneously oxidised and reduced
83
Chlorine plus water
Application
Cl2 + H20 ⇌ H+ +Cl- + HOCl
Disproportionation
HOCl is a strong disinfectant - used in water treatment
Bromine reacts but position of eq more to left
Iodine does not react
84
Chlorine aq sodium hydroxide
COLD
Cl2 + 2OH- -> Cl- + H20 + ClO-
ClO- bleach
HOT
3Cl2 + 6OH- -> 5Cl- + 3H2O ClO3^-1
85
NaCl + H2SO4
NaHSO4 + HCl
Steamy fumes which turn blue litmus red and give white fumes with gaseous ammonia
No reduction
NaCl is a poor oxidising agent
86
KBr + H2SO4
HBr + HKS04
2HBr + H2SO4 -> SO2 + H20 +Br2
Brown Br2 gas reduction to plus four
87
KI + H2SO4
HI + HKSO4
6HI + H2SO4-> 3I2 + S + 4H20
Violet gas and yellow solid
Reduced to 0
Or
8HI + H2SO4-> 4I2 + H2S + 4H2O
Rotten egg smell
Reduce to -2
88
HCl and water
| And other hydrogen halides
HCl + H2O -> H3O+ + Cl-
All soluble to form acidic solutions
89
HCl plus Ammonia
| And other hydrogen halides
HCl(g) + NH3 (g) -> NH4Cl (s)
Reaction is a test for gaseous hydrogen halides, but does not distinguish between them
90
Test for carbonate ions and hydrogen carbonate ions
Use aqueous acid
CO3^2- +2H+ -> H20 + CO2
Or HCO3^-1 + H+ -> H20 + CO2
Turns limewater cloudy
Effervescence
91
Test for sulphate ions
SO4^2- + BaCl2 -> BaSO4 + 2Cl^-
Milky white ppt
92
Test for ammonium ions
Sodium hydroxide solution Warm
NH4+ + OH^- -> H20 + NH3
Gas observed which
Turns damp red litmus blue
When a glass rod is dipped in conc HCl, a white smoke of ammonia colourised observed
NH3(g) + HCl(g) -> NH4Cl(s)
93
Test for halide ions
Add nitric acid and silver nitrate ethanol solvent
Chloride white ppt
Bromide cream ppt
Iodine pale yellow ppt
Then add dil ammonia
Chloride soluble
Then add conc
Bromine in conc
Iodine never
Ag+ + Cl- -> AgCl (s)
Then AgCl (s) + 2NH3(aq) -> [Ag(NH3)2]+(aq) + Cl- (aq)
94
Atom economy
100 x Molar mass of desired produced/ sum of molar mass of all products
95
Standard conditions
100kPa
| Specified temp usually 298K
96
What is enthaply change
Heat energy change measured at a constant pressure
97
Enthalpy level diagrams
Activation energy is not shown in enthalpy level diagrams but is shown on reaction profiles
98
Standard enthalpy change of reaction
Enthalpy change when the number of moles as written react under STP
99
Standard enthalpy of formation
Enthalpy change when one mole of a substance is formed from its elements in their standard states under STP
100
Standard enthalpy change of combustion
Enthalpy change when one mole of a substance is burnt in excess oxygen under STP
101
Standard enthalpy change of neutralisation
Enthalpy change when one mole of water is produced by the neutralisation of a solution of acid by excess base under STP
102
Errors in measuring temperature change
• Heat loss - insulate container
Heat loss is greater for slow reactions
•Therefore extrapolate the graph backwards to time when Bunsen was lit
• heat absorbed by container or thermometer - use a copper calorimeter if heating
•if not using then use a polystyrene cup to insulate
•use power rather than lumps to ensure it all reacts
•stir
• incomplete combustion
•alcohol evaporates
•water vapour not liquid produced
•some heats up the air
•non standard conditions
•beaker absorbs some heat
103
Assumptions in measuring temperature change
* density of solution = density of water
* specific heat capacity of solution = that of water
* negligible heat loss
104
Hess Law
The enthalpy change for any reaction is independent of the route taken from the reactants to products
105
Bond enthalpy
Enthalpy change when one mole of a bond in a gaseous molecule is broken
106
Mean bond enthalpy
Average enthalpy change to break one mole of a bond of that type over a wide variety of elements in their gaseous state
When doing calculations bare in mind a limitation is that it’s not the gaseous state
107
Conditions for dynamic equilibrium
* The rate of forward reaction is equal to the rate of backward reaction
* the concentration of reactants and products must remain constant
108
Effect of change of pressure
if pressure increases Eq shifts to the side with less gaseous molecules
No effect on Kp
109
Effect of change of concentration
The position of eq will shift to the right if conc products decreases and to the left if conc reactants decreases
No effect on Kc
110
Effect of temperature on equilibrium
Exothermic
T increases equilibrium shifts to the left
k down
endothermic
t increases equilibrium to the right
K up
Explained by change in equilibrium constant
Always increases date of rescuing equilibrium
111
kp
First find mole fraction
Then times by pressure to find partial pressure
Same as Kc but with partial pressures
(Use dimensional analysis to find units)
Kp and Kc are not affected by a catalyst
112
Brønsted Lowry bases and acids
Acid -> proton donor
| Base -> proton acceptor
113
What is the difference between a strong acid and a weak acid
Strong acid is totally ionised in aqueous solution forming hydrated hydrogen ions H30+
A weak acid is only very slightly ionised in aqueous solution <10%
114
Assumptions for weak acid calculations
[H+]=[A-] there is no other source of A-
[HA]eq = [HA]inital
Ka = x^2/[HA]
115
Calculation of PH of a strong base
pH + pOH =14
Or [H+][OH-] = Kw
Remember to account for multiple OH in same compound
116
pKa
PKw
- logKa
| - logKw
117
How much does the pH change for a tenfold dilution
1 unit strong acid
| 1/2 weak
118
features of a strong acid strong base titration curve
Vertical range
3-11
Equivalence point ph=7
(Calc vol equivalence point and final and initial pH)
119
Features of a weak acid strong base titration curve
Vertical range 7-11
Equivalence point pH 9
Initial buffer region
Calc vol equivalence point and initial and final pH
120
Features of a strong acid weak base titration curve
Vertical range 3-7
Equivalence point pH 5
Calc vol of equivalence point and final and initial pH
121
Buffer solution
A buffer solution is one that resists a change in pH when a small amount of acid or base is added.
It consists of a weak acid and its conjugate base in similar concentrations, less than factor of ten but more than 0.5 mol difference was
Or consists of a weak base and it’s conjugate acid
122
Action of a buffer solutions
* salt totally ionised
* acid partially ionised (suppressed by A- ions)
The reservoirs are large relative to the added H+ or OH-
When a small amount of H+ is added H+ + A- -> HA therefore the conc of H+ doesn’t change much. Conc of A- decreases slightly but is large relatively and conc of HA increases slightly but they are insignificant changes
When a small amount of OH- is added
OH- + HA -> H20 + A-
Conc A- increased slightly and HA Decreases slightly but that is insignificant relatively ( drives eq to the right).
Hydrogen ion conc has not changed significantly so pH will not change greatest
123
Assumptions for calculating pH of a buffer
```
[A-] = [salt]
[HA] = weak acid
```
(Remember to account for used up acid if some is used producing salt)
Ka = [H+][salt]/[weak acid]
124
How to determine Ka from a weak acid-strong base titration curve
[HA]=[A-]. Therefore Ka = [H+]
Ka = 10^-pH
At half equivalence point
125
Why is there a difference in enthalpy of neutralisation values for strong and weak acids
Strong acid you only have one enthalpy change
For a weak acid the enthalpy of ionisation of the acid must be considered
The weaker the acid the more endothermic the ionisation
Therefore the enthalpy of neutralisation is more endothermic
126
How is the pH in our blood controlled
Carbonic acid in the blood and the conjugate base hydrogencarbonate act as a buffer
127
Lattice energy
Energy change when one mole of an ionic solid is formed from its gaseous ions
(Is always negative)
Provides a measure of ionic bond strength - higher for larger magnitude
MAGNITUDE which decreases for a large ionic radii and increases for a large charge
128
Enthalpy change of atomisation
Enthalpy change when one mole of gaseous atoms is made from an element in its standard state
129
Electron affinity
Energy change when one mole of gaseous ions is formed from one mole of gaseous atoms by the addition of one mole of electrons EXOTHEMRIC
second electron affinity is the energy change when an electron is added to a gaseous 1- ion to form a 2- ion ENDOTHERMIC
130
Born Haber cycle
Hess’s cycle for finding lattice energy
131
Comparing the theoretical lattice energy value with experimental value
The more different the values are the higher the degree of covalent bonding
132
Polarisation
A positive action exerts and attraction on electrons in the negative anion distorting the electron cloud and weakening bonds
133
When is polarisation the strongest
polarising Small cation and high charge
Polarisable Large anion and large charge
134
Enthalpy change of solution
Enthalpy change when one mole of the solid is dissolved in sufficient solvent to give an infinitely dilute solution
135
Enthalpy change of hydration
The enthalpy change when one mole of gaseous ions is dissolved to give an infinitely dilute solution
136
Evidence that enthalpy alone doesn’t control whether reactions occur
Some endothermic reactions can occur at room temperature
If a reaction produces a gas for example, and the entropy increases massively, then it will be spontaneous at room temp
137
What is entropy (chemistry)
Measure of disorder of a system - the natural direction of change is increasing total entropy
The second law of thermodynamics states that spontaneous changes result in an increase in disorder or entropy
138
Change of state
Effect on entropy
Or dissolving
Ordered crystal structure become dispersed through a liquid -> entropy increases
Solid more ordered than liquid than gases
139
Number of moles effect on entropy
More moles in products than reactants means that entropy increases
140
Entropy change in dissolving ammonium nitrate crystals
Dissolve therefore less ordered entropy increases
141
Reacting ethanoic acid with ammonium carbonate
CO2 gas produced
CO2 + 2ammoniumethanoate + h20
No mores increases 3-> 5
Entropy increases
142
Burning magnesium ribbon
Mg (s) + O2 (g) -> 2MgO
Less moles
gas to solid
Entropy decreases
143
Mixing solid barium hydroxide with solid ammonium chloride
Forms ammonia BlC2 and H20
Solid to liquid
Entropy increases
144
Total entropy
ΔS total = ΔS system + ΔS surroundings
Δs total = Rlnk
145
Entropy of surroundings
ΔS surroundings = -ΔH/T
146
Gibbs free energy
A measure of the chemical potential that a mole of a substance has when on its own G=H-TS
ΔG = ΔH - TΔS system
Feasible if ΔG is less than zero
ΔG = -RTlnK reactions which are feasible have large values for the equilibrium constant
Reactions are feasible if it’s negative as it’s losing potential. In equilibrium reactions, the position of eq is where the system has minimum free energy.
147
Why might a - ΔG value occur in practice
If the reaction is kinetically controlled (high activation energy)
Or departure from standard conditions such as ppt forming or conc acid
148
Standard electrode potential
Individual potential of a reversible electrode relative to the standard hydrogen electrode
•all ions in the concentration are at concentration of 1.0moldm^-3
•all gases are at a pressure of 100kPA (1.0atm)
•the system is at a stated temp usually at 298
Salt bridge
149
Features of a standard hydrogen electrode
Hydrogen gas at 100kPa pressure bubbling over a Platinum plate which is dipped into a solution that is 1.0moldm^-3 solution of Hydrogen ions (Eg HCl) at a temp of 298K
Salt bridge
Hydrogen should always be drawn on the left diagram or written in left in cell diagrams
150
How to set up an electrode where the substance is gaseous
Platinum plate dipping into 1.0 moldm^-3 of the solution of ions of the element with the gaseous element, at 100kPa bubbling over the surface of the platinum
151
Cell diagram
Anode|ions || ions | cathode
Eg
Pt(s)|MnO4^-1(aq), 8H+ (aq), Mn2+ (aq)|| Fe3+ (aq), Fe2+ (aq) | Pt(s)
152
Predicting the thermodynamically feasibility of a reaction using standard electrode potentials
Predict when Ecell > 0 thermodynamically feasible
153
Limitations of predictions made using standard electrode potentials
If it is kinetically controlled, Ea is too high, even thermodynamically spontaneous reactions may not occur
If a ppt is formed, it derives the equilibrium to the right, conditions are non standard
154
Electrochemical series
The higher the E value the better the oxidising agent
155
Methanol and hydrogen rich cells
Anode CH3OH + H20 ⇌ HCOOH + 4H+ + 4e-
Cathode 02 + 4H+ + 4e- ⇌ 2H2O
Energy released is utilised in a fuel cell to generate a voltage
156
Application of electrode potentials to storage cells
Rechargeable battery
two lead plates, one solid lead (IV oxide coating, sulphuric acid as the electrolyte
Anode Pb(s) + SO4^2- (aq) ⇌ PbSO4 (s) + 2e-
Cathode PbO2 + 4H+ + SO4^2- + 2e- ⇌ PbSO4 + 2H20
Overall
Pb + PbO2 + 4H+ + 2SO42- -> 2PbSO4 + 2H20
When all the lead oxide is reduced, the battery is flat.
Ecell 2V
Six cells 12V
When a reverse external voltage of 12V is applied the reaction is driven backwards and the cell is recharged
157
Hydrogen-oxygen fuel cell
Acidic electrolyte (H+ ions through solid polymer electrolyte)
Anode 2H2 ⇌ 4H+ + 4e-
Cathode O2 + 4H+ 4e- ⇌ 2H20
Alkaline (KOH electrolyte , porous platinum electrodes)
Anode 2H2 + 4OH- ⇌ 4H20 + 4e-
Cathode O2 + 2H20 + 4e- ⇌ 4OH-
Overall for both 2H2 + O2 -> 2H2O
Alkaline fuel cells have their efficiency reduced by carbon dioxide in the air and corrosive potassium hydroxide solution may leak. They are more efficient however
158
Iodine titration
Used to estimate conc oxidising agent
* known vol of oxidising agent from a pipette to excess KI (May need to acidify) in conical flask
* titrate the liberated I2 with standard solution sodium thiosulphate
* iodine faded to pale straw from brown add starch
* continue adding sodium thiosulphate until blue black starch colour disappears
END POINT COLOURLESS
Why don’t you add starch initially - forms irreversibly a blue black starch iodine complex
I2 + 2Na2S2O3 -> 2NaI + Na2S4O6
Repeat, concordant results
159
Application of electrode potentials to storage cells
Rechargeable battery
two lead plates, one solid lead (IV oxide coating, sulphuric acid as the electrolyte
Anode Pb + SO4^2- -> PbSO4 + 2e-
Cathode PbO2 + 4H+ + SO4^2- + 2e- -> PbSO4 + 2H20
Ecell 2V
Six cells 12V
When a reverse external voltage of 12V is applied the reaction is driven backwards and the cell is recharged
A storage cell must be reversible for both reactions
Oxidised and reduced forms of the anode and cathode must be solid
160
Hydrogen-oxygen fuel cell
Acidic electrolyte (H+ ions through porous polymer)
Anode 2H2 ⇌ 4H+ + 4e-
Cathode O2 + 4H+ 4e- ⇌ 2H20
Alkaline (KOH electrolyte, porous platinum electrode)
Anode 2H2 + 4OH- ⇌ 4H20 + 4e-
Cathode O2 + 2H20 + 4e- ⇌ 4OH-
•alkaline fuel cell efficiency is reduced by carbon dioxide in the air and corrosive potassium hydroxide solution may leak
Overall for both 2H2 + O2 -> 2H2O
161
Configuration of d-block elements Sn-Zn
[Αr]3dx 4s2
Other than chromium and copper due to half filled stability
[Ar] 3d5 4s1 - chromium
[Ar] 3d10 4s1 - copper
Electrons are lost from the 4s subshell before the 3d subshell
162
What is a transition metal
D-block elements that form one or more stable ions with partially filled d orbitals
Has one or more unpaired d electron in one of its ions
163
Why do transition metals show variable oxidation number
The successive ionisation energies increase steadily. At higher oxidation states, the election can be promoted just to unpair them for covalent bonding
164
Ligand
Ion or molecule which is bonded via a dative covalent bond to a central metal ion
165
What is a complex ion
A central metal ion surrounded by Luganda
166
Why do transition metals form coloured ions in solution
The ligands split the energy levels of the d-orbitals. Electronic transitions take placePhotons of a particular frequency are absorbed by electrons if they have the exact energy of the band gap. E=Hf. In the visible range This promotes the electrons You see the complementary colour of the absorbed light
167
Why is there a lack of colour in some aq ions and other complex ions
Ions with no d-electrons are not coloured, or with a full d subshell as there is either nothing to promote or nowhere for them to go. Transitions cannot take place
Or if not full, the band gap may be so large the light is UV not visible
168
What causes colour changes in transition metal ions
* coordination number - no of atoms bonded to the central metal ion changes the colour because the splitting of the d orbitals is different for an octahedral, tetrahedral and planar field
* oxidation number - ions with a higher charge density attract ligand more strongly - splitting of the d- orbitals is greater
* ligand - some ligands interact more strongly with the d subshells, causing a greater splitting, changing the colour
The more the splitting the more violet
169
What is the coordination number
No of atoms bonded to the central metal ion
170
Monodentate ligands
One lone pair therefore can form one dative bond eg h20 or OH-
171
Shape of 6 fold coordination number
Repeal to maximum separation to minimise repulsion
| Octahedral shape
172
Shape of four coordination number
Larger ligands May form tetrahedral complexes eg Cl-
109.5
Could also be square planar eg cisplatin
173
Use of transition metals in cancer treatment
Cisplatin, a platinum2+ ion complex with two ammonia and two chloride ions, is used as a single isomer the complex bonds to adjacent guanine molecules in one strand of DNA In cancer cells by ligand exchange preventing replication. The chloride ligand is repacked by a nitrogen atom. The cancer cell is then destroyed by the bodies immune system.
The transplatin form is ineffective because it is kinetically unstable and it is inaffective at bonding adjacently
Therefore it is supplied as a single isomer and not in a mixture with the trans form
174
Bidentete ligands And multi-dentate ligands
Bidentate have two lp and are sufficiently long to bend round
175
What is haemoglobin
An iron(II) complex containing a multidentate ligand
Planar
Ligand exchange occurs when an oxygen molecule bound to haemoglobin is replaced by a carbon monoxide molecule this is irreversible
176
Colour of vanadium 5+
Yellow VO2+ (colourless VO3-)
177
Colour of vanadium 4+
Blue
178
Colour of vanadium 3+
Green
179
Colour of vanadium 2+
Lavender
180
How do reduce dichromate (VI)
Cr2O7^2- reduced to Cr3+ and Cr2+ using zinc in 50% hydrochloride acid acidic conditions
Final colour blue
181
How to produce dichromate (VI)
Oxidation of Cr3+ ions using hydrogen peroxide in alkaline conditions followed by acidification
182
Equilibrium between Chromate and dichromate
2CrO4^2- + 2H+ ⇌ Cr2O7^2- + H2O
Convert between by altering pH to shift equilbirum
183
Transition metal plus aq NaOH
Deprotonation - number deprotonated is equal to charge on cation
In excess further deprotonation only for ampotheric hydroxides
Cr3+
184
Transition metal plus ammonia
Deprotonation
Excess
Some undergo ligand exchange
Cr3+ 6NH3 exchanged
Co2+ 6NH3 exchanged
Cu2+ 4NH3 exchanged
185
Cr3+ NaOH
(Grey) Green ppt
Excess
(Dark) green solution
186
Fe2+ NaOH
(Dark) green ppt. (Goes brown on exposure as oxidised to Iron(III))hydroxide)
Excess
Insoluable
187
Fe3+ NaOH
Red brown ppt
Excess
Insolvable
188
Cobalt + NaOH
Blue ppt (goes pink on standing)
Excess insoluable
189
Cu2+ NaOH
Blue ppt
Excess
Insoluable
190
Cr3+ aqueous
Green
191
Fe2+ aqueous
Pale green
192
Fe3+ aqueous
Yellow
193
Co2+ aqueous
Pink/red
194
Cu2+ aqueous
Turquoise blue
195
Cr3+ NH3
(Green) grey ppt
Excess
green solution forms slowly
196
Fe2+ NH3
(Dark )green ppt (again darkens in air)
Excess
Insoluable
197
Fe3+ NH3
Red brown ppt
Excess
Insoluable
198
Co2+ NH3
Blue ppt
Excess
brown solution
199
Cu2+ NH3
Blue ppt
Excess
Deep blue solution
200
Ligand exchange with Cl-
[CuCl4]2- from [Cu(H2O)6]2+
| [CoCl4]2- from [Co(H2O)6]2+
201
Ligand exchange affect on entropy and stability
Large positive increase in entropy of the system when a monodentate ligand is substituted a bidentate ligand or multidentate ligand leads to a more stable complex ion
202
Ammonia ligand exchange
[Cr(H20)6]3+ + 3NH3 ⇌ [Cr(H2O)3(OH)3] + 3NH4+
[Cr(H20)6]3+ + 6NH3 ⇌ [Cr(NH3)6]3+ + 6H20
As the ligand exchanges the conc of aq ion decreases, pulling top eq to the left, so the ppt disappears
203
V2O5 as a catalyst
Contact process makes H2SO4 but first makes SO3
Heterogenous
vanadium (V) oxide is the catalyst to convert sulphur dioxide to sulphur trioxide
SO2 (g) + V2O5 (s) + SO3 (g) + 2VO2 (s)
1/2O2 (g) +2VO2 (s) -> V2O5(s)
Therefore overall SO2 + 1/2O2 -> SO3
204
Amphoteric behaviour of complex ions
Acting as an acid
[Cr(H20)3(OH)] + 3OH- ⇌ [Cr(OH)6]3- + 3H20
Acting as a base
[Cr(H20)3(OH)3] + 3H20⇌ [Cr(H20)6]3+ + 3OH-
Cr(OH)3 + 3H+ -> Cr3+ + 3H20
Excess of strong base
Only Cr are amphoteric
205
Catalytic converter
Reduced carbon monoxide and nitrogen monoxide emissions by
• adsorbing the gases onto the surface of a catalyst
•weakening the bonds and chemical reaction
•desorption of the products
2CO + 2NO - 2CO2 + N2
Catalyst is platinum with rhodium on a ceramic base
206
Heterogenous catalyst
In a different phase from the reactants and the reaction occurs on the surface
207
Homogenous catalyst
In the same phase as the reactants and the reaction occurs via and intermediate species
208
Iron as a catalyst
Homogenous
2Fe^2+ + S2O8^2- -> 2Fe^3+ 2SO4^2-
2Fe^3+ + 2I- -> 2Fe2+ + I2
2I^-(aq) + S2O8^2- (aq) -> I2 + 2SO4^2-
209
Autocatalysis
2MnO4 ^- + 16H+ +5C2O4 ^2- -> 2Mn^2+ + 10CO2 + 8H20
Reaction speeds up as the Mn^2+ is produced as it catalyses the reaction
210
What affects the hydration enthalpy
High charge larger magnitude
| Less exothermic for larger radius (the greater the force)
211
What is a homogenous reaction
One in which all the reactants and products are in the same phase
212
What phase is two immiscible layers
Two phases
213
What phase is a mixture of solids
Usually two phase
214
Effect of addition of an inert gas equilibrium
Pressure is increased, there is no effect on the concentrations of the reactants and products. The no of moles of reacting species has not been altered and neither has volume.
The mole fraction decreases but the total pressure increases by the same factor. So the quotient is unaltered.
System is still in equilibrium
215
Effect of adding a catalyst to equilibrium
No affect on equilibrium constant
But the reaction is speed up so equilibrium is reached faster
Unless it is a gaseous reaction catalysed by a solid catalyst. The number of active sites on the surface are all already used up therefore no effect
216
Aims of industrial processes in equilibrium
* maximise yield
* make it quickly
* keep cost low
* have high atom economy
217
How can you improve yield in industrial processes
•Reactants are added continuously at one end and products removed at the other end- this is not an equilibrium system as it is not closed
218
What is a strong base
Something which is totally ionised in aqueous solution forming hydroxide ions OH-
A weak base is protonated to only a small degree in a solution and so only forms a small proportion of hydroxide ions
219
What is a neutral solution
[H+][OH-] are the same
Ph is not necessarily 7
220
pH of strong acid and salt of strong acid and strong base
7
221
pH of weak acid and salt of strong acid and strong base
5.1
222
pH of strong base and salt of weak acid and strong base
8.9
223
pH of weak base and salt of weak acid and weak base
7.0
224
How does an indicator work
HInd ⇌ H+ + Ind-
Kind = [H+][Ind-] / [Hind]
In acid eq is driven to each side and the colour changes
225
What affects lattice energy
* The magnitude of the charges on the ions
* The sum of the radius of the cation and anion
* The arrangement of the ions in the lattice
* The relative sizes of the ions
* The extent of the covalencey
226
What is the standard free energy of formation
The change in free energy that occurs when a compound is formed from its elements in their most thermodynamically stable stables under standard conditions
=sum of Gibbs of products - reactants
227
What’s ΔSsystem of dissolving a gas
Negative because the system is more ordered. So therefore it will dissolve exothermically the surroundings so that the surroundings become more disordered
Eq is driven to left by increase in temp so gases are less soluble in hot water than cold one
228
ΔS system of dissolving a solid
The solute always becomes more disordered, but the dilute can become more ordered due to the forces of attraction between solute and solvent
Eg high charge dense ions
Extent is determined by the total entropy change for one mole of a solid
Dissolving group 1 compounds - only a small change in entropy of water therefore Δsystem is always pos
Doubly charged cations causes a large decrease of entropy water so Δs system is negative
229
Why do hydroxides get more soluble down and group and sulphates get more soluable up the ground
* depends on hydration enthalpy which depends on charge density
* lattice energy depends on the charges of two ions multiplied together divided by the sum of the two ionic radio
* OH ion is small l, therefore the value of the sum of the radii increases considerably as the value of the radius of the cation increases
* the sulphate ion is much larger therefore the sum of radii is less significant than the decrease in magnitude of the hydration enthalpy. This makes enthalpy of solution less exothermic
230
Which is the anode and which is the cathode
The anode is where oxidation occurs
The cathode is where reduction occurs
Anions ions move to anode
Cations move to cathode
231
What does a salt bridge consist of
A conc solution of an inert electrolyte such as potassium nitrate or potassium chloride
232
Features of an electrode with two cations
Electrode consists of a platinum rod dipping into a solution containing 1 moldm^-3 of both ions, each with that conc
233
Relationship between equilibrium constant and ECell
lnK α Ecell
234
Potassium manganate titrations
Used to estimate conc reducing agent
* titrate a standard solution of potassium manganate (VII) with the solution of the reducing agent in acidic conditions
* puppetry a known vol of reducing agent into a conical flask and acidfiy with dilute sulphuric acid
* fill burette with standard solution of potassium manganate
* add and swirl until purple colour disappears
* then add dropwise until solution is slightly pink
END POINT PINK
Concordant titres
MnO4- + 8 H+ + 5e- -> Mn2+ + 4H20
With ion 1:5 ratio
5Fe2+ + MnO4-...
235
Why are scandium and zinc not transition metals
It’s only ion is Sc3+ which has no d electrons
Zinc does not form + or 3+ Zn2+ has 10 d electrons that are all paired in five full d orbitals
transition metal) forms an ion with incomplete d sub-shell
scandium and zinc are not transition metals
Sc3+ and 1s2 2s2 2p6 3s2 3p6
Zn2+ and 1s2 2s2 2p6 3s2 3p6 3d10
Sc3+ and d sub-shell empty / d-orbitals empty
Zn2+ and d sub-shell full / ALL d-orbitals are full
236
Trend in ionisation energy across period 4 from potassium to zinc
And successive ionisation energies
•nuclear charge increases, the no of shielding 3d electrons increases as well. The ionisation energies are all fairly similar
Big jump does not occur until after all the 4s and 3d electrons are removed because they are so close in energy
237
How to reduced vanadate (v) vanadium (II)
Warm with powdered zinc with ammonium vanadate in presence of 50% hydrochloride acid solution
238
Colour of Cr2+
Blue
239
Colour of Cr3+
Green
240
Colour of Cr in oxidation state of +6
CrO4 2-
Yellow
Cr2O7 2-
Orange
241
Melting temp of metals
Group one have low melting temps
Group two have higher as there are two electrons lost
Melting temp increases across a period because the metallic radius decreases and more electrons are released for bonding
It decreases down a group because metallic radius increases and the force of attraction between metal ions and delocalised electrons becomes less
242
Reaction of metals with less restive metal
Fe(s) + Cu^2+ aq -> Fe^2+ (aq) + Cu (s)
243
What is an oxidation number
The charge that the element would have if the compound were fully ionic
244
What is group one called
Alkali metals
245
What is group two called
Alkaline earth metals
246
Lime water with carbon dioxide
Ca(OH)2 (aq) + CO2 (g) -> CaCO3 (s) + H2O (l)
247
Test to distinguish between a carbonate and a hydrogen carbonate of group one
Calcium hydrogen carbonate is soluble in water
Calcium carbonate is not
Add a solution of the test substance to calcium chloride. A carbonate will form a white ppt of calcium carbonate
Hydrogen carbonate will form no ppt until mixture is heated, decomposing the hydrogencarbonate ions
248
Chlorine plus phosphorus
Limited supply of chlorine
2P + 3Cl2 -> 2PCl3
In excess
2P + 3I2 -> 2PCl5
If damp
PI3 + 3H20 -> 3HI + H3PO3
249
Boiling temps of hydrogen halides
Increases from HCl to HI due to more electrons therefore more London forced
HF is anomalous because it forms hydrogen bonds
250
Solubility of halides
All soluble except silver halides and lead II halides
251
Steps to carry out a titration
* rinse burette with distilled water and then with a little of the solution
* rinse conical flask with distilled water
* the burette is filled with solution and tap is opened so some solution is run out so that the stem below the tap does not contain air bubbles (record vol) bottom of meniscus
* fill pipette so bottom of meniscus is on line. Discharge into conical flask and and indicator
* the solution from the buffet added, swirl constantly, until change of colour
* read volume and calculate titre
Repeat for concordant results (within 0.2cm^3)
252
Examples of poor technique
| In titration stuff
* not rinsing solid from weighing bottle
* not rinsing stirrer and funnel into volumetric flask
* not shaking after making up to 250cm^3
* not rinsing out the burette and pipette with correct solutions
* not ensuring there is no air below tap in burette
* getting air bubbles in stem of pipette
* running in the solution from burette and overshooting
* not swirling the flask after each addition of solution from burette
253
Preparation of a standard solution
* weigh
* pour in beaker and wash our weighing bottle into beaker
* add water and stir until solid dissolved
* pour the solution through a funnel with water so the liquid goes in volumetric flask
* take washings of beaker and funnel
* make volume up to the mark with distilled water
* shake / invert funnel
254
Experimental method to find the enthalpy of combustion of a liquid
* spirit burner containing liquid is weighed
* a known volume of water is added to the copper calorimeter
* temp measured at regular intervals
* burner lit after 4.5 min
* when temp has reacted 20 degrees above room temp, the flame is extinguished and the burner immediately reweighed
255
When to include water in equilibrium constant
* gaseous state INCLUDE
* reactant but not solvent INCLUDE
* when solvent, even if reactant or product, DO NOT INCLUDE because the conc remains constant
256
When can the enthalpy change of reaction not be determined directly
When heat is supplied, you cannot distinguish the heat change due to the reaction
257
Formula of potassium dichromate
K2Cr2O7
258
Test for NO3-
Gives off brown gas of NO2 when heated in Biden and caused glowing splint to relight
259
Why may a student add a pinch of something to a reaction
If finding a volume, the pinch may saturate the solution to stop excess dissolving
260
Is CO2 soluable
Yes a bit
261
Is hydrogen soluable
No
262
Unit of entropy
JK^-1mol^-1
263
Cobalt chloride
Blue
264
Copper chloride
Yellow
265
Metal plus acid
Salt and hydrogen
266
Calculating ECell
Either Eoxidant - E reductant
Or right - left
267
Environmental issues behind hydrogen fuel cells
* hydrogen is produced by electrolysis, the electricity of this comes from fossil fuels
* however they are efficient
* ethanol could be used from fermenting sugar or cereals
* but this used agricultural land
268
Salt bridge
Allows ions to flow
Therefore cannot be unreactive metal wire
Must be inert as they must not form a ppt with the ions in the cells
Potassium nitrate or potassium chloride
269
Why would you want to use a high resistance voltmeter
Little reaction takes place therefore the conc of the ions remains approximately constant
270
What is an electrolyte
electrolyte is a chemical compound that dissociates into ions and hence is capable of transporting electric charge
271
Bonding in a transition metal
In 2+ and 3+ a lot of the time is ionic
+4 or higher covalently bonded
Can also be a complex ion in the lower states too
For each covalent bond to form, the element must have an unpaired electron in its valence shell. - 4s and 3d orbitals
In Mn04- an electron can be promoted from 4s to 4p - giving 7 unpaired electrons
272
Stereoisomerism in a transition metal
Cis-trans -> occurs in octahedral and square planar complexes. Adjacent are cis, opposite are trans
Optical -> bidenetate ligands
273
How to reduce to vanadium(II)
Warm a solution of powered zinc with ammonium vanadate In presence of 50% hydrochloride acid
Conical flask fitted with a Biden valve to exclude air
Final colour lavender
274
Chromium(II) chloride
[Cr(H2O)6(Cl)3]. Grey blue
[Cr(H20)5Cl]Cl2 pale green
[Cr(H2O)4(Cl)2]Cl green
275
Deprotonation by water
[Cr(H20)6] 3+ + H20 ⇌ [Cr(H20)5(OH)]2+ +H30+
Occurs with iron 3+
Does not with copper 2+
276
Deprotonation by stronger bases
Sodium hydroxide or weak ammonia
Uncharged - therefore insoluable
277
Test for iron(III) ions
Deep blue ppt forms when hexacyanoferrate is added
278
en ligand
1,2 diaminoethane
Can fit 3
Neutral charge
279
Why can diaminomethane not act as a mutlidetenate ligand
Bond angle would have to be 90 degrees
Too short
Would have too much strain
280
Colour of solid hydrate copper(II) sulfate
| Or anhydrous
Torquoise blue
Anhydrous is white as there are no ligands
281
Copper(I) complexes in water
Disproportionate
| 2Cu+-> Cu(s) + Cu2+
282
Why are transition metals good catalysts
Energetically avaliable d orbitals can accept electrons or it’s own d electrons can form a bond
283
Why is NaCl formed not NaCl2
Formation NaCl -441kJmol-1
Formation NaCl2 +2177kJmol-1
Basically to endothermic
The point is it’s not about Noble gas electronic configuration stability
284
Effect of change of temp on equilibrium using entropy
If exothermic -H/T is positive
Therefore as T increases Δs total decreases
Therefore lnK smaller therefore K smaller
Eq more to left
Reverse arg for other one
285
Effect of temperature on equilibrium in terms of Gibbs free energy
ΔG = -RTlnK
ΔG = ΔΗ - TΔS
lnK = ΔS/R - ΔH/RT
If exothermic then -ΔH/T positive, becomes less positive as T increases
Therefore ln K and K become smaller and eq moves to left
286
Law of mass action
When reactions reach equilibrium, the equilibrium concentrations of the products multiplied together and divided by the equilibrium concentrations of the reactants also multiplied together, with the concentration of each substance raised to the power appropriate to the reaction stoichiometry are constant at a given temperature
(Reaction quotient )
287
What is a partial pressure
The pressure that gas A would exert if it were alone in the container at that particular temperature
288
Phase of a dissolved solid
Single liquid phase with the solvent
289
Why is a solid left out of the reaction quotient
It’s concentration is constant Kc
| It has no vapour pressure Kp
290
Uses of mass spectrometers
Detection of drugs and their metabolites in urine and blood samples
Identity of a new compound in the pharmaceutical industry
Carbon-14 dating
291
Effective nuclear charge
The effective nuclear charge is the net charge on the nucleus after allowing for the electrons in orbit around the nucleus shielding it’s full charge
* increases across a period
* d block hardly alters
292
Energy in dissolving
Hydration enthalpy / enthalpy of solution energy released compensates for energy in bond breaking
293
Proof of ionic bonding
Electron density maps when X-rays are passed through a crystal, the radiation is scattered and diffraction pattern obtained. This is dependant on electron density therefore a contour map can be produced
294
Which octets can expand
Phosphorus
| Sulphur
295
Factors which determine strength of a covalent bond
* sum of atomicr radio, the shorter the stronger
| * the more electrons shared the stronger
296
Dot and cross for hydronium ion
Central O
Lone pair on O
Dative pair from oxygen to hydrogen
Then two normal pairs
Ammonium ion also has one dative pair
297
Aluminium chloride bonding
Al2Cl6
Two AlCl3 molecules bond together, the lone pair on one chloride bonds to empty aluminium orbital of the other
Each Al has four bonds
The Al are bonded together by two chlorines
Each of these chlorine has a dative bond to one of the Al
298
PCl6 ^- ion bonding
One PCl5 molecule looses Cl- ion which uses one of its lone pairs of electrons to form a dative covalent bond with an employ orbital of another PCl5
PCl4+PCl6-
On heating forms PCl5
299
Polarity of molecules
* tetrahedral are not polar if all four groups are the same
* if one is different or a lone pair it is polar
* same with triagonal planar
* linear are non polar
* best cam be polar
300
Heterogenous reaction
The reactants and profits are not all in the same phase
Assume conc of solid is constant - omit it
301
Enthalpy of reaction experiment
* if at room temp the reaction occurs at a reasonable rate, heat change can be measured using an expanded polystyrene cup as calorimeter
* thermal insulator minimised heat loss to the surroundings, and absorbs little heat itself
* measure every 30 seconds for two minutes, then add second reagent and continue recording temp until the max/min temp is reached take 3 more readings
* extrapolate the graph back
302
Charge on sulfide ion
S^2-
303
Charge on nitride and nitrate ions
N^3-
| NO3^-
304
Charge on Manganate (VII)
MnO4 -
305
Charge on hydrogen carbonate and carbonate ions
CO3 2-
| HCO3 -
306
Chlorate (I)
| chlorate (V) ions
OCl -
| ClO3 -
307
Oxide and superoxide and peroxide ions
O 2-
O2 -
O2 2-
308
Sulfate and sulfite and thiosulphate
SO4 2-
SO3 2-
S2O3 2-
309
Chromate and dichromate ions (VI)
CrO4 2-
| Cr2O7 2-
310
Phosphate ions
PO4 3-
311
Oxidation number of aluminium in compounds
+3
312
When does oxygen not have an oxidation state of -2
Peroxide
Superoxide
With fluorine
313
When does hydrogen not have an oxidation number of +1
When combined with a metal
314
Reaction of sodium thiosulphate with chlorine bromine and iodine
Chlorine and bromine
4Cl2 + S2O3^2- + 5H2O -> 8Cl- + 2SO4 ^2- + 10H+
Iodine
I2 + 2S2O3 ^2- -> 2I- + S4O6 ^2-
315
Why does hydrogen bromide appear is misty fumes when in contact with moist air
Dissolved in water
| To form hydrobromic acid
316
EDTA
4-
6 sites
317
Why does hydrogen not belong above lithium
The rest of group one are metals
Hydrogen has different chemical properties
Forms H- ion
