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Chemistry > Paper 1 > Flashcards

Paper 1 Flashcards

1
Q

Relative mass of an electron

A

0.00055

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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

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3
Q

Relative isotopic mass

A

The mass of an atom of that isotope divided by 1/12 of the mass of carbon twelve

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4
Q

First ionisation energy

A

The energy required to remove one mole of electrons from one mile of atoms in its gaseous form

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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)
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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

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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

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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

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9
Q

What is an orbital

A

Region within an atom that can hold up to two electrons with opposite spins

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10
Q

Shapes S and P

A

S- is spherical
P- infinity sign

Draw on axis

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11
Q

Conditions for electrons to fill subhsells

A

Fill singly before pairing up

Two electrons in the same orbital must have opposite spin

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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)

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13
Q

Where are the s p and d block

A

D transition metal area
P right
S left and hydrogen

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14
Q

What is periodicity

A

Repeating pattern across different periods

Recurring trends that are shown in the properties of an element

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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

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16
Q

Ionic bonding

A

Strong electrostatic attraction between oppositely charged ions

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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

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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

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19
Q

Covalent bond

A

Strong electrostatic attraction between two nuclei and the shared pair of electrons between them

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20
Q

Relationship between bond strength and bond length

A

The longer the bond length the weaker the bond

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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

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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

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23
Q

Two bonds

No lone pair

A

Linear
180
Cl-Be-Cl

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24
Q

Three bond pairs

A

Trigonal planar
120
BCl3 Ethene
Sulphur trioxide

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25
Q

Two bond pairs

One lone pair

A

Bent 120

Sulphur dioxide

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26
Q

Four bond pairs

A

Tetrahedral
109.5

Wedges and slashed lines

Methane
(NH4)+

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27
Q

Three bond pairs

One lone pair

A

Trigonal pyramidal
107

Ammonia

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28
Q

Two bond pairs two lone pairs

A

Bent
104.5

Water

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29
Q

Five bond pairs

A

Trigonal bi-pyramidal
120,90
PCl5

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30
Q

6 bond pairs

A

Octahedral
90

SF6

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31
Q

Electronegativity

A

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

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32
Q

Polar molecules

A

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

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33
Q

London forces

A

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

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34
Q

Permanent dipole - dipole

A

Electronegativity differences lead to polarity. The charges attract neighbouring polar molecules

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35
Q

Hydrogen bonding

A

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

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36
Q

How many H bonds in
H20
NH3
HF

A

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

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37
Q

Melting and boiling temperature of water and alcohols

A

Higher due to strong Hydrogen bonds

Alcohols have a lower volatility than alkanes of a similar number of e

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38
Q

Density of ice

A

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

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39
Q

Trends in boiling temperature of alkanes

A

Chain length increases, more electrons, more London forces. Higher bp
More branching, molecules not stacked, less London forces, lower bp

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40
Q

Trends in boiling temperatures of hydrogen halides

A

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

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41
Q

Choice of solvents

A
  • 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)
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42
Q

Metallic bonding

A

Strong electrostatic attraction between metal ions and delocalised electrons

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43
Q

Where are giant lattices present

A
  • 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)

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44
Q

Graphite

A

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

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45
Q

Diamond structure

A

Four sigma bonds to carbons per carbon
3D tetrahedral
Bond angle 109.5

Silicon has a similar structure

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46
Q

Trend in reactivity down group 2

A

Reactivity increase

IE decreases -> group two reacts by loosing electrons

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47
Q

Group two plus oxygen

A

2Ca + O2 -> 2CaO

Burn in air

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48
Q

Group two plus chlorine

A

Ca + Cl2 -> CaCl2

When heated

These ionic chlorides can dissolve in water producing solutions that contain the hydrated cations

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49
Q

Magnesium and water

A

COLD

Mg +2H20 -> Mg(OH)2 + H2

HOT (steam)

Mg + H20 (g) -> MgO + H2

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50
Q

Group two plus water

A

Magnesium different in heat

Ca + 2H20 -> Ca(OH)2 + H2

Reacts rapidly in cold water

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51
Q

Group 2 oxides with water

A

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-

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52
Q

Group 2 oxides with dilute acid

A

MO + 2H+ -> M^2+ + H20

With a sulphuric acid can form MSO4
Or nitrate with nitric acid

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53
Q

Hydroxides with dilute acid

A

M(OH)2 + 2H+ -> M^2+ + 2H20

Forms the salts

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54
Q

Solubility of group 2 hydroxides

A

Increases down a group

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55
Q

Solubility of group two sulphate

A

Decreases down the group

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56
Q

Trend in thermal stability of nitrates and carbonates of groups one and two

A

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

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57
Q

Decomposition of group one nitrates

A

Only lithium

4LiNO3 -> 2Li2O +4NO2 + O2

all Decompose on strong heating in a different way
They melt
2NaNO3 -> 2NaNO2 + O2

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58
Q

Decomposition of group two nitrates

A

2Mg(NO3)2 -> 2MgO + 4NO2 + O2

All of them decompose this way
On heating

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59
Q

Decomposition of group one carbonates

A

Li2CO3 -> Li2O + CO2

Only lithium

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60
Q

Decomposition of group 2 carbonates

A

CaCO3 -> CaO + CO2

All on heating

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61
Q

Formation of flame test colours

A

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

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62
Q

Li flame colour

A

Crimson

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63
Q

Na flame colour

A

Yellow

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64
Q

K flame colour

A

Lilac

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65
Q

Mg flame colour

A

No colour

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66
Q

Ca flame colour

A

Yellow red

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67
Q

Sr flame colour

A

Red

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68
Q

Ba flame colour

A

Apple green

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69
Q

Experimental procedure for thermal decomposition

A
  • 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
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70
Q

Experimental procedure for a flame test

A

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
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71
Q

F state at room temperature

A

Gas

Pale yellow

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72
Q

Cl state at room temperature

A

Gas

Green

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73
Q

Br state at room temperature

A

Volatile liquid

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74
Q

I state at room temperature

A

Solid which sublimes when heated

Grey black solid sublimed into violet vapour

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75
Q

Trend in electronegativity down group 7

A

The atom is larger, therefore attraction to bonding electrons decreases and therefore the electronegativity decreases down a group

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76
Q

Trend in reactivity of group 7

A

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

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77
Q

Cl Aq colour

Organic

A

Pale green

Pale yellow/green

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78
Q

Br aq colour and in organic colour

A

Orange

Orange

(Bromine liquid is brown)

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79
Q

I aq colour and organic colour

A

Pale brown but only slightly soluble

Forms I3^-1 Deep red brown in the presence of I2 and I^-

Violet

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80
Q

Redox reactions of halogens and halides in aq followed by organic

A

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

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81
Q

Halogens plus group 1 and2 metals

A

Heat

Na+1/2Cl2 -> NaCl

Mg + I2 -> MgI2

Or
2Fe + 3Cl2 -> 2FeCl3
Fe + I2 -> FeI2 (less poweful oxidising agent)

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82
Q

Disproportionation

A

A single species is simultaneously oxidised and reduced

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83
Q

Chlorine plus water

Application

A

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

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84
Q

Chlorine aq sodium hydroxide

A

COLD

Cl2 + 2OH- -> Cl- + H20 + ClO-

ClO- bleach

HOT

3Cl2 + 6OH- -> 5Cl- + 3H2O ClO3^-1

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85
Q

NaCl + H2SO4

A

NaHSO4 + HCl

Steamy fumes which turn blue litmus red and give white fumes with gaseous ammonia

No reduction
NaCl is a poor oxidising agent

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86
Q

KBr + H2SO4

A

HBr + HKS04

2HBr + H2SO4 -> SO2 + H20 +Br2

Brown Br2 gas reduction to plus four

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87
Q

KI + H2SO4

A

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

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88
Q

HCl and water

And other hydrogen halides

A

HCl + H2O -> H3O+ + Cl-

All soluble to form acidic solutions

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89
Q

HCl plus Ammonia

And other hydrogen halides

A

HCl(g) + NH3 (g) -> NH4Cl (s)

Reaction is a test for gaseous hydrogen halides, but does not distinguish between them

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90
Q

Test for carbonate ions and hydrogen carbonate ions

A

Use aqueous acid

CO3^2- +2H+ -> H20 + CO2

Or HCO3^-1 + H+ -> H20 + CO2

Turns limewater cloudy
Effervescence

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91
Q

Test for sulphate ions

A

SO4^2- + BaCl2 -> BaSO4 + 2Cl^-

Milky white ppt

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92
Q

Test for ammonium ions

A

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)

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93
Q

Test for halide ions

A

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)

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94
Q

Atom economy

A

100 x Molar mass of desired produced/ sum of molar mass of all products

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95
Q

Standard conditions

A

100kPa

Specified temp usually 298K

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96
Q

What is enthaply change

A

Heat energy change measured at a constant pressure

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97
Q

Enthalpy level diagrams

A

Activation energy is not shown in enthalpy level diagrams but is shown on reaction profiles

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98
Q

Standard enthalpy change of reaction

A

Enthalpy change when the number of moles as written react under STP

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99
Q

Standard enthalpy of formation

A

Enthalpy change when one mole of a substance is formed from its elements in their standard states under STP

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100
Q

Standard enthalpy change of combustion

A

Enthalpy change when one mole of a substance is burnt in excess oxygen under STP

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101
Q

Standard enthalpy change of neutralisation

A

Enthalpy change when one mole of water is produced by the neutralisation of a solution of acid by excess base under STP

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102
Q

Errors in measuring temperature change

A

• 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

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103
Q

Assumptions in measuring temperature change

A
  • density of solution = density of water
  • specific heat capacity of solution = that of water
  • negligible heat loss
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104
Q

Hess Law

A

The enthalpy change for any reaction is independent of the route taken from the reactants to products

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105
Q

Bond enthalpy

A

Enthalpy change when one mole of a bond in a gaseous molecule is broken

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106
Q

Mean bond enthalpy

A

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

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107
Q

Conditions for dynamic equilibrium

A
  • The rate of forward reaction is equal to the rate of backward reaction
  • the concentration of reactants and products must remain constant
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108
Q

Effect of change of pressure

A

if pressure increases Eq shifts to the side with less gaseous molecules

No effect on Kp

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109
Q

Effect of change of concentration

A

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

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110
Q

Effect of temperature on equilibrium

A

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

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111
Q

kp

A

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

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112
Q

Brønsted Lowry bases and acids

A

Acid -> proton donor

Base -> proton acceptor

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113
Q

What is the difference between a strong acid and a weak acid

A

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%

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114
Q

Assumptions for weak acid calculations

A

[H+]=[A-] there is no other source of A-

[HA]eq = [HA]inital

Ka = x^2/[HA]

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115
Q

Calculation of PH of a strong base

A

pH + pOH =14

Or [H+][OH-] = Kw

Remember to account for multiple OH in same compound

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116
Q

pKa

PKw

A
  • logKa

- logKw

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117
Q

How much does the pH change for a tenfold dilution

A

1 unit strong acid

1/2 weak

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118
Q

features of a strong acid strong base titration curve

A

Vertical range
3-11

Equivalence point ph=7

(Calc vol equivalence point and final and initial pH)

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119
Q

Features of a weak acid strong base titration curve

A

Vertical range 7-11
Equivalence point pH 9
Initial buffer region

Calc vol equivalence point and initial and final pH

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120
Q

Features of a strong acid weak base titration curve

A

Vertical range 3-7
Equivalence point pH 5

Calc vol of equivalence point and final and initial pH

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121
Q

Buffer solution

A

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

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122
Q

Action of a buffer solutions

A
  • 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

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123
Q

Assumptions for calculating pH of a buffer

A 
[A-] = [salt] 
[HA] = weak acid

(Remember to account for used up acid if some is used producing salt)

Ka = [H+][salt]/[weak acid]

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124
Q

How to determine Ka from a weak acid-strong base titration curve

A

[HA]=[A-]. Therefore Ka = [H+]

Ka = 10^-pH

At half equivalence point

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125
Q

Why is there a difference in enthalpy of neutralisation values for strong and weak acids

A

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

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126
Q

How is the pH in our blood controlled

A

Carbonic acid in the blood and the conjugate base hydrogencarbonate act as a buffer

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127
Q

Lattice energy

A

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
Q

Enthalpy change of atomisation

A

Enthalpy change when one mole of gaseous atoms is made from an element in its standard state

129
Q

Electron affinity

A

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
Q

Born Haber cycle

A

Hess’s cycle for finding lattice energy

131
Q

Comparing the theoretical lattice energy value with experimental value

A

The more different the values are the higher the degree of covalent bonding

132
Q

Polarisation

A

A positive action exerts and attraction on electrons in the negative anion distorting the electron cloud and weakening bonds

133
Q

When is polarisation the strongest

A

polarising Small cation and high charge

Polarisable Large anion and large charge

134
Q

Enthalpy change of solution

A

Enthalpy change when one mole of the solid is dissolved in sufficient solvent to give an infinitely dilute solution

135
Q

Enthalpy change of hydration

A

The enthalpy change when one mole of gaseous ions is dissolved to give an infinitely dilute solution

136
Q

Evidence that enthalpy alone doesn’t control whether reactions occur

A

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
Q

What is entropy (chemistry)

A

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
Q

Change of state
Effect on entropy
Or dissolving

A

Ordered crystal structure become dispersed through a liquid -> entropy increases

Solid more ordered than liquid than gases

139
Q

Number of moles effect on entropy

A

More moles in products than reactants means that entropy increases

140
Q

Entropy change in dissolving ammonium nitrate crystals

A

Dissolve therefore less ordered entropy increases

141
Q

Reacting ethanoic acid with ammonium carbonate

A

CO2 gas produced

CO2 + 2ammoniumethanoate + h20

No mores increases 3-> 5
Entropy increases

142
Q

Burning magnesium ribbon

A

Mg (s) + O2 (g) -> 2MgO

Less moles
gas to solid
Entropy decreases

143
Q

Mixing solid barium hydroxide with solid ammonium chloride

A

Forms ammonia BlC2 and H20

Solid to liquid
Entropy increases

144
Q

Total entropy

A

ΔS total = ΔS system + ΔS surroundings

Δs total = Rlnk

145
Q

Entropy of surroundings

A

ΔS surroundings = -ΔH/T

146
Q

Gibbs free energy

A

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
Q

Why might a - ΔG value occur in practice

A

If the reaction is kinetically controlled (high activation energy)

Or departure from standard conditions such as ppt forming or conc acid

148
Q

Standard electrode potential

A

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
Q

Features of a standard hydrogen electrode

A

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
Q

How to set up an electrode where the substance is gaseous

A

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
Q

Cell diagram

A

Anode|ions || ions | cathode

Eg
Pt(s)|MnO4^-1(aq), 8H+ (aq), Mn2+ (aq)|| Fe3+ (aq), Fe2+ (aq) | Pt(s)

152
Q

Predicting the thermodynamically feasibility of a reaction using standard electrode potentials

A

Predict when Ecell > 0 thermodynamically feasible

153
Q

Limitations of predictions made using standard electrode potentials

A

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
Q

Electrochemical series

A

The higher the E value the better the oxidising agent

155
Q

Methanol and hydrogen rich cells

A

Anode CH3OH + H20 ⇌ HCOOH + 4H+ + 4e-

Cathode 02 + 4H+ + 4e- ⇌ 2H2O

Energy released is utilised in a fuel cell to generate a voltage

156
Q

Application of electrode potentials to storage cells

A

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
Q

Hydrogen-oxygen fuel cell

A

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
Q

Iodine titration

A

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
Q

Application of electrode potentials to storage cells

A

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
Q

Hydrogen-oxygen fuel cell

A

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
Q

Configuration of d-block elements Sn-Zn

A

[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
Q

What is a transition metal

A

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
Q

Why do transition metals show variable oxidation number

A

The successive ionisation energies increase steadily. At higher oxidation states, the election can be promoted just to unpair them for covalent bonding

164
Q

Ligand

A

Ion or molecule which is bonded via a dative covalent bond to a central metal ion

165
Q

What is a complex ion

A

A central metal ion surrounded by Luganda

166
Q

Why do transition metals form coloured ions in solution

A

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
Q

Why is there a lack of colour in some aq ions and other complex ions

A

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
Q

What causes colour changes in transition metal ions

A
  • 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
Q

What is the coordination number

A

No of atoms bonded to the central metal ion

170
Q

Monodentate ligands

A

One lone pair therefore can form one dative bond eg h20 or OH-

171
Q

Shape of 6 fold coordination number

A

Repeal to maximum separation to minimise repulsion

Octahedral shape

172
Q

Shape of four coordination number

A

Larger ligands May form tetrahedral complexes eg Cl-

109.5

Could also be square planar eg cisplatin

173
Q

Use of transition metals in cancer treatment

A

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
Q

Bidentete ligands And multi-dentate ligands

A

Bidentate have two lp and are sufficiently long to bend round

175
Q

What is haemoglobin

A

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
Q

Colour of vanadium 5+

A

Yellow VO2+ (colourless VO3-)

177
Q

Colour of vanadium 4+

178
Q

Colour of vanadium 3+

179
Q

Colour of vanadium 2+

180
Q

How do reduce dichromate (VI)

A

Cr2O7^2- reduced to Cr3+ and Cr2+ using zinc in 50% hydrochloride acid acidic conditions

Final colour blue

181
Q

How to produce dichromate (VI)

A

Oxidation of Cr3+ ions using hydrogen peroxide in alkaline conditions followed by acidification

182
Q

Equilibrium between Chromate and dichromate

A

2CrO4^2- + 2H+ ⇌ Cr2O7^2- + H2O

Convert between by altering pH to shift equilbirum

183
Q

Transition metal plus aq NaOH

A

Deprotonation - number deprotonated is equal to charge on cation

In excess further deprotonation only for ampotheric hydroxides
Cr3+

184
Q

Transition metal plus ammonia

A

Deprotonation

Excess
Some undergo ligand exchange

Cr3+ 6NH3 exchanged
Co2+ 6NH3 exchanged
Cu2+ 4NH3 exchanged

185
Q

Cr3+ NaOH

A

(Grey) Green ppt

Excess
(Dark) green solution

186
Q

Fe2+ NaOH

A

(Dark) green ppt. (Goes brown on exposure as oxidised to Iron(III))hydroxide)

Excess
Insoluable

187
Q

Fe3+ NaOH

A

Red brown ppt

Excess
Insolvable

188
Q

Cobalt + NaOH

A

Blue ppt (goes pink on standing)

Excess insoluable

189
Q

Cu2+ NaOH

A

Blue ppt

Excess
Insoluable

190
Q

Cr3+ aqueous

191
Q

Fe2+ aqueous

A

Pale green

192
Q

Fe3+ aqueous

193
Q

Co2+ aqueous

194
Q

Cu2+ aqueous

A

Turquoise blue

195
Q

Cr3+ NH3

A

(Green) grey ppt

Excess
green solution forms slowly

196
Q

Fe2+ NH3

A

(Dark )green ppt (again darkens in air)

Excess
Insoluable

197
Q

Fe3+ NH3

A

Red brown ppt

Excess
Insoluable

198
Q

Co2+ NH3

A

Blue ppt

Excess
brown solution

199
Q

Cu2+ NH3

A

Blue ppt

Excess
Deep blue solution

200
Q

Ligand exchange with Cl-

A

[CuCl4]2- from [Cu(H2O)6]2+

[CoCl4]2- from [Co(H2O)6]2+

201
Q

Ligand exchange affect on entropy and stability

A

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
Q

Ammonia ligand exchange

A

[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
Q

V2O5 as a catalyst

A

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
Q

Amphoteric behaviour of complex ions

A

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
Q

Catalytic converter

A

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
Q

Heterogenous catalyst

A

In a different phase from the reactants and the reaction occurs on the surface

207
Q

Homogenous catalyst

A

In the same phase as the reactants and the reaction occurs via and intermediate species

208
Q

Iron as a catalyst

A

Homogenous

2Fe^2+ + S2O8^2- -> 2Fe^3+ 2SO4^2-
2Fe^3+ + 2I- -> 2Fe2+ + I2

2I^-(aq) + S2O8^2- (aq) -> I2 + 2SO4^2-

209
Q

Autocatalysis

A

2MnO4 ^- + 16H+ +5C2O4 ^2- -> 2Mn^2+ + 10CO2 + 8H20

Reaction speeds up as the Mn^2+ is produced as it catalyses the reaction

210
Q

What affects the hydration enthalpy

A

High charge larger magnitude

Less exothermic for larger radius (the greater the force)

211
Q

What is a homogenous reaction

A

One in which all the reactants and products are in the same phase

212
Q

What phase is two immiscible layers

A

Two phases

213
Q

What phase is a mixture of solids

A

Usually two phase

214
Q

Effect of addition of an inert gas equilibrium

A

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
Q

Effect of adding a catalyst to equilibrium

A

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
Q

Aims of industrial processes in equilibrium

A
  • maximise yield
  • make it quickly
  • keep cost low
  • have high atom economy
217
Q

How can you improve yield in industrial processes

A

•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
Q

What is a strong base

A

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
Q

What is a neutral solution

A

[H+][OH-] are the same

Ph is not necessarily 7

220
Q

pH of strong acid and salt of strong acid and strong base

221
Q

pH of weak acid and salt of strong acid and strong base

222
Q

pH of strong base and salt of weak acid and strong base

223
Q

pH of weak base and salt of weak acid and weak base

224
Q

How does an indicator work

A

HInd ⇌ H+ + Ind-

Kind = [H+][Ind-] / [Hind]

In acid eq is driven to each side and the colour changes

225
Q

What affects lattice energy

A
  • 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
Q

What is the standard free energy of formation

A

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
Q

What’s ΔSsystem of dissolving a gas

A

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
Q

ΔS system of dissolving a solid

A

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
Q

Why do hydroxides get more soluble down and group and sulphates get more soluable up the ground

A
  • 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
Q

Which is the anode and which is the cathode

A

The anode is where oxidation occurs
The cathode is where reduction occurs

Anions ions move to anode
Cations move to cathode

231
Q

What does a salt bridge consist of

A

A conc solution of an inert electrolyte such as potassium nitrate or potassium chloride

232
Q

Features of an electrode with two cations

A

Electrode consists of a platinum rod dipping into a solution containing 1 moldm^-3 of both ions, each with that conc

233
Q

Relationship between equilibrium constant and ECell

A

lnK α Ecell

234
Q

Potassium manganate titrations

A

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
Q

Why are scandium and zinc not transition metals

A

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
Q

Trend in ionisation energy across period 4 from potassium to zinc

And successive ionisation energies

A

•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
Q

How to reduced vanadate (v) vanadium (II)

A

Warm with powdered zinc with ammonium vanadate in presence of 50% hydrochloride acid solution

238
Q

Colour of Cr2+

239
Q

Colour of Cr3+

240
Q

Colour of Cr in oxidation state of +6

A

CrO4 2-
Yellow

Cr2O7 2-
Orange

241
Q

Melting temp of metals

A

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
Q

Reaction of metals with less restive metal

A

Fe(s) + Cu^2+ aq -> Fe^2+ (aq) + Cu (s)

243
Q

What is an oxidation number

A

The charge that the element would have if the compound were fully ionic

244
Q

What is group one called

A

Alkali metals

245
Q

What is group two called

A

Alkaline earth metals

246
Q

Lime water with carbon dioxide

A

Ca(OH)2 (aq) + CO2 (g) -> CaCO3 (s) + H2O (l)

247
Q

Test to distinguish between a carbonate and a hydrogen carbonate of group one

A

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
Q

Chlorine plus phosphorus

A

Limited supply of chlorine

2P + 3Cl2 -> 2PCl3

In excess

2P + 3I2 -> 2PCl5

If damp

PI3 + 3H20 -> 3HI + H3PO3

249
Q

Boiling temps of hydrogen halides

A

Increases from HCl to HI due to more electrons therefore more London forced

HF is anomalous because it forms hydrogen bonds

250
Q

Solubility of halides

A

All soluble except silver halides and lead II halides

251
Q

Steps to carry out a titration

A
  • 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
Q

Examples of poor technique

In titration stuff

A
  • 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
Q

Preparation of a standard solution

A
  • 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
Q

Experimental method to find the enthalpy of combustion of a liquid

A
  • 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
Q

When to include water in equilibrium constant

A
  • gaseous state INCLUDE
  • reactant but not solvent INCLUDE
  • when solvent, even if reactant or product, DO NOT INCLUDE because the conc remains constant
256
Q

When can the enthalpy change of reaction not be determined directly

A

When heat is supplied, you cannot distinguish the heat change due to the reaction

257
Q

Formula of potassium dichromate

258
Q

Test for NO3-

A

Gives off brown gas of NO2 when heated in Biden and caused glowing splint to relight

259
Q

Why may a student add a pinch of something to a reaction

A

If finding a volume, the pinch may saturate the solution to stop excess dissolving

260
Q

Is CO2 soluable

261
Q

Is hydrogen soluable

262
Q

Unit of entropy

A

JK^-1mol^-1

263
Q

Cobalt chloride

264
Q

Copper chloride

265
Q

Metal plus acid

A

Salt and hydrogen

266
Q

Calculating ECell

A

Either Eoxidant - E reductant

Or right - left

267
Q

Environmental issues behind hydrogen fuel cells

A
  • 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
Q

Salt bridge

A

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
Q

Why would you want to use a high resistance voltmeter

A

Little reaction takes place therefore the conc of the ions remains approximately constant

270
Q

What is an electrolyte

A

electrolyte is a chemical compound that dissociates into ions and hence is capable of transporting electric charge

271
Q

Bonding in a transition metal

A

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
Q

Stereoisomerism in a transition metal

A

Cis-trans -> occurs in octahedral and square planar complexes. Adjacent are cis, opposite are trans

Optical -> bidenetate ligands

273
Q

How to reduce to vanadium(II)

A

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
Q

Chromium(II) chloride

A

[Cr(H2O)6(Cl)3]. Grey blue
[Cr(H20)5Cl]Cl2 pale green
[Cr(H2O)4(Cl)2]Cl green

275
Q

Deprotonation by water

A

[Cr(H20)6] 3+ + H20 ⇌ [Cr(H20)5(OH)]2+ +H30+

Occurs with iron 3+

Does not with copper 2+

276
Q

Deprotonation by stronger bases

A

Sodium hydroxide or weak ammonia

Uncharged - therefore insoluable

277
Q

Test for iron(III) ions

A

Deep blue ppt forms when hexacyanoferrate is added

278
Q

en ligand

A

1,2 diaminoethane

Can fit 3

Neutral charge

279
Q

Why can diaminomethane not act as a mutlidetenate ligand

A

Bond angle would have to be 90 degrees
Too short
Would have too much strain

280
Q

Colour of solid hydrate copper(II) sulfate

Or anhydrous

A

Torquoise blue

Anhydrous is white as there are no ligands

281
Q

Copper(I) complexes in water

A

Disproportionate

2Cu+-> Cu(s) + Cu2+

282
Q

Why are transition metals good catalysts

A

Energetically avaliable d orbitals can accept electrons or it’s own d electrons can form a bond

283
Q

Why is NaCl formed not NaCl2

A

Formation NaCl -441kJmol-1
Formation NaCl2 +2177kJmol-1

Basically to endothermic

The point is it’s not about Noble gas electronic configuration stability

284
Q

Effect of change of temp on equilibrium using entropy

A

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
Q

Effect of temperature on equilibrium in terms of Gibbs free energy

A

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
Q

Law of mass action

A

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
Q

What is a partial pressure

A

The pressure that gas A would exert if it were alone in the container at that particular temperature

288
Q

Phase of a dissolved solid

A

Single liquid phase with the solvent

289
Q

Why is a solid left out of the reaction quotient

A

It’s concentration is constant Kc

It has no vapour pressure Kp

290
Q

Uses of mass spectrometers

A

Detection of drugs and their metabolites in urine and blood samples

Identity of a new compound in the pharmaceutical industry

Carbon-14 dating

291
Q

Effective nuclear charge

A

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
Q

Energy in dissolving

A

Hydration enthalpy / enthalpy of solution energy released compensates for energy in bond breaking

293
Q

Proof of ionic bonding

A

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
Q

Which octets can expand

A

Phosphorus

Sulphur

295
Q

Factors which determine strength of a covalent bond

A
  • sum of atomicr radio, the shorter the stronger

* the more electrons shared the stronger

296
Q

Dot and cross for hydronium ion

A

Central O
Lone pair on O
Dative pair from oxygen to hydrogen
Then two normal pairs

Ammonium ion also has one dative pair

297
Q

Aluminium chloride bonding

A

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
Q

PCl6 ^- ion bonding

A

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
Q

Polarity of molecules

A
  • 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
Q

Heterogenous reaction

A

The reactants and profits are not all in the same phase

Assume conc of solid is constant - omit it

301
Q

Enthalpy of reaction experiment

A
  • 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
Q

Charge on sulfide ion

303
Q

Charge on nitride and nitrate ions

A

N^3-

NO3^-

304
Q

Charge on Manganate (VII)

305
Q

Charge on hydrogen carbonate and carbonate ions

A

CO3 2-

HCO3 -

306
Q

Chlorate (I)

chlorate (V) ions

A

OCl -

ClO3 -

307
Q

Oxide and superoxide and peroxide ions

A

O 2-
O2 -
O2 2-

308
Q

Sulfate and sulfite and thiosulphate

A

SO4 2-
SO3 2-
S2O3 2-

309
Q

Chromate and dichromate ions (VI)

A

CrO4 2-

Cr2O7 2-

310
Q

Phosphate ions

311
Q

Oxidation number of aluminium in compounds

312
Q

When does oxygen not have an oxidation state of -2

A

Peroxide
Superoxide
With fluorine

313
Q

When does hydrogen not have an oxidation number of +1

A

When combined with a metal

314
Q

Reaction of sodium thiosulphate with chlorine bromine and iodine

A

Chlorine and bromine
4Cl2 + S2O3^2- + 5H2O -> 8Cl- + 2SO4 ^2- + 10H+

Iodine
I2 + 2S2O3 ^2- -> 2I- + S4O6 ^2-

315
Q

Why does hydrogen bromide appear is misty fumes when in contact with moist air

A

Dissolved in water

To form hydrobromic acid

316
Q

EDTA

A

4-

6 sites

317
Q

Why does hydrogen not belong above lithium

A

The rest of group one are metals
Hydrogen has different chemical properties
Forms H- ion

Chemistry (5 decks)

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