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A2 Chem > Physical > Flashcards

Physical Flashcards

(61 cards)

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1
Q

Relative atomic mass

A

Ar is the weighted average mass of an atom of an element, taking into account its naturally occurring isotopes, relative to 1/12 the relative atomic mass of an atom of carbon-12

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2
Q

Relative molecular mass

A

Mr of a molecule is the mass of the molecule compared to 1/12 the relative atomic mass of an atom of carbon-12

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3
Q

Avogadro’s Constant

A

A mole of a substance is the number of particles (6.022x1023) in 12g of carbon-12

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4
Q

m3 -> dm3

A

x1000

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5
Q

dm3 -> cm3

A

x1000

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6
Q

m3 -> cm3

A

x1000000

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7
Q

dm3 -> m3

A

/1000

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8
Q

cm3 -> dm3

A

/1000

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9
Q

cm3 -> m3

A

/1000000

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10
Q

moldm-3 -> gdm-3

A

x Mr

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11
Q

why are atoms ionised in TOF spec

A

acceleration and detection

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12
Q

TOF spec ionisation (2)

A

electron impact ionisation

  • vapourised sample is bombarded by high energy electrons
  • sample loses electron forming M+

electrospray ionisation

  • high voltage is applied to sample dissolved in a polar volatile solvent
  • sample molecule m gains a proton to for MH+
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13
Q

first ionisation energy / Ionisation enthalpy (ΔHiƟ)

A

the enthalpy change when an electron is removed from each atom in a mole of gaseous atoms to form one mole of 1+ gaseous ions

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14
Q

electronegativity

A

the power of an atom to attract a pair of electrons in a covalent bond

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15
Q

trend of 1st ionisation energy down group

A

decreases
more shielding as more energy levels
further distance from nucleus (bigger atomic radius) decreases attraction
despite increase in nuclear charge

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16
Q

trend of 1st ionisation energy across a period

A

generally requires more energy
increase in nuclear charge
atoms become smaller so closer distance
same shielding

2->3 decrease as p orbital is higher in energy than s
5->6 decrease as extra electron-electron repulsion means less energy required to remove electron from p4 than p3

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17
Q

Standard enthalpy change of formation (ΔHfƟ)

A

when one mole of a compound is formed in its elements under standard conditions with all substances in their standard states

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18
Q

Standard enthalpy change of combustion (ΔHcƟ)

A

when one mole of an element or compound reacts completely with oxygen under standard conditions with all elements in their standard states

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19
Q

Hess’s Law

A

the total enthalpy change for a chemical reaction is independent of the route by which the reaction takes place, provided the initial and final conditions are the same

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20
Q

mean bond energy

A

enthalpy needed to break the covalent bonds into gaseous atoms averaged over different molecules

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21
Q

oxidation state

A

charge an atom of a molecule would have if it were an ion

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22
Q

Enthalpy of lattice formation (ΔHLFƟ)

+ or -

A

energy transferred when one mole of a solid ionic lattice is formed from its gaseous ions under standard conditions
always - , exothermic

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23
Q

perfect ionic model

A

assuming all attractions are purely electrostatic and ions are perfectly spherical and all ionic

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24
Q

Enthalpy of lattice dissociation (ΔHLDƟ)

+ or -

A

energy transferred when one mole of a solid ionic lattice is separated into its gaseous ions under standard conditions
always + , endothermic

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25
Enthalpy of atomisation (ΔHatƟ)
enthalpy change when one mole of gaseous atoms is formed from the element in its standard state o = enthalpy of sublimation (ΔHsubƟ): if the element is a solid in standard states
26
Enthalpy of electron affinity (ΔHeaƟ)
enthalpy change when one mole of gaseous atoms gain one mole of electrons to form one mole of gaseous 1- ions 1st is exothermic as more stable 2nd is endothermic as negative ion repels electron
27
Why is calcium chloride CaCl2 (and not CaCl or CaCl3)
Formula with the most exothermic enthalpy of formation will form as it is the most thermodynamically stable
28
Enthalpy of solution (ΔHsolƟ)
enthalpy change when one mole of ionic solid dissolves in a large enough amount of water to ensure the dissolved ions are well separated ad do not interact with one another (soluble is exothermic)
29
Enthalpy of hydration (ΔHhydƟ)
enthalpy change when one mole of aqueous ions is formed from gaseous ions (exothermic as ions and water molecules form bonds)
30
order of reaction
order of reaction with respect to a reactant is the power of the concentration of a species in the rate equation
31
rate determining step
same proportion of species involved in or before the rate determining step also appear in the rate equation slowest step which controls the overall rate of reaction
32
Arrhenius equation
k=Ae^(-Ea/RT)
33
logged Arrhenius equation
ln⁡k = -Ea/RT+ln⁡A
34
activation energy equation
Ea=RT(ln⁡A-ln⁡k ) = -gradient x R
35
entropy
measure of the amount of disorder in a system (JK^-1mol^-1) | increase in disorder is positive ΔS
36
Gibbs free energy change ΔG
∆G= ∆H(kJ)-T∆S(J) When ∆G < 0 reaction is feasible When ∆G ≥ 0 reaction is not feasible When ∆G = 0, T= ΔH⁄ΔS
37
Bronsted Lowry acid
proton donor
38
Bronsted Lowry base
proton acceptor
39
Lewis acid
electron pair acceptor
40
Lewis base
electron pair donor
41
Kw at 25°C
Kw = [H+][OH-] = 1x10-14 mol2dm-6
42
at equivalence point (neutralisation volume added)
moles of H+ = moles of OH- | ؞ M1V1 = M2V2
43
at half equivalence (half neutralisation volume added)
[HA] = [A-] | so Ka = [H+] and pKa = pH = -log[H+]
44
phenolphthalein
8.3 - 10.0
45
methyl orange
3.2 - 4.4
46
buffer
a solution that resists a change in pH when small amounts of acid or alkali are added, and on dilution with water
47
if acid is added to buffer
H+ + A- -> HA | equilibrium shift to the left
48
if alkali is added to buffer
OH- + H+ -> H2O | equilibrium shift to the right to replace reacted H+
49
if water is added to buffer
dilutes buffer and doesn't affect pH is ratio [HA]:[A-] is constant, therefore [H+] is the same
50
why are platinum electrodes used
used if system does not have a metal for electrode Provides a conducting surface for electron transfer Inert and can conduct electricity
51
KCl not used as salt bridge in copper systems
will produce complexes with copper ions
52
cell EMF
``` Cell EMF (E⦵cell) = E⦵RHS - E⦵LHS If E⦵cell is positive, the reaction is feasible ```
53
cell diagram of hydrogen standard electrode
Pt | H2 (g) | H+(aq) ||
54
standard conditions for hydrogen standard electrode (3)
Hydrogen gas at pressure of 101kPa Solution containing hydrogen ions at 1.0moldm-3 (e.g. HCl 1M) Temperature at 298K
55
Non-rechargeable cells advantages (1) disadvantages (1)
+ cheaper | - waste issues
56
Rechargeable cells advantages (3) disadvantages (1)
+ less waste + cheaper in the long run + lower environmental impact - some waste issues (at end of useful life)
57
example of rechargeable cell ``` positive electrode negative electrode overall reaction during discharge overall reaction during recharge cell diagram ```
E.g. Lithium ion Cell Positive electrode: Li+ + CoO2 + e- → Li+[CoO2]- Negative electrode: Li -> Li+ + e- Overall reaction during discharge: Li + CoO2 -> LiCoO2 Overall reaction during recharge: LiCoO2 -> Li + CoO2 Cell diagram: Li | Li+ || Li+ , CoO2 | LiCoO2 | Pt
58
Fuel cells advantages (3) disadvantages (5)
``` + only waste product is water + does not require recharging + very efficient - need constant supply of fuels - hydrogen is highly flammable and explosive - storing hydrogen - hydrogen source from fossil fuels or electrolysis - high cost of fuel cells ```
59
example of fuel cell
hydrogen-oxygen fuel cell
60
hydrogen-oxygen fuel cell in alkaline conditions Negative anode Positive cathode Overall equation Cell EMF
Negative anode: H2 + 2OH- -> 2H2O + 2e- Positive cathode: O2 + 2H2O + 4e- -> 4OH- Overall equation: 2H2 + O2 -> 2H2O Cell EMF = +1.23V
61
hydrogen-oxygen fuel cell in acidic conditions Negative anode Positive cathode Overall equation Cell EMF
Negative anode: H2 -> 2H- + 2e- Positive cathode: O2 + 4H+ + 4e- -> 2H2O Overall equation: 2H2 + O2 -> 2H2O Cell EMF = +1.23V

A2 Chem (5 decks)

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