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Chemistry A level > Year 1 - A level > Flashcards

Year 1 - A level Flashcards

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1
Q

What are the 5 stages of mass spectroscopy

A 
1-Vapourise sample
2-Ionisation(remove an electron)
3-Acceleration(using electric field)
4-Deflection(using magnetic field)
5-Detection
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2
Q

Ionisation energy

A

The energy required to remove 1 mole of electrons from 1 mole of gaseous atoms

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3
Q

What is an orbital?

A

Area of space which can hold a maximum of 2 electrons

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4
Q

How many electrons can be held in s, p and d subshells?

A

S - 1 orbital - 2e
P - 3 orbitals - 6e
D - 5 orbitals - 10e

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5
Q

What is the order of the first 5 shells?

A

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d

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6
Q

What is relative atomic mass?

A

The abundance of an elements isotopes using percentage to calculate an average atomic mass against a carbon 12 atom

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7
Q

What is an isotope?

A

Atoms with the same number of protons and different number of neutrons.

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8
Q

What effects ionisation energy?

A
  • Electron shielding, further subshell lower I.E.
  • No. protons, more protons higher I.E.
  • Electron repulsion
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9
Q

How does I.E. change generally: -down a group, -across a period

A

Down group: 1st I.E. decreases as more shielding
Across period: Increases as more protons.
-Mg –> Al decrease since Al e coming from a higher subshell
-P –> S decrease as S electron is paired = more repulsion

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10
Q

Oxidation number for: Element, fluorine, hydrogen and oxygen(not with F or H)

A 
element = 0
Fluorine = -1
Hydrogen = +1
Oxygen = -2
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11
Q

Steps to making 1/2 ionic equations?

A
  1. Balance all but H and O
  2. Add H2O to balance O
  3. Add H+ to balance H
  4. Add e- to balance charge
  5. Check with oxidation states
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12
Q

What is the reaction between Cl2 and a cold and hot alkali(e.g. NaOH)?

A 
COLD:
Cl2 + 2NaOH --> NaCl + H2O + NaOCl
-Cl2 is reduced and oxidised called: Disproportionation reaction
HOT:
3Cl2 + 6NaOH --> 5NaCl + NaClO3 + 3H2O
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13
Q

What is ΔH⦵?

A

Change in enthalpy

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14
Q

What is enthalpy?

A

Energy per mole

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15
Q

What does ⦵ mean?

A

Standard conditions, 100kPa and 298K

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16
Q

What does it mean when ΔH⦵ is negative or positive?

A 
Negative = exothermic, surroundings gain heat
Positive = endothermic, surroundings get colder
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17
Q

What is Q = mcΔT

A

Energy of surroundings = mass of surroundings x specific heat capacity x temp. change.
(In Joules)

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18
Q

What does energy change of system =?

A

Energy change of system = - energy change of surroundings

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19
Q

ΔH =?

A

ΔH = energy change of system / moles

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20
Q

Definition of bond energy?

A

Amount of energy required to break ‘1 mole of a particular bond’ for gaseous atom

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21
Q

Why aren’t bond energies correct?

A

They are averages

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22
Q

Define ΔH⦵c? Enthalpy of combustion

A

Enthalpy change when 1 mole of a substance is burnt in excess O2 under standard conditions

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23
Q

What are all the diatomic elements?

A

B,O,C,I,H,N,F(Becasue Of Chem I Have No Friends)

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24
Q

Define ΔH⦵f? Enthalpy of formation

A

Enthalpy change when 1 mole of a compound is formed from it’s elements, in their STANDARD STATES under standard conditions

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25
Q

ΔH⦵n? Enthalpy of neutralisation

A

Energy change when ‘1 MOLE OF WATER’ is formedin reduction between acid + base, under standard conditions

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26
Q

What are the 3 types of Hess Cycle?

A
  1. Given ΔH⦵c to work out ΔH⦵f
  2. Given ΔH⦵f to work out ΔH⦵ decomposition
  3. Work out ΔH⦵ solution
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27
Q

ΔH⦵ solution?

A

Difference between lattice energy and hydration energy

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28
Q

What is lattice energy?

A

Energy required to form 1 mole of ionic solid from gaseous ions, under standard states

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29
Q

What affects distortion(how good it is at polarising) in +ve ion?

A

-Small radius and high charge means best at polarising.

More distortion means furthest from ionic model and bigger gap between experimental and theoretical lattic enthalpies

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30
Q

What affects how the -ve ion is distorted?

A

-Close to 0 charge + big radius means easy to distort.

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31
Q

What is Entropy(S)

A

The number of ways of organising particles in a substance. Measured in JK-1mol-1?

  • Many ways(gas) = high entropy
  • Few ways(solid) = low entropy
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32
Q

ΔS =

A

ΔS = S products - S reactants

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33
Q

ΔG =

A

ΔG = ΔH - TΔS sys - change Δs by /1000
If -ve then spontaneous
Measured in KjMol-1

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34
Q

ΔS total =

Therefore ΔSsurr =

A

ΔS total = ΔS sys - ΔH/T - change ΔH x1000 for KJ to J
If +ve then spontaneous
Measured in JK-1Mol-1
ΔSsurr = -ΔH/T

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35
Q

What is the ideal gas equation?

A

pV = nRT

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36
Q

What are the three main gas laws?

A
  1. Boyle’s law
  2. Charles’ law
  3. Avagadro’s principle
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37
Q

What is Boyle’s law?

A
  • Volume of a fixed mass of gas is inversely proportional to it’s pressure
  • V ∝ 1/p
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38
Q

What is Charles’ law?

A
  • The volume of a fixed mass of gas is directly proportional to its thermodynamic temperature
  • V ∝ T
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39
Q

What is Avagadro’s principle?

A
  • Equal volumes of different gases at the same temperature and pressure contain the same No. of molecules
  • V ∝ n
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40
Q

What are features of an ideal gas?

A
  • No IMF’s
  • The kinetics theory explains how their molecules act
  • All collisions are perfectly ‘elastic’
  • The molecules of an ideal gas have no size (they occupy zero volume)
41
Q

What are isoelectric ions?

A

Ions that have the same number of electrons of same electronic configuration. E.g. Li+ and He

42
Q

How does ionic radius change: -down group, -across period?

A
  • Down group it gets larger as more shells

- Across period it gets smaller as more protons E.g. N3- –> Al+

43
Q

Properties of ionic compounds? M.P. and strength

A

Melting point is very high as a large amount of energy is needed to overcome electrostatic attractions.
Strength is brittle as any dislocation causes the layers moving and similar ions repel each other.

44
Q

How do ionic substances dissolve?

A

The negative oxygen in water attracts to positive crystal ion.
Positive hydrogen in water attracts to negative crystal ion.

45
Q

What is the octet rule?

A

The tendency for an atom to acquire a noble gas configuration.
As a general rule, atoms share electrons to get the nearest noble gas configuration.

46
Q

How is bond length and strength linked?

A

The larger the atom the longer the bond and therefore weaker.

47
Q

What is a dative covalent bond?

A

A shared pair of electrons which has been provided by one of the bonding atoms only. Also called co-ordinate bond.

48
Q

What is a sigma bond?

A

A bond where the line of electron density is concentrated symmetrically along the line connecting the two atoms.
Can be formed with 2 S orbitals or an S and a P or 2 P orbitals end to end

49
Q

What is a pi bond?

A

Both P orbitals overlapping side by side so a bond is formed above and below the nucleus.(forms one pi bond)

50
Q

What is hybridisation?

A

When an S and a P orbital produce new orbitals called sp orbitals.
These are called hybrid orbitals, this means is that both the s and one p orbital are involved in bonding to the connecting atoms.

51
Q

What are the five allotropes of carbon and a use for each?

A 
Graphine: -2d, very strong
Nano-tube: -Used for injections
Buckminsterfullerene: -60 carbons in hexagons and pentagons
Diamond: -Drill bits
Graphite: -Lubricants
52
Q

Which bond in more repulsive, lone pair or bonding pair?

A

LP LP > LP BP > BP BP

With lone pair lone pair most repulsive

53
Q

What is the Valence shell electron pair repulsion theory?(VESPR)

A

The ability to predict a simple molecule’s shape due to the ideas that electrons repel and will therefore, sit as far from each other as possible

54
Q

What scale measures electronegativity?

A

The Pauling scale

55
Q

What is electronegativity?

A

The ability of an atom IN A MOLECULE to attract electrons to itself.

56
Q

What are the electronegativity trends in the periodic table?

A

-Decreases down group, increases across period

This is because more shells means more repulsion and more protons means more attraction.

57
Q

What is the difference between non-polar and polar bonds in terms of electronegativity?

A
  • Non-polar(covalent): Have similar atoms with similar electronegativity
  • Polar(ionic): One side will have delta -(negative side) and one delta +(positive side) Dipole formed and bond is said to be polar.
58
Q

What does greater electronegativity difference mean?

A

Greater polarity = more ionic

59
Q

What does polarisation do to bonding?

A

It distorts ionic bonds to give them covalent characteristics. Also distorts electron-cloud.

60
Q

What affects how an anion or cation is polarised/polarises?

A

Anions with charge close to 0 or larger radius are more easily polarised.
Cations with a smaller radius or greater charge have a greater ability to polarise anion.

61
Q

What are the two categories of bonds and what bonds are in them?

A

Chemical(Intramolecular): Covalent, Ionic, Metallic, dative covalent
Physical(Intermolecular): London forces, dipole-dipole interaction, hydrogen.

62
Q

What are london (dispersion) foces?

A

Intermolecular forces between molecules. Moving electrons group instantaneously on one side to create instantaneous charge forming temporary dipole

63
Q

What are dipole-dipole interactions?

A

Between molecules with polar bonds working with London forces meaning more energy needed to separate molecules. Creates higher M.P than expected for mass.

64
Q

What is hydrogen bonding?

A

In molecules containing hydrogen bonded to N, O or F. Increases B.p.

65
Q

Give 5 facts about group II elements?

A

Metal, Good conductivity, grey solids, Oxidation state of 2+, Form basic oxides and basic hydroxides. Less reactive than group I but, more dense and high M.P.

66
Q

What colours are flame tests for group II ions?

A 
SR2+ = Crimson
Ca2+ = Brick red
Ba2+ = Apple green
Mg2+ = White
Electrons jump back down energy levels and energy is given out as light
67
Q

Solubility and PH in group II hydroxides?

A

More soluble down group as ionic bonds weaker. Therefore, more OH- ions makes solution more alkaline

68
Q

What is disproportionation? Happens with halogens

A

Simultaneous oxidation and reduction of a single species during a chemical reaction.

69
Q

Change in energy for bonds breaking?

A

Energy taken in - endothermic

70
Q

Change in energy for bonds formed?

A

Energy released - exothermic

71
Q

What is the definition of rate?

A

The change in concentration of a reactant of product per unit time.

72
Q

On a concentration time graph how do you find rate? Hence what formula gives rate?

A

Rate = gradient. = concentration/time

73
Q

What 5 factors increase rate? And why

A
  1. Higher concentration
  2. Higher temperature
  3. Larger Sa:V = small pieces
  4. Higher pressure
  5. Using catalyst.
74
Q

On concentration time graphs, how do you know the order?

A

Zero order = straight
First order = constant half life
Second order = Half life increases with time

75
Q

What are the two types of catalysts?

A 
Heterogeneous = catalyst and reactants in same phase (state).
Homogeneous = catalyst and reactants in different phases (states).
76
Q

What is the definition of activation energy?

A

The energy a colliding molecule must possess before a collision will result in a reaction.

77
Q

What are the axis on a Maxwell-Boltzmann distribution graph?

A

Fraction of molecules on Y. Energy on X

78
Q

What does the addition of a catalyst do to a Maxwell-Boltzmann distribution graph?

A

Shits the Ea line to the left. More molecules with sufficient energy to react.

79
Q

Name an economic benefit of a catalyst?

A
  • Not much needed to catalyse large amount of reactants.

- Lowered energy demands = less environmental impact due to less CO2 = less production costs.

80
Q

What do the letters stand for in the rate equation, Rate = k[A]^m[B]^n?

A

k = rate constant.
A and B are reactants.
m and n are their orders.

81
Q

What does order tell you?

A

How concentration of reactants affects rate.

82
Q

What is the rate determining step (R.D.S)?

A

The slowest step which determines overall rate. Any step which appears after it will not be in rate equation.

83
Q

What is an intermediate?

A

A species formed in one step of a multi-step reaction which is not seen as a reactant or product in the overall equation.

84
Q

What is the Arrhenius equation?

A 
k = A e^(-Ea/RT)
k= rate constant
A= frequency factor
e= maths e
Ea= activation energy
R= gas constant = 8.314 JK-1mol-1
T= Temp. in K
85
Q

How do you rearrange the Arrhenius equation for a graph?

A 
ln k = -Ea/R x (1/T) + ln A
y = ln k
x = 1/T
Grad = -Ea/R
y intercept = ln A
86
Q

Definition of a base?

A

Hydrogen accepter

87
Q

When should you talk about electronegativity?

A

With covalent molecules

88
Q

Mass of electrons, neutrons and protons?

A

0.0005, 1 and 1 in u(atomic mass unit)

89
Q

What is on the x axis and y axis for mass spectroscopy?

A

x axis: mass/charge aka m/z

y axis:%

90
Q

What does emission spectra prove?

A

Evidence of shells

91
Q

What is VESPR?

A

Valence shell electron pair repulsion

92
Q

Why does H20 have a higher boiling point than HF which is higher than NH3?

A

It has the ‘perfect’ number of lone pairs to hydrogen meaning it can maximise hydrogen bonding.
NH3 has too few lone pairs.

93
Q

What does hydrogen bonding change?

A

Higher bp than expected for size of molecule.

H20 less dense than water as hydrogen bonds go stiff.

94
Q

What is the difference in bp for straight vs brached carbon chains?

A

Straight have higher as fit more compactly meaning better imf attraction as larger surface area.

95
Q

What is difference between alkanes than alcohols bp with similar number of electrons?

A

Alcohols have higher bp as hydrogen bond.

96
Q

What are the trends in hydrogen halide bp?

A

HF highest due to hydrogen bonding.

HCl, HBr, HI increasing in that order due to increasing number of electrons and increasing london forces

97
Q

What type of compound structures are there?

A

Giant ionic lattice, simple covalent molecule, giant covalent lattice, giant metallic lattice.

98
Q

What are the different bond types depending on difference in electronegativity?

A

Zero: Covalent
Low: Polar covalent
High: ionic

99
Q

What is the atom economy equation?

A

% atom economy = molar mass of useful product / total molar mass of starting reactants x 100

Chemistry A level (6 decks)

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