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OCR B Chemistry > Elements Of Life > Flashcards

Elements Of Life Flashcards

1
Q

Describe Dalton’s model of the atom

A

Solid spheres, where different spheres were different elements

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2
Q

Describe and give evidence for Thompson’s atomic model

A

Plum pudding model, negative raisin electrons and positive pudding
Charge and mass measurements proved existence of smaller negative particles - electrons

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3
Q

Describe and give evidence for Rutherford’s atomic model

A

Model of tiny, dense, positive nucleus and cloud of negative electrons around (empty space in between)
Geiger-Marsden experiment, alpha particles fired at gold sheet, most went through but some deflected

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4
Q

Describe and give evidence for Bohr’s atomic model

A

Electrons arranged in (quantised) energy levels around nucleus
Atomic emission/absorption spectra - lines are shown rather than whole spectrum - levels must be present

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5
Q

Define atomic number

A

Number of protons in nucleus of an atom

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6
Q

Define mass number

A

Total number of protons and neutrons in atomic nucleus

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7
Q

Define relative atomic mass

A

Mean atomic mass compared to 1/12th the mass of a C12 atom

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8
Q

Define relative molecular mass

A

Average molecular mass compared to 1/12th of a C12 atom

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9
Q

Define relative isotopic mass

A

Mass of a specific isotope compared to 1/12th of a C12 atom

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10
Q

Define a mole

A

Amount of a substance that contains same number of particles as 12g of C12 (6.02 x 10^23 particles)

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11
Q

Define ionic equation

A

An equation that omits spectator ions

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12
Q

Formula for percentage yield

A

Percentage yield = actual yield / theoretical yield x 100

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13
Q

Formula for percentage composition my mass

A

Mass of desired substance / mass of mixture x 100

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14
Q

How is a standard solution prepared from a conc solution?

A

C1V1 = C2V2 is used to calculate measurements
Calculated volume of conc solution added to volumetric flask
Distilled water added to bro mg solution up to line

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15
Q

What shapes are s and p orbitals?

A

s - spherical
p - dumbbell

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16
Q

Define orbital

A

Regions of space around an atom where electrons are most likely to be found (contain max 2 electrons)

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17
Q

Describe tendencies of electrons as they fill orbitals

A

Lazy + antisocial
Lowest energy orbitals first
One in each orbital if possible

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18
Q

Why does the 4s orbital fill before 3d orbital?

A

4s at lower energy level than 3d

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19
Q

What is fusion?

A

Two lighter nuclei collide, and combine to form a heavier nucleus (and release energy)

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20
Q

Conditions for fusion

A

Very high temp and pressure

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21
Q

What is covalent bonding?

A

Electron pair is shared between 2 atoms.
Strong bond due to strong electrostatic attraction.

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22
Q

Why do giant covalent lattices have greater melting point than simple covalent molecules? (Give examples for each)

A

Covalent bonds in lattice much stronger than intermolecular bonds between simple molecules
Lattice - diamond, silicon dioxide
Simple - oxygen, carbon dioxide

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23
Q

Why are ionic substances only conductive when dissolved?

A

Free movement of ions to carry charge

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24
Q

Why do covalent and ionic lattices have high bp?

A

Both have strong bonds to many atoms, so lots of (heat) energy needed to break many strong bonds

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25
Q

Linear bond angle + conditions

A

180°, two bonded electron groups

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26
Q

Triangular planar bond angle + conditions

A

120° and 3 bonded electron groups

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27
Q

Bent bond angle + conditions

A

104.5° and 2 bonded groups, 2 lone pairs

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28
Q

Trigonal pyramidal bond angle + conditions

A

107° and 3 bonded groups, 1 lone pair

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29
Q

Tetrahedral bond angle + conditions

A

109.5° and 4 bonding groups

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30
Q

Octahedral bond angle + conditions

A

90° and 6 bonded groups

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31
Q

Explain lone pair repulsion theory

A

Electron pairs repel (same negative charge)
To reduce repulsion, they move as far from each other as possible
Lone pairs repel more than bonded pairs, causing diff bond angles/shapes

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32
Q

What is first ionisation enthalpy

A

The enthalpy change from the removal of 1 mol electrons from 1 mol of gaseous atoms

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33
Q

What’s does group number mean in periodic table

A

Number of electrons in atoms outer shell

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34
Q

What does period mean in periodic table

A

Number of electron shells in the atom

35
Q

Why does first ionisation enthalpy increase as you move across a period

A

More attraction between outer shell electrons and positive nucleus, due to increase in protons, so more energy needed to remove them

36
Q

Nitrate ion formula

A

NO3 -

37
Q

Sulphate ion formula

A

SO4 2-

38
Q

Carbonate ion formula

A

CO3 2-

39
Q

Hydroxide ion formula

A

OH -

40
Q

Ammonium ion formula

A

NH4 +

41
Q

Hydrogencarbonate ion formula

A

HCO3 -

42
Q

Charge of most transition metals in period 4

A

2+

43
Q

Why does first ionisation enthalpy decrease down a group

A

More shielding (more electrons between nucleus and outer shell)
Longer distance between nucleus and outer shell
Even though protons have increased, attraction between nucleus and outer shell is lower so less energy needed to remove

44
Q

How do the charge densities of group 2 ions affect the thermal stability of their carbonates?

A

Higher charge density - more polarising - distorts carbonate ion more - lower thermal stability (break down at lower temps)

45
Q

How to test for Fe 2+ ions

A

Add OH -
Green precipitate forms

46
Q

How to test for Fe 3+ ions

A

Add OH -
Orange precipitate forms

47
Q

How to test for Cu 2+ ions

A

Add OH -
Blue precipitate forms

48
Q

How to test for NH4 + ions

A

Add NaOH (aq) to tube with NH4 +
Heat
Vapours from tube will turn damp red litmus paper blue

49
Q

How to test for Al 3+ ions

A

Add OH -
White precipitate forms

50
Q

How to test for halide ions

A

Dissolve substance in HNO3
Add AgNO3
Ag+(aq) + X-(aq) ➡️ AgX(s) (Where X is halide)
Silver chloride is white precipitate (dissolves in dilute ammonia)
Silver bromide is cream precipitate (dissolves in conc ammonia)
Silver iodide is yellow precipitate (doesn’t dissolve)

51
Q

How to test for SO4 2- ions

A

Add Ba 2+ ions
White precipitate forms

52
Q

How to test for CO3 2- ions

A

Add dilute nitric acid
Effervescence occurs

OR

Add dilute HCl
Bubble gas (CO2) through lime water, which will turn cloudy

53
Q

How to test for Li+ , Na+ , K+ and Ca+ ions

A

Use flame tests

54
Q

Give flame test colours

A

Sodium - orange
Calcium - brick red
Lithium - crimson
Potassium - lilac
Barium - apple green
Copper - blue-green

55
Q

How to test for Pb 2+ ions

A

Add OH- ions,
White precipitate forms
adding more OH- ions causes precipitate to dissolve

56
Q

What is an acid

A

Species that donates H+ ions

57
Q

What is a base

A

Species that accepts H+ ions in solution (releases OH- ions)

58
Q

What is an alkali

A

A base that is soluble in water

59
Q

What is a neutralisation reaction

A

Reaction between a base and acid

60
Q

Explain Mash, Basho, cashoco

A

Mash
Metal + acid -> salt + hydrogen
Basho
Base + acid -> salt + hydrogen + oxygen
Cashoco
Carbonate + acid -> salt + H2O + CO2

61
Q

What is metallic bonding

A

Lattice of metal ions + sea of delocalised electrons

62
Q

Describe and explain the trend in reactivity with water as you move down group 2

A

Forms hydroxides (+ H2) with water, outer electrons further from nucleus so are lost more easily

63
Q

Explain the trend in pH and solubility reacting with water as you go down group 2 oxides

A

Forms group 2 hydroxides, which are strongly alkaline
As you go down, they get more alkaline because they also get more soluble

64
Q

Describe the trend in group 2 carbonate solubility

A

As you go down the group, the solubility descreases

65
Q

Compare and contrast the appearances of atomic emission spectra and atomic absorption spectra

A

Emission:
Dark background, colour lines
Absorption:
Black lines, coloured background

Both have more lines, which are closer together, at higher frequencies

66
Q

Explain what causes the lines in an atomic emission spectra, and why they are vary for different elements

A
  • electrons are excited up to higher energy levels
  • they drop back down very quickly, releasing energy in the form of a photon of light
  • the frequency of the photon is proportional to the distance dropped
    ΔΕ=hf
  • each photon corresponds to a line on the spectra - higher energy levels are closer together so higher frequencies have lines closer together
  • each element has a unique pattern of distances between energy levels, so frequency of photons emitted and therefore spectra will also be unique
67
Q

State mass + mol equation

A

moles = mass/Mr

68
Q

State vol mol equation

A

Moles = vol (in cm3)/1000 x concentration

69
Q

State dilution equation

A

C1V1 = C2V2

70
Q

State gas volume and mol equation

A

Moles = volume/24 dm3

71
Q

Define periodicity and give an example

A

The repeating pattern of properties when elements are put in atomic number order (across a period)
Melting point (and bp) increases across a period until simple covalent nonmetals, where it drops down very low and starts to decrease

72
Q

Explain physical properties of ionic substances

A

High mp and bp - lots of energy needed to break strong ionic bonds
Don’t conduct electricity as solid, but do in solution/molten - as ions are free to move and carry charge
Soluble - polar substance, forms ion-dipole bonds with water

73
Q

Explain physical properties of simple covalent structure

A

Relatively low mp and bp - weak id-id bonds between molecules so little energy needed to break
Do not conduct electricity - no charged particles free to move
Not very soluble (usually) - as often nonpolar

74
Q

Explain properties of giant covalent lattices

A

Very high mp and bp - many strong covalent bonds, lots of energy needed to break
Do not conduct electricity usually (excluding graphite, which can conduct due to delocalised electrons between carbon sheets)
Not soluble (nonpolar)

75
Q

Explain properties of metals

A

High mp and bp - strong electrostatic attraction (metallic bonding) , lots of energy needed to break
Good conductor, delocalised electrons that are free to move can carry charge
Not soluble (nonpolar)

76
Q

Name two elements that break the pattern with numbers of bonded electrons in outer shell

A

awkward Buggers
Beryllium and Boron

77
Q

Name three groups of ionic compounds that are always soluble

A

All group one salts
All ammonium salts
All nitrate salts

78
Q

What are the only ionic compounds that contain halogens that are insoluble

A

Lead (II) halide and silver halide

79
Q

What are the only 4 sulfate compounds that are insoluble

A

Barium sulfate
Lead (II) sulfate
Silver sulfate
Calcium sulfate

80
Q

What are the only two types of carbonates that are soluble

A

Group one carbonates
Ammonium carbonates

81
Q

What does flash the cash mean

A

FlaSH - for positive ions
Flame test first, then add sodium hydroxide to test for precipitates

CaSH - for negative ions
Test for Carbonate first (nitric acid or HCl)
Then Sulfate (add barium nitrate)
Then Halides (precipitate with silver nitrate)

82
Q

How would you prepare an insoluble salt e.g. barium sulfate, giving examples of reactants

A

Mix sulfate ions (sodium sulfate solution) and barium ions (barium chloride solution)
Filter to collect precipitate that forms, rinse and dry

83
Q

How would you make a soluble salt using a neutralisation reaction
E.g. copper sulfate

A

Use acid that contains anion (sulfuric acid)
And insoluble base that contains cation (copper carbonate)
Mix and allow neutralisation to form copper sulfate solution
Filter to remove excess copper carbonate
Evaporate water by warming until crystals form
Collect crystals when mostly dry by filtration, rinse and then dry

84
Q

State equation to calculate percentage uncertainty

A

Total uncertainty / measurement size x 100

OCR B Chemistry (11 decks)

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