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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
Linear bond angle + conditions
180°, two bonded electron groups
26
Triangular planar bond angle + conditions
120° and 3 bonded electron groups
27
Bent bond angle + conditions
104.5° and 2 bonded groups, 2 lone pairs
28
Trigonal pyramidal bond angle + conditions
107° and 3 bonded groups, 1 lone pair
29
Tetrahedral bond angle + conditions
109.5° and 4 bonding groups
30
Octahedral bond angle + conditions
90° and 6 bonded groups
31
Explain lone pair repulsion theory
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
32
What is first ionisation enthalpy
The enthalpy change from the removal of 1 mol electrons from 1 mol of gaseous atoms
33
What’s does group number mean in periodic table
Number of electrons in atoms outer shell
34
What does period mean in periodic table
Number of electron shells in the atom
35
Why does first ionisation enthalpy increase as you move across a period
More attraction between outer shell electrons and positive nucleus, due to increase in protons, so more energy needed to remove them
36
Nitrate ion formula
NO3 -
37
Sulphate ion formula
SO4 2-
38
Carbonate ion formula
CO3 2-
39
Hydroxide ion formula
OH -
40
Ammonium ion formula
NH4 +
41
Hydrogencarbonate ion formula
HCO3 -
42
Charge of most transition metals in period 4
2+
43
Why does first ionisation enthalpy decrease down a group
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
How do the charge densities of group 2 ions affect the thermal stability of their carbonates?
Higher charge density - more polarising - distorts carbonate ion more - lower thermal stability (break down at lower temps)
45
How to test for Fe 2+ ions
Add OH - Green precipitate forms
46
How to test for Fe 3+ ions
Add OH - Orange precipitate forms
47
How to test for Cu 2+ ions
Add OH - Blue precipitate forms
48
How to test for NH4 + ions
Add NaOH (aq) to tube with NH4 + Heat Vapours from tube will turn damp red litmus paper blue
49
How to test for Al 3+ ions
Add OH - White precipitate forms
50
How to test for halide ions
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
How to test for SO4 2- ions
Add Ba 2+ ions White precipitate forms
52
How to test for CO3 2- ions
Add dilute nitric acid Effervescence occurs OR Add dilute HCl Bubble gas (CO2) through lime water, which will turn cloudy
53
How to test for Li+ , Na+ , K+ and Ca+ ions
Use flame tests
54
Give flame test colours
Sodium - orange Calcium - brick red Lithium - crimson Potassium - lilac Barium - apple green Copper - blue-green
55
How to test for Pb 2+ ions
Add OH- ions, White precipitate forms adding more OH- ions causes precipitate to dissolve
56
What is an acid
Species that donates H+ ions
57
What is a base
Species that accepts H+ ions in solution (releases OH- ions)
58
What is an alkali
A base that is soluble in water
59
What is a neutralisation reaction
Reaction between a base and acid
60
Explain Mash, Basho, cashoco
Mash Metal + acid -> salt + hydrogen Basho Base + acid -> salt + hydrogen + oxygen Cashoco Carbonate + acid -> salt + H2O + CO2
61
What is metallic bonding
Lattice of metal ions + sea of delocalised electrons
62
Describe and explain the trend in reactivity with water as you move down group 2
Forms hydroxides (+ H2) with water, outer electrons further from nucleus so are lost more easily
63
Explain the trend in pH and solubility reacting with water as you go down group 2 oxides
Forms group 2 hydroxides, which are strongly alkaline As you go down, they get more alkaline because they also get more soluble
64
Describe the trend in group 2 carbonate solubility
As you go down the group, the solubility descreases
65
Compare and contrast the appearances of atomic emission spectra and atomic absorption spectra
Emission: Dark background, colour lines Absorption: Black lines, coloured background Both have more lines, which are closer together, at higher frequencies
66
Explain what causes the lines in an atomic emission spectra, and why they are vary for different elements
- 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
State mass + mol equation
moles = mass/Mr
68
State vol mol equation
Moles = vol (in cm3)/1000 x concentration
69
State dilution equation
C1V1 = C2V2
70
State gas volume and mol equation
Moles = volume/24 dm3
71
Define periodicity and give an example
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
Explain physical properties of ionic substances
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
Explain physical properties of simple covalent structure
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
Explain properties of giant covalent lattices
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
Explain properties of metals
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
Name two elements that break the pattern with numbers of bonded electrons in outer shell
awkward Buggers Beryllium and Boron
77
Name three groups of ionic compounds that are always soluble
All group one salts All ammonium salts All nitrate salts
78
What are the only ionic compounds that contain halogens that are insoluble
Lead (II) halide and silver halide
79
What are the only 4 sulfate compounds that are insoluble
Barium sulfate Lead (II) sulfate Silver sulfate Calcium sulfate
80
What are the only two types of carbonates that are soluble
Group one carbonates Ammonium carbonates
81
What does flash the cash mean
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
How would you prepare an insoluble salt e.g. barium sulfate, giving examples of reactants
Mix sulfate ions (sodium sulfate solution) and barium ions (barium chloride solution) Filter to collect precipitate that forms, rinse and dry
83
How would you make a soluble salt using a neutralisation reaction E.g. copper sulfate
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
State equation to calculate percentage uncertainty
Total uncertainty / measurement size x 100

OCR B Chemistry (11 decks)

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