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A LEVEL chemistry > Elements of Life > Flashcards

Elements of Life Flashcards

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1
Q

What is an atomic structure made of?

A
  • 3 sub atomic particles
  • protons and neutrons form the nucleus
  • electrons occupy energy levels/shells orbiting the nucleus
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2
Q

What is the relative atomic mass of each sub-atomic particle?

A

protons: 1
neutrons: 1
electrons: 0.00055 (1/1836)

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3
Q

What is the relative charge of each sub-atomic particle?

A

protons: +1
neutrons: 0 (no charge)
electrons: -1

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4
Q

What is the atomic number?

A

the number of protons in the nucleus of an atom

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5
Q

What is the mass number?

A

the number of protons and neutrons in the nucleus of an atom

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6
Q

What are isotopes?

A

atoms of the same element with different mass numbers

same no. of protons identifying the element with different no. of neutrons

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7
Q

What does mass spectrometry do?

A

measures atomic or molecular mass of different particles in a sample as well as relative abundance of different isotopes in an element

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8
Q

What happens in a mass spectrometer?

A

sample atoms/molecules are ionised into positively charged cations which are separated according to mass m to charge z ratios
separated ions are detected together with relative abundance

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9
Q

How do you calculate relative atomic mass from a mass spectra?

A
  • %abundance of each ion (isotope) calculated by height of peak
  • (relative isotopic mass x relative abundance) … add together each isotope and then divide all by 100
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10
Q

What is nuclear fusion?

A

when 2 light atomic nuclei fuse together to form a single heavier nucleus of a new element, releasing an enormous amount of energy

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11
Q

What are the required conditions for nuclear fusion?

A
  • must be very close together
  • very high temperature (e.g. star) so that nuclei move quickly and collide with more energy to overcome repulsive forces of nucleus
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12
Q

What does the absorption spectra look like?

A
  • a series of black absorption lines on a coloured background
  • missing frequencies of light where they have been absorbed by particles in the chromosphere
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13
Q

What does the emission spectra look like?

A
  • a series of coloured lines on a black background

- correspond to emitted frequencies

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14
Q

What does the emission spectrum show?

A

the frequencies of emitted electromagnetic radiation where electrons of atoms/molecules/ions are raised from their ground state to higher energy states (excited states)
particles lose extra energy when emitting EM radiation

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15
Q

What is the name of the characteristic emission spectrum in the ultraviolet of the EM spectrum called?

A

the lyman series

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16
Q

What is the name of the hydrogen emission spectrum in visible light?

A

the balmer series

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17
Q

Is light seen from stars continuous?

A

no - consists of lines corresponding to absorption or emission of specific frequencies of light - atomic spectra

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18
Q

What are the 2 theories used to describe behaviour of light?

A

wave theory

particle theory

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19
Q

What is the speed of light? c

a constant

A

3.00 x 10(8) ms-1

when travelling in a vacuum

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20
Q

What is the formula relating frequency and wavelength?

A

speed of light c (ms-1) = wavelength λ (m) x frequency v (s-1)

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21
Q

What happens as the wavelength increases?

A

the frequency decreases (c remains the same)

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22
Q

What is the particle theory of light?

A
  • light regarded as steady stream of photons (energy)
  • energy of photons related to position in EM spectrum
  • linked
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23
Q

What is the formula used in the particle theory of light?

A

energy of a photon E (J) = planck constant h (Js-1) x frequency v (s-1)

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24
Q

What is the value of the planck constant?

A

6.63 x 10(-34) Js-1

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25
Q

Bohr’s theory: how is an emission spectrum given off?

A
  • when an atom is excited, electrons jump to higher energy levels
  • later drop back into lower levels emitting extra energy as EM radiation and give off emission spectrum
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26
Q

Bohr’s theory: how is an absorption spectrum given off?

A
  • electrons have been raised to higher levels without then dropping down again
  • correspond to light absorbed by atoms in sample (and with coloured lines of emission spectrum of that element)
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27
Q

What are the main points of Bohr’s theory?

A
  • the electron in the hydrogen atom exists only in certain definite energy levels or electron shells
  • a photon of light is emitted or absorbed when the electron changes from one energy level to another
  • the energy of the photon is equal to the difference between the 2 energy levels (deltaE)
  • since E=hv it follows that the frequency of the emitted or absorbed light is related to deltaE by deltaE=hv
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28
Q

What are quanta?

A

definite quantities of energy possessed by electrons (can only change to specific values and not continuously)

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29
Q

What happens to the energy levels as you move further away from the nucleus of an atom?

A
  • they get higher

- they become closer together (frequency)

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30
Q

What are the maximum numbers of electrons that can be in each shell (up to 7) and what is its principal quantum number n label?

A
  • 1st shell: (n=1) 2 electrons
  • 2nd and 3rd shells: (n=2) (n=3) 8 electrons
  • 4th and 5th shells: (n=4) (n=5) 18 electrons
  • 6th and 7th shells: (n=6) (n=7) 32 electrons
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31
Q

What are the different types of sub-shell and how many electrons can be held in each one?

A

S : 2
P : 6
D : 10
F : 14

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32
Q

Split the number of electrons held in each shell into sub-shells for: n=1

A
  • 2 electrons in s sub-shell

- total = 2

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33
Q

Split the number of electrons held in each shell into sub-shells for: n=2

A
  • 2 electrons in s sub-shell
  • 6 electrons in p sub-shell
  • total = 8
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34
Q

Split the number of electrons held in each shell into sub-shells for: n=3

A
  • 2 electrons in s sub-shell
  • 6 electrons in p sub-shell
  • 10 electrons in d sub-shell
  • total = 18
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35
Q

Split the number of electrons held in each shell into sub-shells for: n=4

A
  • 2 electrons in s sub-shell
  • 6 electrons in p sub-shell
  • 10 electrons in d sub-shell
  • 14 electrons in f sub-shell
  • total = 32
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36
Q

Flame test: colour of Li+ ion flame?

A

bright red

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37
Q

Flame test: colour of Na+ ion flame?

A

yellow

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38
Q

Flame test: colour of K+ ion flame?

A

lilac

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39
Q

Flame test: colour of Ca2+ ion flame?

A

brick red

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40
Q

Flame test: colour of Ba2+ ion flame?

A

apple green

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41
Q

Flame test: colour of Cu2+ ion flame?

A

blue green

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42
Q

How many orbitals does an s sub-shell always contain?

A

1

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43
Q

How many orbitals does an p sub-shell always contain?

A

3

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44
Q

How many orbitals does an d sub-shell always contain?

A

5

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45
Q

How many orbitals does an f sub-shell always contain?

A

7

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46
Q

What is the maximum number of electrons each atomic orbital can hold? (how are they shown?)

A

2

each electron spins at same rate in opposite direction so drawn with opposing arrows

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47
Q

What 4 pieces of information are required when describing an electron?

A
  • the electron shell it’s in
  • the sub-shell
  • its orbital within the sub-shell
  • its spin
48
Q

What is the electronic configuration of an atom?

A

the arrangement of electrons in shells and orbitals

filled in specific order to produce lowest energy arrangement possible

49
Q

How are orbitals filled?

A
  • in order of increasing energy

- orbitals first occupied singly by electrons which keeps electrons as far apart as possible (spins are parallel)

50
Q

What is significant about the 3d sub-shell and the 4s sub-shell?

A

the 4s sub-shell is filled first

still write 3d before 4s

51
Q

How did Mendeleev arrange elements in the periodic table and why did he leave gaps?

A
  • arranged elements in order of increasing atomic mass and so elements with similar properties were in the same vertical group
  • he left gaps for undiscovered elements in order to keep elements with similar properties in the same group
52
Q

How is the periodic table arranged today?

A
  • in order of atomic number
  • in blocks s d p and f (similarities)
  • vertical columns groups = no. of electrons in highest energy level
  • horizontal rows periods = no. of energy levels
53
Q

What is periodicity?

A

the occurrence of periodic patterns

54
Q

What are the arrangements of noble gases called?

A

closed shell arrangements

all sub-shells fully occupied

55
Q

What ions do groups 1 2 6 and 7 form?

A

1: +1
2: +2
6: -2
7: -1

56
Q

What are dot and cross diagrams used for?

A

used to represent the way that atoms bond together

57
Q

What is covalent bonding?

A

when 2 non-metals share electrons to bond together

58
Q

What are lone pairs?

A

pairs of electrons not involved in bonding

59
Q

What are bonding pairs?

A

electron pairs that form bonds

60
Q

What is a dative covalent bond?

A

this is when both bonding electrons come from the same atom (shown by an arrow)

61
Q

What are the physical properties of covalently bonded simple molecules?

A
  • strong intramolecular forces
  • intermolecular forces weak (electrostatic forces between molecules)
  • low melting/boiling points (not much energy required to break intermolecular forces)
  • do not conduct electricity (no charged particles)
  • do not dissolve readily in water
62
Q

For a linear shape molecule…
how many bond pairs are there?
what are the bond angles?
what is an example?

A
  • 2 bond pairs
  • 180 degree bond angles
  • e.g. BeCl2
63
Q

For a trigonal planar shape molecule…
how many bond pairs are there?
what are the bond angles?
what is an example?

A
  • 3 bond pairs
  • 120 degree bond angles
  • e.g. BF3
64
Q

For a tetrahedral shape molecule…
how many bond pairs are there?
what are the bond angles?
what is an example?

A
  • 4 bond pairs
  • 109.5 degree bond angles
  • e.g. CH4
65
Q

For a trigonal bipyramidal shape molecule…
how many bond pairs are there?
what are the bond angles?
what is an example?

A
  • 5 bond pairs
  • 120 degree and 90 degree bond angles
  • e.g. PF5
66
Q

For a octahedral shape molecule…
how many bond pairs are there?
what are the bond angles?
what is an example?

A
  • 6 bond pairs
  • 90 degree bond angles
  • e.g. SF6
67
Q

What is the order of repulsion between lone pairs and bond pairs?

A

lone pair-lone pair > lone pair-bond pair > bond pair-bond pair
(lone pairs have greater repulsion)

68
Q

For a pyramidal shape molecule…
how many bond/ lone pairs are there?
what are the bond angles?
what is an example?

A
  • 3 bond pairs, 1 lone pair
  • 107 degree bond angles
  • e.g. NH3
69
Q

For a angular shape molecule…
how many bond/ lone pairs are there?
what are the bond angles?
what is an example?

A
  • 2 bond pairs, 2 lone pairs
  • 104.5 degree bond angles
  • e.g. H2O
70
Q

For a square planar shape molecule…
how many bond/ lone pairs are there?
what are the bond angles?
what is an example?

A
  • 4 bond pairs, 2 lone pairs
  • (90 degree bond angles?)
  • e.g. XeF4
71
Q

What is relative atomic mass? Ar

A

the mass of an atom relative to carbon-12

Ar have no units

72
Q

What is the formula for amount in moles?

A

amount in moles n (mol) = mass m (g)
————————-
molar mass M (gmol-1)

73
Q

What is a mole?

A

a unit for measuring the amount of each substance

contains as many particles as there are in atoms in 12g of carbon-12

74
Q

What is molar mass?

A

the mass of one mole

75
Q

What is the relative formula mass? Mr

A

the sum of the relative atomic masses of the elements making it up

76
Q

What are formula units?

A

basic units/building blocks atoms/molecules/ion groups that match the formulae of the substances

77
Q

What does the amount in moles of formula units equal?

A

molar mass

78
Q

What is the avogadro constant?

A

the number of formula units in one mole of substance
it’s a constant NA
value: 6.02 x 10(23) formula units per mole

79
Q

What does the molecular formula tell you?

A

the actual numbers of different types of atom

80
Q

What is the empirical formula?

A

the simplest ratio for the moles of atoms

81
Q

What is water of crystallisation?

A

when crystals of ionic lattices include molecules of water which are fitted into attic regularly

82
Q

What is the difference between hydrous and anhydrous?

A
  • hydrous means water molecules are present

- anhydrous means water molecules are removed

83
Q

What factors can reduce %yield?

A
  • loss of products from reaction vessels
  • side-reactions occurring, producing unwanted by-products
  • impurities in reactants
  • changes in temperature and pressure
  • if reaction is reversible in an equilibrium system
84
Q

How do you calculate % yield?

A

experimental yield
—————————– x100
theoretical yield

85
Q

What are the state symbols for a solid, liquid, gas and aqueous solution?

A

solid : s
liquid: l
gas: g
aqueous solution: aq

86
Q

What is ionic bonding?

A

when a metal and a non-metal bond by transferring electrons (from M to NM)
cations and anions formed held together by electrostatic forces

87
Q

What are the 4 salt making equations?

A
  • acid + alkali –> salt + water
  • acid + base –> salt + water
  • acid + carbonate –> salt + water + carbon dioxide
  • acid + metal –> salt + hydrogen
88
Q

What ionic substances are soluble and which ones aren’t?

A

soluble:

  • all G1 compounds
  • ammonium compounds
  • nitrates

insoluble:

  • barium, calcium, lead and silver sulfates
  • silver and lead halides (chlorides/bromides/iodides)
  • all metal carbonates
  • metal hydroxides (excluding G1 and aluminium)
89
Q

How can ionic substances conduct electricity?

A
  • when in a solution or molten
  • ions become differentiated, acting independently of each other and spread out
  • ions can carry charge and therefore conduct
90
Q

What are spectator ions?

A

ions not involved in the reaction

91
Q

Why must ionic equations always have state symbols?

A

help identify precipitates etc e.g. ionic precipitation (a suspension of solid particles produced by a chemical reaction in a solution)

92
Q

What do you use sodium hydroxide to test for (when it forms a precipitate)?

A

metal ions e.g. Cu2+ (blue) , Fe(ii) 2+ (dirty green), Fe(iii) 3+ (orange/brown)

93
Q

What do you use potassium iodide to test for (when it forms a precipitate)?

A

Pb2+ ions (bright yellow)

94
Q

What do you use silver nitrate to test for (when it forms a precipitate)?

A

halide ions e.g. Cl- (white), Br- (cream), I- (pale yellow)

95
Q

What do you use barium chloride to test for (when it forms a precipitate)?

A

sulfate ions (white)

96
Q

What are the physical properties of ionic compounds?

A
  • solid at room temperature and pressure
  • have regular lattices of positive and negative ions in 3D (regularly shaped crystals)
  • high melting/boiling points because of strong electrostatic forces requiring lots of energy to be broken
  • conduct electricity when dissolved or molten
97
Q

How does metallic bonding work?

A

a giant lattice structure of metal cations is held together by a sea of delocalised valence electrons which are free to move

98
Q

What are the properties of metallic lattices?

A
  • conduct electricity due to delocalised electrons free to carry charge
  • relatively high melting/boiling points
  • insoluble
  • relatively hard but malleable
99
Q

What are covalent networks and what properties do they have?

A

e.g. diamond and silicon dioxide
large covalent networks formed from strong covalent bonds between networks
- high melting/boiling points because it takes a lot of energy to break intramolecular forces but intermolecular forces still weak
- insoluble
- do not conduct electricity (exception graphite with a free delocalised electron each carbon atom)

100
Q

What is first ionisation enthalpy?

A

the first ionisation enthalpy of an element is the energy needed to remove one electron from every atom in one mole of isolated gaseous atoms of the element - a mole of gaseous ions with one positive charge are formed

101
Q

How does reactivity change across the s block of metals?

A
  • elements become more reactive down a group

- become less metallic across a period left to right (G1 more reactive than G2 of same period)

102
Q

What are the general patterns of first ionisation energies of elements 1-20?

A
  • as you cross the period it becomes harder to remove an electron so ionisation enthalpy increases
  • going down a group, it becomes easier to remove an electron (more shielding from nucleus) so next period increases less
  • group 1 have the lowest ionisation enthalpies whilst the noble gases have the highest ionisation enthalpies
103
Q

Why is there a decrease in ionisation enthalpy between beryllium and boron?

A

Be) 1s2 2s2
B) 1s2 2s2 2p1
the s sub-shell is lower in energy than the p sub-shell therefore less energy is needed to remove the outer electron of boron despite increased nuclear charge because the p electron can more strongly resist the nucleus (the s sub-shell is stable)

104
Q

Why is there a decrease in first ionisation enthalpies between nitrogen and oxygen?

A

the extra repulsion from the paired electron sub-shell in oxygen means less energy is needed to remove one of the paired electrons despite increased number of protons

105
Q

As you go down group 2, what is the trend of reactivity with water? What does it form?

A
  • increases

- form hydroxides and hydrogen

106
Q

As you go down group 2, what is the trend in thermal stability of carbonate? What does it form?

A
  • decomposes at increasingly higher temperature

- form oxides and carbon dioxide

107
Q

As you go down group 2, what is the trend in pH of hydroxide in water? What does it form?

A
  • increasing pH

- form alkaline solutions (not very soluble)

108
Q

As you go down group 2, what is the trend in solubility of hydroxide and carbonate? What does it form?

A
  • hydroxide - increase in solubility

- carbonate - decrease in solubility

109
Q

What is thermal stability?

A

how much energy / temperature needed to start decomposition

measured in terms of charge density (smaller=higher density and less stable)

110
Q

What is an acid?

A

a compound that disassociates in water to produce hydrogen H+ ions

111
Q

What is a base?

A

a compound that reacts with an acid - a proton acceptor- to produce water and a salt

112
Q

What is an alkali?

A

a base that dissolves in water to produce hydroxide OH- ions

113
Q

When a solution is made, what will its concentration depend on?

A
  • the amount of solute

- the final volume of solution

114
Q

How do you find the concentration in MOLdm-3?

A

molar mass (g mol-1)

115
Q

How do you work out the amount of solute in a particular volume?

A

amount (mol) = concentration of solution (mol dm-3) x volume of solution

116
Q

How many cm3 is in 1dm3?

A

1000cm3

117
Q

What is the equation relating amount of moles, concentration and volume?

A

amount n (moles) = concentration c (mol dm-3) x volume v (dm3)

A LEVEL chemistry (2 decks)

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