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elements of life Flashcards

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1
Q

what is an elements mass number

A

the sum of the protons and neutrons

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2
Q

what is an elements atomic number

A

the number of protons

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3
Q

what is an isotope

A

atoms of the same element with the same atomic number, but a different number of neutrons, so a different mass number

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4
Q

how to work out number of particles

A

number of particles = number of moles x avogadro’s number

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5
Q

mole equation for mass

A

mass = Mr x moles

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6
Q

mole equation for concentration

A

moles = concentration x volume

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7
Q

mole equation for volume of gases

A

moles = volume x 24 dm^3

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8
Q

ionic equations

A

show the ions that are formed in solution and show which particles are reacting

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9
Q

empirical formula definition

A

the simplest whole number ratio of elements in a compound

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10
Q

empirical formula process

A
  1. write out elements
  2. write out %s as masses
  3. divide these by element Mr to get number of moles
  4. divide by smallest number of moles
  5. multiply all by 2 if a 0.5 value
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11
Q

working out molecular formula process

A

work out Mr of empirical formula
divide by Mr of molecular formula
use this number to multiple all the atoms in the empirical formula

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12
Q

calculating water of crystallisation

A
  1. write out 2 molecules involved
  2. write out masses of each molecule
  3. divide these by relative molecular mass to get moles
  4. divide all of these by the smallest number of moles
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13
Q

percentage yield equation

A

percentage yield = ( actual yield / theoretical yield ) x 100

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14
Q

measuring the mass of a solid

A
  1. Zero the balance
  2. Place a weighing boat onto balance, add in the solid
  3. Record the weight of both the weighing boat and solid
  4. Empty the solid into the container where it will be used
  5. Reweigh the weighing bottle
  6. Subtract this mass from the combined weight to find the mass of solid used
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15
Q

measuring volumes of liquid in a volumetric pipette

A
  1. Dip the pipette into the solution and use a pipette filler to draw liquid into it
  2. Stop once the bottom of the meniscus is touching the line on the pipette
  3. Transfer to the glassware the solution is being transferred to
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16
Q

measuring volumes of liquid with a burette

A
  1. Use a funnel to empty the solution into the burette
  2. Run a small amount into a beaker until there are no air bubbles
  3. Record the volume to the nearest 0.05 cm3
  4. Slowly open the tap, letting the solution out until the end point is reached
  5. At the end point, close the tap and read again to the nearest
    0.05 cm3
  6. Subtract this from the first reading to get the titre
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17
Q

calculating standard solutions

A
  1. work out number of moles of solution required
  2. use number of moles to calculate the mass of the solid
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18
Q

calculating volumes to use when making solutions

A

volume to use = ( final concentration / initial concentration ) x volume required

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19
Q

acid - base titrations process

A
  1. Rinse and fill a burette with the acid
  2. Measure the start volume
  3. Fill a 25 cm3 pipette with the alkali solution and empty into a conical flask
  4. Add 2-3 drops of an indicator
  5. Run the acid into the flask and swirl the flask until a colour change is seen
  6. Record the end volume and note this as a trial titration
  7. Keep repeating this until three titres are reached which are concordant
  8. Concordant results must be within 0.1 cm3 of each other
  9. ind the mean of these - the mean titre
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20
Q

indicators to use for acid - base titration

A
  • phenolphthalein
  • methyl orange
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21
Q

phenolphthalein indicator

A

acid = colourless
base = pink

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22
Q

methyl orange indicator

A

acid = yellow
base = red

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23
Q

what are electrons arranged in

A

electrons are arranged in energy levels

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24
Q

each electron shell is made up of

A

subshells

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25
Q

what are the electron subshells

A

s, p, d, f

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26
Q

what are orbitals

A

clouds of negative charge that hold electrons

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27
Q

dispersion of electrons within orbitals

A

two electrons are located in each orbital

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28
Q

dispersion of orbitals within subshells

A

each subshell has a different number of orbitals

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29
Q

number of electrons in S subshell

A

2

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30
Q

number of electrons in P subshell

A

6

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31
Q

number of electrons in D subshell

A

10

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32
Q

number of electrons in F subshell

A

14

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33
Q

area of periodic table with S block elements

A

groups 1 and 2

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34
Q

area of periodic table with D block elements

A

transition metals

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35
Q

area of periodic table with P block elements

A

groups 3 to 8

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36
Q

area of periodic table with F block elements

A

actinides and lanthanides

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37
Q

order of which electron subshells are filled

A

s, then p, then d, then f
exception: 3d subshell has higher energy than 4s subshell, so 4s subshell is filled before 3d subshell

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38
Q

s orbital shape

A

sphere

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39
Q

p orbital shape

A

dumb bell

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40
Q

development of the model of the atom

A
  1. Dalton - atoms are small solid spherical particles
  2. Thomson - plumb pudding model (negative electrons in positive atom), discovery of electron
  3. Rutherford - discovered positively charged nucleus
  4. Bohr - electrons were in fixed energy shells
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41
Q

number of orbitals s subshells contain

A

1 orbital

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42
Q

number of orbitals p subshells contain

A

3 orbitals

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43
Q

number of orbitals d subshells contain

A

5 orbitals

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44
Q

number of orbitals f subshells contain

A

7 orbitals

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45
Q

nuclear fusion definition

A

the forcing together of 2 light atomic nuclei to make a heavier nuclei, and thus a new element, releasing energy as gamma radiation

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46
Q

nuclear fusion conditions

A

very high temperature and pressure

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47
Q

electrostatic attraction definition

A

the attraction between a positive and negative particle

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48
Q

electrostatic repulsion definition

A

the repulsion between two particles of the same charge

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49
Q

covalent bonding definition

A

sharing of electrons between non-metals

50
Q

covalent bond features

A
  • insoluble
  • doesn’t conduct
  • low melting and boiling point (due to weak intermolecular forces)
51
Q

dative covalent bond definition

A

covalent bonding where both electrons are from the same atom

52
Q

ionic bonding definition

A

bonding through transference of electrons between a metal and a non-metal

53
Q

ionic bond features

A
  • conduct electricity when molten or dissolved in water
  • arranged in a positive-negative pattern to reduce repulsions and maximise attractions
  • high melting and boiling points due to high electrostatic forces
54
Q

metallic bonding definition

A

bonding through delocalised electrons between metals

55
Q

metallic bonding features

A
  • insoluble in water
  • conducts heat (as vibrations pass between closely packed particles)
  • high melting point (as ions packed close together)
56
Q

shape and angle for 2 bonding pairs 0 lone pairs

A

linear, 180 degrees

57
Q

shape and angle for 2 bonding pairs 1 lone pair

A

non-linear, 117.5 degrees

58
Q

shape and angle for 2 bonding pairs 2 lone pairs

A

non-linear, 104.5 degrees

59
Q

shape and angle for 3 bonding pairs 0 lone pairs

A

trigonal planar, 120 degrees

60
Q

shape and angle for 3 bonding pairs 1 lone pair

A

pyramidal, 107 degrees

61
Q

shape and angle for 4 bonding pairs 0 lone pairs

A

tetrahedral, 109.5 degrees

62
Q

shape and angle for 4 bonding pairs 2 lone pairs

A

square planar, 107 degrees

63
Q

shape and angle for 5 bonding pairs 0 lone pairs

A

trigonal bipyramidal, 120 and 90 degrees

64
Q

shape and angle for 6 bonding pairs 0 lone pairs

A

octahedral, 90 degrees

65
Q

for each lone pair remove ? off bond angle

A

2.5 degrees off bond angle

66
Q

the periodic table is ordered by

A

proton number

67
Q

periodic table groups

A
  • going down columns
  • have the same number of electrons in the outershell
  • all have similar properties
68
Q

periodic table periods

A
  • going across rows
  • have the same number of electron shells
  • have trends within these (periodicity)
69
Q

periodicity trends in melting point - groups 1 to 3

A

metals
- have metallic bonding with delocalised electrons
- general increase in melting points going left to right as metal ions have an increasing positive charge, increasing the number of delocalised electrons and smaller ionic radius, meaning a stronger metallic bond

70
Q

periodicity trends in melting point - group 4

A

giant covalent structure
- has the highest melting point in the period
- as many strong covalent bonds hold the atoms together and thus a large amount of energy is needed to overcome these strong covalent bonds

71
Q

periodicity trends in melting point - groups 5 to 7

A

simple molecular structure
- big drop in melting point
- has a weaker structure, thus weaker imf
- for group 6, melting poitn increases due to larger molecular structure thus it has larger imf
- for group 7 (diatomic). melting point lowers due to smaller simple molecular structure, meaning smaller imf

72
Q

periodicity trends in melting point group 8

A

atom
- lowest melting point in the period as it is only an individual atom, thus it has smaller intermolecular forces and hence a lower melting point

73
Q

group 2 reactions with water

A

group 2 + 2 water –> metal hydroxide + H2

74
Q

group 2 reactions with steam

A

group 2 + H2O (g) –> metal oxide + H2

75
Q

group 2 trends

A

reactivity with water increases down the group

76
Q

group 2 reaction with oxygen

A

2 group 2 + oxygen –> 2 metal oxide

77
Q

thermal decomposition of group 2 carbonates

A

XCO3 –> XO + CO2

78
Q

thermal stability trend down group 2 carbonates

A

thermal stability increases down group 2 carbonates because
as you go down the group, the group 2 metals have the same charge but a lower charge density (as radius increases)
thus as you go down, they distort the negative carbonate ion less
so more energy is required to break down XCO3 into XO and CO2

79
Q

charge density definition

A

the ratio of an ions charge to its volume (size)

80
Q

trends in group 2 carbonates

A
  • solubility decreases down group
81
Q

trends in group 2 hydroxides

A
  • solubility increases down the group
82
Q

first ionisation definition

A

the minimum energy required to remove one mole of electrons from one mole of atoms in a gaseous state (in kJmol-1)

83
Q

trends ionisation energy across periods

A
  • ionisation enthalpy increases across each period as nuclear charge increases so more energy is needed to remove an electron from the outer shell
  • drops in ionisation occur between periodsa s electrons fill a new shell, menaing these valence electrons are further from the nucleus and experience greater shielding from inner shields of electrons
84
Q

successive ionisation energy definition

A

energy required to remove one mole of electrons from a cation

85
Q

trends of successive ionisation energies within atoms

A
  • successive ionisation energies increase because atomic radius decreases and there is a greater attraction between the outer shell electrons and the nucleus
  • large jumps in successive ionisation energy when crossing shells as lower shells are closer to the nucleus meaning less shielding and thus more energy required to remove electrons due to stronger attraction between nucleus and the outershell
86
Q

reactions for formation of salts

A

acid + base –> salt + water
acid + carbonate –> salt + water + carbon dioxide
acid + metal –> salt + hydrogen

87
Q

testing for cations with sodium hydroxide (OH-)

A

Ag+ - brown
Ca 2+ - white
Cu 2+ - blue
Pb 2+ - white
Zn 2+ - white (in excess colourless)
Fe 2+ - green
Fe 3+ - red-brown
Al 3+ - white (in excess colourless)

88
Q

testing for halides

A

add dilute nitric acid
- chloride - white
- bromine - cream
- iodine - yellow

89
Q

testing for lead ions

A

add potassium iodide
forms lead iodide which is yellow

90
Q

Testing for sulfates

A

Add dilute HCl and barium chloride
Sulfates (SO42-) form a white barium sulfate precipitate

91
Q

Litmus paper test

A
  • red litmus paper turns blue in alkaline solutions
  • blue litmus paper turns red in acidic solutions
92
Q

test for ammonium ions

A

add sodium hydroxide (OH- ions) and heat
this forms ammonia gas, which is alkaline and will turn damp red litmus paper blue

93
Q

testing for hydroxides

A

dip red litmus paper into solution, will turn blue if hydroxide present

94
Q

testing for nitrates

A

warm solution and add NaOH and aluminium
forms ammonia gas
turns damp red litmus paper blue

95
Q

flame test for Li+

A

crimson

96
Q

flame test for Na+

A

yellow-orange

97
Q

flame test for K+

A

lilac

98
Q

flame test for Ca 2+

A

dark red

99
Q

flame test for Ba 2+

A

green

100
Q

flame test for Cu 2+

A

blue green

101
Q

acid definition

A

a proton donor, and a compound that produces hydrogen ions (H+) in water

102
Q

base definition

A

proton acceptor

103
Q

alkali definition

A

a base that is soluble in water

104
Q

base + acid reaction

A

base + acid –> salt + water

105
Q

neutralisation reaction definition

A

when a base reacts with an acid to form a salt

106
Q

making water-insoluble salts with precipitation reactions

A
  1. Add the desired solutions into a beaker to form a precipitate of the salt
  2. Filter the precipitate
  3. Wash the precipitate with deionised water
  4. Transfer to a watch glass and dry in an oven (below the melting point of the salt), or in air
  5. At regular intervals, weigh
  6. Once the solid has dried to a constant mass, you can stop heating
107
Q

making salt from an acid and alkali

A
  1. Carry out an acid-base titration to find out how much acid is required to neutralise
    25 cm3 of the alkaline solution
  2. Measure another
    25 cm3 of alkali. Using the burette, add the amount of acid that is required to neutralise the alkali
  3. Transfer the neutralised solution to a clean evaporating basin and heat lightly (to avoid spitting) to evaporate the water
  4. Leave the mixture to cool in the evaporating basin
  5. Filter the mixture, and wash the solid with cold distilled water
  6. Transfer the residue to a watch glass and leave it to dry. Measure its mass at regular intervals; you can stop once it remains constant
108
Q

making salt from an acid and base

A
  1. In a beaker, warm excess of the insoluble base in dilute acid
  2. Continue to heat, then add universal indicator to see if the solution is neutral. If required, more solid base can be added
  3. Leave to cool
  4. Filter off the excess base and transfer the filtrate to a clean, dry evaporating basin
  5. Heat the evaporating basin until salt crystals begin to appear on the sides of the basin
109
Q

reactions of group 2 oxides with acids

A

MO (s) + 2HCl (aq) –> MCl2 (aq) + H2O (l)

110
Q

reaction of group 2 hydroxides with acids

A

M(OH)2 (s) + H2SO4 (aq) –> MSO4 (aq) + 2 H2O (l)

111
Q

group 2 oxide trends in pH

A

group 2 oxides become increasingly more basic as you move down the group

112
Q

electromagnetic spectrum from lowest to highest wavelength

A

ultraviolet
visible
infrared

113
Q

wave theory of light

A

light can be described from its wavelength and frequency

114
Q

wave theory of light equation

A

speed of light = wavelength (m) x frequency (Hz)

115
Q

particle theory of light

A

light can be described as tiny packets of energy called photons

116
Q

particle theory of light equation

A

quantum energy of a photon (J) = planck’s constant x frequency (Hz)

117
Q

Bohr’s light theory

A

when heated, the atom gains energy and its electrons become ‘excited’
- This causes electrons to jump to higher energy levels, further from the nucleus than at their neutral states
- Later the electrons will drop back to lower levels, giving off a single photon of energy, often as visible light
Bohr’s theory is that electrons exist in fixed, quantised energy levels. The ground state is n = 1 is closest to the nucleus.

118
Q

atoms and spectra fingerprinting

A

each atom releases and absorbs photons at different frequencies, and these frequencies can be plotted on absorption / emission spectra and are unique to each element.

119
Q

emission spectra

A

vertical lines of colour from the emissions on a black background
Frequency increases from right to left, and wavelength decreases from right to left

120
Q

absorption spectra

A

vertical black lines from absorption on a coloured background
frequency increases from right to left, and wavelength decreases from right to left

121
Q

flame test method

A
  1. nichrome wire loop dipped in HCl until it produces no colour when placed in a blue bunsen burner flame (to cleanse it)
  2. wire dipped into solid sample of compound being tested
  3. loop is put in blue flame
  4. colour flame produced is recorded
122
Q

calculating relative atomic mass from mass spectra

A

relative atomic mass can be calculated by multiplying the relative isotopic mass by relative abundance (%) for each ion mass detected (for each bar on the graph). The sum of these should then be divided by 100 to get the average mass of each atom of this type

chemistry (10 decks)

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