elements of life Flashcards
(122 cards)
1
Q
what is an elements mass number
A
the sum of the protons and neutrons
2
Q
what is an elements atomic number
A
the number of protons
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
4
Q
how to work out number of particles
A
number of particles = number of moles x avogadro’s number
5
Q
mole equation for mass
A
mass = Mr x moles
6
Q
mole equation for concentration
A
moles = concentration x volume
7
Q
mole equation for volume of gases
A
moles = volume x 24 dm^3
8
Q
ionic equations
A
show the ions that are formed in solution and show which particles are reacting
9
Q
empirical formula definition
A
the simplest whole number ratio of elements in a compound
10
Q
empirical formula process
A
- write out elements
- write out %s as masses
- divide these by element Mr to get number of moles
- divide by smallest number of moles
- multiply all by 2 if a 0.5 value
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
12
Q
calculating water of crystallisation
A
- write out 2 molecules involved
- write out masses of each molecule
- divide these by relative molecular mass to get moles
- divide all of these by the smallest number of moles
13
Q
percentage yield equation
A
percentage yield = ( actual yield / theoretical yield ) x 100
14
Q
measuring the mass of a solid
A
- Zero the balance
- Place a weighing boat onto balance, add in the solid
- Record the weight of both the weighing boat and solid
- Empty the solid into the container where it will be used
- Reweigh the weighing bottle
- Subtract this mass from the combined weight to find the mass of solid used
15
Q
measuring volumes of liquid in a volumetric pipette
A
- Dip the pipette into the solution and use a pipette filler to draw liquid into it
- Stop once the bottom of the meniscus is touching the line on the pipette
- Transfer to the glassware the solution is being transferred to
16
Q
measuring volumes of liquid with a burette
A
- Use a funnel to empty the solution into the burette
- Run a small amount into a beaker until there are no air bubbles
- Record the volume to the nearest 0.05 cm3
- Slowly open the tap, letting the solution out until the end point is reached
- At the end point, close the tap and read again to the nearest
0.05 cm3 - Subtract this from the first reading to get the titre
17
Q
calculating standard solutions
A
- work out number of moles of solution required
- use number of moles to calculate the mass of the solid
18
Q
calculating volumes to use when making solutions
A
volume to use = ( final concentration / initial concentration ) x volume required
19
Q
acid - base titrations process
A
- Rinse and fill a burette with the acid
- Measure the start volume
- Fill a 25 cm3 pipette with the alkali solution and empty into a conical flask
- Add 2-3 drops of an indicator
- Run the acid into the flask and swirl the flask until a colour change is seen
- Record the end volume and note this as a trial titration
- Keep repeating this until three titres are reached which are concordant
- Concordant results must be within 0.1 cm3 of each other
- ind the mean of these - the mean titre
20
Q
indicators to use for acid - base titration
A
- phenolphthalein
- methyl orange
21
Q
phenolphthalein indicator
A
acid = colourless
base = pink
22
Q
methyl orange indicator
A
acid = yellow
base = red
23
Q
what are electrons arranged in
A
electrons are arranged in energy levels
24
Q
each electron shell is made up of
A
subshells
25
what are the electron subshells
s, p, d, f
26
what are orbitals
clouds of negative charge that hold electrons
27
dispersion of electrons within orbitals
two electrons are located in each orbital
28
dispersion of orbitals within subshells
each subshell has a different number of orbitals
29
number of electrons in S subshell
2
30
number of electrons in P subshell
6
31
number of electrons in D subshell
10
32
number of electrons in F subshell
14
33
area of periodic table with S block elements
groups 1 and 2
34
area of periodic table with D block elements
transition metals
35
area of periodic table with P block elements
groups 3 to 8
36
area of periodic table with F block elements
actinides and lanthanides
37
order of which electron subshells are filled
s, then p, then d, then f
exception: 3d subshell has higher energy than 4s subshell, so 4s subshell is filled before 3d subshell
38
s orbital shape
sphere
39
p orbital shape
dumb bell
40
development of the model of the atom
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
41
number of orbitals s subshells contain
1 orbital
42
number of orbitals p subshells contain
3 orbitals
43
number of orbitals d subshells contain
5 orbitals
44
number of orbitals f subshells contain
7 orbitals
45
nuclear fusion definition
the forcing together of 2 light atomic nuclei to make a heavier nuclei, and thus a new element, releasing energy as gamma radiation
46
nuclear fusion conditions
very high temperature and pressure
47
electrostatic attraction definition
the attraction between a positive and negative particle
48
electrostatic repulsion definition
the repulsion between two particles of the same charge
49
covalent bonding definition
sharing of electrons between non-metals
50
covalent bond features
- insoluble
- doesn't conduct
- low melting and boiling point (due to weak intermolecular forces)
51
dative covalent bond definition
covalent bonding where both electrons are from the same atom
52
ionic bonding definition
bonding through transference of electrons between a metal and a non-metal
53
ionic bond features
- 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
metallic bonding definition
bonding through delocalised electrons between metals
55
metallic bonding features
- insoluble in water
- conducts heat (as vibrations pass between closely packed particles)
- high melting point (as ions packed close together)
56
shape and angle for 2 bonding pairs 0 lone pairs
linear, 180 degrees
57
shape and angle for 2 bonding pairs 1 lone pair
non-linear, 117.5 degrees
58
shape and angle for 2 bonding pairs 2 lone pairs
non-linear, 104.5 degrees
59
shape and angle for 3 bonding pairs 0 lone pairs
trigonal planar, 120 degrees
60
shape and angle for 3 bonding pairs 1 lone pair
pyramidal, 107 degrees
61
shape and angle for 4 bonding pairs 0 lone pairs
tetrahedral, 109.5 degrees
62
shape and angle for 4 bonding pairs 2 lone pairs
square planar, 107 degrees
63
shape and angle for 5 bonding pairs 0 lone pairs
trigonal bipyramidal, 120 and 90 degrees
64
shape and angle for 6 bonding pairs 0 lone pairs
octahedral, 90 degrees
65
for each lone pair remove ? off bond angle
2.5 degrees off bond angle
66
the periodic table is ordered by
proton number
67
periodic table groups
- going down columns
- have the same number of electrons in the outershell
- all have similar properties
68
periodic table periods
- going across rows
- have the same number of electron shells
- have trends within these (periodicity)
69
periodicity trends in melting point - groups 1 to 3
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
periodicity trends in melting point - group 4
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
periodicity trends in melting point - groups 5 to 7
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
periodicity trends in melting point group 8
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
group 2 reactions with water
group 2 + 2 water --> metal hydroxide + H2
74
group 2 reactions with steam
group 2 + H2O (g) --> metal oxide + H2
75
group 2 trend with water reactivity
reactivity with water increases down the group
76
group 2 reaction with oxygen
2 group 2 + oxygen --> 2 metal oxide
77
thermal decomposition of group 2 carbonates
XCO3 --> XO + CO2
78
thermal stability trend down group 2 carbonates
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
charge density definition
the ratio of an ions charge to its volume (size)
80
trends in group 2 carbonates
- solubility decreases down group
81
solubility trends in group 2 hydroxides
- solubility increases down the group
82
first ionisation definition
the minimum energy required to remove one mole of electrons from one mole of atoms in a gaseous state (in kJmol-1)
83
trends ionisation energy across periods
- 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
successive ionisation energy definition
energy required to remove one mole of electrons from a cation
85
trends of successive ionisation energies within atoms
- 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
reactions for formation of salts
acid + base --> salt + water
acid + carbonate --> salt + water + carbon dioxide
acid + metal --> salt + hydrogen
87
testing for cations with sodium hydroxide (OH-)
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
testing for halides
add dilute nitric acid
- chloride - white
- bromine - cream
- iodine - yellow
89
testing for lead ions
add potassium iodide
forms lead iodide which is yellow
90
Testing for sulfates
Add dilute HCl and barium chloride
Sulfates (SO42-) form a white barium sulfate precipitate
91
Litmus paper test
- red litmus paper turns blue in alkaline solutions
- blue litmus paper turns red in acidic solutions
92
test for ammonium ions
add sodium hydroxide (OH- ions) and heat
this forms ammonia gas, which is alkaline and will turn damp red litmus paper blue
93
testing for hydroxides
dip red litmus paper into solution, will turn blue if hydroxide present
94
testing for nitrates
warm solution and add NaOH and aluminium
forms ammonia gas
turns damp red litmus paper blue
95
flame test for Li+
crimson
96
flame test for Na+
yellow-orange
97
flame test for K+
lilac
98
flame test for Ca 2+
dark red
99
flame test for Ba 2+
green
100
flame test for Cu 2+
blue green
101
acid definition
a proton donor, and a compound that produces hydrogen ions (H+) in water
102
base definition
proton acceptor
103
alkali definition
a base that is soluble in water
104
base + acid reaction
base + acid --> salt + water
105
neutralisation reaction definition
when a base reacts with an acid to form a salt
106
making water-insoluble salts with precipitation reactions
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
making salt from an acid and alkali
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
making salt from an acid and base
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
reactions of group 2 oxides with acids
MO (s) + 2HCl (aq) --> MCl2 (aq) + H2O (l)
110
reaction of group 2 hydroxides with acids
M(OH)2 (s) + H2SO4 (aq) --> MSO4 (aq) + 2 H2O (l)
111
group 2 oxide trends in pH
group 2 oxides become increasingly more basic as you move down the group
112
electromagnetic spectrum from lowest to highest wavelength
ultraviolet
visible
infrared
113
wave theory of light
light can be described from its wavelength and frequency
114
wave theory of light equation
speed of light = wavelength (m) x frequency (Hz)
115
particle theory of light
light can be described as tiny packets of energy called photons
116
particle theory of light equation
quantum energy of a photon (J) = planck's constant x frequency (Hz)
117
Bohr's light theory
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
atoms and spectra fingerprinting
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
emission spectra
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
absorption spectra
vertical black lines from absorption on a coloured background
frequency increases from right to left, and wavelength decreases from right to left
121
flame test method
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
calculating relative atomic mass from mass spectra
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
