Chemistry Paper 1 Flashcards
(129 cards)
1
Q
Physical separation methods (3) and equipment
A
- Filtration (filter paper, conical flask, funnel)
- crystallisation (bunsen burner, tripod stand, wire gauze, evaporation basin)
- simple distillation (flask, bunsen burner, cooling jacket, beaker)
2
Q
what is the name of the liquid that passes through filter paper?
A
the filtrate
3
Q
describe the fractional distillation of crude oil
A
- Crude oil is heated in a fractionating column
- There is a gradient of temperature (cooler going up)
- The substances with lower boiling points evaporate first
- substances cool and re-evaporate on the beads
- different substances cool at different fractions/levels
- The unevaporated substances are emptied through the bottom
4
Q
paper chromatography definitions:
1. stationary phase
2. mobile phase
3. Rf value
A
- chromatography paper
- solvent
- distance travelled by pigment / distance travelled by solvent
5
Q
describe paper chromatography
A
- draw a pencil start line on chromatography paper
- apply inks on the pencil line
- place paper vertically in solvent (solvent must start below pencil line)
6
Q
Plum Pudding model:
1. describe
2. who proposed it?
A
- negative electrons embedded in a cloud of positive charge
- JJ Thomson
7
Q
Alpha scattering experiment:
1. describe
2. conclusions drawn
A
- Alpha particles were fired at an piece of gold foil, 1-2 atoms thick. Most alpha particles passed straight through, some were deflected some bounced directly backwards
- Most of the atom is empty space
The atoms has a positive centre
The centre of the atom contains most of the mass of the atom
8
Q
Describe the nuclear model
A
Mostly empty space
Positive nucleus made up of protons and neutrons
Electrons orbiting in shells (discovered by Niels Bohr)
9
Q
What are isotopes?
A
atoms of the same element with different numbers of neutrons
10
Q
Relative atomic mass equation
A
(mass of isotope 1 x percentage abundance 1) + (mass of isotope 2 x percentage abundance 2) divided by
100
11
Q
Table of octaves:
1. describe
2. who proposed it?
3. problems
A
- elements were ordered by atomic mass in groups of 8: he noticed every 8th element had similar properties
- John Newlands
- Some elements were grouped together with completely different properties
Some were in the some boxes as each other
12
Q
Mendeleev’s periodic table:
1. describe
2. why it was accepted
A
- ordered in order of atomic weight, left some gaps for undiscovered elements, switched the order of some elements to match properties
- the predicted properties of the elements that he left gaps for matched the elements that were later discovered.
13
Q
do metals lose or gain electrons?
A
metals lose electrons to form positive ions
14
Q
Describe group 1 elements) :
1. hard/soft
2. density?
3. reaction with oxygen
4. how does reactivity change down the table?
5. reaction with water
A
- soft
- low density
- they react rapidly with oxygen to form metal oxides
- reactivity increases going down the table
- effervescence, alkaline solution is created
15
Q
Describe group 0 elements:
1. commonly known as?
2. boiling points going down the table
3. reactivity? and why?
A
- the noble gases
- increase down the table
- highly unreactive due to their full & stable outer electron shell
16
Q
Describe group 7 elements:
1. commonly known as?
2. how many outer shell electrons?
3. what is special about them?
4. melting/boiling points down the group?
5. mass down the group?
6. what type of bond when reacting with other non-metals?
7. what type of bond when reacting with metals?
8. reactivity down the group?
A
- The Halogens
- 7
- they are covalently bonded to another atom of the same element
- increases
- increases
- covalent
- ionic
- decreases
17
Q
Describe transition metals:
1. hard/soft
2. density?
3. reactivity?
4. what is special about them?
5. can be used as?
6. melting points?
A
- hard and strong
- high density
- much less reactive than group 1
- they form coloured compounds
- catalysts
- high
18
Q
describe the particles in a solid
A
particles vibrate around a fixed position
19
Q
describe the particles in a liquid
A
particles are close together but they can flow over each other
20
Q
describe the particles in a gas
A
particles move in random directions with random speeds
21
Q
limitations of the simple particle model
A
- Assumes that particles are solid spheres
- Does not show the forces between particles
22
Q
Ionic bonding between groups 1 and 7
A
group 1 elements lose their single outer electron to the group 7 element, causing both atoms to have full outer electron shells.
group 1 forms positive ions and group 7 forms negative ions
23
Q
Ionic bonding between groups 2 and 6
A
group 2 elements lose 2 outer shell electrons to the group 6 elements, causing both to have full outer shells.
24
Q
Properties of ionic compounds
A
- ionic compounds form giant structures (giant ionic lattices). In a giant ionic lattice, every positive ion is surrounded by negative ions and every negative ion is surrounded by positive ions
- Very high boiling/melting points : strong electrostatic forces of attraction
- They can conduct electricity when molten or in aqueous solution but not when solid.
25
properties of small covalent molecules
1. low melting and boiling points : weak intermolecular forces
2. they do not conduct electricity
26
what is covalent bonding
1. the bonding between two non-metal elements where they share electron (s)
27
diamond
- giant covalent molecule
- very hard
-very high melting point
- cannot conduct electricity
- each atom forms covalent bonds with four others
28
silicon dioxide
- giant covalent molecule
- very high melting/boiling point
- many covalent bonds take a lot of energy to overcome
29
graphite
- giant covalent molecule
-soft and slippery : made up of hexagonal layers that can slide over each other due to no covalent bonds between them
-very high melting/boiling point
- each carbon atom forms covalent bonds with 3 others
- good conductor of both electricity and heat: each carbon atom has 1 electron not covalently bonded. they are released from the carbon atoms: they are delocalised electrons.
30
Graphene
- conductor of electricity: delocalised electrons
- strong
- high melting/boiling point
31
Fullerenes
- hollow shapes
- hexagonal rings of carbon atoms (usually)
- buckminsterfullerene : 60 carbon atoms arranged in a hollow sphere.
uses:
to deliver drugs into toe body
lubricants
catalysts
32
Bonding in polymers
- polymers are many monomers joined together
- monomers: double carbon bond
- polymers: single carbon bond
- most polymers are solid at room temperature : strong intermolecular forces
33
DRAW A POYLMER REPEATING UNIT
CANNOT PUT IN FLASHCARDS
34
Metallic bonding
In metals: outer electrons are not attached to a specific atom --> they form a sea of free delocalised electrons. therefore, the metal atoms are positive metal ions.
- There is a strong electrostatic attraction between the sea of electrons and the positive metal ions
- metals can be bent/shaped since atoms can slide over each other
35
alloys
- mixtures of metals
- sifferent sized atoms makes it harder for layers to slide over each other: making alloys harder than pure metals
36
Nanoparticles
- coarse particles: contains thousands of atoms
- fine particles: can contain several thousands atoms
- nanoparticles: contain a few hundred atoms
-as the size decreases, the SA : V ration increases by the same factor. (e.g. 10x smaller, 10x larger surface area : vol. ratio)
uses:
sunscreen,catalysts, deoderants, cosmetics etc.
37
limitations of bonding diagrams
dot and cross - do not tell us the shape of the molecule
stick diagram - cannot tell which electron came from which atom. No idea of outer electrons that are not in bonds
ball and stick diagram - atoms look like solid spheres, ions are shown as spaced apart but they are really closely pact. Only shows a small part of the giant ionic lattice : gives a mistaken impression of the size.
38
Law of conservation of mass
no atoms are lost or made during a chemical reaction so the total mass of the products is equal to that of the reactants
39
sulfate ion formula
SO4 ^2-
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nitrate ion formula
NO3 ^1-
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carbonate ion formula
CO3 ^2-
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Formula of ionic compounds
- there must be no overall charge
- the charges must cancel each other out
- e.g. Na+ + O^2- --> Na2O
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Relative atomic mass
- Ar
- the average mass of the isotopes of that element, weighted to take the abundance of each isotope into account
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Relative formula mass of a compound
the sum of the relative atomic masses of the atoms in the formula
45
Relative atomic mass formula
(Ar1 x abundance 1) + (Ar2 x abundance 2) + (Ar3...
/ 100
(all divided by 100)
46
percentage by mass of an element in a compound
(total relative atomic mass of atoms of that element / relative formula mass of compound (Mr) x 100
47
calculating mass of an element/compound
MASS = MR X MOLE
m = mr x n
(mass is Mr. Mole!)
48
Avogadro's constant
The atomic mass of an element in grams = 1 mole of that element.
6.02 × 10^23
49
using moles to balance equations
using the Mr. mole grid, the lowest possible ratio of the moles = the co-efficient of the substance
50
limiting reactant
the reactant that is fully used up and is stopping more product from being formed.
51
concentration (simple definition)
the mass of a solute in a given solution
52
solute
a substance that is dissolved in a solvent
53
unit of concentration
g/dm^3
54
percentage yield
(actual yield / total possible yield) x 100
- this shows how good a process is (yield can be 100%, but there may still be waste products)
55
reasons why percentage yield is not 100%
- some of the product can be lost when separated from the reaction mixture
- there may be unexpected products produced
- reversible reactions may not go to completion / products may react to reform the reactants
56
Atom economy
(mass of useful product(s) / total mass of reactants (or total mass of all products) x 100
These tells you how efficient a reaction. Atom economy is a measure of how much of the starting materials end up as useful products.
57
importance of high atom economy
- lower price
- higher sustainability / less waste
58
concentration definition (higher)
the number of moles of a solute in a given volume of solution
59
concentration unit (higher)
mol/dm^3
60
moles in terms of concentration
n = c (mol/dm^3) x V (dm^3)
61
volume of 1 mol. of a gas
24 dm^3
(at room temperature and pressure)
temp: 20 degrees Celsius
pressure: 1 atmosphere
62
volume in terms of moles
dm^3 = n x 24
63
Metals + Oxygen
The metal is oxidised
this forms: METAL OXIDES
- This is an oxidation reaction
64
reduction reaction
the loss of oxygen
65
oxidation reaction
the gain of oxygen
66
metal + water
metal hydroxide + hydrogen
67
Reactivity series
Potassium
Sodium
Lithium
Calcium
Magnesium
Aluminium
Carbon
Zinc
Iron
Tin
Lead
Hydrogen
Copper
Silver
Gold
68
comparing the reactivity of metals
potassium/sodium/lithium/calcium can be ordered by their reaction with water
less reactive metals can be ordered by their reaction with dilute acids
69
metal + acid
metal + acid --> Salt + hydrogen
70
metals in reactions
metals lose electrons to form positive ions
the reactivity of a metal is based on its tendency to lose electrons.
71
Displacement reactions
A more reactive element will push out and replace (displace) a less reactive element from a compound
72
Extraction of metals
carbon is used in a displacement reaction to obtain metals that are less reactive than it
73
Oxidation
- the loss of electrons
- the gain of oxygen
74
Reduction
- The gain of electrons
- The loss of oxygen
75
Acids in aqueous solutions
produce Hydrogen ions (H+)
76
Base + Acid
Base + acid --> Salt + Water
77
base v alkali
bases are insoluble in water whilst alkalis are soluble
78
Alkalis in aqueous solutions
produce hydroxide ions (OH-)
79
Sulfuric acid
H2SO4
80
Nitric acid
HNO3
81
Metal carbonates + acid
Metal carbonates + acid --> Salt + water + carbon dioxide
82
RP1: Making soluble salts
Equipment:
- beaker
- wire gauze
- tripod stand
- bunsen burner
- evaporating basin
-heatproof mat
method:
1. start with a fixed volume of sulfuric acid (this must be the limiting reactant or it will contaminate the salt)
2. gently heat the acid with a Bunsen Burner until it is almost (but not) boiling
3. Add small amounts of copper oxide to the acid and stir using a glass rod
4. continue adding copper oxide until in excess (no more dissolves)
5. Filter excess copper oxide out
6. pour copper sulfate solution into an evaporating basin until roughly 1/2 the solution remains
7. leave the basin overnight until copper sulfate crystals appear
83
strong acid
fully ionises in aqueous solution
includes: Hydrochloric acids, sulfuric acid, nitric acid
84
weak acid
partially ionise in aqueous solutions
includes: carbonic acid, ethanoic acid, citric acid
85
strong vs weak acids
- Lower pH = stronger acid
- As the pH decreases by 1 unit, the concentration of hydrogen ions increases by 10x
86
concentration of acids
dilute acids: fewer molecules per unit of volume than concentrated acids
87
RP2: Titration
1. use a pipette to transfer 25cm^3 base/alkali into a conical flask (reduces chance of splashing)
2. Add 5 drops of phenolphthalein to the alkali in the conical flask
3. place the flask onto a white tile to make colour change clear
4. fill a beuret with acid
5. Add the acid slowly, until you start to see a colour change (from pink to colourless) and then add it drop by drop (SWIRLING THE WHOLE TIME)
6. Read the volume of acid added : the eye must be level with the bottom of meniscus
7. repeat the titration until two readings are within 0.1cm^3 of each other (concorrdant results)
88
Titration calculation
moles = concentration x volume
89
cm^3 to dm^3
divide by 1000
90
mol/dm^3 to g/dm^3
multiply by Mr
91
Why is electrolysis done with aqueous or molten solutions (+ what are they called)
Forces of attraction are broken, allowing the charged particles to move freely as they could not before. This means they can conduct electricity.
These solutions are called electrolytes
92
negative electrode in electrolysis
cathode
93
positive electrode in electrolysis
anode
94
What forms at the cathode in electrolysis?
- Positive metal ions undergo reduction at the Cathode to form metal atoms.
-
95
Electrolysis of aluminium oxide
1. aluminium oxide is mixed with cryolite to reduce the boiling point
- This reduces the energy required which saves mone
2. an electric current is applied to the molten aluminum oxide
96
aluminium oxide electrolysis half equations
Al3+ + 3e- --> Al
2O2- --> O2 + 4e-
97
Whys is electrolysis expensive
- lots of energy to produce an electric current
- lots of energy to melt the aluminium oxide
98
Electrolysis of aqueous solutions
- solution contains 2 positive and 2 negative ions : H+ and a metal ion, OH- and a non-metal ion
- unreactive anode + cathode (e.g. platinum)
99
what forms at the cathode in an aqueous electrolysis
- Hydrogen forms when the metal is more reactive than it
- otherwise the metal form s
100
what forms at the anode in an aqueous electrolysis
- Oxygen forms at the anode
- UNLESS halide ions are present - the the halogen forms
101
RP3 : Electrolysis
1. pour approx. 50cm^3 copper chloride solution into a beaker
2. a plastic petri dish is paced over the beaker - insert carbon graphite rods into each hole of the petri dish (these are electrodes)
3. attach crocodile leads to the rods of and then connect the rods to the terminals of a low-voltage power supply (4V)
4. cathode will be coated in copper
5. anode will have bubbles of chlorine gas - test for this by seeing if litmus paper becomes bleached
102
chlorine gas test
if litmus paper becomes bleached, chlorine is present
103
Exothermic reactions (and example)
energy of the reactants decreases
energy of the surroundings increases
- neutralisation / certain oxidation reactions
104
endothermic reactions
energy of the reactants increases
energy of the surroundings decreases
105
exothermic reaction uses
- hand warmers
- self-heating cans for food/drink
106
bond energy calculations
all bonds have an energy
( KJ/mol.)
1. work out the total bond energy to break the bonds in the reactants (provided in the question, simply use addition). Breaking these bonds means the bond energy is positive (putting in energy)
2. work out the total bond energy to make bonds in the products. Making bonds means the bond energy is negative (releasing energy)
3. add the positive bond energy of the reactants to the negative bond energy of the products (subtract the product energy from the reactant energy)
4. Negative bond energy = exothermic (energy is released)
Positive bond energy = endothermic (energy is required)
107
RP4 : Temperature change
independent variable : volume of sodium hydroxide solution
dependent variable : maximum temperature reached
control variables: HCl conc., sodium hydroxide AND HCl conc.
1. measure 30cm^3 dilute HCl (measuring cylinder)
2. transfer the acid into a polystyrene cup and place this cup in a beaker so it doesn't tip over
3. measure and record the temperature of the acid (thermometer)
4. measure 5cm^3 sodium hydroxide solution into the polystyrene cup
5. fit a plastic lid with a thermometer through it - uses the thermometer to stir gently
6. record temperature rise when it stops changing / record highest temperature reached
7. repeat with increasingly large volumes of sodium hydroxide solution increasing by 5cm^3 each time until 40cm^3
108
RP4: temperature change
Why does the temperature decreases after the maximum temperature is reached?
The HCl becomes the limiting factor, so there is excess sodium hydroxide solution meaning it cannot all react. This meas the energy produced is no greater than the maximum yet it is spread out over a wider volume.
109
chemical cell basic set up
2 different metals are placed in an electrolyte to produce electricity
110
chemical cells
- 2 different metals placed in an electrolyte (setup)
- cells can only produce electricity for a certain amount of time
- cells only produce electricity if the 2 metals have different reactivities : the greater the difference in reactivity of the metals, the greater the potential difference of the cell
111
battery (ies)
- multiple cells connected in series to produce a higher voltage
-
112
Alkaline batteries
- non-rechargeable : the reactants run out and no more electricity is produced
- there is no way to reverse the reaction
113
rechargeable batteries
- the chemical reactions are reversible, so batteries can be used again
114
Hydrogen fuel cells
- in the fuel cell, hydrogen is reacted with either pure oxygen or air
- a chemical reaction takes place inside the fuel cell, producing an electric current
- hydrogen is oxide, this creates a waste product of water
115
hydrogen fuel cell half equations
At the negative electrode: 2H2 + 4OH- → 4H2O + 4e-
At the positive electrode: O2 + 2H2O + 4e- → 4OH-
116
advantages of hydrogen fuel cells compared to rechargeable batteries
- they will produce electricity for as long as hydrogen is provided whereas batteries need to be recharged
- they do not get less efficient with use/time however batteries do
- they can be a source of drinkable water e.g. on aircraft
117
disadvantages of hydrogen fuel cells vs rechargeable batteries
- hydrogen is an explosive gas that is difficult to store safely but no dangerous fuels are required for batteries
- they produce a relatively low p.d. so many need to be used together however batteries can produce a greater p.d.
118
carbon nanotubes
carbon nanotubes: high length:diameter ratio. they have a high tensile strength. they have delocalised electrons: good conductors of heat/electricity.
They can be used to reinforce materials e.g. tennis rackets
- they are cylindrical fullerenes
119
uses of nanoparticles
- catalysts
- suncream
- cosmetics
- medicine
- electronics
120
homologous group
a family of compounds that share similar chemical properties and the same functional group
121
what is crude oil
Crude oil is a finite resource found in rocks. Crude oil is the remains of an ancient biomass consisting mainly of plankton that was buried in mud.Crude oil is a mixture of a very large number of compounds. Most of the compounds in crude oil are hydrocarbons
122
general formula for homologous series of alkanes
CnH2n+2
123
general formula for homologous series of alkenes
CnH2n
124
factors that affect the rate of chemical reactions
- concentration of reactants in solution
- temperature
- catalysts
- surface area of reactants
- pressure of reacting gases
125
RP5: Rates of reaction
1. use a measuring cylinder to put 10cm^3 sodium thiosulfate in a conical flask
2. place the conical flask onto a printed black cross
3. add 10cm^3 HCl to the flask
4. swirl the solution and start a stopwatch
5. when the you can no longer see the cross stop the clock
6. repeat with lower conc. of sodium thiosulfate
7. repeat these and calculate means
126
RP5: rates of reaction (gas produced)
1. use a measuring cylinder to place 10cm^3 HCl in a conical flask
2. attach a bung and delivery tube to the flask
3. place the delivery tube into a container filled with water
4. place a measuring cylinder filled with water over the delivery tube
5. add a 3cm strip of Mg to the HCl
6. every 10s measure the vol. of hydrogen gas in the measuring cylinder until the increase stops
7. repeat with different concs and calculate a mean
127
repeatable
the experiment can be done again under the same conditions and the same equipment and the results will be the same
128
reproducible
the same results can be achieved with a different experiment / set up/ equipment
129
catalysts
catalysts increase the rate of reaction by providing a pathway for the reaction with a lower activation energy
