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november mock Flashcards

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1
Q

what are the 3 states of matter

A

solid, liquid, gas

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2
Q

arrangement of a solid

A

has least energy – particles are not moving/are just vibrating and they are arranged regularly and very closely together and all touching

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3
Q

arrangement of a liquid

A

particles have more energy than those in a solid, but less than those in a gas and the particles are closer together but have a random arrangement

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4
Q

arrangement of a gas

A

particles have the most energy – as the particles are the most spread apart with a random arrangement

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5
Q

what is melting

A

solid -> liquid

increase in particle energy

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6
Q

what is boiling

A

liquid -> gas

increase in particle energy

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7
Q

what is freezing

A

liquid -> solid

decrease in particle energy

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8
Q

what is evaporation

A

liquid -> gas (at surface)

increase in energy

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9
Q

what is condensation

A

gas -> liquid

decrease in energy

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10
Q

what is sublimation

A

solid -> gas

increase in energy

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11
Q

what is deposition

A

gas -> solid

decrease in energy

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12
Q

how does melting occur

A

The process requires heat energy which transforms into kinetic energy, allowing the particles to move.
It occurs at a specific temperature known as the melting point which is unique to each pure solid

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13
Q

how does boiling occur

A

This requires heat which causes bubbles of gas to form below the surface of a liquid, allowing for liquid particles to escape from the surface and from within the liquid

It occurs at a specific temperature known as the boiling point which is unique to each pure liquid

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14
Q

how does freezing occur

A

This is the reverse of melting and occurs at exactly the same temperature as melting, hence the melting point and freezing point of a pure substance are the same

It requires a significant decrease in temperature (or loss of thermal energy) and occurs at a specific temperature which is unique for each pure substance

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15
Q

how does evaporation occur

A

Evaporation occurs only at the surface of liquids where high energy particles can escape from the liquid’s surface at low temperatures, below the boiling point of the liquid

The larger the surface area and the warmer the liquid/surface, the more quickly a liquid can evaporate

Evaporation occurs over a range of temperatures, but heating will speed up the process as particles need energy to escape from the surface

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16
Q

how does condensation occur

A

When a gas is cooled its particles lose energy and when they bump into each other, they lack energy to bounce away again, instead grouping together to form a liquid

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17
Q

what is diffusion

A

The random movement of particles from a high to low concentration

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18
Q

why does diffusion occur slower in a liquid than a gas

A

Diffusion in liquids is slower than in gases because the particles in a liquid are closely packed together and move more slowly

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19
Q

why does the colour get paler when a solution is diluted

A

when you add water the evenly spaced particles become much more spread out making the colour paler

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20
Q

solvent

A

the liquid in which something is dissolved in

eg the water in sea water

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21
Q

solute

A

the substance which is dissolved in something to form a solution

eg salt in sea water

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22
Q

solution

A

the mixture of solvent and dissolved solute

eg sea water

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23
Q

saturated solution

A

a solution with the maximum concentration of solute dissolved in the solvent

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24
Q

soluble

A

something that will dissolve in a specific liquid

eg salt is soluble in water

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25
Q

insoluble

A

something that will not dissolve in a specific liquid

eg sand in water

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26
Q

what is solubility normally expressed as

A

Solubility can be expressed in g per 100 g of solvent

for example 10g of salt can dissolve (before it gets saturated) in 100g of water which means the solubility is 10g/100g

if something is given not out of 100g then it must be converted
eg 10g/50g = 20g/100g

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27
Q

what is solublility

A

Solubility is a measurement of how much of a substance (solute) will dissolve in a given volume of a liquid (solvent)

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28
Q

what do solubility curves represent

A

Solubility graphs or curves represent solubility in g per 100 g of solvent with different temperature

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29
Q

what does above the solubility line mean

A

above the line means the solute will no longer dissolve

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30
Q

what does below the solubility line mean

A

the solute will dissolve

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31
Q

1.7C practical - investigate the solubility of a solid in water at a specific temperature

A
  1. set a water bath to a specific temperature (eg 50c)
  2. measure out 100g of water into a beaker
  3. place the beaker in the water bath
  4. add a thermometer to the beaker and once it has reached the same temperature as the water bath
  5. add a known mass of solid and stir until it has dissolved
  6. keep adding known masses in increments until not more dissolves
  7. write down what mass of solid that was added before solution got saturated
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32
Q

whats an element

A

a substance made up of the same atoms

eg a beaker containing all carbon

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33
Q

whats a compound

A

a substance made up of 2 or more elements chemically joined

eg a beaker containing water (H2O)

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34
Q

what a mixture

A

a substance made up of 2 or more elements not chemically bonded

eg a beaker containing O2 and N2 but they aren’t joined to each other

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35
Q

why does a pure substance have a fixed melting point

A

there is only 1 type of thing in it so it will all have the same melting point

eg water boils at 100c

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36
Q

how to distinguish between a pure substance and a mixture

A

slowly heat the substance up as a pure substance will boil at one temperature whereas the mixture will boil of a range of temperatures

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37
Q

what does simple distillation separate

A

a liquid and soluble solid from a solution (e.g., water from a solution of salt water) or a pure liquid from a mixture of liquids

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38
Q

simple distillation method

A
  1. heat the solution
  2. the thing with the lower boiling point will evaporate first leaving the other substance

eg in a water and ethonal mixture, the ethanol would evaporate first leaving the water

  1. if you want to keep both substances use a simple distillation set up (a delivery tube surrounded with cold water, to recondense the evaporated substance, leading into a beaker)
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39
Q

what does fractional distillation separate

A

This is used to separate two or more liquids that have different boiling points which are mixed together

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40
Q

fractional distillation method

A
  1. place liquid in boiling tube
  2. attach a fractionating column above it with a thermometer at the top
  3. attach a delivery tube to the fractionating column with a constant supply of cold water around it
  4. have a beaker after the delivery tube to colloct the liquid
  5. slowly heat until the boiling point of the lowest substance
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41
Q

what does filtration separate

A

an undissolved solid from a mixture of the solid and a liquid / solution (eg sand from a mixture of sand and water)

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42
Q

filtration method

A
  1. place a piece of filter paper in a funnel
  2. place the funnel above a beaker
  3. pour the substance through the funnel

the liquid will be in the beaker and the solid will be left in the filter paper

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43
Q

what does crystallisation separate

A

a dissolved solid from a solution, when the solid is much more soluble in hot solvent than in cold solvent

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44
Q

crystallisation method

A
  1. heat the solution which allows any excess solvent (liquid) to evaporate
  2. once the ruffly half the solution has evaporated
  3. remove from heat to allow cooling
  4. crystals will appear of the solute that was dissolved
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45
Q

chromatography method

A

Use a ruler to draw a horizontal pencil line (as a pens ink would run into the other samples) 2 cm from the end of the chromatography paper
Use a different capillary tube to put a tiny spot of each colouring A, B, C and D on the line
Use the fifth tube to put a small spot of the unknown mixture U on the line
Make sure each spot is no more than 2-3 mm in diameter and label each spot in pencil
Pour water into the beaker to a depth of no more than 1 cm (to avoid the samples washing into the solvent container) and clip the top of the chromatography paper to the wooden spill. The top end is the furthest from the spots
Carefully rest the wooden spill on the top edge of the beaker. The bottom edge of the paper should dip into the solvent

Allow the solvent to travel undisturbed at least three quarters of the way up the paper
Remove the paper and draw another pencil line on the dry part of the paper as close to the wet edge as possible. This is called the solvent front line
Measure the distance in mm between the two pencil lines. This is the distance travelled by the water solvent
For each of food colour A, B, C and D measure the distance in mm from the start line to the middle of the spot
calculate the Rf values for each one. Compare the Rf values from the known samples to the unknown dye to see what it is composed of

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46
Q

how to calculate Rf value

A

Rf = distance travelled by substance / distance traveled by solvent

These values are used to identify the components of mixtures
The Rf value of a particular compound is always the same but it is dependent, however, on the solvent used
If the solvent is changed then the value changes
Calculating the Rf value allows chemists to identify unknown substances because it can be compared with Rf values of known substances under the same conditions
These values are known as reference values

The Rf value will always lie between 0 and 1; the closer it is to 1, the more soluble is that component in the solvent

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47
Q

what does a chromatogram do

A

This technique is used to separate substances that have different solubilities in a given solvent (e.g., different coloured inks that have been mixed to make black ink)

An impure substance will show up with more than one spot, a pure substance should only show up with one spot

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48
Q

where is a proton found

A

in the nucleus

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49
Q

where is a neutron found

A

in the nucleus

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50
Q

where is an electron found

A

orbiting the nucleus in shells

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51
Q

relative mass of a proton

A

1

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52
Q

relative mass of a neutron

A

1

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53
Q

relative mass of a electron

A

1/1840

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54
Q

charge of a proton

A

+1

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55
Q

charge of a neutron

A

0

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56
Q

charge of an electron

A

-1

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57
Q

what is the atomic number

A

number of protons an atom has

eg the atomic number of carbon is 6

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58
Q

what is the mass number

A

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

eg carbons mass is 12

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59
Q

what is an isotope

A

an atom with the same number of protons and electrons but a different number of neutrons

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60
Q

what is the RAM (relative atomic mass or Ar)

A

the weighted average mass of an atom of an element, compared to 1/12th the mass of a carbon 12 atom

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61
Q

how to calculate RAM (Ar) from isotope abundance

A

( % of isotope A x mass of isotope A ) + ( % of isotope B x mass of isotope B ) / 100

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62
Q

how are the elements arranged in the periodic table

A

in the order of atomic number

in groups and periods

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63
Q

how to tell a metal or non-metal based on electrical conductivity

A

metals are good electrical conductors

non-metals are poor electrical conductors

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64
Q

how to split the periodic table into metals and nonmetals

A

zigzag line going down right between boron and carbon

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65
Q

what forms when metals react with water

A

a base

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66
Q

what forms when nonmetals react with water

A

an acid

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67
Q

how to calculate the Mr of a molecule from the Ar’s of the atoms in the molecule

A

add up the masses of each of the atoms

the big numbers do not apply
small numbers apply only to the element that they are behind
everything in a bracket is multiplied by the small number outside the bracket

for water (h2O)

H= 1 x 2 = 2 (1 is the mass of hydrogen and times by 2 because there is 2 hydrogen’s in water)

0= 16 x 1 = 16

2+16 = 18

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68
Q

mass, moles and Ar/Mr formula

A

moles = mass / Ar or Mr

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69
Q

how to calculate reacting masses

A
  1. balance the symbol equation if not alr
  2. calculate the moles from mass given
  3. work out the ratio of substances (using the
    large numbers)
  4. convert moles to mass
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70
Q

how to calculate percentage yield

A

actual yield / theoretical yield x 100

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71
Q

how to experimentally find the formula for metal oxides

A
  1. measure mass of crucible and lid
  2. measure a mass of the metal and add
    into the crucible
  3. heat strongly and lift the occasionally
  4. every 5 minutes remove from heat and
    reweigh
  5. once mass stops increasing stop heating
  6. calculate metal oxide weight by:
    final mass - mass of crucible and lid

to find the formula:

  1. work out the mass of the metal and the
    oxygen individually
  2. work out the moles of the metal and the
    oxygen individually
  3. divide both moles by the smallest mole
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72
Q

how to experimentally find the formula for water / hydrated salts

(water of crystallisation)

A
  1. measure mass of evaporating basin
  2. measure a mass of hydrated salt and add
    into the evaporating basin
  3. heat strongly until the salt turns completely
    white
  4. reweigh the crucible and white salt
  5. calculate anhydrous salt weight by:
    the basins mass and white salt - just the
    basin
  6. calculate the mass of water lost by:
    hydrated salt mass - unhydrated salt mass

to find the formula:

  1. find mass of the water and the mass of the
    unhydrated salt
  2. work out the Mr of the salt and water
    individually
  3. work out the moles of the salt and the
    water individually
  4. divide both moles by the smallest mole
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73
Q

what is the molecular formula

A

shows the number and type of each atom in a molecule

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74
Q

what is the empirical formula

A

simplest whole number ratio of the atoms of each element of a compound

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75
Q

how to calculate the empirical formula

A
  1. work out the moles of each atom in the
    molecule
  2. divide both sets of moles by the smallest
    mole
  3. if the answer contains a 0.5 then times all
    by 2
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76
Q

how to determine the formula of a metal oxide by combustion

A
  1. measure mass of crucible and lid
  2. measure a mass of the metal and add
    into the crucible
  3. heat strongly and lift the lid occasionally
  4. every 5 minutes remove from heat and
    reweigh
  5. once mass stops increasing stop heating
  6. calculate metal oxide weight by:
    final mass - mass of crucible and lid

then calculate the empirical formula

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77
Q

how to determine the formula of a metal
oxide by reduction

A
  1. Measure mass of the empty boiling tube
  2. Place metal oxide into a horizontal boiling tube and measure the mass again
  3. Support the tube in a horizontal position held by a clamp
  4. A steady stream of natural gas(methane) is passed over the copper(II)oxide and the excess gas is burned off
  5. The copper(II)oxide is heated strongly using a Bunsen burner
  6. Heat until metal oxide completely changes colour, meaning that all the oxygen has been removed
  7. Measure mass of the tube with remaining metal powder and subtract the mass of the tube to work out mass of metal

to work out empirical formula:

  1. work out mass of oxygen by
    mass of metal oxide - metal
  2. then work out empirical formula
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78
Q

how are ions formed

A

by the loss or gain of electrons

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79
Q

what is the charge of a Ag ion

A

1+

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80
Q

what is the charge of a Cu(II) ion

A

2+

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81
Q

what is the charge of a Fe(II) ion

A

2+

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82
Q

what is the charge of a Fe(III) ion

A

3+

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83
Q

what is the charge of a Pb ion

A

2+

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84
Q

what is the charge of a Zn ion

A

2+

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85
Q

what is the charge of a hydrogen ion

A

1+

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86
Q

what is the charge of a hydroxide (OH) ion

A

1-

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87
Q

what is the charge of a ammonium (NH4) ion

A

1+

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88
Q

what is the charge of a carbonate (CO3) ion

A

2-

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89
Q

what is the charge of a nitrate (NO3) ion

A

1-

90
Q

what is the charge of a sulfate (SO4) ion

A

2-

91
Q

ionic bonding definition

A

the electrostatic force of attraction between 2 oppositely charged ions

92
Q

why do compounds with giant ionic lattices have high melting and boiling points

A

strong electrostatic forces acting between the oppositely charged ions
These forces act in all directions and a lot of energy is required to overcome them giving them a high melting and boiling point

93
Q

when do ionic compounds conduct electricity

A

in an aqueous solution or when molten as the ions have more energy and are able to move

not as a solid as there are no free ions to move

94
Q

covalent bond definition

A

the electrostatic force of attraction between the nuclei and a shared pair of electrons

95
Q

why are substances with simple molecular formulas liquids or gases or solids with low boiling points

A

the covalent forces between atoms are strong but the inter-molecular forces connecting each molecule are weak

the weak inter-molecular forces are easy to overcome (require a low temperature to break) meaning at room temperature the bonds are broken and therefore are liquids or gases or solids with a low boiling points

96
Q

why does the melting points of simple molecular structures increase as the relative molecular mass increases

A

the increase in mass means there will also be more electrons and therefore there are more intermolecular forces of attraction that need to be overcome which increases the melting point

97
Q

why do giant covalent structures have a higher melting point then simple molecular ones

A

to melt giant covalent bonds you must break the strong covalent bonds whereas melting a simple molecular structure only requires overcoming the weak intermolecular forces

98
Q

do covalent compounds conduct electricity

A

no as there are no free electrons to move

99
Q

how does the structure of diamond influence its physical properties including electrical conductivity and hardness

A

giant covalent structure

very hard as in order to break you have to break the strong covalent bonds

does not conduct electricity because there are no free electrons or ions

high melting and boiling point

100
Q

how does the structure of graphite influence its physical properties including electrical conductivity and hardness

A

giant covalent structure

arranged in layers where there are weak intermolecular forces in between the layers however there are layers themselves are bonded covalently which means the layers can slide over each other

high melting and boiling point

soft and slippery

can conduct as it has declocalized electrons

used as a lubricant

101
Q

how does the structure of bucminster fullerene influence its physical properties including electrical conductivity and hardness

A

simple molecular structure

low melting and boiling points as it has weak intermolecular forces so easy to overcome

it is also soft for the same reason

it can conduct as there are delocalised electrons

used as a lubricant or a drug delivery system

102
Q

definition of a metallic bond

A

the electrostatic force of attraction between the delocalised free moving electrons and the positive metal ions

103
Q

why do metals conduct so well

A

have a large sea of delocalised electrons that are free to move and conduct

104
Q

why do ionic compounds only conduct when molten or in an aqueous solution

A

They cannot conduct electricity in the solid state as the ions are in fixed positions within the lattice and are unable to move but when molten or in solution the particles become free to move as the lattice is broken down.

105
Q

what % of nitrogen is there in the air

A

78%

106
Q

what % of oxygen is there in the air

A

21%

107
Q

what % of CO2 is there in the air

A

0.04%

108
Q

what % of argon is there in the air

A

0.9%

109
Q

how to determine the percentage by volume of oxygen in air using experiments involving the reactions of metals (e.g. iron) with air

A
  1. place an excess of wet iron filings into a conical flask (100cm3) full of air
  2. place a bung on the top of the conical flask connected to a syringe

over time the iron will rust (react with oxygen) until all of the oxygen is used up. this will move the syringe reading from 100 to 79ish showing oxygen makes up 21% of the air

110
Q

how to determine the percentage by volume of oxygen in air using experiments involving the reactions of non-metals (e.g. phosphorus) with air

A
  1. place a small amount of phosphorus onto a evaporating dish in a bell jar which is sitting in a trough of water
  2. set the phosphorus on fire

As the phosphorus burns it uses up the oxygen inside the bell jar and the water level rises

By making careful measurements of water levels before and after the experiment you can determine the percentage of oxygen in the air

111
Q

combustion of magnesium in oxygen

A

intense white flame
white powder produced

112
Q

combustion of hydrogen in oxygen

A

exothermic
water is produced

113
Q

combustion of sulpher in oxygen

A

blue flame
colourless, poisonous gas produced

114
Q

what is the formation of carbon dioxide from the thermal decomposition of metal carbonates

A

metal carbonate → metal oxide + carbon dioxide

115
Q

what is the happens when copper(II) carbonate is thermally decomposed

A

copper(II) carbonate → copper(II) oxide + carbon dioxide

green -> black

116
Q

practical: determine the approximate percentage by volume of oxygen in air using a metal or a non-metal

A
  1. place an excess of wet iron filings into a conical flask (100cm3) full of air
  2. place a bung on the top of the conical flask connected to a syringe

over time the iron will rust (react with oxygen) until all of the oxygen is used up. this will move the syringe reading from 100 to 79ish showing oxygen makes up 21% of the air

117
Q

what is the order of the reactivity series

A

potassium, sodium, lithium, calcium, magnesium, aluminium, carbon, zinc, iron, lead, hydrogen, copper, silver, gold

118
Q

how can metals be arranged into a reactivity series based on their reaction with water

A

whichever reacts most violently with water is the most reactive

for example potassium is the most reactive with water because it is the top of the reactivity series

metal + water ⟶ metal hydroxide + hydrogen

119
Q

how can metals be arranged into a reactivity series based on their reaction with dilute acids

A

whichever reacts most violently with dilute acid is the most reactive

for example potassium is the most reactive with dilute acid because it is the top of the reactivity series

only elements more reactive than copper will react with acid

metal + acid ⟶ salt + hydrogen

120
Q

how can metals be arranged in a reactivity series based on their displacement reactions between metal and metal oxides

A

a more reactive metal will displace a less reactive metal from a compound

to do this in metal oxides you can heat it

for example more reactive zinc will displace less reactive copper
zinc + copper(II) oxide → zinc oxide + copper

we can repeat this reaction with different metals to determine which is more / less reactive
so potassium will displace everything as it as the most reactive metal

121
Q

how does displacement work with the reactivity series

A

a more reactive metal will displace a less reactive metal from a compound

122
Q

how can metals be arranged in a reactivity series based on their displacement reactions between metal and aqueous solutions of metal salts

A

a more reactive metal will displace a less reactive metal from a compound so

This is easily seen as the more reactive metal slowly disappears from the solution, displacing the less reactive metal

For example, magnesium is a reactive metal and can displace copper from copper(II)sulfate solution:
Mg + CuSO4→ MgSO4 + Cu

The blue colour of the CuSO4 solution fades as colourless magnesium sulfate solution is formed

Copper coats the surface of the magnesium and also forms solid metal which falls to the bottom of the beaker

123
Q

what is oxidation

A

Oxidation is any reaction in which a substance gains oxygen

oxidation is loss of electrons

124
Q

what is reduction

A

Reduction is a reaction in which a substance loses oxygen

reduction is gain of electrons

125
Q

what is a redox reaction

A

a reaction where reduction and oxidation happen (loss and gain of electrons)

Oxidation cannot occur without reduction happening simultaneously, hence these are called redox reactions

126
Q

what is a oxidising agent

A

makes oxidation happen

so it is the thing that is being reduced as it will take the electrons from the thing being oxidised allowing oxidation to happen

the thing that is reduced is the oxidising agent

127
Q

what is a reducing agent

A

makes reduction happen

so it is the thing that is being oxidised as it will gain the electrons from the thing being reduced allowing reduction to happen

the thing that is oxidated is the reducing agent

128
Q

practical: investigate reactions between dilute hydrochloric and sulfuric acids and metals

A
  1. place 3 test tubes in a rack
  2. Using a small measuring cylinder, add 5 cm3 of dilute hydrochloric acid to each of three test tubes
  3. Add about 1 cm length of magnesium ribbon to the first tube, observe and note down what you see
  4. Use a lit splint to test for any gases given off
  5. To the second test tube add a few pieces of iron filings and to the third some zinc turnings
  6. Observe what happens, test for any gases and note down your observations
    7.Repeat the experiment with dilute sulfuric acid

The metals can be ranked in reactivity order Mg > Zn > Fe

129
Q

what are the uses of aluminium

A

aeroplane bodies - very strong and low density (so it doesn’t weigh a lot)

power cables - good electrical conductor

food cans - non-toxic, resistant to corrosion and acidic food

window frames - resistant to corrosion

130
Q

what are the uses of copper

A

wires - great electrical conductor

cooking pans - great heat conductor

water pipes - unreactive and malleable

131
Q

what are the uses of iron

A

making steel - steel is more useful

132
Q

what are the uses of low-carbon steel

A

ships, cars, bridges - strong and malleable

133
Q

what are the uses of high-carbon steel

A

tools (knives + screwdrivers) - strong and less malleable so won’t change shape

134
Q

what are the uses of stainless steel

A

cutlery, kitchen sinks - Cr forms oxide layers that is resistant to corrosion and stays shiny and clean

135
Q

what is an alloy

A

a mixture of a metal and one or more elements, usually other metals or carbon

136
Q

why are alloys harder than pure metals

A

the different-sized atoms/ions which distorts the regular arrangement therefore preventing the layers of metal ions from sliding over each other making it harder

137
Q

universal indicator colours

A

0-3 strongly acidic - dark red/red/orange
4-6 weakly acidic - orange/yellow/light green
7 neutral - green
8-10 weakly alkaline - dark green/light blue
11-14 strongly alkaline - dark blue/purple

add a few drops to a solution and match the colour it turns to the universal indicator chart to find out the pH

138
Q

whats an alkali

A

a OH- donor
bases that are soluble in water

139
Q

what is a acid

A

a proton (H+ ion) donor

140
Q

what is a base

A

a proton (H+ ion) acceptor

141
Q

what happens if an acid is added to a alkali

A

a neutralisation reaction

the H+ in the acid reacts with the OH- in a base forming water

142
Q

how to carry out an acid-alkali titration

A

Titrations are a method of analysing the concentration of solutions
They can determine exactly how much alkali is needed to neutralise a quantity of acid – and vice versa

Use the pipette and pipette filler and place exactly 25 cm3 known acid solution into the conical flask
fill the burette with an alkali
Place the conical flask on a white tile so the tip of the burette is inside the flask
Add a few drops of phenolphthalein indicator to the solution in the conical flask
Perform a rough titration by taking the burette reading and running in the solution in 1 – 3 cm3 portions, while swirling the flask vigorously
Quickly close the tap when the end-point is reached (sharp colour change) and record the volume, placing your eye level with the meniscus
Now repeat the titration with a fresh batch of acid
As the rough end-point volume is approached, add the solution from the burette one drop at a time until the indicator just changes colour
Record the volume to the nearest 0.05 cm3
Repeat until you achieve two concordant results (two results that are within 0.1 cm3 of each other) to increase accuracy

143
Q

solubility of common sodium, potassium, ammonium compounds

A

they are soluble with everything

144
Q

solubility of common chlorides

A

soluble with everything except silver and lead(II)

145
Q

solubility of nitrates

A

they are soluble with everything

146
Q

solubility of common sulfates

A

soluble with everything except barium, calcium, lead(II)

147
Q

solubility of common carbonates

A

insoluble except with sodium, potassium, and ammonium

148
Q

solubility of common hydroxides

A

insoluble except with sodium, potassium, ammonium and calcium (calcium hydroxide is slightly soluble)

149
Q

reaction between metals and acids

A

Acid + Metal → Salt + Hydrogen

a salt is formed because the H is displaced from the acid

observations:
the metal will disappear as it is used up in the reaction
there is fizzing because a gas (H) is produced

150
Q

reaction between bases and acids

A

Acid + Base → Salt + Water

observations:
the base disappears as it is used up in the reaction

151
Q

reaction between metal carbonates and acids

A

Acid + Carbonate → Salt + Water + Carbon dioxide

observations:
the carbonate disappears because it is used up in the reaction
there is fizzing because a gas (CO2) is being produced

152
Q

examples of bases

A

a metal oxide, metal hydroxide and ammonia

153
Q

describe an experiment to prepare a pure, dry sample of a soluble salt, starting from an insoluble reactant

A

solid metal, insoluble base or insoluble carbonate

metal/metal oxide/metal carbonate + acid –> salt + water

  1. gently heat acid in a beaker until warm
  2. add excess insoluble metal/oxide/carbonate
  3. to ensure all acid is reacted stir
  4. stop heating and filter out any excess undissolved metal
  5. filter it into an evaporating basin
  6. gently heat again until crystals first appear
  7. as soon as they appear stop heating
  8. leave evaporating basin in a warm place in the sun for a few days to allow water to evaporate
154
Q

describe an experiment to prepare a pure, dry sample of a soluble salt, starting from an acid and alkali

A

acid + alkali (soluble base) –> salt + water

titration
1. the alkali and a few drops of indicator to a conical flask
2. add acid to a burette and note the starting volume
3. slowly add the acid to the alkali until the indicator changes colour
4. record the final volume of acid and calculate the total volume of acid added
5. get the same amount of alkali without the indicator and add the total volume of acid added to the other one (this creates pure crystals without the indicator in them)
5. heat the salt solution until crystals just start to appear
6. as soon as they appear, remove from heat
7. leave evaporating basin in a warm place in the sun for a few days to allow water to evaporate

155
Q

describe an experiment to prepare a pure, dry sample of an insoluble salt, starting from two soluble reactants

A

soluble salt 1(aq) + soluble salt 2(aq) –> insoluble salt(s) + soluble salt 3(aq)

  1. mix the 2 soluble salts together in water using a stirring rod
  2. filter the solution using a funnel and filter paper
  3. pour distilled water through the filter paper to wash the insoluble salt
  4. place the insoluble salt in a warm dry sunny place to dry the salt
156
Q

practical: prepare a sample of pure, dry hydrated copper(II) sulfate crystals starting from copper(II) oxide

A

soluble salt from insoluble reactant method

  1. Add 50 cm3 dilute acid into a beaker and warm gently using a Bunsen burner
  2. add the copper(II) oxide slowly to the hot dilute acid and stir until the base is in excess (i.e. until the base stops
    dissolving and a suspension of the base forms in the acid)
  3. Filter the mixture into an evaporating basin to remove the excess base
  4. Gently heat the solution in a water bath or with an electric heater to evaporate the water and to make the
    solution saturated
  5. Check the solution is saturated by dipping a cold glass rod into the solution and seeing if crystals form on the
    end
  6. Leave the filtrate in a warm sunny place to dry and crystallise
157
Q

practical: prepare a sample of pure, dry lead(II) sulfate

A

insoluble salt from 2 soluble reactants

  1. Measure out 25 cm3 of 0.5 mol dm3 lead(II)nitrate solution and add it to a small beaker
  2. Measure out 25 cm3 of 0.5 mol dm3 of potassium sulfate add it to the beaker and mix together using a stirring
    rod
  3. Filter to remove precipitate from mixture
  4. Wash filtrate with distilled water to remove traces of other solutions
  5. leave in sunny warm place to dry
158
Q

test for hydrogen

A

place a lit splint into hydrogen and you will hear a squeaky pop

159
Q

test for carbon dioxide

A

bubble through limewater and it will turn cloudy

160
Q

test for oxygen

A

place a glowing splint into oxygen and it will relight

161
Q

test for chlorine

A

place a damp blue litmus paper into chlorine and it will bleach (turn white)

162
Q

test for ammonia

A

hold damp red litmus paper in the NH3
will turn blue

163
Q

method to carry out the flame test

A
  1. turn the bunsen burner onto non-luminous
  2. get a nichrome wire loop and dip it into HCl
  3. Dip the nichrome wire into the substance
  4. hold the wire into the blue cone in the bunsen
  5. observe the colour change and repeat
164
Q

colour of Li+ flame

A

(bright) red

165
Q

colour of Na+ flame

A

yellow

166
Q

colour of K+ flame

A

lilac

167
Q

colour of Ca2+ flame

A

orange/red (brick red)

168
Q

colour of Cu2+ flame

A

blue/green

169
Q

test for ammonium cation

A
  1. place 1cm of ammonia solution into a test tube
  2. add 2cm of sodium hydroxide solution into the test tube
  3. place the test tube into a beaker containing warm water
  4. hold damp red litmus over the test tube

turns from red to blue

170
Q

test for Cu2+ cation

A
  1. place 1cm of the salt (Cu) solution in a test tube
  2. slowly add drops of 0.5mol sodium hydroxide into test tube until precipitate is visible

light blue gelantinous precipitate is formed
floats on surface

171
Q

test for Fe2+ cation

A
  1. place 1cm of the salt (Fe) solution in a test tube
  2. slowly add drops of 0.5mol sodium hydroxide into test tube until precipitate is visible

green gelantinous precipitate is formed
floats on the surface

172
Q

test for Cl-, Br-, I-

A

silver nitrate solution

dissolve the sample (which you are testing for the anion) in water
add nitric acid (prevents false positives)
add silver nitrate solution

if Cl, Br, I are present then a precipitate will form

Cl - white precipitate
Br - cream precipitate
I - yellow precipitate

173
Q

test for carbonates

A

hydrochloric acid

add the solution into a test tube with a delivery tube connected to another test tube
add hydrochloric acid to the solution
if the solution effervesces then the gas produced will be captured
test the captured gas for CO2 in limewater
if carbonate was present then the limewater will turn cloudy

174
Q

test for sulphates

A

barium chloride

add hydrochloric acid to the solution and look for effervescence
if there’s not effervescence then add barium chloride (stops false positives)
if theres sulphates present then a white precipitate will form

175
Q

test for water

A

anhydrous copper(II) sulfate

add the solution to anhydrous copper(II) sulfate
if water is present then the solution will turn from white to blue

176
Q

physical test to show water is pure

A

check boiling point

gently heat the solution to 100’c
if the water is pure then it will all boil at 100’c

177
Q

method to work out the effect in change of surface area of a solid on the rate of reaction

A
  1. Add dilute hydrochloric acid to a conical flask
  2. Use a delivery tube to connect this flask to an inverted measuring cylinder upside down in a water trough
  3. Add calcium carbonate chips into the conical flask and close the bung
  4. Measure the volume of gas produced in a fixed time using the measuring cylinder
  5. Repeat with different sizes of calcium carbonate chips (use 3g of small, medium, large chips)

results:
more gas will be produced from the smaller chips as they have a larger surface area

178
Q

method to work out the effect in change of concentration of a solution on the rate of reaction

A
  1. Measure 50 cm3 of sodium thiosulfate solution into a conical flask
  2. Measure 5 cm3 of dilute hydrochloric acid into a measuring cylinder
  3. Draw a cross on a piece of paper and put it underneath the flask
  4. Add the acid into the flask and immediately start the stopwatch
  5. Look down at the cross from above and stop the stopwatch when the cross can no longer be seen
  6. Repeat using different concentrations of sodium thiosulfate solution (mix different volumes of sodium thiosulfate
    solution with water to dilute it)

results:
with an increased of the concentration the rate of reaction will increase meaning the cross will disappear quicker

179
Q

method to work out the effect in change of temperature of a solution on the rate of reaction

A
  1. Dilute hydrochloric acid in a conical flask is heated to a set temperature using a water bath
    2 .Add a strip of magnesium to the HCl and start the stopwatch
  2. Stop the time when the magnesium fully dissolves
  3. Repeat at different temperatures and compare results

results:
With an increase in the temperature the magnesium will dissolve quicker

180
Q

method to work out the effect of using a catalyst has on the rate of reaction

A
  1. Add hydrogen peroxide into a conical flask
  2. Use a delivery tube to connect this flask to a measuring cylinder upside down in water trough
  3. Add the catalyst (manganese(IV) oxide) into the conical flask and close the bung
  4. Measure the volume of gas produced in a fixed time using the measuring cylinder
  5. Repeat experiment without the catalyst of manganese(IV) oxide and compare results

results:
the catalyst will speed up the rate of reaction therefore the volume of gas produced will be greater when the catalyst is used

181
Q

how does a catalyst affect the rate of reaction

A

the catalyst lowers the activation energy providing an alternate pathway for the reaction to take place
meaning the percentage of collisions with enough energy is higher
so there are more successful collisions per unit of time
increasing the rate of reaction

182
Q

how does surface area affect the rate of reaction

A

with a larger SA rate of reaction increases because more surface area of the particles will be exposed to the other reactant
therefore more particles collide more frequently per unit of time
meaning there are more successful collisions per unit of time
and an increased rate of reaction

183
Q

how does the temperature affect the rate of reaction

A

with higher temperature rate of reaction increases because the particles will have more kinetic energy
So the percentage of collisions with enough energy is higher
meaning more successful collisions per unit of time
increasing the rate of reaction

184
Q

how does the concentration/pressure affect the rate of reaction

A

with a higher concentration/pressure the rate of reaction increases because there are more particles per given volume of mass
there for particles collide more frequently per unit of time
meaning there are more successful collisions per unit of time
increasing the rate of reaction

185
Q

does the catalyst change the reaction

A

no the catalyst does not
the reaction is chemically unchanged at the end of the reaction

186
Q

practical: investigate the effect of changing the surface area of marble chips and of changing the concentration of hydrochloric acid on the rate of reaction between marble chips and dilute hydrochloric acid

A
  1. Add dilute hydrochloric acid to a conical flask
  2. Use a delivery tube to connect this flask to an inverted measuring cylinder upside down in a water trough
  3. Add calcium carbonate chips into the conical flask and close the bung
  4. Measure the volume of gas produced in a fixed time using the measuring cylinder
  5. Repeat with different sizes of calcium carbonate chips (use 3g of small, medium, large chips)

results:
more gas will be produced from the smaller chips as they have a larger surface area

187
Q

practical: investigate the effect of different solids on the catalytic decomposition of hydrogen peroxide solution

A
  1. Add hydrogen peroxide into a conical flask
  2. Use a delivery tube to connect this flask to a measuring cylinder upside down in water trough
  3. Add the catalyst (manganese(IV) oxide) into the conical flask and close the bung
  4. Measure the volume of gas produced in a fixed time using the measuring cylinder
  5. Repeat experiment with different catalysts and compare results
188
Q

what is a hydrocarbon

A

a compound made up of only hydrogens and carbons

189
Q

what is the empirical formula

A

the simplest possible ratio of the atoms in a molecule
For example: Hydrogen peroxide is H2O2 but the empirical formula is HO

190
Q

what is the molecular formula

A

the actual number of atoms in a molecule

191
Q

what is the general formula

A

A ratio of atoms in a family of compounds in terms of ‘n’ where n is a varying whole number
For example, the general formula of a molecule that belong to the alkane family is CnH2n+2

192
Q

what is the structural formula

A

In a structural formulae enough information is shown to make the structure clear, but most of the actual covalent bonds are omitted
Only important bonds are always shown, such as double and triple bonds
Identical groups can be bracketed together
Side groups are also shown using brackets
Straight chain alkanes are shown as follows:

CH3CH2CH2CH2CH3 or CH3(CH2)3CH3

Displayed formula:
H H H H
I I I I
H-C-C-C-C-H
I I I I
H H I H
I
H-C-H
I
H

193
Q

what is the displayed formula

A

how the molecule is drawn

H H H H
 I   I  I   I H-C-C-C-C-H
 I   I   I  I
H  H I  H
         I
    H-C-H
         I
        H
194
Q

what is a homologous series

A

a series or family of organic compounds that have similar features and chemical properties due to them having the same functional group

All members of a homologous series have:
The same general formula
Same functional group
Similar chemical properties
Gradation in their physical properties
The difference in the molecular formula between one member and the next is CH2

195
Q

how to name compounds

A

first part of the name second part of the name
NAME Number of C atoms name functional group Family

meth.. 1 ..ane none alkane
eth.. 2 ..ene C = C double bond alkene
prop.. 3 ..anol R-OH alcohol
but.. 4 ..anoic acid R-C=O-OH carboxylic acid
pent.. 5 ..amine R-NH2 amine
hex.. 6 ..yl ..anoate R-C=O-O-R ester

When there is more than one carbon atom where a functional group can be located it is important to distinguish exactly which carbon the functional group is on
Each carbon is numbered and these numbers are used to describe where the functional group is
When 2 functional groups are present di- is used as a prefix to the second part of the name

196
Q

what is a functional group

A

A group of atoms bonded in a specific arrangement that influences the properties of the homologous series (what is attached to the carbon)

197
Q

what is isomerism

A

Isomers are compounds that have the same molecular formula but different displayed formulae

198
Q

how to classify reactions of organic compounds as substitution, addition and combustion

A

A substitution reaction takes place when one functional group is replaced by another
CH4 + Br2 → CH3Br + HBr

An addition reaction takes place when two or more molecules combine to form a larger molecule with no other products
C2H4 + Br2 → C2H4Br2

This is the scientific term for burning. In a combustion reaction, an organic substance reacts with oxygen to form carbon dioxide (or carbon monoxide if incomplete combustion) and water.

199
Q

what is crude oil

A

a mixture of different hydrocarbons

200
Q

what is fuel

A

a substance that when it is burned it releases heat energy

201
Q

how does fractional distillation work

A

Fractional distillation is carried out in a fractionating column which is very hot at the bottom and cool at the top
Crude oil enters the fractionating column and is heated so vapours rise
Vapours of hydrocarbons with very high boiling points will immediately condense into liquid at the higher temperatures lower down and are tapped off at the bottom of the column
Vapours of hydrocarbons with low boiling points will rise up the column and condense at the top to be tapped off
The different fractions condense at different heights according to their boiling points and are tapped off as liquids
The fractions containing smaller hydrocarbons are collected at the top of the fractionating column as gases
The fractions containing bigger hydrocarbons are collected at the lower sections of the fractionating column

202
Q

trend in colour, boiling point and viscosity of the fractions

A

boiling point and viscosity increase as the boiling point increases
colour gets darker as the boiling points increase

203
Q

refinery gases

A

1-4 carbons
boiling point less then 25’c
fuel for home cooking

204
Q

gasoline

A

4-12 carbons
boiling point between 40-100’c
fuel for cars

205
Q

kerosene

A

12-16 carbons
boiling point between 150-240’c
fuel for aircrafts

206
Q

diesel

A

14-18 carbons
boiling point between 220-300’c
fuel for trains

207
Q

fuel oil

A

19-25 carbons
boiling point between 250-320’c
fuel for ships

208
Q

bitumen

A

more than 70 carbons
boiling point more than 350’c
making roads

209
Q

possible products of complete combustion of hydrocarbons with oxygen in air

A

Complete Combustion happens when there is enough oxygen available, producing carbon dioxide (CO2) and water (H2O)

eg CH4 + 2O2 -> CO2 + 2H2O

210
Q

possible products of incomplete combustion of hydrocarbons with oxygen in air

A

Incomplete Combustion happens when there is not enough oxygen available, with possible products being carbon monoxide (CO), carbon (C, soot), carbon dioxide (CO2) and water (H2O)

eg ethane + oxygen -> carbon monoxide + water

211
Q

effects of carbon monoxide on the capacity of blood to carry oxygen

A

Carbon monoxide is a toxic and odourless gas which can cause dizziness, loss of consciousness and eventually death
The CO binds well to haemoglobin which therefore cannot bind oxygen meaning less oxygen and CO2 can be transported to and from organs and working muscles

212
Q

what do the high temperatures in car engines allow

A

the temperature goes high enough to allow and oxygen from the air and nitrogen to react, forming oxides of nitrogen (NO and NO2)

213
Q

how does the combustion of some impurities in hydrocarbon fuels results in the formation of sulfur dioxide

A

Fossil fuels are often contaminated with small amounts of sulfur impurities
When these contaminated fossil fuels are combusted, the sulfur in the fuels get oxidised to sulfur dioxide

214
Q

how does sulfer dioxide and and oxides of nitrogen contribute to acid rain

A

The sulfur dioxide produced from the combustion of fossil fuels dissolves in rainwater droplets to form sulfuric acid
2SO2 (g) + O2 (g) + 2H2O (l) → 2H2SO4 (aq)

Sulfuric acid is one of the components of acid rain which has several damaging impacts on the environment
Nitrogen dioxide produced from car engines reacts with rain water to form a mixture of nitrous and nitric acids, which contribute to acid rain:
2NO2 (g) + H2O (l) → HNO2 (aq) + HNO3 (aq)

215
Q

what is the general formula for alkanes

A

CnH2n+2

216
Q

why are alkanes classified as saturated hydrocarbons

A

they only have single carbon-carbon bonds, there are no double bonds

217
Q

reactions of alkanes with halogens in the presence of ultraviolet radiation

A

Alkanes undergo a substitution reaction with halogens in the presence of ultraviolet radiation

eg CH4 + Br2 → CH3Br + HBr

methane + bromine → bromomethane + hydrogen bromide

218
Q

what is the functional of alkenes

A

> C=C <

219
Q

what is the general formula for alkenes

A

CnH2n

220
Q

why are alkenes classified as unsaturated hydrocarbons

A

Compounds that have a C=C double bond are also called unsaturated compounds
That means they can make more bonds with other atoms by opening up the C=C bond and allowing incoming atoms to form another single bond with each carbon atom of the functional group
Each of these carbon atoms now forms 4 single bonds instead of 1 double and 2 single bonds
This makes them much more reactive than alkanes

221
Q

reactions of alkenes with bromine

A

Alkenes undergo addition reactions in which atoms of a simple molecule add across the C=C double bond
The reaction between bromine and ethene is an example of an addition reaction
The same process works for any halogen and any alkene in which the halogen atoms always add to the carbon atoms across the C=C double bond

eg

H H H H
I I I I
C = C + Br2 —–> H - C - C - H
I I I I
H H Br Br

222
Q

how can bromine water be used to distinguish between an alkane and an alkene

A

Alkanes and alkenes have different molecular structures
All alkanes are saturated and alkenes are unsaturated
The presence of the C=C double bond allows alkenes to react in ways that alkanes cannot
Bromine water is an orange coloured solution
When bromine water is added to an alkane, it will remain as an orange solution as alkanes do not have double carbon bonds (C=C) so the bromine remains in solution
But when bromine water is added to an alkene, the bromine atoms add across the C=C bond, hence the solution no longer contains free bromine so it loses its colour

Chemistry (5 decks)

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