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Chemistry papers 1 and 2 AQA GCSE Content Summarised
Topics 1-10 Simplified
Providing you with content, so you can do the active recall
Topic 1:
❖ The nucleus of an atom has a radius 10,000 times smaller than an atom and
atoms have a radius of 0.1nm.
❖ On the periodic table, the smaller bottom number is the atomic number which
is the number of protons and this is the same as the number of electrons. The
larger top number instead refers to the mass number and this is the total
number of protons and neutrons in an atom.
❖ Isotopes of an element have the same number of protons but a different
number of neutrons - they have the same atomic number but a different mass
number.
❖ Relative atomic mass = sum of (isotope abundance x isotope mass number) ÷
sum of abundances of all isotopes.
❖ Compounds contain elements in fixed proportions that are held together by
chemical bonds: nonmetals and metals form ionic bonds, metals and metals
from metallic bonds and nonmetals and nonmetals form covalent bonds.
❖ Mixtures are comprised of separate substances not chemically bonded to one
another - these can easily be separated out by physical methods.
❖ Paper chromatography method: can be used to separate out mixtures of
liquids of different different colours and can be used to separate out the
different dyes within an ink: 1. Draw a line at the bottom of the chromatography
paper (the stationary phase) in pencil as it will be insoluble 2. Place spots of inks
on the line leaving space between each one and put the sheet upright in a
beaker of a solvent (such as water), making sure the solvent isn’t touching the
ink initially so that it isn’t simply washed away 3. Place a lid on the top of the
beaker to avoid the evaporation of the solvent. 4. The solvent will begin to seep
up the paper and carry the ink with it. 5. The different dyes in the ink will move
up the paper at different rates causing the dyes to separate out and form
spots in different places (any insoluble inks will remain as they were before on
the baseline) 6. The point at which the solvent has reached as it moves of the
stationary phase of the paper is known as the solvent front and when it has
nearly reached the very top of the paper, the paper should be taken out of the
beaker, the line of the solvent front marked in pencil and it left to dry. This is the
completed chromatogram.
❖ Sometimes, but not always, the number of spots on the chromatogram are
indicative of where the ink has separated out into its dyes and therefore how
many dyes the ink is composed of - however, contrary to belief this is not always
true as two types may travel the same distance up a chromatogram and only
show one spot.
❖ Separating an insoluble solid and a liquid - filtration: (it is also a useful
technique for purification), 1. A piece of filter paper should be folded into a
cone, 2. The corner of filter paper placed point down into a filter funnel sitting
in the neck of a container such as a conical flask, 3. The mixture containing the
insoluble solid should be poured into the filter paper lined funnel being careful
to avoid spilling any of the mixture down the side of the filter paper, 4. The
liquid will pass through the filter paper through the tiny pores and the solid will
not and will instead be left behind in the funnel.
,❖ Separating out a soluble solid and a solution - evaporation or crystallisation;
evaporation: can only be used for soluble salts that do not decompose when
heated and the steps are 1. Pouring the solution into an evaporating dish, 2.
Placing the evaporating dish on top of a tripod and gauze - placing a bunsen
burner underneath, 3. Slowly heating the solution and letting the solvent
evaporate and as the solution becomes more concentrated, a solid will begin to
form 4. The evaporating dish should continue to be heated until all that is left is
dry solid. Crystallisation: it is more time consuming but can produce crystals
that would otherwise decompose if heated… 1. Place an evaporating dish on the
top of a tripod with a gauze and mat and a bunsen burner beneath the tripod
2. Pour the solution into the evaporating dish and gently heat it and some of
the solvent will evaporate whilst the solution becomes more concentrated 3.
Once some of the solvent has evaporated, or when crystals begin to form, the
dish should be removed from the heat and left to cool 4. The salt should begin
to crystallise and it becomes insoluble in the cold and highly concentrated
solution 5. The crystals should be filtered out of the solution left in a warm
place to dry or an oven or desiccator. The slower the rate of cooling, the
crystals the larger the crystals produced.
❖ Simple distillation: separates out a liquid from its mixture and can be used to
separate out substances with boiling points significantly far apart from one
another. 1. The equipment is set up and the mixture heated 2. The component of
the mixture with the lowest boiling point evaporates first 3. As the vapour rises it
passes into the condenser and is cooled, condensed and collected in a
container below the condenser 4. Components of the mixture with higher
boiling points are left behind in the flask.
❖ Fractional distillation: 1. set up the equipment - placing the mixture in the flask
and attaching a fractionating column on top and heating it, 2. The different
liquids will begin to evaporate at different points in time due to their different
boiling points, 3. The liquid with the lowest boiling point will evaporate first and
when the temperature on the thermometer matches the boiling point of this
liquid the vapour has reached the top of the column and passed into the
condenser where it will then cool and condense, running out of the end where
the pure liquid can then be collected, 4. Liquids with higher boiling points may
also begin to evaporate but the column is cooler towards the top and so they
will only reach part of the way up the condenser before they run down to the
flask, 5. Following collection of the first liquid, the temperature should be raised
to the next lowest boiling point of the liquids in the mixture.
❖ History of the atom: initially a man named Democritus believed that all matter
was made up of smaller particles that could not be destructed, were
incompressible and were hard and solid and uniform. In 1803 John Dalton
published his ideas regarding ‘atoms’ and he imagined them as tiny spheres
that could not be divided. Nearly 100 years later, JJ Thompson carried out
several experiments - discovering the electron which meant that atoms must
contain smaller negatively charged particles and he created the ‘plum pudding
model’ - presenting atoms as balls of positive charge embedded with negative
electrons. Next came Ernest Rutherford and Marsden who in 1909 who designed
an alpha particle scattering experiment when testing the plum pudding model -
his findings were that although most positively charged alpha particles passed
straight through when fired at a sheet of thin gold foil, some were deflected
much more than expected; Rutherford created a new idea to cope with the new
evidence and suggested that there was a very small positively charged nucleus
at the centre of atoms and this is where the mass of the atom is concentrated.
He also concluded that a ‘cloud’ of negative electrons surrounded the nucleus
and that most of an atom is comprised of empty space and this was an early
, nuclear model. Neils Bohr then adapted the Rutherford model after new
evidence came to light that the atom would simply ‘collapse in on itself’ if the
nucleus was surrounded by a cloud of negative electrons as Rutherford had
suggested so therefore this could not be true; consequently, Bohr proposed
that electrons were contained within shells or energy levels and that these are
at fixed distances from the nucleus. Further experiments led to the discovery of
the proton - and lastly, in 1942, James Chadwick discovered the neutron with
mass but no charge.
❖ Elements today are arranged in order of increasing atomic number and used to
be arranged by atomic weight.
❖ Mendeleev originally worked with the 57 known elements in 1869. They were
ordered mainly by atomic weight but did change some if the properties
appeared not to fit. Gaps were left to ensure that elements of similar properties
could remain in the same groups and undiscovered elements could be filled in
here - whose properties he had already begun to predict.It was the discovery of
isotopes (same number of protons and different numbers of neutrons) that
conformed that Mendeleev was correct to change up the order of some
elements originally ordered by atomic weight due to their properties as they
have different masses but the same chemical properties and so are in the same
position.
❖ Vertical columns are called groups and rows are called periods. The reason that
some mass numbers are not integers is because of isotopes affecting their Ar
values.
❖ Group number indicates number of electrons in the outer shell.
❖ Metals and non metals = covalent, metals and metal = metallic and non metal
and nonmetals = covalent.
❖ Metals form positive ions when they react.
❖ Group 1: alkali metals: silver solids at room temperature, stored in oil to prevent
reactions with the air due to them being highly reactive and are handled with
forceps and often gloves due to the risk of chemical burns that they pose. They
have a single electron in their outer shells and they are very reactive with similar
properties - the first 3 are less dense than water and all are low density.
Reactivity increases down the group and further down the group the outer shell
electron becomes further and further from the positively charged nucleus
meaning the attraction decreases and it's more easily lost. Melting and boiling
points decrease down the group and relative atomic mass increases down the
group. They react with nonmetals to form ionic compounds and easily form 1+
ions, never share covalently - there's quite frankly no point. Compounds
produced when alkali metals react with nonmetals are most usually always
white solids that’ll dissolve in water forming colourless solutions. Alkali metals
also react with water in order to form a metal hydroxide and hydrogen gas.
Alkali metal + water → metal hydroxide + hydrogen. Group 1s also react with
water vigorously, lithium, sodium and potassium float and move about the
surface and effervescent seriously as hydrogen gas is produced - in some
cases even igniting the hydrogen. Elements below potassium react explosively
with water -i.e very dangerous. Hydroxides formed when alkali metals react with
water will dissolve to give alkaline solutions and hence the name - universal
indicator will show up as purple as a result. Remember squeaky pop tests will
also confirm the presence of hydrogen. Alkali metals will also react vigorously if
heated in chlorine gas to form metal chloride salts. Groups 1s and oxygen
forming different metal oxides.
❖ Group 7: the halogens: all of the halogens exist as diatomic molecules and they
are non metals that have coloured vapours. Fluorine is a poisonous and yellow
gas, chlorine is a poisonous and dense green gas, bromine is a dense and
, poisonous and red - brown volatile liquid or as an orange vapour, iodine is a
poisonous dark grey crystalline solid or a purple vapour. All group 7 elements
react in a similar manner and they all have 7 electrons in their outer shell,
reactivity decreases down the group as the greater distance between the
positive nucleus and the negative electrons results in less attraction and a
lower likelihood of the atom gaining an electron. Melting and boiling points of
the halogens increase down the group. Relative atomic mass also increases
down the halogens. Halogen atoms can share electrons through covalent
bonding with other non - metals and the compounds that form when this
occurs all have simple molecular structures. Halogens are also able to gain
another electron in order to form a halide ion. A more reactive halogen may
also displace a less reactive halogen from an aqueous solution of its salt.
Astatine is the least reactive halogen and therefore can displace no other one.
❖ Group 0 - the noble gases: group 0 elements are practically inert and have a full
outer shell of electrons and hence are not particularly reactive at all. They are
energetically stable and therefore neither gain or lose electrons. They do not
react and nor are they flammable. They exist at room temperature as
monatomic and colourless gases. Boiling points of the noble gases increase
down the group and this goes the same for their relative atomic mass. Down the
group, the number of electrons - and hence the intermolecular forces between
atoms increases - greater energy is required to overcome these and so the
boiling point also increases.
❖ Transition metals: located in the centre of the periodic table they are typical
metals all with similar properties; they’re all good conductors of heat and
electricity and are very dense, strong and shiny. When compared to the
transition metals: the transition metals are much less reactive than the alkali
metals, the transition metals are much denser, harder and stronger when
compared to group 1 metals and have much higher melting point when
compared to group 1 metals. Most transition metals are able to form ions of
different charges and different ions usually form different coloured compounds
(this is identifiable in gemstones for example). Transition metals and their
compounds also all make good catalysts. (for example iron is used in the haber
process).
Topic 2:
❖ Ions are charged particles produced when electrons are lost or gained - this all
happens in the name of trying to obtain a full outer shell and a stable
electronic structure. Elements in the same group all have the same number of
outer shell electrons and therefore will also form the same ions with the same
charges.
❖ Dot and cross diagrams are used to represent ionic bonding.
❖ Ionic compounds have a structure known as a giant ionic lattice and have a
very closely packed and regular arrangement that has very strong electrostatic
forces of attraction between oppositely charged ions and these act in all
directions. Ionic compounds also have high melting and boiling points due to
their strong electrostatic forces of attraction between ions - which it takes large
amounts of energy to overcome. Ionic compounds do not conduct electricity in
their solid form due to ions being held in fixed positions, however, they do
conduct electricity when molten or dissolved as ions become free to move and
carry charge. Most ionic compounds will also dissolve easily in water. Ionic
compounds can of course be represented by dot and cross diagrams, however,
these do not show the structure of the compound nor the relative size or
arrangement of the ions. 3D models on the contrary show the relative sizes of
Chemistry papers 1 and 2 AQA GCSE Content Summarised
Topics 1-10 Simplified
Providing you with content, so you can do the active recall
Topic 1:
❖ The nucleus of an atom has a radius 10,000 times smaller than an atom and
atoms have a radius of 0.1nm.
❖ On the periodic table, the smaller bottom number is the atomic number which
is the number of protons and this is the same as the number of electrons. The
larger top number instead refers to the mass number and this is the total
number of protons and neutrons in an atom.
❖ Isotopes of an element have the same number of protons but a different
number of neutrons - they have the same atomic number but a different mass
number.
❖ Relative atomic mass = sum of (isotope abundance x isotope mass number) ÷
sum of abundances of all isotopes.
❖ Compounds contain elements in fixed proportions that are held together by
chemical bonds: nonmetals and metals form ionic bonds, metals and metals
from metallic bonds and nonmetals and nonmetals form covalent bonds.
❖ Mixtures are comprised of separate substances not chemically bonded to one
another - these can easily be separated out by physical methods.
❖ Paper chromatography method: can be used to separate out mixtures of
liquids of different different colours and can be used to separate out the
different dyes within an ink: 1. Draw a line at the bottom of the chromatography
paper (the stationary phase) in pencil as it will be insoluble 2. Place spots of inks
on the line leaving space between each one and put the sheet upright in a
beaker of a solvent (such as water), making sure the solvent isn’t touching the
ink initially so that it isn’t simply washed away 3. Place a lid on the top of the
beaker to avoid the evaporation of the solvent. 4. The solvent will begin to seep
up the paper and carry the ink with it. 5. The different dyes in the ink will move
up the paper at different rates causing the dyes to separate out and form
spots in different places (any insoluble inks will remain as they were before on
the baseline) 6. The point at which the solvent has reached as it moves of the
stationary phase of the paper is known as the solvent front and when it has
nearly reached the very top of the paper, the paper should be taken out of the
beaker, the line of the solvent front marked in pencil and it left to dry. This is the
completed chromatogram.
❖ Sometimes, but not always, the number of spots on the chromatogram are
indicative of where the ink has separated out into its dyes and therefore how
many dyes the ink is composed of - however, contrary to belief this is not always
true as two types may travel the same distance up a chromatogram and only
show one spot.
❖ Separating an insoluble solid and a liquid - filtration: (it is also a useful
technique for purification), 1. A piece of filter paper should be folded into a
cone, 2. The corner of filter paper placed point down into a filter funnel sitting
in the neck of a container such as a conical flask, 3. The mixture containing the
insoluble solid should be poured into the filter paper lined funnel being careful
to avoid spilling any of the mixture down the side of the filter paper, 4. The
liquid will pass through the filter paper through the tiny pores and the solid will
not and will instead be left behind in the funnel.
,❖ Separating out a soluble solid and a solution - evaporation or crystallisation;
evaporation: can only be used for soluble salts that do not decompose when
heated and the steps are 1. Pouring the solution into an evaporating dish, 2.
Placing the evaporating dish on top of a tripod and gauze - placing a bunsen
burner underneath, 3. Slowly heating the solution and letting the solvent
evaporate and as the solution becomes more concentrated, a solid will begin to
form 4. The evaporating dish should continue to be heated until all that is left is
dry solid. Crystallisation: it is more time consuming but can produce crystals
that would otherwise decompose if heated… 1. Place an evaporating dish on the
top of a tripod with a gauze and mat and a bunsen burner beneath the tripod
2. Pour the solution into the evaporating dish and gently heat it and some of
the solvent will evaporate whilst the solution becomes more concentrated 3.
Once some of the solvent has evaporated, or when crystals begin to form, the
dish should be removed from the heat and left to cool 4. The salt should begin
to crystallise and it becomes insoluble in the cold and highly concentrated
solution 5. The crystals should be filtered out of the solution left in a warm
place to dry or an oven or desiccator. The slower the rate of cooling, the
crystals the larger the crystals produced.
❖ Simple distillation: separates out a liquid from its mixture and can be used to
separate out substances with boiling points significantly far apart from one
another. 1. The equipment is set up and the mixture heated 2. The component of
the mixture with the lowest boiling point evaporates first 3. As the vapour rises it
passes into the condenser and is cooled, condensed and collected in a
container below the condenser 4. Components of the mixture with higher
boiling points are left behind in the flask.
❖ Fractional distillation: 1. set up the equipment - placing the mixture in the flask
and attaching a fractionating column on top and heating it, 2. The different
liquids will begin to evaporate at different points in time due to their different
boiling points, 3. The liquid with the lowest boiling point will evaporate first and
when the temperature on the thermometer matches the boiling point of this
liquid the vapour has reached the top of the column and passed into the
condenser where it will then cool and condense, running out of the end where
the pure liquid can then be collected, 4. Liquids with higher boiling points may
also begin to evaporate but the column is cooler towards the top and so they
will only reach part of the way up the condenser before they run down to the
flask, 5. Following collection of the first liquid, the temperature should be raised
to the next lowest boiling point of the liquids in the mixture.
❖ History of the atom: initially a man named Democritus believed that all matter
was made up of smaller particles that could not be destructed, were
incompressible and were hard and solid and uniform. In 1803 John Dalton
published his ideas regarding ‘atoms’ and he imagined them as tiny spheres
that could not be divided. Nearly 100 years later, JJ Thompson carried out
several experiments - discovering the electron which meant that atoms must
contain smaller negatively charged particles and he created the ‘plum pudding
model’ - presenting atoms as balls of positive charge embedded with negative
electrons. Next came Ernest Rutherford and Marsden who in 1909 who designed
an alpha particle scattering experiment when testing the plum pudding model -
his findings were that although most positively charged alpha particles passed
straight through when fired at a sheet of thin gold foil, some were deflected
much more than expected; Rutherford created a new idea to cope with the new
evidence and suggested that there was a very small positively charged nucleus
at the centre of atoms and this is where the mass of the atom is concentrated.
He also concluded that a ‘cloud’ of negative electrons surrounded the nucleus
and that most of an atom is comprised of empty space and this was an early
, nuclear model. Neils Bohr then adapted the Rutherford model after new
evidence came to light that the atom would simply ‘collapse in on itself’ if the
nucleus was surrounded by a cloud of negative electrons as Rutherford had
suggested so therefore this could not be true; consequently, Bohr proposed
that electrons were contained within shells or energy levels and that these are
at fixed distances from the nucleus. Further experiments led to the discovery of
the proton - and lastly, in 1942, James Chadwick discovered the neutron with
mass but no charge.
❖ Elements today are arranged in order of increasing atomic number and used to
be arranged by atomic weight.
❖ Mendeleev originally worked with the 57 known elements in 1869. They were
ordered mainly by atomic weight but did change some if the properties
appeared not to fit. Gaps were left to ensure that elements of similar properties
could remain in the same groups and undiscovered elements could be filled in
here - whose properties he had already begun to predict.It was the discovery of
isotopes (same number of protons and different numbers of neutrons) that
conformed that Mendeleev was correct to change up the order of some
elements originally ordered by atomic weight due to their properties as they
have different masses but the same chemical properties and so are in the same
position.
❖ Vertical columns are called groups and rows are called periods. The reason that
some mass numbers are not integers is because of isotopes affecting their Ar
values.
❖ Group number indicates number of electrons in the outer shell.
❖ Metals and non metals = covalent, metals and metal = metallic and non metal
and nonmetals = covalent.
❖ Metals form positive ions when they react.
❖ Group 1: alkali metals: silver solids at room temperature, stored in oil to prevent
reactions with the air due to them being highly reactive and are handled with
forceps and often gloves due to the risk of chemical burns that they pose. They
have a single electron in their outer shells and they are very reactive with similar
properties - the first 3 are less dense than water and all are low density.
Reactivity increases down the group and further down the group the outer shell
electron becomes further and further from the positively charged nucleus
meaning the attraction decreases and it's more easily lost. Melting and boiling
points decrease down the group and relative atomic mass increases down the
group. They react with nonmetals to form ionic compounds and easily form 1+
ions, never share covalently - there's quite frankly no point. Compounds
produced when alkali metals react with nonmetals are most usually always
white solids that’ll dissolve in water forming colourless solutions. Alkali metals
also react with water in order to form a metal hydroxide and hydrogen gas.
Alkali metal + water → metal hydroxide + hydrogen. Group 1s also react with
water vigorously, lithium, sodium and potassium float and move about the
surface and effervescent seriously as hydrogen gas is produced - in some
cases even igniting the hydrogen. Elements below potassium react explosively
with water -i.e very dangerous. Hydroxides formed when alkali metals react with
water will dissolve to give alkaline solutions and hence the name - universal
indicator will show up as purple as a result. Remember squeaky pop tests will
also confirm the presence of hydrogen. Alkali metals will also react vigorously if
heated in chlorine gas to form metal chloride salts. Groups 1s and oxygen
forming different metal oxides.
❖ Group 7: the halogens: all of the halogens exist as diatomic molecules and they
are non metals that have coloured vapours. Fluorine is a poisonous and yellow
gas, chlorine is a poisonous and dense green gas, bromine is a dense and
, poisonous and red - brown volatile liquid or as an orange vapour, iodine is a
poisonous dark grey crystalline solid or a purple vapour. All group 7 elements
react in a similar manner and they all have 7 electrons in their outer shell,
reactivity decreases down the group as the greater distance between the
positive nucleus and the negative electrons results in less attraction and a
lower likelihood of the atom gaining an electron. Melting and boiling points of
the halogens increase down the group. Relative atomic mass also increases
down the halogens. Halogen atoms can share electrons through covalent
bonding with other non - metals and the compounds that form when this
occurs all have simple molecular structures. Halogens are also able to gain
another electron in order to form a halide ion. A more reactive halogen may
also displace a less reactive halogen from an aqueous solution of its salt.
Astatine is the least reactive halogen and therefore can displace no other one.
❖ Group 0 - the noble gases: group 0 elements are practically inert and have a full
outer shell of electrons and hence are not particularly reactive at all. They are
energetically stable and therefore neither gain or lose electrons. They do not
react and nor are they flammable. They exist at room temperature as
monatomic and colourless gases. Boiling points of the noble gases increase
down the group and this goes the same for their relative atomic mass. Down the
group, the number of electrons - and hence the intermolecular forces between
atoms increases - greater energy is required to overcome these and so the
boiling point also increases.
❖ Transition metals: located in the centre of the periodic table they are typical
metals all with similar properties; they’re all good conductors of heat and
electricity and are very dense, strong and shiny. When compared to the
transition metals: the transition metals are much less reactive than the alkali
metals, the transition metals are much denser, harder and stronger when
compared to group 1 metals and have much higher melting point when
compared to group 1 metals. Most transition metals are able to form ions of
different charges and different ions usually form different coloured compounds
(this is identifiable in gemstones for example). Transition metals and their
compounds also all make good catalysts. (for example iron is used in the haber
process).
Topic 2:
❖ Ions are charged particles produced when electrons are lost or gained - this all
happens in the name of trying to obtain a full outer shell and a stable
electronic structure. Elements in the same group all have the same number of
outer shell electrons and therefore will also form the same ions with the same
charges.
❖ Dot and cross diagrams are used to represent ionic bonding.
❖ Ionic compounds have a structure known as a giant ionic lattice and have a
very closely packed and regular arrangement that has very strong electrostatic
forces of attraction between oppositely charged ions and these act in all
directions. Ionic compounds also have high melting and boiling points due to
their strong electrostatic forces of attraction between ions - which it takes large
amounts of energy to overcome. Ionic compounds do not conduct electricity in
their solid form due to ions being held in fixed positions, however, they do
conduct electricity when molten or dissolved as ions become free to move and
carry charge. Most ionic compounds will also dissolve easily in water. Ionic
compounds can of course be represented by dot and cross diagrams, however,
these do not show the structure of the compound nor the relative size or
arrangement of the ions. 3D models on the contrary show the relative sizes of
