TOPIC 1: Key Concepts in
Chemistry
FORMULAE, EQUATIONS &
HAZARDS
During a chemical reaction, bonds between atoms break as they rearrange
themselves to form different chemicals. The starting products are known as
reactants & they become products.
Word equations are used to show what happens in a chemical reaction
using the full names (e.g. magnesium + oxygen –> magnesium oxide).
Otherwise, symbol equations are used to show reactions using chemical
symbols & formulae (e.g. 2Mg + O2 –> 2MgO).
Equations must always be balanced so the equation shows the same
number of atoms on each side. This creates a ratio.
State symbols can be added to tell you the physical state of the reactants &
products, including solid (s), liquid (l), gas (g) & aqueous (aq – dissolved in
water). If the colour changes, a precipitate is being formed – this is solid. If
there is fizzing, a gas is being formed.
Common ions (charged groups of atoms): Ammonium (NH4), Hydroxide
(OH-), Nitrate (NO3-), Carbonate (CO32-) & Sulfate (SO42-).
HOW TO WRITE IONIC EQUATIONS:
1. Write full balanced symbol equation for reaction.
2. Split aqueous compounds into ions (HNO3(aq) –> H+(aq) + NO3-(aq)).
3. Remove anything that appears on both sides (known as spectator ions).
4. Rewrite equation with ionic charges & balancing numbers.
Hazards are anything that can cause harm – whereas risks are the
probability associated of someone being harmed after exposure.
When you plan an experiment, you must do a risk assessment, including
identifying all hazards & risks, likelihood of something going wrong,
seriousness of result & how to reduce risks.
Hazards symbols are found on chemical containers to highlight dangers &
allow you to work safely.
HOW HAZARD SYMBOLS WORK:
1. Oxidising: provides oxygen allowing other material to catch fire burn
more fiercely SO they must be kept away from flammable substances,
skin & clothes.
2. Highly flammable: catches fire easily SO they must be kept away from
open fires, skin & clothes.
3. Environmental hazard: harmful to organisms & environment SO it must
be disposed of properly & not into environment.
, 4. Corrosive: destroying living material, including living tissue, SO never let
it meet skin, eyes or clothes.
5. Toxic: could cause death by absorption, swallowing or inhaling SO never
let it meet skin or inhale it.
6. Harmful: can cause irritation, blistering or reddening of skin SO keep
away from skin, eyes & clothes.
ATOMIC STRUCTURE & PERIODIC
TABLE
The structure of an atom has a developed by theories over many years,
including:
1. JOHN DALTON: described atoms as solid spheres, where different
spheres are made up of different elements.
2. JJ THOMSON: (concluded atoms were not solid spheres) his
measurements of charge & mass lead to the discovery of the smaller,
negatively charged particles – electrons. New theory was known as ‘plum
pudding model’ with atom as ball of positive charge with electrons stuck
on.
3. ERNEST RUTHERFORD: (disproved ‘plum pudding model’ from gold foil
experiment – where they fired positively charged alpha particles at an
extremely thin sheet of gold. They expected to particles to pass or be
slightly deflected, since the positive charge was spread out. However,
while most passes through, more than expected by deflected & a small
number were even deflected backwards) he developed the nuclear model
of the atom with a tiny positively charged nucleus at the centre, where
most of the mass is concentrated. He also believed there was a ‘cloud’ of
electrons surrounding it, so most of the atom was empty space.
4. NIELS BOHR: (realised a ‘cloud’ of electrons would be attracted to the
nucleus & cause the atom to collapse) he proposed a new nuclear model
with electrons orbiting the nucleus with fixed shells, where each shell
has a fixed energy.
5. Further experiments by Rutherford & others showed that the nucleus
could be divided into smaller particles – they were referred to as protons.
6. JAMES CHADWICK: carried out experiment to provide evidence for
neutral particles in the nucleus – they became known as neutrons.
Atoms have three subatomic particles: protons (heavy & positive charged
(+1)), neutrons (heavy & neutral (0)) & electrons (extremely light &
negatively charged (-1)) – these are their relative charges.
Subatomic particles have a relative mass measured in mass number units
(amu) where protons & neutrons have a mass of 1 & an electron have a
mass of 1/1840.
Atoms have a radius of around 1 x 10-10 m. The nucleus (made of neutrons &
protons) is in the centre & it is orbited by electrons in shells.
, Nuclei have a positive charge overall & it is where most of the atoms mass
is concentrated; however, it is tiny comparatively at only 1 x 10-14 m.
Electrons are tiny & negatively charged but their shells cover a lot of space
– their size is determined by the size of the atom.
In an atom, the number of protons & electrons is always equal; this means
they have no charge/ are neutral because the charges are equal & opposite
(cancelled out). However, if electrons are added or removed, they become
ions because they become charged.
Nuclear symbol tells you can atom’s atomic number (number of protons – &
thus, electrons) & mass number (number of protons + number of neutrons);
therefore, to find the number of neutrons = mass number – atomic number.
Isotopes are atoms of the same element with the same number of protons
but different number of neutrons (e.g. carbon-12 where the mass number is
12). This means they have the same atomic number but different mass
numbers. As neutrons as neutrally charged, the atoms remain neutral.
The periodic table the top number represents the relative atomic mass (Ar).
This is the average mass of one atom of the element compared to 1/12 the
mass of one atom of carbon-12. If an element only has one isotope, its
relative atomic mass will be the same as the mass number.
When an element has different isotopes, they occur in different abundances
or percentages so you must use an equation to calculate the A r.
CALCULATING Ar:
relative atomic mass = sum of (isotope abundance x isotope mass number)
/ sum of
abundances (i.e. 100)
In 1869, Dmitri Mendeleev arranged 50 known elements into the Table of
Elements by sorting them into groups, based on their properties. He
realised, by putting them in order of atomic mass, a pattern appeared, &
element of similar chemical properties could be put in columns.
A few elements ended up in the wrong columns since the wrong atomic
mass was calculated from the presence of isotopes. Wherever this
happened, he switched the order of elements.
To keep element of the similar properties together, he left gaps by using the
properties of other elements in the column. When they were found & fitted,
it confirmed his ideas.
Once protons & electrons were discovered, the atomic number could be
found. Modern periodic tables show elements in order of ascending atomic
number (same pattern as Mendeleev).
The periodic table gives information about the elements, including: group
number (tells you the number of electrons in the outer shells & elements
with similar chemical properties), period number (tells you how many
electron shells it has) & metals vs non-metals (separated by zigzag)
Group 0 is an exception as all elements have a full outer shell of 8 (but
Helium has 2).
Electrons occupy shells/ energy levels with the lowest fill first (closet to
nucleus). Each shell has a maximum number of electrons (1st: 2, 2nd: 8 &
3rd: 8). When the shells are filled, they are stable.
, Each electron shell can hold up to 2n2 electrons – where 2 is the number of
the shell.
Electronic configurations show the arrangement of electrons in an atom. It
can be represented in a 2D diagram or using number (e.g. 2.8.8). Using the
period & group number from the periodic table can ensure accuracy from
the number of shells to the number of electrons on the last shell.
The first four electron sub-shells (orbitals) are s, p, d & f (in this order);
they hold up to 2 electrons.
BONDING & TYPES OF
SUBSTANCE
Ions are charged particles formed when atoms lose or gain electrons. They
are either anions (negative – more electrons than protons) or cations
(positive – more protons than electrons).
The number of electrons lost or gained is the same as the charge on the ion.
When they lose or gain electrons, they end up with a full outer shell &
stability.
The number of electrons they transfer is dependent on the number of the
electrons on the outer shell – which is also their group number. Metals
(group 1 & 2) lose electrons & become cations; non-metals (group 6 & 7)
gain electrons & become anions.
Metals have less electrons on their outer shell, so it is easier for them to
lose than gain electrons.
Ionic compounds are made of a negative & positive part, but the overall
charge is 0 (so the negative & positive charges should balance each other
(e.g. Ca2+ + NO3- -> Ca(NO3)2).
After ionic bonding, some chemicals change their name, including:
polyatomic ions (Hydroxide (OH-), Nitrate (NO3-), Carbonate (CO32-),
Sulfate (SO42-)), part of group 6 & all of group 7.
When ions end with ‘-ate’, they are negative ions with oxygen & at least one
other element. If they end in ‘-ide’, they are negative ions with one element.
Ionic bonding is when a metal reacts with a non-metal & electrons are
transferred by the metal to the non-metal atoms.
The strong electrostatic attraction between the oppositely charged ions
holds the ions together.
Ionic bonding is represented by a dot & cross diagram of before & after the
transferral of electrons – where on type of atom is represented by dots &
the other by crosses. You show the charges as + or – next to the big square
bracket for every labelled ion involved.
Ionic compounds have a structure called a giant ionic lattice (closely
packed regular arrangement of ions) held together by strong electrostatic
attraction in all directions.
Chemistry
FORMULAE, EQUATIONS &
HAZARDS
During a chemical reaction, bonds between atoms break as they rearrange
themselves to form different chemicals. The starting products are known as
reactants & they become products.
Word equations are used to show what happens in a chemical reaction
using the full names (e.g. magnesium + oxygen –> magnesium oxide).
Otherwise, symbol equations are used to show reactions using chemical
symbols & formulae (e.g. 2Mg + O2 –> 2MgO).
Equations must always be balanced so the equation shows the same
number of atoms on each side. This creates a ratio.
State symbols can be added to tell you the physical state of the reactants &
products, including solid (s), liquid (l), gas (g) & aqueous (aq – dissolved in
water). If the colour changes, a precipitate is being formed – this is solid. If
there is fizzing, a gas is being formed.
Common ions (charged groups of atoms): Ammonium (NH4), Hydroxide
(OH-), Nitrate (NO3-), Carbonate (CO32-) & Sulfate (SO42-).
HOW TO WRITE IONIC EQUATIONS:
1. Write full balanced symbol equation for reaction.
2. Split aqueous compounds into ions (HNO3(aq) –> H+(aq) + NO3-(aq)).
3. Remove anything that appears on both sides (known as spectator ions).
4. Rewrite equation with ionic charges & balancing numbers.
Hazards are anything that can cause harm – whereas risks are the
probability associated of someone being harmed after exposure.
When you plan an experiment, you must do a risk assessment, including
identifying all hazards & risks, likelihood of something going wrong,
seriousness of result & how to reduce risks.
Hazards symbols are found on chemical containers to highlight dangers &
allow you to work safely.
HOW HAZARD SYMBOLS WORK:
1. Oxidising: provides oxygen allowing other material to catch fire burn
more fiercely SO they must be kept away from flammable substances,
skin & clothes.
2. Highly flammable: catches fire easily SO they must be kept away from
open fires, skin & clothes.
3. Environmental hazard: harmful to organisms & environment SO it must
be disposed of properly & not into environment.
, 4. Corrosive: destroying living material, including living tissue, SO never let
it meet skin, eyes or clothes.
5. Toxic: could cause death by absorption, swallowing or inhaling SO never
let it meet skin or inhale it.
6. Harmful: can cause irritation, blistering or reddening of skin SO keep
away from skin, eyes & clothes.
ATOMIC STRUCTURE & PERIODIC
TABLE
The structure of an atom has a developed by theories over many years,
including:
1. JOHN DALTON: described atoms as solid spheres, where different
spheres are made up of different elements.
2. JJ THOMSON: (concluded atoms were not solid spheres) his
measurements of charge & mass lead to the discovery of the smaller,
negatively charged particles – electrons. New theory was known as ‘plum
pudding model’ with atom as ball of positive charge with electrons stuck
on.
3. ERNEST RUTHERFORD: (disproved ‘plum pudding model’ from gold foil
experiment – where they fired positively charged alpha particles at an
extremely thin sheet of gold. They expected to particles to pass or be
slightly deflected, since the positive charge was spread out. However,
while most passes through, more than expected by deflected & a small
number were even deflected backwards) he developed the nuclear model
of the atom with a tiny positively charged nucleus at the centre, where
most of the mass is concentrated. He also believed there was a ‘cloud’ of
electrons surrounding it, so most of the atom was empty space.
4. NIELS BOHR: (realised a ‘cloud’ of electrons would be attracted to the
nucleus & cause the atom to collapse) he proposed a new nuclear model
with electrons orbiting the nucleus with fixed shells, where each shell
has a fixed energy.
5. Further experiments by Rutherford & others showed that the nucleus
could be divided into smaller particles – they were referred to as protons.
6. JAMES CHADWICK: carried out experiment to provide evidence for
neutral particles in the nucleus – they became known as neutrons.
Atoms have three subatomic particles: protons (heavy & positive charged
(+1)), neutrons (heavy & neutral (0)) & electrons (extremely light &
negatively charged (-1)) – these are their relative charges.
Subatomic particles have a relative mass measured in mass number units
(amu) where protons & neutrons have a mass of 1 & an electron have a
mass of 1/1840.
Atoms have a radius of around 1 x 10-10 m. The nucleus (made of neutrons &
protons) is in the centre & it is orbited by electrons in shells.
, Nuclei have a positive charge overall & it is where most of the atoms mass
is concentrated; however, it is tiny comparatively at only 1 x 10-14 m.
Electrons are tiny & negatively charged but their shells cover a lot of space
– their size is determined by the size of the atom.
In an atom, the number of protons & electrons is always equal; this means
they have no charge/ are neutral because the charges are equal & opposite
(cancelled out). However, if electrons are added or removed, they become
ions because they become charged.
Nuclear symbol tells you can atom’s atomic number (number of protons – &
thus, electrons) & mass number (number of protons + number of neutrons);
therefore, to find the number of neutrons = mass number – atomic number.
Isotopes are atoms of the same element with the same number of protons
but different number of neutrons (e.g. carbon-12 where the mass number is
12). This means they have the same atomic number but different mass
numbers. As neutrons as neutrally charged, the atoms remain neutral.
The periodic table the top number represents the relative atomic mass (Ar).
This is the average mass of one atom of the element compared to 1/12 the
mass of one atom of carbon-12. If an element only has one isotope, its
relative atomic mass will be the same as the mass number.
When an element has different isotopes, they occur in different abundances
or percentages so you must use an equation to calculate the A r.
CALCULATING Ar:
relative atomic mass = sum of (isotope abundance x isotope mass number)
/ sum of
abundances (i.e. 100)
In 1869, Dmitri Mendeleev arranged 50 known elements into the Table of
Elements by sorting them into groups, based on their properties. He
realised, by putting them in order of atomic mass, a pattern appeared, &
element of similar chemical properties could be put in columns.
A few elements ended up in the wrong columns since the wrong atomic
mass was calculated from the presence of isotopes. Wherever this
happened, he switched the order of elements.
To keep element of the similar properties together, he left gaps by using the
properties of other elements in the column. When they were found & fitted,
it confirmed his ideas.
Once protons & electrons were discovered, the atomic number could be
found. Modern periodic tables show elements in order of ascending atomic
number (same pattern as Mendeleev).
The periodic table gives information about the elements, including: group
number (tells you the number of electrons in the outer shells & elements
with similar chemical properties), period number (tells you how many
electron shells it has) & metals vs non-metals (separated by zigzag)
Group 0 is an exception as all elements have a full outer shell of 8 (but
Helium has 2).
Electrons occupy shells/ energy levels with the lowest fill first (closet to
nucleus). Each shell has a maximum number of electrons (1st: 2, 2nd: 8 &
3rd: 8). When the shells are filled, they are stable.
, Each electron shell can hold up to 2n2 electrons – where 2 is the number of
the shell.
Electronic configurations show the arrangement of electrons in an atom. It
can be represented in a 2D diagram or using number (e.g. 2.8.8). Using the
period & group number from the periodic table can ensure accuracy from
the number of shells to the number of electrons on the last shell.
The first four electron sub-shells (orbitals) are s, p, d & f (in this order);
they hold up to 2 electrons.
BONDING & TYPES OF
SUBSTANCE
Ions are charged particles formed when atoms lose or gain electrons. They
are either anions (negative – more electrons than protons) or cations
(positive – more protons than electrons).
The number of electrons lost or gained is the same as the charge on the ion.
When they lose or gain electrons, they end up with a full outer shell &
stability.
The number of electrons they transfer is dependent on the number of the
electrons on the outer shell – which is also their group number. Metals
(group 1 & 2) lose electrons & become cations; non-metals (group 6 & 7)
gain electrons & become anions.
Metals have less electrons on their outer shell, so it is easier for them to
lose than gain electrons.
Ionic compounds are made of a negative & positive part, but the overall
charge is 0 (so the negative & positive charges should balance each other
(e.g. Ca2+ + NO3- -> Ca(NO3)2).
After ionic bonding, some chemicals change their name, including:
polyatomic ions (Hydroxide (OH-), Nitrate (NO3-), Carbonate (CO32-),
Sulfate (SO42-)), part of group 6 & all of group 7.
When ions end with ‘-ate’, they are negative ions with oxygen & at least one
other element. If they end in ‘-ide’, they are negative ions with one element.
Ionic bonding is when a metal reacts with a non-metal & electrons are
transferred by the metal to the non-metal atoms.
The strong electrostatic attraction between the oppositely charged ions
holds the ions together.
Ionic bonding is represented by a dot & cross diagram of before & after the
transferral of electrons – where on type of atom is represented by dots &
the other by crosses. You show the charges as + or – next to the big square
bracket for every labelled ion involved.
Ionic compounds have a structure called a giant ionic lattice (closely
packed regular arrangement of ions) held together by strong electrostatic
attraction in all directions.
