CHEMISTRY – PAPER 1 1.3 Recall the relative charge and relative mass of:
KEY CONCEPTS - ATOMIC STRUCTURE Particle Relative Mass Relative Charge
1.1 Describe how the Dalton model of an atom has changed over time PROTON 1 +1
because of the discovery of subatomic particles NUETRON 1 0
1) John DALTON described atoms as solid spheres and that different types of ELECTRON Negligible (0.0005) -1
spheres made up different elements. 1.4 Explain why atoms contain equal numbers of protons & electrons
2) In 1897 JJ THOMSON carried out experiments that showed that atoms have - Atoms are neutral – they have no overall charge. This is because the number
smaller negatively charged particles – electrons. The atom was then of protons (positive) is equal to the number of electrons (negative)
described with the plum pudding model. This shows an atom to be a ball of 1.5 Describe the nucleus of an atom as very small compared to the
positive charge, with electrons spread through overall size of the atom
3) In 1909 RUTHERFORD conducted the gold foil experiment. They fired alpha 1.6 Recall that most of the mass of an atom is concentrated in the
particles at an extremely thin sheet of gold. The plum pudding model should nucleus
mean the particles should pass straight through or be deflected slightly. 1.7 Recall the meaning of the term mass number of an atom
- However, while some did go straight through, some deflected more than ATOMIC NUMBER: The number of protons (unique to each element). In a
expected and some were deflected backwards. This meant the plum pudding neutral atom, this is also the number of electrons
model was incorrect, and the positive charge must be in a very concentrated MASS NUMBER: Protons + neutrons (only these have mass)
point (the nucleus)
1.8 Describe atoms of a given element as having the same number of
- He then came up with the nuclear model. This shows positively charged protons in the nucleus and that this number is unique to that element
particles in the nucleus, surrounded by a ‘cloud’ of negative electrons
Every atom of an element has the same number of protons, which is unique
4) However, the ‘cloud’ means nothing stopped the electrons from rushing to the element, changing this will make a new element
towards the nucleus, causing the atom to collapse
1.9 Describe isotopes as different atoms of the same element containing
- The BOHR model shows electrons have fixed orbits (shells). Each shell has a the same number of protons but different numbers of neutrons in their
fixed energy level nuclei
5) New experiments showed that the positive charge in the nucleus is made ISOTOPES: Different forms of the same element, which have the same
up of small discrete particles (PROTONS) as well as neutral particles number of protons but a different number of neutrons
(NEUTRONS)
- Thus, they have the same atomic number, but different mass number
1.2 Describe the structure of an atom as a nucleus containing protons
1.10 Calculate the numbers of protons, neutrons and electrons in atoms
and neutrons, surrounded by electrons in shells
given the atomic number and mass number
NUCLEUS: Middle of atom, containing the protons & neutrons. It has an
1.11 Explain how the existence of isotopes results in relative atomic
overall positive charge & most of the mass of an atom is concentrated here. It
masses of some elements not being whole numbers
is very tiny compared to the atom. The electrons surround nucleus in shells
,- Isotopes have different masses so elements need an ‘average mass’. This is THE PERIODIC TABLE
the RELATIVE ATOMIC MASS (Ar) 1.13 Describe how Mendeleev arranged the elements, known at that time,
- If an element only has 1 isotope than its Ar will be the mass number. But if in a periodic table by using properties of these elements and their
there are multiple isotopes the Ar will be an average, taking into account how compounds
much there is of each. Thus, it might not be a whole number - Mendeleev arranged the ~50 elements known into a periodic table. He did
- Relative atomic mass is the average mass of one element compared to this by sorting elements into groups based on their properties
1/12th the mass of Carbon-12 - Doing so, he realised putting elements into order of atomic mass created a
1.12 Calculate the relative atomic mass of an element from the relative pattern where elements with similar properties were put into columns
masses and abundances of its isotopes - He switched the order of specific elements if they did not fit the properties of
- To calculate the Ar you need to use the formula: their group, so their atomic number would be in reverse order
ABUNDANCE: How common/rare the isotope is 1.14 Describe how Mendeleev used his table to predict the existence and
If abundance is in PERCENTAGE %: Ar= properties of some elements not then discovered
- He left some gaps to keep elements with similar properties together. He
then used the properties of the other elements in the column to predict the
new element. When they were found, they fitted the pattern.
1.15 Explain that Mendeleev thought he had arranged elements in order of
increasing relative atomic mass but this was not always true because of
the relative abundance of isotopes of some pairs of elements in the
Otherwise it is divided by: (SUM of ISOTOPIC ABUNDANCE)
periodic table
- Mendeleev arranged the atoms using the mass number, however the mass
number he had was sometimes wrong due to isotopes
- Some elements didn’t fit the pattern so he reversed their order to keep them
in the same groups
1.16 Explain the meaning of atomic number of an element in terms of
position in the periodic table and number of protons in the nucleus
The modern periodic table is ordered by the ATOMIC NUMBER
, IONIC BONDING
1.17 Describe that in the periodic table 1.21 Explain how ionic bonds are formed by the transfer of electrons
a elements are arranged in order of increasing atomic number, in rows between atoms to produce cations and anions, including the use of dot
called periods and cross diagrams
b elements with similar properties are placed in the same vertical - Ionic Bonding occurs between a METAL & NON-METAL. The metal loses the
columns called groups electron (turning into a CATION) & the non-metal gains it (turning into an
- Rows are called PERIODS. Each new period represents another full shell of ANION). Electrons are TRANSFERED. These oppositely charged ions are
electrons, thus the period corresponds to the number of shells. strongly attracted to each other by electrostatic forces
- Columns are called GROUPS. Elements that have similar properties are - Dot & Cross diagrams are used to show where the electrons originally came
arranged in vertical groups. This corresponds to the number of electrons in from, as one atom uses dots and the other uses crosses
the outer shell 1.22 Recall that an ion is an atom or group of atoms with a positive or
1.18 Identify elements as metals or non-metals according to their negative charge
position in the periodic table, explaining this division in terms of the ION: Charged particles (single atom/group). They are formed when an atom
atomic structures of the elements loses/gains an electron to get a full outer shell (stable electronic structure)
Zig-Zag line separates metals (lose electrons) & non-metals (lose electrons) CATION: positive ion ANION: Negative ion
1.19 Predict the electronic configurations of the first 20 elements in the 1.23 Calculate the numbers of protons, neutrons and electrons in simple
periodic table as diagrams and in the form, for example 2.8.1 ions given the atomic number and mass number
In an atom, electrons are placed in fixed shells according to the rules: 1.24 Explain the formation of ions in ionic compounds from their atoms,
limited to compounds of elements in groups 1, 2, 6 and 7
1st shell: 2 electrons. 2nd shell: 8 electrons. 3rd shell: 8 electrons...
Elements most readily forming ions are in Groups 1,2 6, 7
1.20 Explain how the electronic configuration of an element is related to
its position in the periodic table Group 1 & 2 are metals that lose their electrons to form cations
The electronic configuration of an element shows how many electrons in Group 6 & 7 are non-metals that gain electrons to form anions
each shell. So, 2.8.7 shows 2 electrons in 1st shell, 8 in the 2nd shell, & 7 in the 1.25 Explain the use of the endings –ide and –ate in the names of
3rd shell. This can also be found by its position in the periodic table as shown compounds
in 1.17 The name of ionic compounds end in:
- IDE: When only the metal and non-metal are present
Eg: Potassium Iodide
- ATE: When oxygen is present and at least 1 other element + metal
Eg: Magnesium Sulphate
, 1.26 Deduce the formulae of ionic compounds (including oxides, 1.31 Explain the formation of simple molecular substances, using dot and
hydroxides, halides, nitrates, carbonates and sulfates) given the cross diagrams, including: hydrogen, hydrogen chloride, water, methane,
formulae of the constituent ions oxygen, carbon dioxide
In ionic compound, charges must cancel out Hydrogen - ELEMENT H2
Oxide O2- Nitrate NO3-
Hydroxide OH- Carbonate CO32-
Halide Grp. 7 (F-, Cl-, Sulphate SO42- Hydrogen Chloride - HCl
Br-, i- etc.) COMPOUND
1.27 Explain the structure of an ionic compound as a lattice structure
a consisting of a regular arrangement of ions
b held together by strong electrostatic forces (ionic bonds) between
Water - COMPOUND H2O
oppositely-charged ions
- When an ionic bond is made, an ionic compound is formed. These have
giant ionic lattice structures with millions of ions closely packed together,
held by strong electrostatic forces due to the oppositely charged ions.
Methane - CH4
- The ionic compound consists of ions in a regular arrangement
COMPOUND
COVALENT BONDS
1.28 Explain how a covalent bond is formed when a pair of electrons is
shared between two atoms
Covalent bonds are formed when a PAIR of electrons is SHARED between 2 Oxygen - ELEMENT O2
NON-METAL atoms
- Each atom needs to GAIN electrons to get a full outer shell. Thus, they share
electrons so both gain an electron
1.29 Recall that covalent bonding results in the formation of molecules Carbon Dioxide - CO2
COMPOUND
- Molecules are 2 or more atoms chemically bonded together. They can be of
the same element (O2) or different elements.
- There are 2 types of molecules: Simple molecular & Giant covalent. (These
are described later on)
1.30 Recall the typical size (order of magnitude) of atoms and small
molecules
Average size of simple molecular substances: 10-10m (0.1nm)
KEY CONCEPTS - ATOMIC STRUCTURE Particle Relative Mass Relative Charge
1.1 Describe how the Dalton model of an atom has changed over time PROTON 1 +1
because of the discovery of subatomic particles NUETRON 1 0
1) John DALTON described atoms as solid spheres and that different types of ELECTRON Negligible (0.0005) -1
spheres made up different elements. 1.4 Explain why atoms contain equal numbers of protons & electrons
2) In 1897 JJ THOMSON carried out experiments that showed that atoms have - Atoms are neutral – they have no overall charge. This is because the number
smaller negatively charged particles – electrons. The atom was then of protons (positive) is equal to the number of electrons (negative)
described with the plum pudding model. This shows an atom to be a ball of 1.5 Describe the nucleus of an atom as very small compared to the
positive charge, with electrons spread through overall size of the atom
3) In 1909 RUTHERFORD conducted the gold foil experiment. They fired alpha 1.6 Recall that most of the mass of an atom is concentrated in the
particles at an extremely thin sheet of gold. The plum pudding model should nucleus
mean the particles should pass straight through or be deflected slightly. 1.7 Recall the meaning of the term mass number of an atom
- However, while some did go straight through, some deflected more than ATOMIC NUMBER: The number of protons (unique to each element). In a
expected and some were deflected backwards. This meant the plum pudding neutral atom, this is also the number of electrons
model was incorrect, and the positive charge must be in a very concentrated MASS NUMBER: Protons + neutrons (only these have mass)
point (the nucleus)
1.8 Describe atoms of a given element as having the same number of
- He then came up with the nuclear model. This shows positively charged protons in the nucleus and that this number is unique to that element
particles in the nucleus, surrounded by a ‘cloud’ of negative electrons
Every atom of an element has the same number of protons, which is unique
4) However, the ‘cloud’ means nothing stopped the electrons from rushing to the element, changing this will make a new element
towards the nucleus, causing the atom to collapse
1.9 Describe isotopes as different atoms of the same element containing
- The BOHR model shows electrons have fixed orbits (shells). Each shell has a the same number of protons but different numbers of neutrons in their
fixed energy level nuclei
5) New experiments showed that the positive charge in the nucleus is made ISOTOPES: Different forms of the same element, which have the same
up of small discrete particles (PROTONS) as well as neutral particles number of protons but a different number of neutrons
(NEUTRONS)
- Thus, they have the same atomic number, but different mass number
1.2 Describe the structure of an atom as a nucleus containing protons
1.10 Calculate the numbers of protons, neutrons and electrons in atoms
and neutrons, surrounded by electrons in shells
given the atomic number and mass number
NUCLEUS: Middle of atom, containing the protons & neutrons. It has an
1.11 Explain how the existence of isotopes results in relative atomic
overall positive charge & most of the mass of an atom is concentrated here. It
masses of some elements not being whole numbers
is very tiny compared to the atom. The electrons surround nucleus in shells
,- Isotopes have different masses so elements need an ‘average mass’. This is THE PERIODIC TABLE
the RELATIVE ATOMIC MASS (Ar) 1.13 Describe how Mendeleev arranged the elements, known at that time,
- If an element only has 1 isotope than its Ar will be the mass number. But if in a periodic table by using properties of these elements and their
there are multiple isotopes the Ar will be an average, taking into account how compounds
much there is of each. Thus, it might not be a whole number - Mendeleev arranged the ~50 elements known into a periodic table. He did
- Relative atomic mass is the average mass of one element compared to this by sorting elements into groups based on their properties
1/12th the mass of Carbon-12 - Doing so, he realised putting elements into order of atomic mass created a
1.12 Calculate the relative atomic mass of an element from the relative pattern where elements with similar properties were put into columns
masses and abundances of its isotopes - He switched the order of specific elements if they did not fit the properties of
- To calculate the Ar you need to use the formula: their group, so their atomic number would be in reverse order
ABUNDANCE: How common/rare the isotope is 1.14 Describe how Mendeleev used his table to predict the existence and
If abundance is in PERCENTAGE %: Ar= properties of some elements not then discovered
- He left some gaps to keep elements with similar properties together. He
then used the properties of the other elements in the column to predict the
new element. When they were found, they fitted the pattern.
1.15 Explain that Mendeleev thought he had arranged elements in order of
increasing relative atomic mass but this was not always true because of
the relative abundance of isotopes of some pairs of elements in the
Otherwise it is divided by: (SUM of ISOTOPIC ABUNDANCE)
periodic table
- Mendeleev arranged the atoms using the mass number, however the mass
number he had was sometimes wrong due to isotopes
- Some elements didn’t fit the pattern so he reversed their order to keep them
in the same groups
1.16 Explain the meaning of atomic number of an element in terms of
position in the periodic table and number of protons in the nucleus
The modern periodic table is ordered by the ATOMIC NUMBER
, IONIC BONDING
1.17 Describe that in the periodic table 1.21 Explain how ionic bonds are formed by the transfer of electrons
a elements are arranged in order of increasing atomic number, in rows between atoms to produce cations and anions, including the use of dot
called periods and cross diagrams
b elements with similar properties are placed in the same vertical - Ionic Bonding occurs between a METAL & NON-METAL. The metal loses the
columns called groups electron (turning into a CATION) & the non-metal gains it (turning into an
- Rows are called PERIODS. Each new period represents another full shell of ANION). Electrons are TRANSFERED. These oppositely charged ions are
electrons, thus the period corresponds to the number of shells. strongly attracted to each other by electrostatic forces
- Columns are called GROUPS. Elements that have similar properties are - Dot & Cross diagrams are used to show where the electrons originally came
arranged in vertical groups. This corresponds to the number of electrons in from, as one atom uses dots and the other uses crosses
the outer shell 1.22 Recall that an ion is an atom or group of atoms with a positive or
1.18 Identify elements as metals or non-metals according to their negative charge
position in the periodic table, explaining this division in terms of the ION: Charged particles (single atom/group). They are formed when an atom
atomic structures of the elements loses/gains an electron to get a full outer shell (stable electronic structure)
Zig-Zag line separates metals (lose electrons) & non-metals (lose electrons) CATION: positive ion ANION: Negative ion
1.19 Predict the electronic configurations of the first 20 elements in the 1.23 Calculate the numbers of protons, neutrons and electrons in simple
periodic table as diagrams and in the form, for example 2.8.1 ions given the atomic number and mass number
In an atom, electrons are placed in fixed shells according to the rules: 1.24 Explain the formation of ions in ionic compounds from their atoms,
limited to compounds of elements in groups 1, 2, 6 and 7
1st shell: 2 electrons. 2nd shell: 8 electrons. 3rd shell: 8 electrons...
Elements most readily forming ions are in Groups 1,2 6, 7
1.20 Explain how the electronic configuration of an element is related to
its position in the periodic table Group 1 & 2 are metals that lose their electrons to form cations
The electronic configuration of an element shows how many electrons in Group 6 & 7 are non-metals that gain electrons to form anions
each shell. So, 2.8.7 shows 2 electrons in 1st shell, 8 in the 2nd shell, & 7 in the 1.25 Explain the use of the endings –ide and –ate in the names of
3rd shell. This can also be found by its position in the periodic table as shown compounds
in 1.17 The name of ionic compounds end in:
- IDE: When only the metal and non-metal are present
Eg: Potassium Iodide
- ATE: When oxygen is present and at least 1 other element + metal
Eg: Magnesium Sulphate
, 1.26 Deduce the formulae of ionic compounds (including oxides, 1.31 Explain the formation of simple molecular substances, using dot and
hydroxides, halides, nitrates, carbonates and sulfates) given the cross diagrams, including: hydrogen, hydrogen chloride, water, methane,
formulae of the constituent ions oxygen, carbon dioxide
In ionic compound, charges must cancel out Hydrogen - ELEMENT H2
Oxide O2- Nitrate NO3-
Hydroxide OH- Carbonate CO32-
Halide Grp. 7 (F-, Cl-, Sulphate SO42- Hydrogen Chloride - HCl
Br-, i- etc.) COMPOUND
1.27 Explain the structure of an ionic compound as a lattice structure
a consisting of a regular arrangement of ions
b held together by strong electrostatic forces (ionic bonds) between
Water - COMPOUND H2O
oppositely-charged ions
- When an ionic bond is made, an ionic compound is formed. These have
giant ionic lattice structures with millions of ions closely packed together,
held by strong electrostatic forces due to the oppositely charged ions.
Methane - CH4
- The ionic compound consists of ions in a regular arrangement
COMPOUND
COVALENT BONDS
1.28 Explain how a covalent bond is formed when a pair of electrons is
shared between two atoms
Covalent bonds are formed when a PAIR of electrons is SHARED between 2 Oxygen - ELEMENT O2
NON-METAL atoms
- Each atom needs to GAIN electrons to get a full outer shell. Thus, they share
electrons so both gain an electron
1.29 Recall that covalent bonding results in the formation of molecules Carbon Dioxide - CO2
COMPOUND
- Molecules are 2 or more atoms chemically bonded together. They can be of
the same element (O2) or different elements.
- There are 2 types of molecules: Simple molecular & Giant covalent. (These
are described later on)
1.30 Recall the typical size (order of magnitude) of atoms and small
molecules
Average size of simple molecular substances: 10-10m (0.1nm)
