CHEMISTRY
PAPER 1 CONTENT!!
, CHEMISTRY
1.1.1 A SIMPLE MODEL OF THE ATOM, SYMBOLS,
RELATIVE ATOMIC MASS, ELECTRONIC CHARGE
AND ISOTOPES
1.1.1 ATOMS, ELEMENTS AND COMPOUNDS
All substances are made of atoms.
An atom is the smallest part of an element that can exist.
A chemical symbol represents atoms of each element, e.g. O represents an atom of oxygen, Na
represents an atom of sodium.
There are about 100 different elements.
Elements are shown in the periodic table.
Compounds are formed from elements by chemical reactions.
Chemical reactions always involve forming one or more new substances, often involving a
detectable energy change.
Compounds contain two or more elements chemically combined in fixed proportions and can be
represented by formulae using the symbols of the atoms from which they were formed.
Compounds can only be separated into elements by chemical reactions.
Chemical reactions can be represented by word equations or equations using symbols and
formulae.
1.1.2 MIXTURES
A mixture consists of two or more elements or compounds not chemically combined together. ?
The chemical properties of each substance in the mixture are unchanged. Mixtures can be
separated by physical processes such as filtration, crystallisation, simple distillation, fractional
distillation and chromatography.
These physical processes do not involve chemical reactions and no new substances are made.
1.1.3 THE DEVELOPMENT OF THE MODEL OF THE
ATOM
Before the discovery of the electron, atoms were thought to be tiny spheres that could not be
divided.
The discovery of the electron led to the plum pudding model of the atom which suggested that
the atom is a ball of positive charge with negative electrons embedded within
The results from the alpha particle scattering experiment led to the conclusion that the mass of
an atom was concentrated at the centre (nucleus) and that the nucleus was charged.
This nuclear model replaced the plum pudding model. Niels Bohr adapted the nuclear model by
suggesting that electrons orbit the nucleus at specific distances.
The theoretical calculations of Bohr agreed with experimental observations.
Later experiments led to the idea that the positive charge of any nucleus could be subdivided
into a whole number of smaller particles, each particle having the same amount of positive
charge.
The name proton was given to these particles.
The experimental work of James Chadwick provided the evidence to show the existence of
neutrons within the nucleus.
This was about 20 years after the nucleus became an accepted scientific idea.
, CHEMISTRY
1.1.4 RELATIVE ELECTRICAL CHARGES OF
SUBATOMIC PARTICLES
The relative electrical charges of the particles in atoms are:
NAMEOF RELATIVE
PARTICLE CHARGE In an atom, the number of electrons is equal to the number of
protons in the nucleus.
Atoms have no overall electrical charge.
PROTON
+1 The number of protons in an atom of an element is its atomic
number.
NEUTRON All atoms of a particular element have the same number of
0 protons.
Atoms of different elements have different numbers of
ELECTRON -1 protons.
NAME OF RELATIVE
PARTICLE MASS
1.5 SIZE AND MASS OF ATOMS PROTON 1
Atoms are very small, having a radius of about 0.1 nm
(1 x 10-10 m). The radius of a nucleus is less than 1/10
NEUTRON 1
000 of that of the atom (about 1 x 10-14 m). Almost all
of the mass of an atom is in the nucleus. The relative ELECTRON
VERY
SMALL
masses of protons, neutrons and electrons are:
The sum of the protons and neutrons in an atom is its mass number.
Atoms of the same element can have different numbers of neutrons;
these atoms are called isotopes of that element.
Mass number : 23
Atoms can be represented as shown in this example: Na
Atomic number: 11
1.1.6 RELATIVE ATOMIC MASS
The relative atomic mass of an element is an average value that takes account of the
abundance of the isotopes of the element.
, CHEMISTRY
1.1.7 ELECTRONIC STRUCTURE
The electrons in an atom occupy the lowest available
energy levels (innermost available shells).
The electronic structure of an atom can be
represented by numbers or by a diagram.
For example, the electronic structure of sodium is
2,8,1 or :
showing two electrons in the lowest energy level,
eight in the second energy level and one in the third
energy level.
(1.2) 1.2.1 THE PERIODIC TABLE
The elements in the periodic table are arranged in order of atomic (proton) number and so that
elements with similar properties are in columns, known as groups.
The table is called a periodic table because similar properties occur at regular intervals. Elements
in the same group in the periodic table have the same number of electrons in their outer shell
(outer electrons) and this gives them similar chemical properties.
1.2.2 DEVELOPMENT OF THE PERIODIC TABLE
Before the discovery of protons, neutrons and electrons, scientists attempted to classify the
elements by arranging them in order of their atomic weights.
The early periodic tables were incomplete and some elements were placed in inappropriate
groups if the strict order of atomic weights was followed.
Mendeleev overcame some of the problems by leaving gaps for elements that he thought had not
been discovered and in some places changed the order based on atomic weights.
Elements with properties predicted by Mendeleev were discovered and filled the gaps.
Knowledge of isotopes made it possible to explain why the order based on atomic weights was
not always correct.
1.2.3 METALS AND NON-METALS
Elements that react to form positive ions are metals.
Elements that do not form positive ions are non-metals.
The majority of elements are metals.
Metals are found to the left and towards the bottom of the periodic table.
Non-metals are found towards the right and top of the periodic table.
1.2.4 GROUP 0
The elements in Group 0 of the periodic table are called the noble gases. They are unreactive and
do not easily form molecules because their atoms have stable arrangements of electrons. The
noble gases have eight electrons in their outer shell, except for helium, which has only two
electrons. The boiling points of the noble gases increase with increasing relative atomic mass
(going down the group).
PAPER 1 CONTENT!!
, CHEMISTRY
1.1.1 A SIMPLE MODEL OF THE ATOM, SYMBOLS,
RELATIVE ATOMIC MASS, ELECTRONIC CHARGE
AND ISOTOPES
1.1.1 ATOMS, ELEMENTS AND COMPOUNDS
All substances are made of atoms.
An atom is the smallest part of an element that can exist.
A chemical symbol represents atoms of each element, e.g. O represents an atom of oxygen, Na
represents an atom of sodium.
There are about 100 different elements.
Elements are shown in the periodic table.
Compounds are formed from elements by chemical reactions.
Chemical reactions always involve forming one or more new substances, often involving a
detectable energy change.
Compounds contain two or more elements chemically combined in fixed proportions and can be
represented by formulae using the symbols of the atoms from which they were formed.
Compounds can only be separated into elements by chemical reactions.
Chemical reactions can be represented by word equations or equations using symbols and
formulae.
1.1.2 MIXTURES
A mixture consists of two or more elements or compounds not chemically combined together. ?
The chemical properties of each substance in the mixture are unchanged. Mixtures can be
separated by physical processes such as filtration, crystallisation, simple distillation, fractional
distillation and chromatography.
These physical processes do not involve chemical reactions and no new substances are made.
1.1.3 THE DEVELOPMENT OF THE MODEL OF THE
ATOM
Before the discovery of the electron, atoms were thought to be tiny spheres that could not be
divided.
The discovery of the electron led to the plum pudding model of the atom which suggested that
the atom is a ball of positive charge with negative electrons embedded within
The results from the alpha particle scattering experiment led to the conclusion that the mass of
an atom was concentrated at the centre (nucleus) and that the nucleus was charged.
This nuclear model replaced the plum pudding model. Niels Bohr adapted the nuclear model by
suggesting that electrons orbit the nucleus at specific distances.
The theoretical calculations of Bohr agreed with experimental observations.
Later experiments led to the idea that the positive charge of any nucleus could be subdivided
into a whole number of smaller particles, each particle having the same amount of positive
charge.
The name proton was given to these particles.
The experimental work of James Chadwick provided the evidence to show the existence of
neutrons within the nucleus.
This was about 20 years after the nucleus became an accepted scientific idea.
, CHEMISTRY
1.1.4 RELATIVE ELECTRICAL CHARGES OF
SUBATOMIC PARTICLES
The relative electrical charges of the particles in atoms are:
NAMEOF RELATIVE
PARTICLE CHARGE In an atom, the number of electrons is equal to the number of
protons in the nucleus.
Atoms have no overall electrical charge.
PROTON
+1 The number of protons in an atom of an element is its atomic
number.
NEUTRON All atoms of a particular element have the same number of
0 protons.
Atoms of different elements have different numbers of
ELECTRON -1 protons.
NAME OF RELATIVE
PARTICLE MASS
1.5 SIZE AND MASS OF ATOMS PROTON 1
Atoms are very small, having a radius of about 0.1 nm
(1 x 10-10 m). The radius of a nucleus is less than 1/10
NEUTRON 1
000 of that of the atom (about 1 x 10-14 m). Almost all
of the mass of an atom is in the nucleus. The relative ELECTRON
VERY
SMALL
masses of protons, neutrons and electrons are:
The sum of the protons and neutrons in an atom is its mass number.
Atoms of the same element can have different numbers of neutrons;
these atoms are called isotopes of that element.
Mass number : 23
Atoms can be represented as shown in this example: Na
Atomic number: 11
1.1.6 RELATIVE ATOMIC MASS
The relative atomic mass of an element is an average value that takes account of the
abundance of the isotopes of the element.
, CHEMISTRY
1.1.7 ELECTRONIC STRUCTURE
The electrons in an atom occupy the lowest available
energy levels (innermost available shells).
The electronic structure of an atom can be
represented by numbers or by a diagram.
For example, the electronic structure of sodium is
2,8,1 or :
showing two electrons in the lowest energy level,
eight in the second energy level and one in the third
energy level.
(1.2) 1.2.1 THE PERIODIC TABLE
The elements in the periodic table are arranged in order of atomic (proton) number and so that
elements with similar properties are in columns, known as groups.
The table is called a periodic table because similar properties occur at regular intervals. Elements
in the same group in the periodic table have the same number of electrons in their outer shell
(outer electrons) and this gives them similar chemical properties.
1.2.2 DEVELOPMENT OF THE PERIODIC TABLE
Before the discovery of protons, neutrons and electrons, scientists attempted to classify the
elements by arranging them in order of their atomic weights.
The early periodic tables were incomplete and some elements were placed in inappropriate
groups if the strict order of atomic weights was followed.
Mendeleev overcame some of the problems by leaving gaps for elements that he thought had not
been discovered and in some places changed the order based on atomic weights.
Elements with properties predicted by Mendeleev were discovered and filled the gaps.
Knowledge of isotopes made it possible to explain why the order based on atomic weights was
not always correct.
1.2.3 METALS AND NON-METALS
Elements that react to form positive ions are metals.
Elements that do not form positive ions are non-metals.
The majority of elements are metals.
Metals are found to the left and towards the bottom of the periodic table.
Non-metals are found towards the right and top of the periodic table.
1.2.4 GROUP 0
The elements in Group 0 of the periodic table are called the noble gases. They are unreactive and
do not easily form molecules because their atoms have stable arrangements of electrons. The
noble gases have eight electrons in their outer shell, except for helium, which has only two
electrons. The boiling points of the noble gases increase with increasing relative atomic mass
(going down the group).
