Chapter 2: Properties of matter
2.1 – Types of matter
Compounds
 Compounds contain different types of atom chemically bonded together in definite
proportions
Homogeneous and heterogeneous mixtures
 Matter can be either classified as pure substance (elements and compounds) or a
mixture
 Mixtures are combinations of either two or more elements, two or more
compounds or two or more elements and compounds
 A mixture has no definite composition, the proportions of each substance can vary
 Difference between a compound and mixture is that compound is two or more
elements chemically bonded but mixture is two or more pure substance physically
combined
 Mixtures can be classified as homogenous or heterogeneous
o Homogenous mixtures form a uniform composition
o Heterogeneous mixture has a composition that varies within the mixture eg.
Sand in water
2.5 – Elements and the periodic table
Metals
Typically metals are:
 Lustrous
 Malleable
 Ductile
 Silver coloured
 Dense
 Of high melting point
 High boiling point
 High tensile strength
 Good conductors of electricity
 Good conductors of heat
Non-metals
Typically, non-metals are:
 Not malleable
 Not ductile
 Dull in colour, not shiny
 Not dense
 Lower in melting and boiling points
 Poor conductors of electricity
 Poor conductor of heat
Metalloids
 Metalloids have some properties of both metal and non-metals
 Tend to be shiny but brittle
 Can conduct heat and electricity but not well
 Known as semi conductors
Distinguishing between physical and chemical properties
 Physical properties describe features that can be observed or measured without
changing the element
 Chemical properties relate to how easily an element undergoes chemical change
o Changing from one substance to another
Chapter 3: Atomic structure and atomic mass
3.1 – Inside atoms
Atomic models
 all atoms are made up of small, positively charged nucleus surrounded by a larger
cloud of negatively charged electron
 nucleus Is made from positive protons and no charged neutrons
Electrons
 electrons are bound to the nucleus due to electrostatic attraction with the proton
The nucleus of an atom
Particles
Symbol
Proton
P
Neutron
n
Electron
e
Charge
+1
0
-1
Size relative to a proton
1
1
1/1800
Mass (kg)
1.673 X 10^-27
1.675 X 10^-27
9.109 X 10^-31
3.2 – Classifying atoms
Different types of atoms
 elements are determined by the number of protons in the atom
 the number of protons is known as the atomic number and represented by a Z
 the number of protons and neutrons is known as the mass number represented by A
Representing atomic structure
 Standard way of representing an atom of a particular element is with its atomic and
mass numbers

Isotopes
 Not all elements have the same mass number as they have more or less neutrons
 Atoms that have the same number of protons but different number of neutrons are
known as isotopes
 Isotopes have identical chemical properties but different physical such as weight and
density
Radioisotopes
 Isotopes which are radioactive are referred to as radioisotopes
 If a nucleus lies within the band of stability then they're stable
 If not, they will decay to form stable isotopes through the emission of alpha
particles, beta particles or gamma radiation

3.3 – Masses of particles
Relative masses
 All in relation to carbon-12
Relative isotopic masses
 Relative isotopic mass, symbol I (subscript r)
Relative isotopic abundance

The percentage abundance of an isotope in the natural environment is called its
relative isotopic abundance
The mass spectrometer
 Relative isotopic masses of elements are their isotopic abundance are determined by
using an instrument called a mass spectrometer
 It separates the individual isotopes in a sample of an element and determines the
mass of each isotope and relative abundance
Mass spectra
 Output of a mass spectrometer is called a mass spectrum and it shows:
o The number of peaks indicates the number of isotopes
o The position of each peak on the horizontal axis indicates the relative isotopic
mass
o The relative height of the peaks corresponds to the relative abundance of the
isotopes
Relative atomic mass
 Average mass of all isotopes of an element is called the relative atomic mass (A
subscript r)
Relative molecular mass
 For elements and compounds that exist as molecules relative molecular mass can be
determined (M subscript r)
 It is equal to the sum of the relative atomic masses of the atoms in the molecule
Relative formula mass
 For ionic compounds the term relative formula mass is used
 Calculated by taking the sum of the relative atomic masses of the elements in the
formula
3.4 – Electronic structure of atoms
Flame test
 Used to identify metals in a sample
 When metal atoms are heated they give off light of a characteristic colour
Emission spectra
 When atoms are heats they give off electromagnetic radiation or light
 If the light passes through a prism, it produces a spectrum with a black background
and a number of coloured lines
 These spectra are known as line spectra or emission spectra and are related to the
electronic structure within the atoms
 Each emission spectrum is unique
 Each line in the spectrum corresponds to light of a different energy
Information from emission spectra

Emission spectra gives clues about the electronic structure of atoms, two being:
o Atoms of the same element produce identical line spectra
o Each element has a unique line spectrum and therefore a unique electronic
structure
The Bohr model of the atom
 The Bohr model proposed the following
o Electrons revolve around the nucleus in fixed, circular orbits
o The electrons orbit corresponds to specific energy levels in the atom
o Electrons can only occupy fixed energy levels and cannot exist between two
energy level
o Orbits of larger radii correspond to higher energy levels
 It is possible for electrons to move between energy levels by absorbing or emitting
energy in the form of light
Electron shells
 Electrons are grouped into different energy levels called electron shells
 Electrons with low energy are in orbits close to the nucleus while high energy
electrons are in outer orbits
 Shells with higher values of n correspond to high energy levels
 As the values of n increase, the energy levels get closer together
Emission spectra and the shell model
 Heating an element can cause an electron to absorb energy and jump to higher
energy levels
 The lowest energy state of an atom is called the ground state
 When electrons absorbs energy and jumps to a higher energy level it is known as an
excited state
 Shortly afterwards, the electron returns to the ground state, releasing a fixed
amount of energy as light
 As electrons falls to a lower energy level it emits energy in the form of light
 This energy is exactly equal to the energy difference between the two energy levels
 Each transition corresponds to a specific energy of light and specific line in the line
spectrum
3.5 – Electronic configuration and the shell model
Ionisation energy
 Ionisation energy is the energy needed to remove an electron from an atom
 Electrons in the same shell
o Are all about the same distance from the nucleus
o Have about the same energy
 Arrangement of electrons around the nucleus is called the electronic configuration
Electronic configuration in shells
 Electrons will generally occupy inner shells before outer shells

Basic electronic configuration of any atoms follow three rules
1. Each shell can contain a maximum number of electrons shown below
2. Lower energy shells fill before higher energy shells
3. Electron shells fill in a particular order
Electron shell (n)
Max number of
electrons
1
2
2
8
3
18
4
32
n
2n^2
Valence electrons
 The outmost shell is called a valence shell
 Electrons in the valence shell are called valence electrons and require the least
energy to remove
3.6 – The Schrödinger model of the atom
A quantum mechanical view of atoms
Quantum mechanics
 The energy of the electrons is said to be quantised
 Schrödinger proposed that electrons behave as waves around the nucleus and
developed a model of the atom called quantum mechanics
 Quantum mechanics describe the behaviour of extremely small particles like
electrons
The Schrödinger model and electron properties
 Fundamental difference between Bohr model and Schrödinger model is the way the
electrons are considered
o Bohr viewed electrons as tiny, hard particles that revolve around the nucleus
In circular orbits
o Schrödinger viewed electrons to have wave like properties, occupying a three
dimensional space known as orbitals
 Schrödinger determined the following
o There are major energy levels in an atom that, for historical reasons, are
called shells
o These shells contain separate energy levels of similar energy, called subshells,
and he labelled them s,p,d,f
Pauli exclusion principle
 States that each orbital can contain maximum of two electrons

Each having a different spin
Electronic configuration and the Schrödinger model
Aufbau principle
 States that the lowest energy orbitals are always filled with electrons first

Hund’s rules
 States that every orbital in a subshell must first be filled with one electron with that
same spin before an orbital is filled with a second electron
Chapter 4: Periodicity
4.2 – Trends in the periodic table: Part 1
State of matter at room temperature
 Melting point increases across periods from group 1 to 14 and then drastically drops
for groups 15-18
Core charge
 Core charge is the measure of the attractive force felt by the valence shell electrons
towards the nucleus
 A shielding effect is displayed due to repulsion of electrons in inner shells
 Core charge = number of protons in the nucleus – number of total inner shell
electrons
 As you look down a group, you can see that:
o The core charge remains constant, but the number of electron shells increase
o The valence electrons become farther from the nucleus thus, will have less
pull towards the nucleus
Trend in core charge
Trend in attraction between
the nucleus and valence
electrons
Left to right across a period increases
Core charge experienced by
the valence shell electrons
in atoms of elements
increase from left to right
across a period
Down a group
Remains constant
Core charge stays constant
down a group, but the
valence electrons are held
less strongly because they
are farther from the nucleus
Electronegativity
 It is the ability of an atom to attract electrons towards itself
 More strongly the valence electrons of an atom are attracted to the nucleus the
greater the electronegativity
Trend in electronegativity
Explanation
Down a group
decreases
Core charge stays constant
while atom radius is
increasing making valence
electrons less strongly
attracted, hence lower
electronegativity
Left to right across a group
increases
Core charge increases
across a period. Therefore,
the valence electrons
become more strongly
attracted to the nucleus. As
a result, electronegativity
increase
Atomic radius
 A measurement used for the size of atoms, distance from the nucleus to the valence
electrons
Trend in atomic radius
Explanation
Down a group
Increases
Core charge stays constant
and number of shell
increases making atomic
radius larger
Left to right across a
decreases
Core charge increases,
period
valence electrons become
more strongly attracted to
the nucleus so the atomic
radius decreases across a
period
4.3 – Trends in the periodic table: Part 2
First ionisation energy
 Process of removing an electron from an atom and forming an ion is called ionisation
 Energy required to remove one electron from an atom of an element in the gas
phase is called first ionisation energy
Down a group
Trend in ionisation energy
Decreases
Left to right across a period
Increases
Explanation
Core charge is constant and
atomic radius increases
therefore valence shells are
less attracted. Energy
required to overcome the
attraction between nucleus
and valence electron is less
Core charge increases and
the number of shell are
constant. Thus the valence
electrons are more strongly
attracted to the nucleus
making the energy required
larger
Reactivity
 Reactivity of an element is an indication of how easily an atom of that element loses
or gains electrons
Reactivity of metals
 Metals lose electrons in reactions
 Weaker the attraction of the valence electrons to the nucleus, the more easily the
electrons can be lost
 Reactivity of metals:
o Increases down a Group because it is easier for a metal atom with a greater
number of shells to lose electrons
o Decreases across the period because the increasing core charge makes it
more difficult for a metal atom to lose electrons
Reactivity of non metals
 Non-metal elements gain electrons in reactions
 Reactivity of non-metals
o Decreases down a group because it is harder for a non-metal atom to attract
electrons into its valence shell with a greater number of shells
o Increases across a period because the increasing core charge makes it easier
for a non-metallic atom to attract electrons
Chapter 5: Bonding
5.1 Metallic bonding
Properties of metals
Metals have the following characteristics:
 Malleable
 Ductile
 Lustrous or reflective when freshly cut or polished
 Often hard with high tensile strength
 Low ionisation energy and electronegativity
 Good conductors of electricity
 High melting and boiling points
 Good conductors of heat
 High density
The metallic bonding model
 In a solid sample of a material
o Positive ions, cations, are arranged in a closely packed three-dimensional
network structure or lattice. They occupy a fixed position in the lattice
o Negatively charged electrons move freely throughout the lattice. These
electrons are called delocalised electrons because they belong to the lattice
as a whole instead of staying in the shell of a particular atom
o Delocalised electrons come from the outer shells of the atoms
o Positive cations are held in the lattice by the electrostatic force of attraction
between these cations and the delocalised electrons. This attraction extends
through the lattice and is called metallic bonding.
Explaining the properties of metals
Property
Explanation
Diagram
Metals are hard and have
Strong electrostatic forces
relatively high boiling points of attraction between the
positive cations and sea of
delocalised of electrons
holds the metallic lattice
together
Metals are good conductors Free moving delocalised
of electricity
electrons will move towards
a positive electrode and
away from a negative
electrode in an electric
circuit
Metals are malleable and
ductile
Metals have high density
Metals are good conductors
of heat
Metals are lustrous
When a force causes layers
of metal ions to move past
one another, the layers are
still held together by their
electrostatic attraction to
the delocalised electrons
between them
Cations in a metal lattice
are closely packed
When the delocalised
electrons bump into one
another and into the metal
ions, they transfer energy to
their neighbours. They
vibrate more quickly with
more energy thus rapidly
transmitting the energy
throughout the lattice
Presence of free electrons
in the lattice reflect light of
all wavelengths and appear
shiny
Metallic bonding model
In a solid sample of a metal:
 Cations are in a closely packed 3-dimensional lattice
 Negatively charged electrons move freely through this lattice called delocalised
electrons
 Delocalised electrons come from the outer shell of the atoms
 Electrostatic force of attraction between cations and delocalised electrons hold the
cations in place
 This attraction extends throughout the lattice and is called metallic bonding
Strength of metallic bonds
 Transition metals are harder, denser and have higher melting points than group 1
and 2 metals due to the smaller size of transitions metals as they have greater core
charge


Outer s subshell and inner d subshell electrons in transition metals are delocalised
Greater number of delocalised electrons in transition metals results in a greater
attraction between electrons and cations resulting in a stronger metallic bond
5.2 – Ionic bonding
Properties of ionic compounds
 Generally ionic bonds:
o Have high melting and boiling points
o Hard but brittle
o Do not conduct electricity at the solid state
o Good conductors of electricity in liquid or molten state
o Vary in solubility In water, insoluble in non-polar solvents such as oil
The ionic bonding model
 Forces between the particles are strong
 No free moving electrons present
 Ions present but in solid state they are not free to move
 When melted ions are free to move so can conduct electricity
 When metallic and non-metallic atoms react, the following steps occur:
o Metal atoms lose electron to non-metal due to lower electronegativity
becoming positively charged (cations)
o Non-metals gain electron, become anions
 Cations and anions then arrange themselves in the following way:
o Large numbers of cations and anions combine to form a three-dimensional
network or lattice
o The lattice is held together strongly by electrostatic forces of attraction, ionic
bonding
Explaining the properties of ionic compounds
High melting points
 Strong electrostatic attraction and lots of energy is required to break it
Hardness and brittleness
 Strong electrostatic forces of attraction between ions so a strong force is needed to
disrupt the crystal lattice
 Will shatter under strong force since it causes layers of ions to move relative to one
another
 They will shift so the anions and cations are in line creating a repulsion force
Electrical conductivity
 Ions in crystal lattice aren’t free to move
 When in liquid state or dissolved ions are free to move and conduct electricity
5.3 – Covalent bonding
Properties of covalent substances
Properties of covalent elements and compounds
 Very low melting and boiling points
o Some of the bonding in covalent substances must be weak
 Absence of electrical conductivity in any phase
o Covalent compounds do not contain ions or delocalised electrons
 Bonds holding atoms together within a molecule is intramolecular bonds
 Bonding holding molecules together is intermolecular bonds
 Intramolecular forces are stronger than intermolecular forces
Covalent bonding
Single covalent bonds
 When atoms share two electrons, one from each atom, the covalent bond formed is
called single covalent bond
o Eg. Hydrogen gas
Double covalent bonds
 When atoms share four electrons, two from each atom, the covalent bond formed is
called double covalent bond
o Eg. Oxygen molecule
Triple covalent bonds
 When atoms share six electrons, three from each atom, the covalent bond formed is
called triple covalent bond
o Eg. Nitrogen molecule
Molecular compounds
Polyatomic molecules
 Molecules made up with more than two atoms are called polyatomic molecules
Allotropes
 Elements exist with their atoms in several different structural arrangement called
allotropes
5.4 – Intermolecular forces
 Covalent molecular substance has much greater range of properties than ionic or
metallic substances
Valence-shell electron-pair repulsion (VSEPR) theory


Valence-shell electron-pair repulsion (VSEPR) theory predicts the shape of small
molecules
Based on the principle that negatively charged electron pairs in the outer shell of an
atom repel each other
Electron-pair repulsion
 VSEPR theory states that the electron pairs around a central atom tends to orient
themselves so that they are as far apart as possible
Lone pairs of electrons
 Lone pairs of electrons are treated I the same way as electron pairs in covalent
bonds in order to determine the shape of a molecule
Double bond, triple bonds and valence-shell electron-pair repulsion theory
 VSEPR theory treats double and triple bonds in the same way that it treats single
bonds and lone pairs
Application of VSEPR
 To apply the VSEPR model correctly, we need to answer two questions
o Which atom is the central atom?
 This is the least electronegative atom that is not hydrogen
o How many (effective) electron pairs are there surrounding the central atom?
 You already know this from electron dot formula
AXE notation
 To keep track of the number and type of (effective) electron pairs are there
surrounding the central atom we denote each bonding pairs with an X and each lone
pair with an E
 The sum of the X and E is known as the steric number
o Eg. H2O has two bonding pairs and two lone pairs therefore steric number of
4
o CO2 has two double bonds and no lone pairs therefore steric number of 2
Steric number 2
 2 electron pairs around the central atom
 Electron pair geometry is called linear
 Molecular shape is linear
Steric number 3
 3 electron pairs around the central atom
 Electron pair geometry is called trigonal planar
 AX3
o Electron pair geometry is called trigonal planer
o Molecular shape is trigonal planar
Steric number 4
 4 electron pairs around the central atom
 Electron pair geometry is called tetrahedral
 AX4
o Molecular shape is called tetrahedral
 AX3E
o Molecular shape is called trigonal pyramidal
 AX2E2
o Molecular structure is called bent
 AXE3
o Molecular shape is linear
Intermolecular forces
Strength of intermolecular forces
 Intermolecular forces are 100 times weaker than bonds found in ionic, metallic and
covalent bonds
Electronegativity and polarity
 Are an example of electrostatic forces
Polarity of diatomic molecules
 Electronegativity determines the electron distribution in diatomic molecules
 Non-polar diatomic molecules
o Covalent bonds with an equal distribution of valence electrons are said to be
non-polar because there is no charge on either end of the molecule
o This happens between atoms with very similar electronegativity
 Polar diatomic molecules
o If the covalent bond between atoms of two different elements, then the
electrons will stay closer to the more electronegative atom because it has a
stronger pull on the electrons in the bond
o Molecules with imbalanced electron distribution is called polar
o The separation of the positive and negative charges is known as an electric
dipole or dipole
Polarity of polyatomic molecules
 Generally symmetrical molecules are non-polar and asymmetrical molecules are
polar
Non-polar molecules
 Even molecules with polar covalent bonds can be non-polar if the molecule is
symmetrical
Polar molecules
 In asymmetrical molecules, the individual dipoles of the covalent bonds do not
cancel each other out, resulting in a net dipole and the overall molecule is polar
Types of intermolecular forces
 There are three main types
o Dipole-dipole forces
o Hydrogen bonding
o Dispersion forces
Dipole-dipole forces
 Only occurs In polar molecules
 Force is a result from electrostatic attraction between negative and positive ends of
the polar molecules
 Dipole-dipole forces are relatively weak however the more polar a molecule the
stronger the force
 Stronger dipole-dipole forces require more energy to break
Hydrogen bonding
 Special form of dipole-dipole forces
 Only occurs between molecules where the hydrogen is covalently bonded to an
oxygen, nitrogen or a fluorine
 Nitrogen, oxygen and fluorine are small and highly electronegative and strongly
attract the electrons within a covalent bond
 The exposed positive hydrogen is attracted to a negative lone pair on a nitrogen,
oxygen or fluorine atom in another molecule
 This results in a relatively strong bond being ten times stronger than dipole-dipole
but only one tenth of ionic and covalent bonds
 Stronger bond leads to higher melting and boiling points
Dispersion forces
 They are forces of attraction between non-polar molecules
 Are caused by a temporary dipole in the molecule that are the result of random
movement of the electrons surrounding the molecule
 These temporary dipoles are called instantaneous dipoles
 Strength of dispersion forces increase as the size of the molecule increases because
larger molecules have a larger number of electrons, making it easier to produce
temporary dipoles
 Stronger dispersion forces means higher melting and boiling
 The shape of a molecule also influences the strength, molecules forming long chains
will tend to have higher dispersion forces than more compact molecules with similar
number of electrons

Large molecules may have stronger dispersion forces than even hydrogen bonding
and dipole-dipole
5.5 – Covalent network structures
 Three dimensional structures are called covalent network structures
 Generally high melting points or decomposition temperature and very hard
