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Chemistry Module One

Preliminary Chemistry
Module One: Properties and Structure of Matter
Properties of Matter
Inquiry Question: How do the properties of substances help us to classify and
separate them?
• Explore homogeneous mixtures and heterogeneous mixtures through practical
- Pure Substances: Simplest substances made up of the same type of atom- cannot be
decomposed by physical means- fixed properties
Molecule: Two or more atoms chemically bonded- can be the same (O2) or different (NH3)
Compound: Pure substances consisting of two or more different elements (OH)
Note: Compounds are molecules but not all molecules are compounds- Oxygen gas is a molecule
as it contains two atoms bonded together. However, since atoms are identical- not a compound.
- Mixture- two or more substances physically combined- able to be separated based on
physical properties of individual components
Ratio of composition and physical properties may vary
Homogeneous Mixture- appears as uniform composition- can be made of variable
substances. e.g. air, blood (fixed ratio)
Heterogeneous Mixture- appears as non uniform composition- individual components can be
distinguished. e.g. salt and sand, oil and water
Mass Units: Grams or kilograms. Volume Units: mL, L or cm3
• Using separation techniques based on physical properties
Separation Technique
Property Exploited
Extra Notes
Solubility of Substances
Large Difference in MP/BP
Simple Distillation
Small Difference in MP/BP
Fractional Distillation
Small Difference in MP/BP
Separates many components at
Separating Funnel
Difference in Densities
Separate immiscible liquids
Difference in Particle Size
Decanting and Sedimentation
Difference in Densities
Insoluble solid from liquid/ gas
Simple Chromatography
Difference in Absorption Rates
Mobile phase (ink) absorbs into
stationary phase (paper)
Separation Technique
Property Exploited
Extra Notes
Difference in Densities
Denser objects pushed outwards
(centripetal acceleration)
• Calculation percentage composition of weight of component elements and/or
Gravimetric Analysis
- Used to determine the composition of a mixture by separating the mixture into its components
and weighing them separately
- Uses include finding mineral contents of ores (whether it is economic to mine an ore body) or
determine % of dissolved solids in water (suitability of drinking/ irrigation), % of alcohol
%X =
• Investigate the nomenclature of inorganic substances using International Union of Pure
and Applied Chemistry (IUPAC) naming conventions
Ionic Compounds
- Formed when there is a direct transfer of electrons from (usually) metals to non-metals
- Metals form cations (positively charged ions) as they lose electrons
- Non-metals form anions (negatively charged ions) as they gain electrons
Representation Using Lew Dot Diagrams
- Metals should not have any electrons inside
- Non-metals should have all eight electrons inside
Naming Ionic Compounds
1. Write down the full name of the metal (or cation) first. For transition metals, add roman
numerals to indicate charge
2. Write down the name of the non-metal (or anion), using suffix “ide”
E.g. BaS= Barium Sulfide, PbO= Lead (II) Oxide
Covalent Compounds
- Formed when pairs of atoms (usually non-metals) share electrons to gain a stable octet electron
Representation Using Lewis Dot Diagrams
- Draw out valence electrons
- Represent covalent bonds with a straight line, pairing all lone electron pairs
Naming Covalent Compounds
1. Write down the name of the ‘first element’ and add “ide” to the second element
First element is the furthest to the left (for same period) or furthest down (for same group)
2. Add prefixes to indicate number of atoms
(Mono, Di, Tri, Tetra, Penta, Hexa, Hepta, Oct)
E.g. N2O5= Dinitrogen Pentoxide
Extra Notes:
- All transition elements have a valency of +2 unless stated otherwise
- Exception is Silver which has +1 unless stated otherwise
- Only use roman numeral when writing name of transition metal and not chemical formula
- E.g. Lead (II) Oxide and LbO
• Classify the elements based on their properties and position in the periodic table
through their:
Physical and Chemical Properties of Groups:
Group Number
Name of Group
Physical Properties
Chemical Properties
Alkali Metals
Soft, silvery coloured
Extremely Reactive
Alkaline Earth Metals
Silvery-white Metals
Found in rocks of Earth
Fairly Reactive
Fairly Reactive (with
alkali metals)
Noble Gases
Gaseous at RTP
Properties of Metals vs. Non-Metals vs. Semi-Metals
Metals (Left-Hand Side)
Non-Metals (Right-Hand Side)
Semi-Metals (Staircase)
High Lustre (shiny)
Low Lustre (dull)
Variable Appearance
High electrical and heat
Poor electrical and heat
Good electrical and heat
High malleability and ductility
Poor malleability and ductility
High MP/BP
Very high MP/BP
Atomic Structure and Mass
Investigate the basic structure of stable and unstable isotopes by examining:
• Their position in the periodic table
- Group Number: Number of valence electrons
- Period Number: Number of electron shells
• The distribution of electrons, protons and neutrons in the atoms
- Nucleus: location of protons and neutrons
- Orbitals: location of electrons (NB: electrons exist within sub shells of these orbitals)
• Representation of the symbol, atomic number and mass number (nucleon number)
A: Mass Number (number of neutrons and protons)
Z: Atomic Number (number of protons)
Weight: Average mass of all isotopes of an element
Model the atom’s discrete energy levels, including electronic configuration and spdf
Bohr’s Model and Schrodinger’s Model
Bohr’s Model
Schrodinger’s Model
- Electrons orbit the nucleus in fixed energy shells
- Electrons are bound to discrete energy shells
- Electrons are considered tiny particles
- Electron shells consist of smaller sub shells i.e.
shell > sub shell > orbitals > electrons
- spdf corresponds to orbitals- region of space
occupied by the electron 90% of the time
- Electrons behave as waves
(Note: A “quantum” of energy is the amount of energy absorbed by the electron when excited to
a higher energy level or dropping to a lower energy level)
Drawbacks of Bohr’s Model
- Unable to accurately predict emission spectra of atoms with more than one electron
- Unable to explain why atoms can hold a maximum number of electrons of 2n2
- Does not explain why fourth shell accepts two electrons before filling third shell
Shells, Sub shells, Orbitals
Principal Quantum Number: energy level or shells
Sub shells: Each sub shell describes the orbitals of the electrons within the energy level
Atomic Orbital (spdf): likely location, about 90%, of an electron- region of space of geometric
shape- fits only two electrons
There are four sub-shells/ orbitals: s, p, d and f
- S Orbital: One orbital- spherical shape (max. 2 e)
- P Orbital: Three orbitals- dumbbell shape (max. 6 e)
- D Orbital: Five orbitals- complex shape (max. 10 e)
- F Orbital: Seven orbitals- complex shape (max. 7 e)
Energy Shell/ Level: 1, 2, 3, 4
Sub-shells: s, p, d, f
Number of orbitals: 1, 3, 5, 7
Number of e- : 2, 6, 10, 14
Each energy level (principal quantum number) consists of sub shells which describe the specific
atomic orbitals of that level. E.g. n=2 has 2 sub shells (orbitals s and p)
Electron Configuration (Four Rules)
1. Aufbau Principle: electrons enter the lowest energy orbitals first
2. Pauli Exclusion Principle: atomic orbitals hold a maximum of 2 electrons (represented by an
up and down arrow), within each orbital electrons have a different spin
3. Hund’s Rule: Electrons fill degenerates (meaning orbitals in the same sub-shell with similar
energy) half way first- fill with one spin first throughout, then other spin
4. 2n2 Rule: Maximum number of electrons able to be placed in shell is 2n2 E.g. Third shell- max
is 18 electrons- 3s2, 3p6, 3d10 = 18 electrons (NB: even though it can hold a max of 18
electrons, fill shells with lower energy I.e. 4s)
Atomic Orbitals (spdf)
Aufbau Principle
Hund’s and Pauli Exclusion
Investigate energy levels in atoms and ions through:
• Collecting primary data from a flame test using different ionic solutions of metals
Flame Test
Flame test is a procedure that identifies the presence of certain metal ions
Electrons absorb the energy sourced from the flame
Quantum of energy allows electron to transition to higher energy levels or their excited state
When returning to ground state, this energy is emitted as photons observed as light- emission
spectrum of each element
- Each element produces a different flame colour dependent on their specific energy gaps
Metal Ion
Flame Colour
Intense Yellow
Apple Green
Brick Red
Lilac (Pale Violet)
• Examining spectral evidence for the Bohr Model and introducing the Schrodinger
Emission Spectrum
- Electrons are bound to discrete energy levels- each element has different energy gaps
between their energy shells
- When electrons absorb a ‘quantum of energy’ (through flames), they are excited to a higher
state- excited state
- Upon returning to ground state, they release this quantum of energy as photons with a specific
frequency- called relaxation energy
- Each atom absorbs a unique amount of energy due to having varying energy gaps between
electron shells, hence releases specific frequencies of photons observed as light
- More energy= higher frequency = different light spectra
- Experiment: Electricity passes through gas tube with an element — element absorbs quantum
of energy — releases as photons with fixed energy — energy correlates to frequency and
wavelength — pass through glass prism — coloured spectral lines appears on black
background — emission spectrum Absorption Spectrum
- Similarly the absorption spectrum of each element produces specific black spectral lines
superimposed on visible light spectrum
- Experiment: Light passes through gas-filled tube with an element — element absorbs quantum
of energy — element releases relaxation energy; photons — photons have a specific
wavelength — pass through glass prism — absorption spectrum obtained as visible light
E.g. Emission spectrum
Emission Spectrum
Emission Spectrum: Appear as coloured spectral lines
superimposed on a black background
• Calculate the relative atomic mass from isotopic composition
Atomic weight is the average of all isotopes and its relative abundance in nature
Relative proportions determined experimentally by a mass spectrometer
E.g Distribution of isotopes of Neon is shown
Abundantly 20 amu (mass atomic units) - atomic weight close to 20
Relative Atomic Mass (RAM)- know percentage of isotopic compositions
E.g. Cl-35 has abundance 75.78%, Cl-37 has abundance 27.22%
Mr = 0.7578 x 35 + 0.2722 x 37 = 35.48 amu
Mr = (75.78 x 35 + 27.22 x 37)% divided by 100% = 35.48 amu
Investigate the properties of unstable isotopes using natural and human-made
radioisotopes as examples, including but not limited to:
• Types of radiation
Emissions of radiation occur when a nucleus is unstable
Factors: mass of nucleus and neutron : proton ratio
Neutron to Proton Ratio: nucleus held together by strong nuclear force
Nuclear force > electrostatic repulsion of protons
As proton numbers increase, so does neutron numbers if nucleus is stable
Mass of Nucleus: Greater than 83 (bismuth)
Use of radioisotopes:
- Diagnosis: non-invasive assessment of organ functionality
- Treatment: targeted cancer therapy (chemotherapy)
Technetium- 99
- Emits game radiation able to be detected past tissue and muscles
- Short half life of 6 hours and can assess function of almost every organ
- Used to scan for damage to organs
- Used particularly in treatment of thyroid cancer
- Injected through IV or orally- concentrates in thyroid gland
- Undergoes beta decay and kills cancer cells
Alpha Decay: Helium nuclei (4He2) ejected from a heavy unstable nuclei
Beta Decay: Fast moving electrons ejected from nucleus-surplus of neutrons (high n:p ratio)
Gamma Decay: High energy electromagnetic radiation emitted from a nucleus in an excited
state (often accompanies alpha and beta decay)
Type of Radiation
Penetrating Power
Ionising Power
Low (stopped by piece of paper)
Moderate (stopped by 0.5mm Pb) Moderate
High (stopped by 1m concrete)
Note: Although gamma rays have the low ionising power, it is still considered dangerous because
of its high penetrating power (as compared to alpha particles)
Half life refers to the time taken for half of the material to decay to stable form
• Types of balanced nuclear reactions
Alpha decay:
Beta Decay:
Gamma Decay:
Demonstrate, explain and predict the relationships in the observable trends in the
physical and chemical properties of elements in periods and groups in the periodic
table, including but not limited to:
• Properties/ Trends
Boiling Point
The temperature at which a liquid boils and turns into a vapour
Melting Point
The temperature at which a solid melts and turns into a liquid
Electron Configuration
The distribution of electrons in an atom’s orbitals
Atomic radius
The distance between an atom’s nucleus to its valance shell
1st Ionisation Energy
The energy required to remove the first valence electron from an atom
A measure of an atom’s tendency to attract a shared pair of electrons (covalent)
Reactivity (with water)
How reactive an atom is when placed in water
Core Charge
Expresses the net positive attractive force experienced by the valence electrons
Calculated by ‘atomic number’ minus all electrons (except valence)
e.g. Cl- 17 : 17 - 10 = +7 —> electron configuration is 2, 8, 7
Inner electrons are referred to as shielding - valence electrons experience the ‘core charge’
• Electronic configurations and atomic radii
Atomic Radii: distance from nucleus to valence
electron/ shell
Electron Configuration: Distribution of electrons
in an atom’s orbitals
- Down a Group: Increase in atomic radius
- Across a Period: Decrease in atomic radius
- Going down group- core charges remain constant - increase in electron shell- increase in atomic
radius e.g. Li= +1 and Na= +1 (K has one more shell - larger atomic radius)
- Going across a period- core charges increase - shielding remains the same - decrease in atomic
radius e.g. C= +4 and N= +5 (higher core charge, stronger attraction, small atomic radius)
• State of matter at room temperature (Melting and Boiling Points)
Trend for MP/ BP
Increases from G1-G4
Sudden Drop from G4-G5
Same from G5-G8
Metals - Semi-metals- Non-metals
- Melting and Boiling Point is dependant on the chemical structure of an element (not electronic)
- Across Period: Metals —> semi-metals —> non-metals
- Metallic Lattices in metals (G1-G3) have relatively strong bonds between fixed positive
cations and sea of delocalised electrons (electrostatic forces)
- Covalent Network have extensive covalent bonds which are even stronger (G4)
- Covalent Molecules in non-metals are only held together by weak intermolecular bonds
(dispersion, hydrogen bonds, dipole-dipole) (G5-G8)
- Down Group: MP/BP decreases for metals, increases for non-metals
- Metals- valence electrons are further from nucleus, metallic bonds weaken
- Non-metals- increase in molecular mass, increase strength of dispersion force (uneven e-)
• First ionisation energy and electronegativity
First Ionisation Energy and Electronegativity:
- Down a group: atomic radius increases — valence electrons are further from nucleus —
experience weaker force of attraction — easier to remove — 1st ionisation energy decreases —
electronegativity decreases (unable to attract electrons)
- Across a period: core charge increases — atomic radius decreases — valence electrons closer
and experience stronger force of attraction — 1st ionisation energy increases —
electronegativity increases (stronger force of attraction- easier to attract electrons)
Metals have low electronegativity. Non-metals have high electronegativity.
As valence electrons are removed, atoms become increasingly positively charged, thus stronger
electrostatic forces between electrons and nucleus and an increase in ionisation energy.
Large Jumps in energy occur when last valence electron on outermost shell is removed.
For e.g. Element of Aluminium: 3rd IE is 2754J and 4th IE 11 577J (Group 3- 3 electrons)
• Reactivity with Water
A metal’s reactivity is dependent on its ability to lose electrons (ionisation). The lower
its ionisation energy is, the easier for metals to lose its valence electrons (react)
Ionisation energy for metals increases down a group, reactivity increases.
Ionisation energy for metals decreases across a period, reactivity decreases.
A non-metal’s reactivity is dependent on its ability to gain electrons
(electronegativity). The higher its electronegativity, the easier for non-metals to gain
electrons and hence react.
Electronegativity for non-metals decreases down a group, reactivity decreases.
Electronegativity for non-metals increases across a period, reactivity increases.
Down a Group
Across a Period
Atomic Radius
Ionisation Energy
Reactivity for Metals
Reactivity for Non-Metals
Chemical Bonds
Intermolecular Forces
Dispersion Forces
Dipole-Dipole Forces
Intramolecular Forces
Ionic Bonding
Hydrogen Bonding
Metallic Bonding
Covalent Bonding
Investigate the role of electronegativity in determining the ionic or covalent nature of
bonds between atoms
Electronegativity is a measure of an atom’s ability to attract electrons toward itself when forming
chemical bonds. This can be used to predict the whether a compound is ionic or covalent based
on the difference in electronegativity between its elements.
Ionic Bonds
- Formed when there is a large difference in EN (large difference means one element attracts an
electron completely and hence the transfer of electrons)
- Lower EN cannot keep electron and higher EN can steal electron — complete transfer
- Can be regarded as the extreme limit of covalent bonds
Covalent Bonds
- Formed when there is a small difference in EN (small difference means both elements attract
electrons equally or similar pull on electrons and hence shares the electrons)
In general:
- Difference in EN > 1.7, bond is ionic
- Difference in EN < 1.7, bond is covalent
Investigate the differences between ionic and covalent compounds through:
Using the nomenclature, valency and chemical formulae (including Lew Dot Diagrams)
Ionic Compounds
Covalent Compounds
• Examining the spectrum of bonds between atoms with varying degrees of polarity with
respect to their constituent elements’ positions on the periodic table
• Modelling the shapes of molecular substances
VSEPR Theory
1. Draw the Lewis structure to determine the amount of electron pairs around the central atom
(NB: double or triple bonds count as ONE pair)
2. Determine the optimum geometry (number of electron pairs)
3. Analyse only bonded electron pairs to determine molecular shape (cases within geometry)
4. Find net dipole to determine overall polarity
E.g. H2O
- 4 electron pairs (tetrahedral)
- 2 lone pairs, 2 bonded pairs
- Bent shape
E.g. NH3
- 4 electron pairs (tetrahedral)
- 1 lone pairs, 3 bonded pairs
- Trigonal Pyramidal
Investigate elements that possess the physical property of allotropy
- Different forms of the same element
- Different atomic arrangements resulting in different chemical and physical properties
• Carbon
- Rigid tetrahedral structure with each carbon atom bonded to four other atoms
- Not capable of conducting electricity (electrons ‘tied-up’ in covalent bonds)
- Orderly structure disperses light effectively (brilliant lustre), hard but brittle
- Hexagonal layers or planar structure with each carbon atom bonded to three other atoms
- Excess electrons from each carbon atom allows free, mobile electrons — electrical conductivity
- Dull, black colour, soft and ‘slippery’
Investigate the different chemical structures of atoms and elements, including but not
limited to:
• Ionic networks
- Consists of positive and negative ions arranged in a 3D structure- ionic lattice
- Strong ionic bonds due to oppositely charged ions (strong electrostatic force of attraction)
E.g. Sodium Chloride
• Covalent networks lattice (including diamond and silicon dioxide)
- Covalent bonds exist throughout the lattice structure
- Typically occurs in Group IV where number of valence electrons (4) allow atoms to bond with 4
other atoms, creating a repeating array of atoms
E.g. Carbon
• Covalent molecular substance
- Discrete molecules where atoms are held together by covalent bonds
- Bonding is dependent on interaction of these molecules (intermolecular forces)
- The stronger these interactions (e.g. stronger the dipole-dipole forces), the stronger the force
of attraction
E.g. Oxygen Gas
• Metallic structure
- Fixed positive ions surrounded by a sea of delocalised electrons
- Electrostatic force of attraction (bond) exists between ions and electrons
E.g. Sodium compared with Aluminium
- Trend going across a period: Valency of metal ions increase (more valence electrons lost during
chemical bonding) - Increase in valency means an increase in the number of delocalised
electrons as well as higher core charge —> stronger force of attraction
- Trend going down a group: Number of delocalised electron and core charge remain constantatomic radius is bigger —> weaker force of attraction
Explore the similarities and differences between the nature of intermolecular and
intramolecular bonds and the strength of each forces associated with each, in order to
• Physical properties of elements
• Physical properties of compounds
• Dipole- Dipole
- Electrostatic force of attraction between permanently polar molecules
- Partially negative end (more electronegative) attracts partially positive end (less electronegative)
E.g. HCl
• Hydrogen Bonds
- Special case of dipole-dipole between highly positive nuclei of hydrogen and a small, highly
electronegative element (oxygen, fluorine, nitrogen)
- Large difference in electronegativity; when hydrogen ‘loses’ electron it becomes a positively
charged ‘proton’ nuclei with no shielding
- Highly concentrated and strong partial charge, resulting in strong dipole-dipole forces
E.g. H2O
• Dispersion Forces
- Weak intermolecular forces caused by temporary uneven dipoles
1. Neutral Molecules
2. Uneven distribution of electrons (electron clouds) causes temporary dipoles to form
3. One side becomes positive, other negative, temporary dipole attracts electrons from
neighbouring molecules causing induced dipoles to form — dispersion forces
Electrical Conductivity
Hardness and
Ionic Compound
Relatively High
Low in solid state
High in molten state
Hard but brittle
Relatively High
Hard and malleable
Covalent Network
Extremely High
Hard but brittle
Covalent Molecular
Extremely Low
Soft and malleable
Ionic Compound
- High MP/BP: Strong ionic bonds holding compound together
- Low electrical conductivity (solid): Ions are held in fixed positions within lattice — electrons
are unable to move freely
- High electrical conductivity (molten): Ionic bonds are broken, causing ions to dissociate.
Electrons become mobile and compound can conduct electricity.
- Hard but brittle: Difficult to distort strong ionic bonds, however if distorted like-charges may
align or are brought closer — like charges repel, shattering lattice
- High MP/BP: Strong bonds between fixed positive cations and sea of delocalised electrons
- High electrical conductivity: Sea of delocalised electrons are mobile and hence metal can
conduct electricity
- Hard, malleable and ductile: Difficult to distort as electrostatic force of attraction is relatively
strong. However, if distorted the delocalised electrons will shift to accommodate the distortion.
Lattice does not shatter as electrons prevent cations from repelling.
Covalent Network
- High MP/BP: Strong covalent bonds between atoms are needed to be broken
- Low electrical conductivity: Electrons are immobile as they are ‘tied-up’ in covalent bonds
- Hard but brittle: Hard to break covalent bonds apart. If broken, lattice will shatter
Covalent Molecular
- Low MP/BP: When melting or boiling, only breaking weak intermolecular forces (covalent
bonds are not broken)
- Low electrical conductivity: Electrons are ‘tied-up’ in covalent bonds- cannot flow
- Soft and malleable: Easy to distort weak intermolecular forces which hold molecules together