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Key
Bold Underline: Module/Topic (heading 1)
Red: Inquiry question (heading 2)
Green: Subheading/title (heading 3) Bold: definition/sub subheading/significant word Blue: syllabus points (title) Normal: content
- Module 1 - Properties & Structure of Matter Table of Contents
- Inquiry Question 1: How do the properties of matter help us to classify and separate them?
- Definitions
- Polyatomic Ions
- Separation Techniques
- Properties
- Types of Physical Properties
- Calculating Percentage Composition (Gravitational Analysis)
- Particle Nature of Matter
- Naming Substances
- Valency
- History
- Inquiry Question 2: Why are atoms of elements different from one another?
- Structure of the Atom
- Isotopes
- Radioisotopes
- Half Lives
- Electronic Structure: Schrödhinger
- Flame Tests
- Inquiry Question 3: Periodicity. Are there patterns in the properties of elements?
- Inquiry Question 4: Bonding. What binds atoms together in elements and compounds?
- Metallic Bonding
- Ionic Bonding
- Covalent Bonding
- Types of Bonding
- Shape of Molecules
- Intermolecular forces
- Allotropy
- Lewis Dot Structure
- Module 2 - Introduction to Quantitative Chemistry - Chemical Reactions & Stoichiometry - Mole Concept - Calculations Using Moles - Limiting reagents:
- reacted/used up of the other substance Definition: the smallest amount of substance in a reaction which limits the amount that will be
- How to do each type of mole calculations:
- Concentration and Molarity
- Dilution
- Pressure
- Molar Volume
- Properties of Gases
- Ideal Gas Law
- Module 3 - Reactive Chemistry
- Types of reactions
- Indications of Chemical change (assumed knowledge)
- Titration:
- Acids and Bases:
- Redox Reactions
- Rates of Reactions:
- Collision Theory:
- Galvanic Cells
- Module 4 - Drivers of Reactions
- Enthalphy
- Thermochemical Equations
Molecule: An element or compound containing non-metal atoms covalently bonded
Ion: Charged atom, pos - cation , neg - anion
Covalent bond: ‘Sharing’ of electron(s) between two non-metal atoms Ionic bond: ‘Sharing’ of electron(s) between a metal and a non-metal
Validity: variables; reliability; accuracy; if the experiment tested the aim; good method Accuracy: correct use of equipment; correct calibration of equipment Reliability: repetition; consistent results
Independent variable: What you are changing; x Dependent variable: What you are measuring; dependent on x; f(x) or y Aim: What you aim to achieve/figure out in the experiment Hypothesis: What you predict; ‘If _____ then ____ Method: Numbered steps on how to conduct the experiment Conclusion: What did you infer from the results Discussion: The significance of your findings; what could you have changed Graph: line/curve of best fit; numbered; title on the side; title; independent on the x axis; points plotted with crosses
Law of conservation of mass: Matter cannot be created nor destroyed only changed; mass of the reactants must balance the mass of the products.
Alloy: Mixture predominantly metal with other substances, e.g. bronze or rose gold
Polyatomic Ions
Polyatomic ion (radical) IUPAC Name Charge
SO 3 2-^ sulfite ion -
SO 4 2-^ sulfate ion -
NO 3 2-^ nitrite ion -
CO 3 2-^ carbonate ion -
PO 4 3-^ phosphate ion -
NH 4 +^ ammonium ion +
● explore homogeneous mixtures and heterogeneous mixtures through practical investigations:
- using separation techniques based on physical properties (ACSCH026)
Separation Techniques
Decantation: Separation of mixtures based on density, for heterogeneous and immiscible mixtures composed in layers. Physically pouring the less dense layer off the top of the mixture.
Separating funnel: Separating mixtures based on density, for heterogeneous, immiscible mixtures. Immiscible mixture goes in the separating funnel and the denser substance is separated out.
Centrifuge: Separating mixtures based on density. The use of centrifugal force forces the denser substance to the end of the tube being able to be decanted off after.
Magnetic separation: Based on the physical property of magnetism, using a magnet to separate the magnetic substance from the one that is not.
Evaporation/Crystallisation: Separation based on evaporation point in a mixture, could be immiscible or miscible mixtures. Evaporation allows for the solute to evaporate in a natural environment from the solute and is slower. Crystallisation is the direct application of heat to drive off the solvent.
Filter Funnel: Separation based on particle size of a mixture, could be suspension. Insoluble particles in a substance are trapped by the filter paper leaving the liquid in the beaker.
Distillation: Separation based on boiling points. Two miscible substances with different boiling points are heated up and the substance with the lower boiling point evaporates and is collected. If there are multiple substances or with very close boiling points, use fractional distillation.
Paper Chromatography: Separation based on solubility. Tests if a reaction has occurred if the two substances create a new dot in a new position, if still two dots then reaction did not occur. Also tests purity of a substance.
Sieve: Separation due to particle size.
Properties ● classify the elements based on their properties and position in the periodic table through their:
- physical properties
- chemical properties
Metal properties:
- Lustre
- Malleable
- Ductile
- Good conductors of heat and electricity
- Delocalised electrons
- Hard
- High melting point
- Usually solid
Naming Substances ● investigate the nomenclature of inorganic substances using International Union of Pure and Applied Chemistry (IUPAC) naming conventions
If a compound contains a cation : the cation will be first, then the anion with the suffix ‘ide’ Cations : keep their name Anions : end in ide If a compound contains oxygen : the compound will end in ate or ite Two oxygen atoms : end in ‘ate’ Three oxygen atoms : end in ‘ite’
For Ionic compounds:
- Write down the name of the compound
- Write down the symbols for the ions in the compound
- Write down the combining numbers (valency).
Valency
Valence shell: The outermost shell of the atom, containing valence electrons that usually dictate the reactivity of the atom
Valency: The ability of an atom to combine with other elements or combinations of elements (polyatomics)
Number of valence electrons: Increases left to right across the periodic table
History ● examining spectral evidence for the Bohr model and introducing the Schrödinger model Bohr: Created the Bohr model of the atom in 1913 with distinct electron shells. One of the limitations of Bohr’s model was that electrons with the same energy were in the same shell
Schrödhinger: Created the most recent atomic model in 1926 where the electrons exist in energy waves and their locations are based on averages where they most likely will be, also SPDF. Also called electron cloud model
Inquiry Question 2: Why are atoms of elements different from one another?
Structure of the Atom ● investigate the basic structure of stable and unstable isotopes by examining:
- their position in the periodic table
- the distribution of electrons, protons and neutrons in the atom
- representation of the symbol, atomic number and mass number (nucleon number)
Structure: Protons and neutrons tightly packed together in the nucleus, protons have a positive charge and usually repel each other but nuclear force compacts them. Electrons are on the outside of the nucleus in orbitals, held together by electrostatic forces of attraction, the positive protons attract the negative electrons.
● model the atom’s discrete energy levels, including electronic configuration and spdf notation Orbitals: S, P, D and F orbitals, can be used in SPDF notation to describe which shells fill up and how many electrons an atom has
Isotopes
Isotope: Atoms of an element that have a differing number of neutrons in the nucleus. Have the same atomic number but different mass numbers. Have identical or very similar properties to their other isotopes. Relative abundance (atomic mass) is the average of all isotopes mass in nature
● calculate the relative atomic mass from isotopic composition Calculating the relative atomic mass: Adding the number of protons to the number of neutrons for the specific isotope of the element you're considering. For example, a carbon-12 atom has 6 protons and 6 neutrons, and so has a relative atomic mass of 12.
Calculating relative abundance: sum of (percentage abundance x isotope mass number) / 100 (or sum of all isotopes)
Radioisotopes
Radioisotope: When a nuclei has an unstable number of neutrons they will emit radiation ● investigate the properties of unstable isotopes using natural and human-made radioisotopes as examples, including but not limited to:
- types of radiation
- types of balanced nuclear reactions
Why are some isotopes radioactive: They may be too large, (the mass attractions overcome the electrostatic forces making it unstable). Or if their ratio from protons to neutrons is too high
Flame Tests ● investigate energy levels in atoms and ions through:
- collecting primary data from a flame test using different ionic solutions of metals
Heating of an atom: When an atom is heated electrons jump up an energy level, then jump back down as that is unstable but due to the law of conservation of mass they emit a photon of light specific to the atom when jumping back down.
To identify atom: because each atom emits a specific colour
Light emission spectra: When the light the atoms create passes through a prism, it produced a number of coloured lines, related to the electronic structure within atoms
Inquiry Question 3: Periodicity. Are there patterns in the properties of elements?
Shielding Effect
The attractive force of the nucleus on the outermost electrons is lessened by the shielding of the force from other closer orbiting electrons
Trends ● 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:
- state of matter at room temperature
- electronic configurations and atomic radii
- first ionisation energy and electronegativity
- reactivity with water
Electronegativity: The ability of an atom to attract electrons Ionisation energy: The energy required to remove electrons from the outermost shell of an atom
Atomic radius: increases top to bottom (each period adds an orbital) and decreases left to right (because as the atomic number increases so does the attractive nucleus force pulling the atom together)
Melting point: Metals and metalloids generally have higher melting points (as metals have high intermolecular forces making it harder for heat to separate them into a liquid)
Electronegativity: Increases across a period as the atoms have a higher charge and a smaller radius increasing their force of attraction
Ionisation energy: Increases as it is harder to remove an electron across a period, making the energy required far higher as well
Group 1: Highly reactive with water as they only have 1 electron in the valence shell so when they move that electron all the energy goes into that electron creating a highly reactive reaction. More reactive as you move down
Group 2: react with water but less vigorously than group 1
Group 8: Noble gases, are unreactive as full valence shell, also have a high ionisation energy
- Always name the metal first. It keeps its whole name, eg magnesium, sodium, copper, etc.
- If there is 1 non-metal, its ending is changed to ‘ide’. So we have chlorine/chloride, oxygen/oxide, bromine/bromide.
- If the are 2 non-metals with a metal, one of which is oxygen, the other
- non-metal gets an ‘ate’ ending, eg nitrogen + oxygen = nitrate, carbon + oxygen = carbonate and so on.
Covalent Bonding
- formed between two atoms sharing electrons (discreet molecules)
- no electrostatic attraction between the atoms
- inter molecular forces: occurs between molecules
- intra molecular forces: occurs between atoms or ions
- covalent molecular substances: forces holding the molecules together
- covalent lattices : extends indefinitely through a crystal
- Properties associated with covalent bonding: - Low melting points: Weak attractive forces between molecules - Non conductors: Lack of delocalised electrons and mobile charged species - Soft: Weak forces existing between molecules
Types of Bonding ● explore the similarities and differences between the nature of intermolecular and intramolecular bonds and the strength of the forces associated with each, in order to explain the:
- physical properties of elements
- physical properties of compounds
- All atoms in matter are held together by forces
- The size of the force determines the bonding and type of bonding
- The size of forces holding atoms together is related to the difference in electronegativity
- the ability of an atom to attract an electron to itself. The more strongly the valence electrons are attracted to the nucleus of the atom the greater the electronegativity. Across the period - increases, down the group - decreases
- There is a continuum of bonding types, ranging from non-polar covalent to polar covalent to ionic
- The level of electronegativity determines ionic or covalent nature of bonds between atoms
Dispersion force : Temporary dipoles, exist in all molecules but are so weak that they are irrelevant when there are other forces active Dipole - dipole : Formed when molecules have atoms with different electronegativity, stronger than dispersion, for example HCl, more elections would sit around the Cl atom rather than the H, making it more electronegative. Therefore the Cl pole is more negative, and the H is more positively charged
Hydrogen bonding : A special case of dipole-dipole, only looks at molecules that have the most electronegative atoms in them, N, O or F, e.g. H2O, ammonia or HFNH
Shape of Molecules
Turn out different because electrons repel each other
Formula BeCl 2 BCl 3 CH 4 NH 3 H 2 O
Name of molecule
Beryllium Chloride
Boron trichloride
Methane Ammon ia Water
Bonding Pairs
Valence Electrons
Lone Pairs 0 0 0 1 2
Angles between bonding pairs
Name of shape
Linear Trigonal Planar
Tetrahedral Triagonal Pyramid
Bent
Image
Intermolecular forces
- Different to intramolecular forces as those are between aspects between molecules
- An attractive force between neighbouring molecules
- Examples: these are far weaker than covalent or ionic bonds
- Permanent dipole-dipole forces
- Hydrogen forces
- Van der Waals force
Allotropy ● Made up of the same element with different 3D structure. The arrangement of the atoms ○ Covalent bonding, hence covalent network lattices
Module 2 - Introduction to Quantitative Chemistry
Writing chemical equations (types of chemical reactions assumed knowledge) Balancing chemical equations Conservation of mass Mole concept and related calculations Concentration and Molarity Gas Laws
Chemical Reactions & Stoichiometry ● conduct practical investigations to observe and measure the quantitative relationships of chemical reactions, including but not limited to:
- masses of solids and/or liquids in chemical reactions
- volumes of gases in chemical reactions (ACSCH046) ● relate stoichiometry to the law of conservation of mass in chemical reactions by investigating:
- balancing chemical equations (ACSCH039)
- solving problems regarding mass changes in chemical reactions (ACSCH046)
Stoichiometry : quantities and ratios of chemical species in a chemical reaction, enabling us to predict the yield of products, whether a reactant is limiting or in excess
- Using a balanced chemical equation to calculate amounts of reactants and products
- Allows for the prediction of the yield of products and the limiting reagent
Yield: amount produced
Limiting reagent:
Aqueous (aq): dissolved in water
Writing Chemical Reactions:
- Identify reactants and products and place them in a word equation
- Convert the chemical names into chemical formulas and write the state symbols
- BALANCE the equation, there must be the same amount of atoms on each side of the equation, this is because of conservation of mass. Total mass of reactants = total mass of products
How to conduct stoichiometry:
- Write the balanced equation
- List the initial amounts in the equation before the reaction
- List the changes in the chemical reactants
- Work out the change in mass, being equal to the sum of the mass lost by the reactants
- Is basically simple algebra
Excess reactant: When there is too much reactant in a chemical reaction
- Minus the smaller mass of reactant from the larger one to work out how much is excess?
Mass-Volume Stoichiometry:
- d = m/v , d is density in g mL or gL, m is mass in g, V is volume in mL or L
- Use this formula for determining the mass of substances
Mole Concept ● conduct a practical investigation to demonstrate and calculate the molar mass (mass of one mole) of:
- an element
- a compound (ACSCH046) ● conduct an investigation to determine that chemicals react in simple whole number ratios by moles ● explore the concept of the mole and relate this to Avogadro’s constant to describe, calculate and manipulate masses, chemical amounts and numbers of particles in: (ACSCH007, ACSCH039)
- moles of elements and compounds n =
- Name for a number, no logical reason similar to the word ‘dozen’ being used for 12 bread loaves
- New way of defining a certain amount of something
- Basically just a set amount, a mole is this set amount: 6.022x10^23
- unit used to measure the amount of chemical
Avogadro’s Number (Avogadro’s Constant):
- Avogadro’s number (NA) = 1mol = 6.022x10^23 units
Relative Atomic Mass: All atomic weights are relative to mass of carbon-
Molecular mass: The mass of a molecule that is the sum of all the masses of all the atoms contained in the molecule
Mole set mass:
- 1mol of any element/compound = its atomic, molecular or formula weight
- Is calculated in grams, 12g of C-12 contains 6.022x10^23 atoms
- n = m / Mw , n is moles (mol), m is mass (g), Mw is molar mass (g/mol), to find n to use in the equation below to find the amount of molecules of an element or compound in a certain weight
n (moles, mol) =
number of molecules/atoms
Avogadro’s number
Calculations Using Moles C 6 H 12 O 6 = molecular formula CH 2 O = empirical formula → the lowest ratio To find mass of unknown:
- Write equation
- Balance
- Find molecular mass (Mw)
- Put into n = m / Mw to find n using given m mass
- Use ratio (put substance you are finding on the top) from the balanced numbers to find unknown mass then put into n = m / Mw formula
Chemical concentrations are expressed in terms of molarity (symbol C ) which expresses the number of atoms, ions or molecules in a certain volume rather than simply the mass per volume.
UNIT DESCRIPTION EXAMPLE OF USE
gL-^
Mass of solute, in g, per litre of solution A bottle of bleach gives the concentration of sodium, hypochlorite as 56gL-
mgL-^
Mass of solute, in mg, per litre of solution Sydney’s water supply has calcium ion concentration typically around 5-10mgL-
molL-^
Moles per litre of solution Hydrochloric acid used in analysis in a laboratory is labelled as 1. mol L-
%(w/w) Percentage by mass; describes the mass of solutes, in g, per 100g of solution
A bottle of hydrogen peroxide is labelled 10%(w/w)
%(w/v) Percentage mass by volume; describes the mass of solute in g, per 100mL solution
Saline drips are a solution of 0.9%(w/v) sodium chloride
%(v/v) Percentage by volume; describes the volume of solute in mL, per 100mL solution
Alcohol content of wine might be 13.5% v/v
ppm Parts per million, e.g. mass, in g, per million equivalent to mg kg-1, or mgL-
The legal limit of mercury, for safe human consumption is 0.5ppm for most species of Australian fish
Ppb Parts per billion, e.g. mass in g, per billion equivalent to μgkg-1, or μgL-
The world Health organisations guideline value for a safe level of arsenic in drinking water is 10ppb
Dilution
- When a solution is diluted more solvent is added, therefore the amount of starting solution stays the same n 1 =n 2
- To calculate the effect of dilution on concentration: C 1 V 1 =C 2 V 2 ( note: only use this for dilution) C can also be represented as M
Pressure
- The force exerted on a unit area of a surface by the particles of the gas as they collide on a surface
- Altitude impacts oxygen levels
- Total pressure: combined pressure of all components of gases in the atmosphere
- Partial pressure: pressure of a specific substance such as oxygen
- One newton per square metre is equivalent to a pressure of one pascal (Pa)
- Pressure formula = force/area P=F/A
- Total pressure formula: Ptotal= P 1 +P 2 +P 2 …
- Pressure and volume formula: V ∝ 1/P (Pressure and volume are inversely proportional)
- Pressure changes with change in temperature - P 1 V 1 =P 2 V 2
Boyle’s Law
- (Temperature is constant)
- Volume and pressure are inversely proportional at a constant temperature
- V∝1/P at constant T
- For example a weather balloon with helium gas at a volume of 40L at a pressure of 1 atm, will increase to 200L at a pressure of 0.2 atm.
- PV=PV
Charles’s Law
- (Pressure is constant)
- The volume of an ideal gas at a constant pressure is proportional to the temperature
- As the temperature increases the volume increases, the density decreases
- There is an energy conversion of heat to kinetic energy, this has the particles vibrate more therefore increasing the volume as the particles need more space to move around - V 1 /T 1 =V 2 /T 2
- V∝T at constant P
Gay Lussac’s Law
- (Volume is constant)
- The pressure changes with change in the temperature
- The higher the temperature, the higher the pressure-
- P 1 /T 1 =P 2 /T 2
- P∝T at constant V
Combined Gas Law!
- P 1 V 1 /T 1 = P 2 V 2 /T 2
- PV∝T or PV/T = constant
