Chemistry Exam Revision Notes

Chemical Properties

How a substance changes its chemical identity.
Examples:
- Flammability: How quickly a material catches fire.
- Toxicity: A substance's ability to harm organisms

Physical Properties

Characteristics observed without changing chemical identity.
Examples:
- Melting and Boiling Point: Temperature where phase changes occur.
- Endothermic Processes: Require heat (melting, boiling).
- Exothermic Processes: Release heat (freezing, condensing).
- Solubility: Mass of solute in a solvent at a given temperature.
- Thermal Conductivity: Rate of heat transfer.
- Electrical Conductivity: Rate of charge flow; depends on free-moving charged particles.
- Density: Mass per unit volume, e.g., g/cm\textsuperscript{3}.

Separating Mixtures

Pure substances have consistent chemical composition.
Mixtures combine substances physically, not chemically.
- Homogenous: Uniform composition (e.g., seawater).
- Heterogeneous: Non-uniform composition.
- Solution: Solute dissolved in solvent.
Techniques:
- Filtration: Separates insoluble material from solution.
- Evaporation: Separates solid solute from liquid solvent.
- Distillation: Separates liquids with different boiling points using heat transfer and condensation.
- Sieving: Separate solids by size.
- Magnetism: Separate solids using magnet.
- Decanting: separate using density.

Nanomaterials

Particles range from 1-100 nanometres (nm) in size.
High surface area to volume ratio creates unique properties.
Example: Silver nanoparticles in bandages to kill bacteria.

Atomic Structure

Atoms are composed of subatomic particles:
- Protons: positively charged particle.
- Neutrons: electrically neutral particle.
- Electrons: negatively charged particle.
Mass number: $\text{A = Protons + Neutrons}$
Atomic number: $\text{Z = Protons}$
Isotopes: Atoms of the same element with different mass numbers due to varying neutron numbers.
- Chemical properties of isotopes are the same.
- Example: $\begin{array}{l}6^{12} \mathrm{C}\end{array}$ , $\begin{array}{l}6^{13} \mathrm{C}\end{array}$ , $\begin{array}{l}_6^{14} \mathrm{C}\end{array}$

Electron Arrangement

Electrons are arranged in shells at different energy levels.
Shell closest to the nucleus has the lowest energy.
Each shell has a principal quantum number and a maximum number of electrons.
Subshells contain orbitals, each holding up to 2 electrons.
Filling subshells:
- Exceptions: Cr and Cu due to lower energy when filling d orbitals (partially or fully).
- Cr: $\text{1s}^2\text{ 2s}^2\text{ 2p}^6\text{ 3s}^2\text{ 3p}^6\text{ 4s}^1\text{ 3d}^5$
- Cu: $\text{1s}^2\text{ 2s}^2\text{ 2p}^6\text{ 3s}^2\text{ 3p}^6\text{ 4s}^1\text{ 3d}^{10}$
Ions: Metals become cations by losing electrons, non-metals become anions by gaining electrons.

Emission and Absorption Spectra

Electrons transition between energy levels by absorbing or emitting photons.
Atomic absorption: Atom absorbs energy and excites an electron.
- Line absorption spectrum: Continuous spectrum with dark absorption lines.
Atomic emission: Atom emits energy as a photon when an electron transitions to a lower energy level.
- Line emission spectrum: Shows lines for specific wavelengths of emitted photons.

Periodic Table

Elements are arranged by increasing atomic number in periods (rows) and groups (columns).
Groups have similar chemical properties.
Groups 1, 2, and 13-18 are main group elements (I-VIII).
Metallic character: Elements with 1, 2, or 3 valence electrons tend to lose electrons, forming cations.
- Group 1: Alkali metals.
- Group 2: Alkaline earth metals.
- Transition metals.
- Group 13: Metals (except boron).
Metalloids: Exhibit both metal and non-metal characteristics (B, Si, Ge, As, Sb, Te, Po).
Non-metals: Elements with 4, 5, 6, 7, or 8 valence electrons, prone to gaining electrons and forming anions.
- Group 7: Halogens.
- Group 18: Noble gases (full outer shell, generally non-reactive).

Chemical Reactivity

*   Determined by number of valence electrons and how readily atoms accept or donate them.

Metals lose electrons (low ionization energy); non-metals gain electrons.
Ionization energy: Energy required to remove an electron; increases with electronegativity.

Atomic Radius

Distances between nucleus and valence shell boundary.
Core charge affects atomic radius.
Increases down a group, decreases from left to right across a period.

Electronegativity

Tendency of an atom to attract a shared pair of electrons.
Non-metals are more electronegative than metals.
Increases from left to right and decreases down a group.

Types of Materials Bonding

Intramolecular: Within a molecule (primary bonding).
Intermolecular: Between molecules (secondary bonding/interactions).

Metallic Bonding

Attractive force between delocalized electron cloud and metal nuclei.
Electrically conductive due to free-moving electrons.

Ionic Bonding

Metal donates electrons to non-metal, forming cations and anions.
Attractive force between cations (+) and anions (-).
Forms a crystalline lattice; solids are not electrically conductive unless molten or dissolved.

Covalent Bonding

Non-metals share electrons.
Can form crystalline lattices or discrete molecular structures.
Most compounds are not electrically conductive (exceptions: graphite, graphene).

Ionic Bonding

Attraction between positively charged cations and negatively charged anions.
Ions arranged in a regular repeating pattern throughout the crystal.

Cations

Metals lose electrons to form cations.
Some atoms can lose multiple valence electrons.

Anions

Non-metals gain electrons to form anions.
Addition of an electron requires energy to overcome the repulsive forces of the electrons already present in the atom.

Monoatomic Ions

Ions composed of one atom

Polyatomic ions

Ions composed of more than one atom

Ionic Formulae

Total Negative charge = total positive charge
Chemical (Empirical) Formulae (In an ionic compound the total negative charge is equal to the total positive charge)
- Swap the number from the charge of the metal ion to the subscript for the non-metal ion and the number from the charge of the non-metal ion to the subscript for the metal ion

Characteristics of ionic compounds

Form crystal structures (crystalline lattice)
High melting/ boiling points
Hard and brittle
Only conduct electricity once dissolved in water.

Metallic Bonding

Metals generally have high thermal conductivity.
Metals generally have high electrical conductivity

Solution of Ionic Substances

For a solute to dissolve in a solvent and form a solution, the particles (molecules, atoms, ions) in both components must be able to separate and mix together into a homogenous mixture.

Precipitation Reactions

When solutions containing these ions are combined, then the insoluble ionic compounds are produced in a precipitation reaction.
The reaction that occurs between two ionic compounds is called a double displacement reaction.
Net ionic equations
Complete ionic equation: precipitation is expanded to show all the dissociated ions

Molecular Polarity

Valence Shell Electron Pair Repulsion (VSEPR) Theory (Shapes of molecules are determined by the spatial arrangement of electron pairs surrounding the central atom)
Pairs of electrons repel each other. Lone pairs of electrons have a similar repulsion to bonding electrons.
Linear: has two regions of electron density around a central atom.

Interactions between Molecules

Secondary Bonding
*Dispersion forces (The weakest forces of attraction which exist between most atoms and molecules)

Hydrocarbons

Organic compounds: Compounds which are found within organisms. They are generally composed of large carbon chains.
Hydrocarbons: organic molecules composed of only carbon and hydrogen.
- $\text{General formula: C}{n}\text{H}{2n+2}$
Trend in melting point and boiling point for alkanes: As the number of carbon atoms in an alkane increases, both the melting point and boiling point increase because larger molecules have stronger London dispersion forces, which require more energy to overcome.

Alkenes

(Homologous series of hydrocarbons (contains elements C and H), with the functional group being a carbon-carbon double bond
$\text{General formula: C}{n}\text{H}{2n}$
Vegetable oils tend to be liquid at room temperature. Animal fasts tend to be solid at room temperature.
Alkenes can undergo addition reactions.

Combustion

Hydrocarbons release large amounts of energy when they combust
Complete combustion: suGicient oxygen molecules present to react with the hydrocarbon and produce carbon dioxide and water.

Polymer

Large molecules constructed from smaller molecules called monomers which are linked via covalent bonds
Thermoplastic Polymers (Polymers where there is secondary bonding between chains)
Thermoset Polymers (Polymers that have crosslinks between chains. These crosslinks are made by primary bonds such as covalent bonds)