Chemistry: Atomic Structure, Rutherford Model

An atom consists of three sub-atomic particles: protons, neutrons, and electrons.
Protons are positively charged with a mass of 1, found inside the nucleus.
Neutrons have no charge with a mass of 1, found inside the nucleus.
Electrons are negatively charged with almost zero mass, orbiting the nucleus.
An atom is neutral, so the number of electrons equals the number of protons.
Atomic Number = Number of protons (also the number of electrons).
Mass Number = Number of protons + Number of neutrons.

Ions are charged particles formed from atoms that donate or accept electrons.
Metal atoms lose electrons to become positive ions.
Non-metal atoms gain electrons to become negative ions.
Ions have an electron arrangement like a Noble Gas, with an outer energy level filled with electrons.

Metal elements have strong metallic bonding between positive ions and delocalized electrons.
The greater the number of delocalized electrons, the greater the charge on the metal ion, the stronger the metallic bond.
Metals have properties such as:
- Conducting electricity due to delocalized electrons.
- High melting and boiling points.
- Range of densities.
- Strength.
- Electronegativity.

Electronegativity is the tendency of an atom to attract electrons in a covalent bond.
Non-metal elements have higher electronegativity values than metal atoms.
Electronegativity is measured using the Pauling scale.

Compounds and molecules are formed when atoms interact and form bonds.
Electronegativity values affect where electrons lie within the bond.
If atoms are equally electronegative, electrons are shared equally between the two atoms.

If atoms have slight differences in electronegativity, electrons are shared unequally between the two atoms.
The atom with the higher electronegativity value attracts electrons more, making the bond polarized.

If atoms in a bond have large differences in electronegativity, electrons are shared unequally between the two atoms.
Electrons are attracted to one side of the bond more than the other, forming an ionic bond and charged particles or ions.

Types of intermolecular bonding include:
- London Dispersion.
- Permanent dipole-Permanent dipole.
- Hydrogen bonding.

Trends in melting and boiling points, atomic radius, and electronegativity downwards a period and across a group.
Reasons for these trends are explained.

Enthalpy is the energy released or taken in when 1 mole of a substance reacts.
Enthalpy (∆H) = Energy reactants - Energy products.
Examples of exothermic and endothermic reactions.
Calculating Enthalpy by Experiment: ∆H = c x mw x ∆T.
Hess's Law: "The enthalpy change for a chemical reaction is independent of the route taken."

Factors affecting the rate of a reaction:
- Temperature of reactants.
- Concentration of reactants.
- Particle size (Surface area) of reactants.
Measuring the rate of a reaction:
- Measuring a change in mass over time.
- Measuring a change in concentration over time.
- Measuring a change in gas volume over time.

The equilibrium constant (Kc) = [C]^c [D]^d / [A]^a [B]^b.
Value of K indicates the extent of the reaction:
- Less than 1: More reactants than products (Effectively no reaction).
- 1: Similar levels of reactants and products.
- Greater than 1: More products than reactants (Effectively complete).

A Bronsted-Lowry acid is a substance that can donate a proton to form a conjugate base.
A Bronsted-Lowry base is a substance that can accept a proton to form a conjugate acid.
Water can act as both an acid and a base when it reacts with actual acids and bases, making it amphoteric.

Transition metal complexes involve ligands that are electron donors.
Coordination number and oxidation number are important in transition metal complexes.
Writing and naming complexes is explained.

Types of formulae: Molecular and Structural.
Naming and drawing alkanes, alkenes, and alcohols.
Electrophiles and nucleophiles are involved in reactions like halogenation (electrophilic addition).