C3

Physical and chemical changes
- Physical change: substance changes physical form
- Chemical change: new substance is formed
  - ex.
    - magnesium + oxygen → magnesium oxide
    - 2Mg(s) + 02(g) → 2MgO(s)
  - All living things need chemical changes to survive
- Reversible reactions
  - Many chemical reactions can’t be reversed easily, but some can
    - ex. metal hydrates such as copper (Il) sulfate pentahydrate (CuSO4.5H20) are heated
      - water of crystallisation → driven away
      - Anhydrous copper (I) sulfate remains
      - Color change takes place (blue → white)

copper (I) sulfate pentahydrate → anhydrous copper (II) sulfate + water

CuSO4.5H20(s) (blue) → CuSO4(s) (white) + 5H20(g)

- - - - Water added to anhydrous copper (II) sulfate → reverse occurs
      - Turns blue
      - Pentahydrate is produced
      - Extremely exothermic process
Elements
- Any substrate that can’T be broken down further into simpler substances
- Each element: made of one kind of atom
- Ex. aluminium
  - Element
  - Only aluminium atoms
  - Can’t get simpler substance chemically
  - Can make more complicated substances
    - Aluminium oxide
    - Aluminium nitrate
    - Aluminium sulfate
  - 118 elements
    - 25 not in nature → have to be artificially made
      - Curium
      - Unnilpentium
    - 91: occur naturally
    - Elements classified by properties
      - Metals

High density

High melting point

High boiling point

- - - - Non-metals

Low density

Low melting point

Low boiling point

Compounds
- Pure substances
- Formed when: 2 or more elements chemically combine
- Water: pure simple compound
  - Hydrogen + oxygen
- Magnesium reacts violently with oxygen → white powder magnesium oxide
  - Releases energy when new chemical bonds are formed
- Difference between mixtures and compounds
  - Mixture
    - Substances don’t undergo chemical reaction
    - Can separate them if they have different physical properties
    - Has properties of all elements
  - Compound
    - New substance formed
    - During reaction heat is lost: chemical bond are formed (exothermic)
Solutions
- Solute dissolves in solvent → solution
- Dissolving involves: separation of added solute particles when they collide with liquid particles
- More solute → more concentrated solution
  - More particles
- Water: not only solvent
Atomic structure
- Everything made of atoms
- Atoms of any element are identical
- Atoms made of sub-atomic particles
Inside atoms
- 3 sub-atomic particles
- Protons + neutrons → in nucleus
  - Small volume
  - Dense
- Protons = positive
- Neutrons = neutral
- Electrons = negative
  - More quickly in electronic shells/ levels
  - Held in atom by electrostatic force of attraction between themselves + positive charge of protons
  - Incredibly small
- Atoms = neutral
  - no. of protons = no. of electrons
- Proton number + nucleon number
  - Proton no./ atomic number = no. of protons
    - Z
    - Elements don’t have same proton no.
  - Neutrons + protons = similar mass
  - Nucleon number/ mass number: no. of protons + neutrons
    - Electron = tiny mass
    - Mass of atom depends on no. of protons + neutrons
    - A
  - number of neutrons = nucleon number (A) - proton number (Z)
- Isotopes
  - Not all atoms in element are identical
  - Some atoms contain different no. of neutrons → different nucleon number
  - Isotope: atoms of same element that have same proton number, but different neutron number
  - Behave same during chemical reaction
    - Reason: same number of valance electrons
  - Different mass + properties
    - Density
  - 2 types
    - Stable
    - Unstable
      - Extra neutrons
      - Radioactive
      - Radioisotopes
Arrangement of electrons in atoms
- Electrons move around nucleus at distance
  - Move fast
- Not possible to give exact position of electron in energy level
- Can only be in energy level
- Each energy level: can hold certain electron number
  - 2
    - Lowest energy
  - 8
  - 8
- Ex. Oxygen
  - 8 electrons
  - 2,6
- Electronic structure/ configuration: way electrons are distributed
Electronic structure and the periodic table
- No. of valance electrons corresponds to group
- Group 0: exception → 2 or 8 electrons
- Valance electrons: responsible for chemical properties
  - Elements in same group = similar properties
- Metallic character of element in group increases moving down group
  - Electrons → easier to lose
    - Less attraction between because of distance
Ions and ionic bonding
- Ions
  - Electrically charged particle
  - Atom loses electrons → positive charge
  - Gains electron → negative
  - Some atoms: multiple atoms
- Ionic bonding
  - Usually metals + non-metals
  - Electrons transferred: metal → non-metal
  - Atoms → more stable (full outer shell) → become like nearest noble gas
  - Ionisation: element → ion
  - Ex.
    - Na(g) → Na+(g) + e-
    - Cl(g) + e- → Cl-(g)
  - Charges are equal, but opposite
  - Oppositely charged ions attract → bonded by strong electrostatic forces
- Ionic structures
  - X ray diffraction: determines arrangement of ionic structures
  - Ions packed together in regular arrangements (lattice)
    - Oppositely charged ions attract
- Properties of ionic compounds
  - Solid at room temperature
  - High melting point
    - Reason: strong electrostatic forced → hold lattice together
    - A lot of energy needed to separate ions + melt substance
  - Hard
  - Can’t conduct electricity when solid
    - Ions can’t move
  - Mainly dissolve in water
    - Water bonds with - and + ions → breaks lattice structure → keeps ions apart
  - Can conduct electricity when molten or in aqueous solution
    - Force between ions → weakened
    - Ions free to move to electrode
    - Electric current passes through molten compound
Covalent bonding
- Different way atoms gain stability of noble gas electron configuration
- Share electrons in outher energy levels
  - Between non-metal atoms
- Covalent bonding: bond formed
  - H2O
    - Each hydrogen atom in molecule = 1 electron
    - Gain 1 electron -> get full outer energy level + gain electron configuration of noble gas (helium)
    - Allow their outer energy levels to overlap
    - Molecule of hydrogen -> formed
    - 2 hydrogen atoms share electrons
    - Single covalent bond (H-H)
  - Other covalent compounds
    - Methane (CH4):
      - Gas
      - Molecules contain carbon + hydrogen atoms

Carbon: 6C 2,4,

Hydrogen 1H 1

- - - - Carbon needs 4 more electrons to get electron configuration of noble gas neon
      - Hydrogen need 1 electron -> helium
    - Water (H2O):
      - Liquid
      - Hydrogen + oxygen
      - Electronic structure

Oxygen: 8O 2,6

Hydrogen: 1H 1

- - - - Oxygen needs two electrons -> neon
      - Hydrogen needs 1 electron -> helium
      - Oxygen + hydrogen share electrons -> 1 water molecules

2 single covalent bonds

- - - Methanol
      - CH3OH
      - Volatile liquid compound
      - Use: biodiesel
      - Elements: carbon, oxygen, hydrogen
      - Electronic structures

Carbon: 6C 2,4

Oxygen: 8O 2,6

Hydrogen: 1H 1

- - - - Carbon needs to share 4 electrons to gain electron configuration of neon
      - Oxygen needs 2 electrons -> neon
      - Hydrogen needs 1 electron -> helium
      - Single covalent bonding

Three carbon-hydrogen

One carbon-oxygen

One oxygen-hydrogen

- - Properties
    - Simple molecular substances
    - Usually low melting + boiling point
      - Reason: weak intermolecular forces of attraction ehich exist between simple molecules
      - Weaker compared to strong covalant bonds
      - Giant molecular substances = higher melting point

Reason: held together by strong covalant bonds

- - - - In ionic compounds -> interionic forces = much stronger than intermolecular forces in simple covalant substances

Melting point = higher

- - - - Do not conduct electricity when molten or dissolved in water

Don’t have ions

However, some molecules react with water -> form ions

Hydrogen chloride gas produces aqueous hydrogen ions + chloride ions when dissolved in water

HCL(g) -> H+(aq) + Cl-(aq)

Usually do not desolve in water

Water = great solvent

Can interat + disolve some covalant molecules better than others

Usually soluble in organic solvents

- - Allotropy
    - Element that can exist in multiple physical forms in same state
    - Allotrope: each of different physical forms
  - Allotropes of carbons
    - Carbon: non-metallic
    - More than one solid structure
    - Allotropes: graphite + diamond
      - Different physical properties
  - Graphite
    - Hexagons flat → Layer structure
      - Each layer: giant molecule
    - Between layers: weak forces of attraction
      - Layers slide
    - Each layer → carbon atom bonded to three others by strong covalent bonds
      - Usually bonded to 4 atoms
      - Has delocalised electrons (free to move around)
      - Conducts electricity
  - Diamond
    - Tetrahedral arrangement
    - Three dimensional structure → extremely hard
    - No free electrons all are covalently bonded
    - Can’t conduct electricity
  - Silicon (IV) oxide
    - Silicon
      - Group 4
      - Non-metals
      - 2,8,4
    - Crystalline silicon: same structure as diamond
      - Add oxygen → make similar macromolecular structure
    - Not as hard as diamond