Atoms, Elements and Compounds
---
### 2.1 Elements, Compounds and Mixtures
* Element = A pure substance made of only one type of atom. Cannot be broken down into simpler substances by chemical means.
* Example: Oxygen (O₂), Iron (Fe), Neon (Ne).
* Compound = A pure substance formed when two or more different elements are chemically combined in a fixed ratio. It has properties different from its elements.
* Example: Water (H₂O), Sodium Chloride (NaCl).
* Mixture = A combination of two or more substances (elements or compounds) that are not chemically combined. The composition can vary.
* Example: Air, Saltwater, Sand and Iron Filings.
* Can be separated by physical methods (filtration, distillation, etc.).
Particle Diagram Summary:
```
Element: [O] [O] [O] (Identical atoms)
Compound: [H-O-H] [H-O-H] (Different atoms bonded)
Mixture: [O] [Na] [Cl] (Different particles, not bonded)
```
---
### 2.2 Atomic Structure and the Periodic Table
* Structure of the atom:
A small, central *nucleus** containing protons and neutrons.
* Electrons orbit the nucleus in shells (energy levels).
* Subatomic Particles:
Proton: Relative Charge = *+1**, Relative Mass = 1
Neutron: Relative Charge = *0**, Relative Mass = 1
Electron: Relative Charge = *-1**, Relative Mass = ~0 (1/1840)
* Definitions:
* Proton (Atomic) Number (Z) = number of protons in the nucleus.
* Mass (Nucleon) Number (A) = total number of protons and neutrons in the nucleus.
* Number of neutrons = A - Z
* Standard Notation:
```
Mass Number (A) → 23
Na
Atomic Number (Z) → 11
```
* Electronic Configuration (for elements Z=1-20):
Shells fill in order: 1st shell holds max *2** electrons, 2nd and 3rd hold max 8.
* Example: Sulfur (Z=16) = 2.8.6
* Links to the Periodic Table:
* Group Number = Number of outer shell electrons (for Groups I-VII).
* Period Number = Number of occupied electron shells.
* Group VIII (Noble Gases) = Have a full outer shell → very unreactive.
ASCII Diagram: Bohr Model of Sodium (Na, Z=11)
```
Nucleus (11p+, 12n⁰)
/ \
(e⁻) / \ (e⁻)
Shell 1: [2 e⁻] → Full
Shell 2: [8 e⁻] → Full
Shell 3: [1 e⁻] → 1 outer electron → Group I
```
Config: 2.8.1
---
### 2.3 Isotopes
* Definition: Atoms of the same element (same number of protons) with different numbers of neutrons.
* Example: Chlorine-35 (17p, 18n) and Chlorine-37 (17p, 20n).
* Chemical Properties: Identical because they have the same electronic configuration.
* Physical Properties: Different (e.g., density, rate of diffusion).
* Calculating Relative Atomic Mass (Aᵣ):
Aᵣ is the *weighted mean mass** of all the isotopes of an element.
* Formula:
```
Aᵣ = ( (% iso1 × mass iso1) + (% iso2 × mass iso2) ) / 100
```
* Worked Example:
* Boron has two isotopes: B-10 (20%) and B-11 (80%).
```
Aᵣ = ( (20 × 10) + (80 × 11) ) / 100
= (200 + 880) / 100
= 1080 / 100 = 10.8
```
---
### 2.4 Ions and Ionic Bonds
* Ions: Charged particles formed when atoms lose or gain electrons.
* Cation (+): Positive ion formed when a metal atom loses electrons.
* Example: Na → Na⁺ + e⁻
* Anion (-): Negative ion formed when a non-metal atom gains electrons.
* Example: Cl + e⁻ → Cl⁻
* Ionic Bond: The strong electrostatic attraction between oppositely charged ions (cations and anions).
* Formation (Dot-and-Cross Diagram for NaCl):
* Sodium (2.8.1) loses 1 electron to become Na⁺ (2.8).
* Chlorine (2.8.7) gains 1 electron to become Cl⁻ (2.8.8).
```
Na⁺ [2.8]⁺ and Cl⁻ [2.8.8]⁻
```
* Use dots (•) for one atom's electrons and crosses (x) for the other.
* Giant Ionic Lattice: Ionic compounds form a regular, repeating 3D structure of alternating positive and negative ions.
```
Na⁺ Cl⁻ Na⁺
Cl⁻ Na⁺ Cl⁻
Na⁺ Cl⁻ Na⁺ ← Strong ionic bonds in all directions
```
* Properties:
* High melting/boiling points: Requires a lot of energy to break the strong electrostatic forces in the giant lattice.
* Conduct electricity when molten or dissolved in water (aqueous): The ions are free to move and carry charge.
* Do NOT conduct electricity when solid: The ions are fixed in position in the lattice.
* Brittle: Layers shift, causing like charges to repel.
---
### 2.5 Simple Molecules and Covalent Bonds
* Covalent Bond: A shared pair of electrons between two non-metal atoms, allowing both to achieve a noble gas configuration.
* Dot-and-Cross Diagrams:
* Hydrogen (H₂): Single bond.
```
H• + •H → H••H or H-H
```
* Water (H₂O): Two single bonds. Oxygen has two lone pairs.
```
H
•
O•••
•
H
```
* Oxygen (O₂): Double bond.
```
O••••O or O=O
```
* Methane (CH₄): Four single bonds.
* Properties of Simple Covalent Molecules:
* Low melting/boiling points: Weak intermolecular forces between molecules require little energy to overcome. (The covalent bonds within each molecule are strong).
* Do NOT conduct electricity: No free ions or delocalised electrons.
* Often gaseous or liquid at room temperature (e.g., CO₂, H₂O).
---
### 2.6 Giant Covalent Structures
* Definition: Huge networks of atoms all bonded together by strong covalent bonds. Very high melting points.
* Diamond:
* Structure: Each carbon atom is covalently bonded to four other carbon atoms in a rigid tetrahedral structure.
```
C
/|\
C C C
\|/
C
```
* Properties:
* Very hard (strong bonds, rigid structure).
* Very high m.p.
* Does not conduct electricity (no delocalised electrons).
* Use: Cutting tools, drill bits.
* Graphite:
* Structure: Each carbon atom is bonded to three others, forming layers of hexagonal rings. Weak intermolecular forces hold the layers together. One delocalised electron per carbon atom.
```
Layer: /-\-/-\ ← Strong covalent bonds
\-/-\-/ ← Delocalised e⁻ move here
```
```
Layer 2: (weak forces hold layers together)
Layer 1: /-\-/-\
```
* Properties:
* Soft & Slippery: Layers can slide over each other.
* High m.p. (strong covalent bonds in layers).
* Conducts electricity and heat: Delocalised electrons can move along the layers.
* Uses: Lubricant, electrodes (in batteries and electrolysis).
* Silicon(IV) Oxide (Silica, SiO₂):
* Structure: Similar to diamond. Each silicon atom is bonded to four oxygen atoms. Each oxygen atom is bonded to two silicon atoms.
* Properties: Hard, high m.p., poor electrical conductor.
* Found in: Sand, quartz.
---
### 2.7 Metallic Bonding
* Definition: The electrostatic attraction between positive metal ions arranged in a giant lattice and a 'sea' of delocalised electrons.
* The delocalised electrons come from the outer shells of the metal atoms.
* Structure:
```
+ + + + + ← Positive metal ions
\ | / \ | /
~~~~~e⁻ sea~~~~~ ← Delocalised electrons
/ | \ / | \
+ + + + +
```
* Properties Explained:
* Good electrical & thermal conductivity: The delocalised electrons are free to move and carry charge/energy through the structure.
* Malleable (can be hammered into shape) & Ductile (can be drawn into wires): Layers of positive ions can slide over each other without breaking the structure because the sea of electrons holds everything together.
* High melting and boiling points: The strong electrostatic attraction between the ions and electrons requires a lot of energy to overcome.