Chemistry

Content Areas:
- Atomic Theory
- The Periodic Table
- Structure and Bonding
- States of Matter
Resources: Knowledge Organisers for the topics attached.

States of Matter

Three Main States:
- Solid
- Liquid
- Gas
Energy Transfer: For a substance to change states, energy must be transferred to the particles, which gains energy leading to the breaking of attractive forces.

Solid

Arrangement: Particles are arranged in a regular pattern.
Movement: Particles vibrate in fixed positions and are tightly packed together.
Kinetic Energy: Solids exhibit a low amount of kinetic energy.
Characteristics:
- Fixed shape and cannot flow like liquids.
- Cannot be compressed since particles are closely packed together.

Liquid

Arrangement: The particles are randomly arranged.
Movement: Particles can move around each other.
Kinetic Energy: Liquids have greater kinetic energy than solids.
Characteristics:
- Can flow and take the shape of the container.
- Cannot be compressed as particles are still close together.

Gas

Arrangement: The particles are randomly arranged and far apart.
Movement: They move very quickly in all directions.
Kinetic Energy: Gas particles have the highest amount of kinetic energy among the three states.
Characteristics:
- Can flow and fill the entire container.
- Can be compressed due to ample space between particles.

Changes in State

Melting: occurs when solid changes to liquid.
Boiling/Evaporation: More energy is needed to overcome the remaining chemical bonds in liquid.
Evaporation: Particles leave the surface of a liquid.
Boiling: Bubbles of gas form throughout the liquid before rising to the surface.

State Determination

Solid: If temperature < melting point.
Liquid: If temperature is between melting point and boiling point.
Gas: If temperature > boiling point.

State Symbols in Chemical Equations

Solid: (s)
Liquid: (l)
Gas: (g)
Aqueous solution: (aq)

Ionic Bonding and Formation of Ions

Definition of Ion: Charged particle, can be positively (cations, e.g., Na^+) or negatively charged (anions, e.g., Cl^-).
Metallic Bonding: Occurs between metals.
- Positive metal ions surrounded by a sea of delocalised electrons.
- Arrangement: Ions are tightly packed and arranged in rows.

Electron Transfer and Ion Formation

Metals: Lose electrons → positively charged ions.
Non-metals: Gain electrons → negatively charged ions.
Electron Configuration:
- Group 1 & 2 elements: lose electrons to form +1 and +2 ions respectively.
- Group 6 & 7 elements: gain electrons to form -2 and -1 ions respectively.

Example of Elements and Their Ions

Element	Group	Ion Formation
Li	1	Li^+
Ca	2	Ca^2+
O	6	O^2-
Br	7	Br^-

Properties and Bonding of Metals & Non-metals

Metals:
- Found on the left-hand side of the periodic table.
- Characteristics: Strong, shiny, malleable, good conductors of heat/electricity.
Non-metals:
- Found on the right-hand side of the periodic table.
- Characteristics: Brittle, dull; not always solids at room temperature; poor conductors.

Alloy Creation

Definition of Alloys: Mixtures of metals for enhanced properties.
Function of Alloys: Different-sized atoms prevent layers from sliding over one another, making them harder than pure metals.

Ionic Compounds

Definition: Form structures called giant lattices held together by strong electrostatic forces of attraction.
Properties of Ionic Compounds:
- High melting and boiling points due to strong attraction.
- Cannot conduct electricity in solid state (ions not mobile).
- Can conduct electricity when molten or in solution (ions can move and carry electrical current).

Covalent Bonding

Definition: Sharing of a pair of electrons between atoms (non-metals).
Characteristics of Simple Covalent Structures:
- Low melting and boiling points due to weak intermolecular forces.
- Do not conduct electricity (no free delocalised electrons).

Dot and Cross Diagrams

Represent the outer electron shell as circles with overlapping circles to indicate covalent bonding; dots or crosses represent electrons.
Molecules to Represent: Chlorine, Oxygen, Nitrogen, Water, Ammonia, Hydrogen Chloride, Methane.

Giant Covalent Structures

Diamond

Structure: Each carbon atom bonded to four others.
Properties: Very strong, high melting/boiling points, does not conduct electricity.

Graphite

Structure: Layers of carbon in hexagons with delocalised electrons.
Properties: High melting point; conducts electricity; layers can slide over each other due to weak forces.

Graphene

Structure: A single layer of graphite.
Properties: Strong due to covalent bonds, conducts electricity, enhances properties of added materials.

Nanoscience

Definition: Structures of size 1-100 nm.
Properties of Nanoparticles: High surface area to volume ratio; effective in fewer amounts.

Types of Particles

Name of Particle	Diameter
Nanoparticle	1-100 nm
Fine Particles (PM2.5)	100-2500 nm
Coarse Particles (PM10)	2500-10000 nm

Polymers

Definition: Long chain molecules from monomers, held by strong covalent bonds.
Intermolecular Forces: Attract polymer chains; longer chains possess stronger attractions.

Fullerenes and Nanotubes

Fullerenes

Molecules shaped like hollow tubes or balls, useful in drug delivery, e.g., Buckminsterfullerene (C60).

Carbon Nanotubes

Tiny cylinders conducting electricity, reinforcing materials with negligible weight.
Applications in electronics and nanotechnology.

Risks of Nanoparticles

Risk of inhalation can lead to potential harmful reactions and toxicity due to the large surface area.
Applications include medicine, cosmetics, and catalysts.

Atomic Structure and the Periodic Table

Atoms

Composition: Nucleus contains protons and neutrons, with electrons surrounding.
Structure:
- Proton: Mass = 1, Charge = +1
- Neutron: Mass = 1, Charge = 0
- Electron: Mass = very small, Charge = -1
Ion Formation: Charged particles do not have equal protons and electrons.

Atomic Number and Mass Number

Definition: Number of protons represents atomic number, while mass number is the total of protons and neutrons.
Isotopes: Same atomic number but different mass number due to different neutrons.

Compounds

Definition: Chemical combinations of two or more elements. Difficulty in separation.
Examples of Formulas: CO₂, NaCl, HCl, H₂O.

Chemical Equations

Formulation of Equations: Reactants on left, products on the right; must be balanced.
Methods of Separation:
- Evaporation: To obtain soluble salts from solutions.
- Filtration: To separate solid from liquids.
- Distillation: To separate liquid mixtures.

Development of the Periodic Table

Historical Context

Early 1800s: Based on atomic mass but incomplete.
Dmitri Mendeleev (1869): Arranged elements in order of mass and left gaps for undiscovered elements.

Modern Periodic Table

Ordered by atomic mass/proton number, metal/non-metal segregation.
Group number indicates the number of valence electrons; periods indicate electron shell completion.

Group Characteristics

Alkali Metals (Group 1)

Characteristics: Soft, very reactive, low density, with one outer electron. Reactivity increases down the group.
Reactions:
- With water: forms hydroxide + hydrogen gas.
- With chlorine: forms metal chloride.

Halogens (Group 7)

Characteristics: Non-metals that become less reactive down the group and have increasing melting/boiling points.

Noble Gases (Group 0)

Characteristics: Non-reactive due to full outer shells; colorless gases.
Boiling Point Trend: Increases down the group due to higher intermolecular forces.

Transition Metals

Block of elements with properties of metals and colored ions. Used as catalysts, e.g., iron in Haber process.

Chemistry Year 10 PPE Schedule

Date of Exam: Tuesday, 27th January
Content Areas: Atomic Theory, The Periodic Table, Structure and Bonding, States of Matter.

States of Matter

Three Main States:
- Solid: Regular pattern, vibrating in fixed positions, low kinetic energy. Fixed shape and incompressible.
- Liquid: Random arrangement, particles can slide over each other. Flow to take shape of container.
- Gas: Far apart, fast movement in all directions, highest kinetic energy. Compressible.
Changes in State:
- Melting: Solid to liquid.
- Boiling/Evaporation: Liquid to gas. Energy is transferred to break attractive forces.
- State Symbols: $(s)$ Solid, $(l)$ Liquid, $(g)$ Gas, $(aq)$ Aqueous solution.

Atomic Structure and The Periodic Table

Atoms: Nucleus (protons and neutrons) surrounded by electrons.
- Proton: Mass $1$ , Charge $+1$ .
- Neutron: Mass $1$ , Charge $0$ .
- Electron: Mass negligible, Charge $-1$ .
Atomic Number: Number of protons. Mass Number: Protons + Neutrons.
Isotopes: Atoms of the same element with different numbers of neutrons.
Development: Mendeleev (1869) arranged by mass and left gaps for undiscovered elements. The modern table is ordered by atomic (proton) number.
Groups:
- Group 1 (Alkali Metals): Highly reactive, reactivity increases down the group.
- Group 7 (Halogens): Non-metals, reactivity decreases down the group.
- Group 0 (Noble Gases): Unreactive, full outer shells.

Ionic Bonding

Definition: The electrostatic attraction between oppositely charged ions. Occurs between metals and non-metals.
Formation:
- Metals: Lose electrons to form positive ions (cations), e.g., $Li^+$ , $Mg^{2+}$ .
- Non-metals: Gain electrons to form negative ions (anions), e.g., $Cl^-$ , $O^{2-}$ .
Ionic Compounds:
- Form giant ionic lattices with strong forces in all directions.
- High melting/boiling points due to strong electrostatic attractions.
- Conductivity: Do not conduct when solid (ions fixed); conduct when molten or aqueous as ions are free to move and carry charge.

Covalent Bonding

Definition: The sharing of pairs of electrons between non-metal atoms.
Simple Molecular Structures:
- Examples: $H<em>2O$ , $CH</em>4$ , $NH<em>3$ , $O</em>2$ .
- Properties: Low melting/boiling points due to weak intermolecular forces; do not conduct electricity (no free electrons or ions).
Giant Covalent Structures:
- Diamond: Each Carbon bonded to 4 others; very hard, high melting point.
- Graphite: Each Carbon bonded to 3 others in layers; has delocalised electrons so it conducts electricity; layers slide (soft).
- Graphene: Single layer of graphite; strong and conductive.
- Fullerenes & Nanotubes: Hollow shapes (e.g., $C_{60}$ ) used for drug delivery and lubricants; nanotubes have high tensile strength.
Polymers: Long chains of monomers held by strong covalent bonds with relatively strong intermolecular forces.

Metallic Bonding

Definition: The attraction between positive metal ions and a "sea" of delocalised electrons.
Structure: Giant lattice of atoms arranged in regular rows.
Properties:
- Conductivity: Delocalised electrons move through the structure to carry charge and heat.
- Malleability: Layers of atoms can slide over each other without the bond breaking.
Alloys: Mixtures of metals where different-sized atoms distort the layers, making it harder for them to slide and thereby making the material stronger than pure metal.

Nanoscience

Scale: Particles sized between $1\text{ nm}$ and U $100\text{ nm}$ .
Properties: High surface area to volume ratio, meaning smaller quantities are needed for effectiveness (e.g., in catalysts or sunscreens).
Risks: Potential toxicity if inhaled or absorbed into the skin due to high reactivity.