Science Semester 1 2025

Balancing Chemical Reactions

• Reactants are the initial substances in a chemical reaction and are always on the left side of the equation.
• Products are the substances formed during a reaction and are always on the right side of the equation, using an arrow to show the change.
• Example word equation: Magnesium + Oxygen → Magnesium oxide.
• A balanced chemical equation has identical numbers for each element on both sides of the equation.
• Atoms cannot be created or destroyed in a reaction, so the atoms present at the start must also be present at the end.
• Balancing an equation does not involve changing a compound or element.

Rules for Balancing Equations

• Write the correct chemical formulas for all reactants and products, linked by "+" and "→".
• Count the atoms of each element on both sides of the equation.
• Use coefficients (numbers in front of compounds) to balance elements one at a time, saving hydrogen (H) and oxygen (O) for last.
• Confirm balance by recounting atoms on each side. Subscripts Must Not Be Changed
• Only coefficients can be changed.
• Changing a subscript alters the chemical formula, indicating a different chemical.
• Coefficients must go in front of chemical formulas, not in the middle.

Balancing Chemical Reactions: Step-by-Step

• List the elements in the compounds for both sides of the reaction (reactants and products)
• Count the atoms of each element on both sides.
• Identify imbalances and decide the order to balance them.
• Add coefficients in front of compounds to balance elements, bearing in mind that one element if by itself should be balanced last.
• Readjust counts and repeat until all elements are balanced.

The Crossover Rule for Ionic Formulas

• Helps to write the formulae
• Identify the metal and nonmetal.
• Write their element symbols.
• Write the charges of each.
• Crisscross the charges so they become subscripts.
• Remove the plus and minus signs from the charges.
• Simplify the numbers and remove the "1s."

Chemical Reactions

• These involve the breaking of bonds in reactant species and the forming of new bonds in product species.

Common Chemical Reaction Types

• Synthesis: A + B → AB
• Decomposition: AB → A + B
• Single Displacement: A + BC → AC + B
• Double Displacement: AB + CD → AD + CB
• Combustion: A + O2 → AO + water
• Neutralization: acid + base (oxide) → salt + water
• Acid + metal: acid + metal → salt + hydrogen
• Acid + metal carbonate: acid + metal carbonate → salt + water + carbon dioxide
• Carbon dioxide + base: carbon dioxide + base → carbonate + water

Synthesis Reactions

• Synthesis reaction combines all reactants to form a single product.
• A synthesis reaction is identified by a single product containing all the reactant atoms.
• Example: 2H2(g) + O2(g) → 2H20(g)
• Steps to predict products: 1. Write word equation, 2. Write chemical symbols, and 3. Balance the equation are.

Synthesis Reactions - Examples

• Synthesis of ammonia
• Commercial production of slaked lime (calcium hydroxide)
• Production of sodium chloride (common salt)
• Preparation of hydrochloric acid and ammonium chloride
• Rusting
• Photosynthesis

Atomic Interactions - Ionic Bonding

• Bonding is an interaction between elements on the periodic table and there are three types: Ionic, Covalent, Metallic.
• Atoms bond to achieve stability.

Ionic Bonding

• It involves electron transfer from on element to another, resulting in charged species or ions.
• Metals lose electrons to become positive cations.
• Nonmetals gain electrons to become negative anions.
• Only valence (outer) electrons take part in bonding.
• Stability is reached when atoms have full outer shells.
• Unfilled outer shells cause atoms to gain/lose electrons and becomes stable.
• An ion's charge is known as its valency.
• An example of an ionic bond is sodium chloride (NaCl) formation from sodium metal and chlorine gas.

Properties of Ionic Substances

• A metal bonded coupled with a nonmetal through electron transfer, and formation of respective ions.
• Properties include: Solid crystals at room temperature, High melting and boiling points, Electrolytic conductivity when mixed with water or melted, Solid & Rigid lattice structures, Hard and brittle, Good insulators as a solid.

Ionic Compound Structure

• The relative sizes of the cations and anions dictates the structure of an ionic compound.
• Ionic compounds include: salts, oxides, hydroxides, sulfides, and the majority of inorganic compounds.
• Electrostatic attration keep ionic solids together between positive and negative ions.
• Large ammounts of energy are needed to separate giant structures with the tightly held ions, resulting in high mpelting points and boiling points.

Ionic Compound Properties

• They possess physical force of attraction between the possitive and negativee ions making compounds very difficult to break/ solids.
• When pressure is applied they tend to come apart from each other hense they are considered "brittle".

Melting and Boiling Points of Ionic Compounds

• High melting and boiling points are attributed to electrostatic forces of attraction between ions.
• Overcoming the strong ionic bonds between the atoms requires a large ammount of energy

The Solubility of Ionic Compounds

• Soluble in polar solvents like water.
• Solubility decreases in nonpolar solvents like petrol and gasoline.

Conduction of Electricity

• Won't conduct electricity in a solid state but will in a molten state & good conductors.
• Electricity involves the movement of charge from one point to another.
• Ionic compounds don't conduct electricity in the solid state due to the immobility of ions.
• Free ions in themolten state facilitate the transfer of electricity as elecrostatic forces of attraction has been released.

Decomposition Reactions

• A compound breaks down into two or more substances.
• This is the opposite of a combination reaction.
• Breakdown of carbonic acid (H2CO3) into carbon dioxide (CO2) and water (H2O) is an example.
• Example Formula: AB → A + B
• Example Decomposition Reactions
• AgCl (s) → Ag (s) + Cl2 (g)
• H2O2 (1) → H2O (1) + 02 (g)
• 2 Fe(OH)3 → Fe2O3 + 3 H₂O
• 2 H₂O →2 H₂ + O₂
• 2 AgBr→ 2 Ag + Br2

Decomposition Reactions - Additional Information

• Requires energy in the form of heat, light, or electricity, making them endothermic.
• Absorption of energy breaks the bonds present in the reaction substances.

Decomposition Reactions - Examples

• Production of calcium oxide (quicklime)
• Production of lithium oxide
• Preparation of oxygen and carbon dioxide
• Mining (extraction of metals)
• When opening fizzy drinks breaks down carbonic acid into both water & dioxide - "fizz"
• Digestion of food via carbohydrates, fats, and proteins.

Single Displacement

• An element replaces another in a compound, resulting in a new element and a compound.
• Requires a metal (solid) and an ionic compound (solution).
• General Form: AB(aq) + C → A + CB(aq)

Single Displacement - Rules

• General rules for writing and prediction single displacement reactions
• Write a word eq. for all species involved
• Write down the chemical symbol for the metal the ionic formula for the solution
• Balance the chemical equation.
• Pure elements, plus an aqueous compound are required as starting point.
• After a reaction, we seee a new aqueous compound and a different pure element are generated .

Single Displacement - More Information

• In predicting whether a single – displacement reaction will occur, compare the reactivity of an element using an activity series table.
• A metal can displace any metal lower in the activity series (cations)
• The metal reactivity series should be consulted when deciding if a single displacement reaction will occur.

Acid & Metal Reactions

• When acid & metal reacts: Metal + Acid → Salt + Hydrogen
• Zinc metal reacts with Hydrochloric acid.

Acid & Metal Carbonate Reactions

• When an acid and metal carbonate are in contact, a salt, water, and carbon dioxide gas form.
• Magnesium Carbonate reacts with Hydrochloric acid.

Neutralization Reactions

• When acid and base are in contact, a salt and water form.
• Sodium Hydroxide reacts with Hydrochloric acid.