Ionic Bonding, Reactions, and Acids & Bases - Vocabulary

Ionic Bonding and Ionic Compounds

• Ionic bonding arises from electrostatic attraction between oppositely charged ions formed when metals lose electrons (cations) and non-metals gain electrons (anions).
• Electron transfer to achieve a stable outer shell leads to positively charged metal ions (cations) and negatively charged non-metal ions (anions).
• If an atom has fewer than 4 electrons in its outer shell, it tends to lose electrons and become positive; if it has more than 4, it tends to gain electrons and become negative.
• A lattice structure is formed as each ion is surrounded by ions of opposite charge, creating a three-dimensional grid of ions.
• Group (column) number equals the number of electrons in the outer shell (valence electrons). Period (row) number equals the number of electron shells.
• Cations are positively charged; anions are negatively charged.
• Common ions (examples):
  • Cations: Li⁺, Na⁺, K⁺, Mg²⁺, Ca²⁺, Al³⁺, etc.
  • Anions: F⁻, Cl⁻, Br⁻, I⁻, O²⁻, S²⁻, N³⁻, OH⁻, CO₃²⁻, NO₃⁻, SO₄²⁻, etc.
• Ionic compounds are formed by combining metal cations with non-metal or polyatomic anions. The metal name stays the same; the non-metal name typically ends with -ide for monoatomic anions (e.g., chloride, oxide) unless a polyatomic ion is involved (e.g., sulfate, nitrate).
• Polyatomic ions keep their names (e.g., NO₃⁻ is nitrate, SO₄²⁻ is sulfate).
• Naming examples:
  • NaCl → Sodium chloride
  • MgCl₂ → Magnesium chloride
  • Ca(OH)₂ → Calcium hydroxide
• Writing ionic formulas using the crossover method (replacement method):
  1) Write the symbols and their charges for the ions.
  2) Write the cation and anion together (fix their charges).
  3) Cross the charges to become subscripts (balance the charges).
  4) Write the formula in standard form with brackets if needed.
• Examples of formula writing using crossover:
  • Mg²⁺ + O²⁻ → MgO
  • Na⁺ + Cl⁻ → NaCl
  • Li⁺ + Br⁻ → LiBr
  • Ca²⁺ + OH⁻ → Ca(OH)₂
  • Cu²⁺ + SO₄²⁻ → CuSO₄
  • Fe³⁺ + S²⁻ → Fe₂S₃
• Practice ionic formulas (polyatomic ions):
  • Na⁺ with NO₃⁻ → NaNO₃
  • Ca²⁺ with CO₃²⁻ → CaCO₃
  • NH₄⁺ with Cl⁻ → NH₄Cl
• Important note: When a polyatomic ion is involved and multiple are needed to balance, use parentheses (e.g., Ca(NO₃)₂).

Writing Ionic Formulas and the Crossover Method

• Steps for writing ionic formulas:
  1) Write the symbols for each ion and their charges in large boxes; write the ion charges in small boxes.
  2) Use the identical ions from the boxes (these cannot change) to balance charges.
  3) Swap the charges (cross them) to give subscripts for the opposite ion.
  4) Write the chemical formula without the boxes; brackets are used if needed.
• Examples shown in notes include:
  • Sodium oxide: Na₂O
  • Calcium chloride: CaCl₂
  • Magnesium chloride: MgCl₂
  • Lithium hydroxide: LiOH
  • Ammonium nitrate: NH₄NO₃
  • Calcium carbonate: CaCO₃
• Polyatomic ion table (examples):
  • Carbonate: CO₃²⁻
  • Hydrogen carbonate (bicarbonate): HCO₃⁻
  • Nitrate: NO₃⁻
  • Phosphate: PO₄³⁻
  • Sulfate: SO₄²⁻
  • Hydroxide: OH⁻
  • Ammonium: NH₄⁺

Writing Ionic Formulas – Practice Problems

• Practice compounds (examples from notes):
  • MgO, NaCl, LiBr, Ca(OH)₂, CuSO₄, CaCO₃, NaNO₃, FeS, CuCO₃
• Note on ionic formulas with polyatomic ions: use parentheses when there are multiple ions (e.g., Ca(NO₃)₂).

Balancing Chemical Equations

• Purpose: to satisfy the conservation of mass; the number and type of atoms on each side must match.
• Historical context: Lavoisier (1788) emphasized the law of conservation of mass.
• Key concepts:
  • Word equations describe reactants and products in words (e.g., Magnesium + Hydrochloric acid → Magnesium chloride + hydrogen).
  • Chemical equations use formulas (e.g., Mg + 2 HCl → MgCl₂ + H₂).
  • Reactants are to the left of the arrow; products to the right; the arrow indicates direction, not an equals sign.
• Subscripts vs coefficients:
  • Subscript indicates the number of atoms of an element within a molecule.
  • Coefficient indicates the number of molecules (or formula units) involved in the reaction.
  • Example: 2 H₂O has a coefficient 2 (two molecules) and a subscript 2 for hydrogen inside each molecule.
• Balancing rules:
  • Do not change subscripts; balance by adjusting coefficients.
  • Ensure the same number and type of atoms for each element on both sides.
• Balanced equation example:
  • Unbalanced: Mg + O₂ → MgO
  • Balanced: 2\,\text{Mg} + \text{O}_2 \rightarrow 2\,\text{MgO}
• Practice balancing problems from notes include:
  • N₂ + 3 H₂ → 2 NH₃
  • 2 KClO₃ → 2 KCl + 3 O₂
  • 2 NaCl + F₂ → 2 NaF + Cl₂
  • 2 H₂ + O₂ → 2 H₂O
  • Pb(OH)₂ + 2 HCl → 2 H₂O + PbCl₂
  • 2 AlBr₃ + 3 K₂SO₄ → 6 KBr + Al₂(SO₄)₃
• Practical notes:
  • You can only change the coefficients at the front of formulas.
  • Use a systematic approach (start with elements that appear only once on each side, then move to other elements).

Word Equations and Phet Simulations

• Word equations are descriptive (e.g., Ammonia formation from N₂ and H₂).
• PhET simulations referenced for balancing, reactants/products, and left-overs concepts (useful for visual learners).

Endothermic and Exothermic Reactions

• Definitions:
  • Exothermic: energy is released to the surroundings (temperature of surroundings increases).
  • Endothermic: energy is absorbed from the surroundings (temperature of surroundings decreases).
• Etymology: exo = outside; endo = inside; thermic = heat.
• Common examples:
  • Exothermic: combustion of fuels, candle flame, rusting of iron, formation of snow in clouds.
  • Endothermic: photosynthesis, baking bread, evaporation of water, melting ice, cooking an egg.
• Energy profile:
  • Exothermic: reactants have higher bond energy; breaking bonds releases energy; the energy difference is emitted as heat/light.
  • Endothermic: products have higher energy; energy is absorbed from the surroundings to form products.
• Example equations:
  • Exothermic: \mathrm{2\,H2 + O2 → 2\,H_2O} (releases energy)
  • Endothermic: photosynthesis (conceptual): 6\,CO2 + 6\,H2O + \text{light energy} → C6H{12}O6 + 6\,O2
• Practical recap questions focus on identifying products and contrasting energy exchange.

Acids and Bases

• Learning goals:
  • Define acids and bases; describe properties; provide examples.
  • Determine acid/base strength and use indicators to predict pH.
  • Understand Aboriginal practices using alkali ash to raise pH.
• Acids:
  • Definition: acids release hydrogen ions (H⁺) into aqueous solution.
  • Common examples: HCl, H₂SO₄, HNO₃, CH₃COOH (ethanoic acid).
  • Properties:
  • Corrosive; react with some metals to produce H₂ gas and a salt; neutralised by bases to form water and salt; turn blue litmus red; conduct electricity; have a sour taste.
  • Strength vs concentration:
  • Strength refers to how many H⁺ ions are released in solution (strong acids release many H⁺; weak acids release few).
  • Concentration refers to the amount of acid per amount of solution (concentrated vs dilute).
  • Names and formulas of common acids (examples):
  • Hydrochloric acid: \text{HCl}
  • Sulphuric acid: \text{H2SO4}
  • Nitric acid: \text{HNO_3}
  • Ethanoic acid: \text{CH_3COOH}
• Bases:
  • Definition: bases release hydroxide ions (OH⁻) into aqueous solution. An alkali is a base that dissolves in water to form an alkaline solution.
  • Common examples: NaOH, Ca(OH)₂, NH₃, NaHCO₃, Na₂CO₃, NH₄OH.
  • Properties:
  • Caustic/silply; soapy/slimy feel; turn red litmus blue; taste bitter; conduct electricity; neutralised by acids to form water and salt.
  • Strength:
  • Strong bases release more OH⁻ in water; weak bases release fewer OH⁻.
  • Names and formulas of common bases (examples):
  • Calcium hydroxide: \text{Ca(OH)_2}
  • Sodium hydroxide: \text{NaOH}
  • Ammonia: \text{NH_3}
  • Sodium hydrogen carbonate: \text{NaHCO_3}
  • Sodium carbonate: \text{Na2CO3}
• pH scale and indicators:
  • pH scale ranges 0-14; acids have pH < 7; bases/alkalis have pH > 7; neutral solutions have pH ~7.
  • Indicators:
  • Litmus paper (blue/red): acids turn blue litmus red; bases turn red litmus blue.
  • Universal indicator: changes color across the pH range to indicate pH value.
  • Quicklime (alkali ash) and Aboriginal use: quicklime (calcium oxide) used to raise pH in medicinal preparations; pyrolysis was used to produce alkali ash for medical uses.
• Concentration vs strength:
  • Concentrated solutions have more solute per litre; dilute solutions have less solute per litre.
• Aboriginal and Torres Strait Islander practices:
  • Alkali salts in wood ashes used as substitutes for sea salt; pyrolysis releases alkali salts (K and Na) from plant matter; used to prepare medicines and aid absorption.

Indicators and pH Testing

• Indicators used to determine acidity/alkalinity:
  • Litmus paper: blue turns red in acids; red turns blue in bases.
  • Universal indicator: color corresponds to pH value on a chart.
  • Other indicators (examples): methyl orange, etc. (synthetic indicators mentioned).
• pH testing and interpretation:
  • A pH of 7 is neutral, below 7 is acidic, above 7 is basic.
  • Titration basics show how pH changes during acid-base neutralization (illustrated in charts and simulations).

Neutralisation Reactions and Acid–Base Chemistry

• Neutralisation: acid reacts with base to form a salt and water (and sometimes additional products).
  • General reaction: ext{acid} + ext{base}
    ightarrow ext{salt} + ext{water}
  • Example: \text{NaOH} + \text{HCl} \rightarrow \text{NaCl} + \text{H_2O}
• Acid and metal carbonate reactions:
  • Acid + carbonate yields salt + water + carbon dioxide (CO₂).
  • Example: \text{H2SO4} + \text{ZnCO3} \rightarrow \text{ZnSO4} + \text{H2O} + \text{CO2}
• Acid and hydrogen carbonate reactions:
  • Acid + bicarbonate yields salt + CO₂ + H₂O.
  • Example: \text{H2SO4} + 2\text{NaHCO3} \rightarrow \text{Na2SO4} + 2\text{H2O} + 2\text{CO_2}
• Acid-metal reactions:
  • Acids react with metals to produce a salt and hydrogen gas.
  • Example (worded): Magnesium + Hydrochloric acid → Magnesium chloride + Hydrogen.
  • Example (chemical equation): \text{Mg} + 2\text{HCl} \rightarrow \text{MgCl2} + \text{H2}
• Environmental context: acid rain results from dissolved SO₂ and NO₂ forming sulfuric and nitric acids; damages materials and ecosystems; acid rain pH can be as low as 3.

Reactions Involving Oxygen and Combustion

• Reactions with oxygen are oxidation reactions and lead to oxide formation (rust is iron oxide).
• Metal and oxygen reactions vary in rate by metal (e.g., Mg oxidizes very fast to MgO; Fe oxidizes to Fe₂O₃).
• Corrosion: general term for deterioration of metals due to environmental interactions (e.g., rust on iron, green patina on copper).
• Combustion: an exothermic reaction with oxygen, producing heat and often light and flame.
• Complete combustion:
  • Occurs with excess O₂; general form for hydrocarbons: \text{CxHy} + \text{O}2 \rightarrow x\text{CO}2 + \frac{y}{2}\text{H_2O}
  • Example: glucose (C₆H₁₂O₆) combusts with O₂ to give CO₂ and H₂O.
• Incomplete combustion:
  • Occurs with limited O₂; can produce carbon monoxide (CO) or carbon (soot) along with CO₂ and H₂O.
  • General idea: hydrocarbon + limited O₂ → CO + C + H₂O (and CO₂ in some cases).
• Practical notes on balancing combustion equations:
  • Write unbalanced formula, balance C first, then H, then O by adjusting O₂ coefficients.
  • Example balancing steps for ethyne C₂H₂:
  • C₂H₂ + O₂ → CO₂ + H₂O
  • Balance C: C₂H₂ + O₂ → 2 CO₂ + H₂O
  • Balance H: C₂H₂ + O₂ → 2 CO₂ + H₂O (already balanced for C and H)
  • Balance O: C₂H₂ + 2.5 O₂ → 2 CO₂ + H₂O
• Aboriginal and Torres Strait Islander practices:
  • Combustion of resins and gums from trees/shrubs for light; torches used for night navigation and fishing.

Key Applications and Examples (Selected Demonstrations)

• Ionic formulas and naming:
  • Example ionic compound names: NaCl (Sodium chloride), MgCl₂ (Magnesium chloride), Ca(OH)₂ (Calcium hydroxide)
• Balanced equation examples:
  • Mg + O₂ → MgO becomes 2\,\text{Mg} + \text{O}_2 \rightarrow 2\,\text{MgO}
  • N₂ + 3 H₂ → 2 NH₃
  • 2 KClO₃ → 2 KCl + 3 O₂
  • 2 NaCl + F₂ → 2 NaF + Cl₂
  • 2 H₂ + O₂ → 2 H₂O
  • Pb(OH)₂ + 2 HCl → 2 H₂O + PbCl₂
  • 2 AlBr₃ + 3 K₂SO₄ → 6 KBr + Al₂(SO₄)₃
• Acid-base indicators and pH testing:
  • Universal indicator color chart indicates pH values; purple generally indicates strongly basic; orange indicates acidic around pH ~3; red indicates strongly acidic; blue indicates basic; color mappings vary by chart.
  • Litmus tests: acids turn blue litmus red; bases turn red litmus blue.
• Environmental and health context:
  • Acids released into the environment can cause corrosion and contribute to acid rain; bases can neutralise acids in environmental contexts.
  • Quicklime (CaO) used historically to raise pH for medicinal absorption via pyrolysis of wood ash.
• Summary recaps useful prompts for studying:
  • What are the products of metal + acid reactions? (Salt + H₂)
  • What are the products of acid + carbonate reactions? (Salt + CO₂ + H₂O)
  • What are the two products of a neutralisation reaction? (Salt + Water)
  • Distinguish complete and incomplete combustion by oxygen availability.
  • Identify acid vs base using pH and indicators.

References to Practice Worksheets and Simulations

• Practice problems and worksheets named in the transcript:
  • Worksheet 1: Naming ionic compounds
  • Worksheet 2: Writing ionic formulas
  • Worksheet 3: Balancing equations
  • Worksheet 4: Endothermic and exothermic reactions
  • Worksheet 5: Acids, bases and indicators
  • Worksheet 6: Types of acid reactions
  • Worksheet 7: Combustion reactions
• PhET simulations mentioned:
  • Reactants, Products and Leftovers
  • Balancing Chemical Equations
  • Acid-Base Solutions
  • PH Scale Basics
  • Carbonate and acid-base activity links provided in notes

Quick Reference: Key Formulas and Equations

• Ionic lattice formation:
  • General lattice concept: an array of alternating cations and anions in 3D, held by electrostatic forces.
• Balanced equations (selected):
  • 2\,\text{Mg} + \text{O}_2 \rightarrow 2\,\text{MgO}
  • \text{N2} + 3\text{H2} \rightarrow 2\text{NH_3}
  • 2\,\text{KClO}3 \rightarrow 2\,\text{KCl} + 3\text{O}2
  • 2\,\text{NaCl} + \text{F}2 \rightarrow 2\text{NaF} + \text{Cl}2
  • 2\,\text{H2} + \text{O}2 \rightarrow 2\text{H_2O}
  • \text{Mg} + 2\text{HCl} \rightarrow \text{MgCl}2 + \text{H2}
  • \text{Pb(OH)2} + 2\text{HCl} \rightarrow 2\text{H2O} + \text{PbCl_2}
  • 2\,\text{AlBr}3 + 3\,\text{K}2\text{SO}4 \rightarrow 6\,\text{KBr} + \text{Al}2(\text{SO}4)3
• Combustion general form:
  • Complete: \text{CxHy} + O2 \rightarrow x\,CO2 + \frac{y}{2}\,H_2O
  • Ethyne balancing example: \text{C}2\text{H}2 + 2.5\,O2 \rightarrow 2\,CO2 + \text{H}_2\text{O}
• Acid-base neutralisation:
  • \text{NaOH} + \text{HCl} \rightarrow \text{NaCl} + \text{H_2O}
• Acid + carbonate/base reactions:
  • \text{H2SO4} + \text{ZnCO}3 \rightarrow \text{ZnSO4} + \text{H2O} + \text{CO}2
  • \text{H2SO4} + 2\text{NaHCO}3 \rightarrow \text{Na2SO4} + 2\text{H2O} + 2\text{CO}_2
• Acid + metal reaction (example):
  • \text{Mg} + 2\text{HCl} \rightarrow \text{MgCl}2 + \text{H2}
• Oxidation/rust example:
  • \text{Fe} + \text{O}2 \rightarrow \text{Fe}2\text{O}_3
• Acid + metal carbonate/incomplete reactions illustrate environmental implications (acid rain) and corrosion effects.