Year 10 Science: Chemical Changes Definitive Study Guide

Key Learning Objectives

• Describe neutrality and relative acidity and alkalinity in terms of $\text{pH}$ measured using Universal Indicator.

• Write word and balanced symbol equations for the reactions of acids with metals, bases (including alkalis), and carbonates.

• Describe the characteristic properties of acids and alkalis, including their effect on litmus paper.

• Describe everyday applications of neutralisation reactions.

• Define electrolysis as the breakdown of an ionic compound when molten or in aqueous solution by the passage of electricity.

• Define the term aqueous.

• Use the terms electrode, electrolyte, anode, and cathode correctly.

• Explain why solid ionic compounds cannot conduct electricity, whereas ionic compounds can conduct electricity when melted or dissolved in water.

• Explain what happens to positive and negative ions during electrolysis and how elements form from their ions.

• Predict the products of electrolysis of molten ionic compounds.

• Recall the reactivity series of metals.

• Give reasons why some metals have to be extracted by electrolysis.

• Know that aluminium is extracted from the ore bauxite by electrolysis.

• Label a diagram and describe the step-by-step process of extracting aluminium by electrolysis.

Part 1: Acids, Bases, and Alkalis

• Aqueous Solutions:

  • An aqueous solution is formed when a substance dissolves in water.

  • The term aqueous ($(aq)$) specifically means "dissolved in water".

  • Aqueous solutions can be acidic, alkaline, or neutral.

• Acids:

  • Definition: A substance that produces hydrogen ions ($H^+$) when dissolved in water.

  • Properties: Acids have a $\text{pH}$ of less than $7$ and taste sharp and sour (e.g., vinegar, lemon juice, citrus fruits).

  • Formulae of common laboratory acids:

  • Hydrochloric acid: $HCl$

  • Sulfuric acid: $H_2SO_4$

  • Nitric acid: $HNO_3$

  • Acid Strength:

  • Strong acids (e.g., $HCl$) fully dissociate in water, producing a high concentration of $H^+$ ions per unit volume.

  • Weak acids (e.g., ethanoic acid in vinegar, citric acid) only partially dissociate.

• Bases and Alkalis:

  • Base: Any substance that reacts with an acid to form a salt and water (typically metal oxides or metal hydroxides).

  • Alkali: A base that is soluble in water, forming an aqueous solution with a $\text{pH}$ greater than $7$

  • Alkalis release hydroxide ions ($OH^-$) when dissolved in water.

  • Formulae of common laboratory alkalis:

  • Sodium hydroxide: $NaOH$

  • Potassium hydroxide: $KOH$

  • Ammonium hydroxide: $NH_4OH$

  • Alkali Strength: Strong alkalis (e.g., $NaOH$ in oven cleaner) produce high concentrations of $OH^-$ ions, whereas weak alkalis (e.g., soap, toothpaste) produce lower concentrations.

• Neutral Solutions:

  • Solutions with a $\text{pH}$ of exactly $7$ (e.g., pure water, sugar solution).

  • Neutral solutions contain equal concentrations of $H^+$ and $OH^-$ ions.

Part 2: Indicators and the pH Scale

• Litmus Indicator:

  • Available as liquid or paper strips (red and blue).

  • Turns red in acidic solutions.

  • Turns blue in alkaline solutions.

  • Remains unchanged in neutral solutions.

  • Limitation: Indicates only whether a solution is acidic or alkaline, not the strength of acidity or alkalinity.

• Universal Indicator and the pH Scale:

  • Universal Indicator is a mixture of several indicators that changes through a continuous spectrum of colours across the $\text{pH}$ scale ($0$ to $14$).

  • The $\text{pH}$ number measures the concentration of $H^+$ ions in solution:

  • $\text{pH}\,0-6$: Acidic (red/orange/yellow; lower numbers mean higher $H^+$ concentration and stronger acidity).

  • $\text{pH}\,7$: Neutral (green).

  • $\text{pH}\,8-14$: Alkaline (blue/purple; higher numbers mean lower $H^+$ concentration / higher $OH^-$ concentration and stronger alkalinity).

• Examples of Substance pH Values:

  • $\text{pH}\,0$: Battery acid

  • $\text{pH}\,1-2$: Hydrochloric acid ($HCl$), stomach acid, lemon juice

  • $\text{pH}\,3$: Vinegar, acid rain

  • $\text{pH}\,4-5$: Black coffee, urine, saliva

  • $\text{pH}\,6$: Milk

  • $\text{pH}\,7$: Pure water, distilled water, blood

  • $\text{pH}\,8$: Sea water

  • $\text{pH}\,9-10$: Baking soda, indigestion tablets, milk of magnesia

  • $\text{pH}\,11-12$: Ammonia solution, soapy water

  • $\text{pH}\,13-14$: Bleach, oven cleaner, drain cleaner ($NaOH$)

Part 3: Neutralisation Reactions and Everyday Applications

• The Neutralisation Process:

  • Reaction between an acid and a base/alkali producing a salt and water:     $\text{acid} + \text{alkali} \rightarrow \text{salt} + \text{water}$

  • Ionic Equation:     $H^+_{(aq)} + OH^-_{(aq)} \rightarrow H_2O_{(l)}$

  • Hydrogen ions ($H^+$) from the acid combine with hydroxide ions ($OH^-$) from the alkali to form neutral water molecules.

• Types of Salts Formed:

  • Metal part of the salt comes from the base/alkali; non-metal part comes from the acid.

  • Hydrochloric acid ($HCl$) produces chloride salts (e.g., $NaCl$).

  • Sulfuric acid ($H_2SO_4$) produces sulphate salts (e.g., $MgSO_4$).

  • Nitric acid ($HNO_3$) produces nitrate salts (e.g., $KNO_3$).

• Neutralisation Equations:

  • Hydrochloric acid + Sodium hydroxide:     $HCl_{(aq)} + NaOH_{(aq)} \rightarrow NaCl_{(aq)} + H_2O_{(l)}$

  • Hydrochloric acid + Potassium hydroxide:     $HCl_{(aq)} + KOH_{(aq)} \rightarrow KCl_{(aq)} + H_2O_{(l)}$

  • Sulfuric acid + Sodium hydroxide:     $H_2SO_{4(aq)} + 2NaOH_{(aq)} \rightarrow Na_2SO_{4(aq)} + 2H_2O_{(l)}$

  • Nitric acid + Sodium hydroxide:     $HNO_{3(aq)} + NaOH_{(aq)} \rightarrow NaNO_{3(aq)} + H_2O_{(l)}$

• Everyday Applications:

  • Indigestion Relief: Excess stomach acid ($HCl$) is neutralised by antacid tablets containing weak bases like sodium bicarbonate ($NaHCO_3$) or calcium carbonate ($CaCO_3$).

  • Bee Stings: Bee stings are acidic and are neutralised by weak alkalis like bicarbonate of soda.

  • Wasp Stings: Wasp stings are alkaline and are neutralised by weak acids like vinegar (acetic acid).

  • Agriculture: Acidic soils are treated with lime (calcium oxide or calcium carbonate) to raise soil $\text{pH}$ for healthy crop growth.

  • Wastewater Treatment: Industrial acidic effluent is neutralised with slaked lime before discharge into rivers.

Part 4: Practical Investigation: Exothermic Neutralisation

• Reaction Energetics: Neutralisation is exothermic, releasing heat energy and raising the temperature of the mixture.

• Experimental Procedure:

  1. Measure $25\,cm^3$ of $NaOH$ into an insulated polystyrene cup and record initial baseline temperature.

  2. Add $5\,cm^3$ portions of $HCl$, stirring continuously.

  3. Record the maximum temperature reached after each addition until the temperature levels off and drops.

• Analysis:

  • Temperature rises as neutralisation releases heat until all $NaOH$ is completely reacted (equivalence point).

  • Adding excess cold acid beyond this point drops the temperature due to dilution and cooling.

• Experimental Improvements:

  • Put a lid on the polystyrene cup to reduce heat loss to surroundings.

  • Use a burette or volumetric pipette instead of a measuring cylinder for precise volume measurement.

  • Repeat trials to calculate a mean maximum temperature and identify anomalies.

Part 5: Reactions of Acids with Metals

• General Reaction:   $\text{ACID} + \text{METAL} \rightarrow \text{SALT} + \text{HYDROGEN}$

• Gas Testing for Hydrogen ($H_2$):

  • Collect gas via upward delivery in an inverted test tube (hydrogen is lighter than air).

  • Hold a lighted splint at the test tube mouth; hydrogen burns rapidly with a squeaky 'pop' sound:     $2H_{2(g)} + O_{2(g)} \rightarrow 2H_2O_{(l)}$

• Reactivity Comparison:

  • Magnesium: Rapid effervescence; vigorous exothermic reaction.     $Mg_{(s)} + 2HCl_{(aq)} \rightarrow MgCl_{2(aq)} + H_{2(g)}$

  • Zinc: Moderate, steady rate of bubbling.     $Zn_{(s)} + 2HCl_{(aq)} \rightarrow ZnCl_{2(aq)} + H_{2(g)}$

  • Iron: Slow, gentle bubbling.     $Fe_{(s)} + 2HCl_{(aq)} \rightarrow FeCl_{2(aq)} + H_{2(g)}$

  • Tin and Lead: Very slow bubbling in cold acid; require gentle warming to react noticeably.

  • Copper: Shows no reaction with dilute acid, even when heated.

  • Potassium and Sodium: Explosive/violent reaction with dilute acids; unsafe for standard laboratory experiments.

• Order of Metal Reactivity with Acids:   $Mg > Zn > Fe > Sn > Pb > Cu$

Part 6: Reactions of Acids with Metal Carbonates

• General Reaction:   $\text{ACID} + \text{METAL CARBONATE} \rightarrow \text{SALT} + \text{WATER} + \text{CARBON DIOXIDE}$

• Gas Testing for Carbon Dioxide ($CO_2$):

  • Bubble the gas through limewater (calcium hydroxide solution).

  • Limewater turns cloudy/milky due to the formation of insoluble calcium carbonate precipitate:     $Ca(OH)_{2(aq)} + CO_{2(g)} \rightarrow CaCO_{3(s)} + H_2O_{(l)}$

• Key Carbonate Equations:

  • Calcium carbonate + Hydrochloric acid:     $CaCO_{3(s)} + 2HCl_{(aq)} \rightarrow CaCl_{2(aq)} + H_2O_{(l)} + CO_{2(g)}$

  • Aluminium carbonate + Sulfuric acid:     $Al_2(CO_3)_3 + 3H_2SO_4 \rightarrow Al_2(SO_4)_3 + 3H_2O + 3CO_2$

  • Iron (II) carbonate + Hydrochloric acid:     $FeCO_3 + 2HCl \rightarrow FeCl_2 + H_2O + CO_2$

  • Sodium carbonate + Nitric acid:     $Na_2CO_3 + 2HNO_3 \rightarrow 2NaNO_3 + H_2O + CO_2$

Part 7: Principles of Electrolysis

• Definition: Breakdown of an ionic compound when molten or in aqueous solution using direct electrical current.

• Electrical Conductivity:

  • Solid Ionic Compounds: Do not conduct electricity because ions are locked in fixed positions within a rigid ionic lattice.

  • Molten or Aqueous Ionic Compounds: Conduct electricity because melting or dissolving breaks the lattice, freeing ions to move and carry charge.

• Key Electrolysis Terminology:

  • Electrolyte: The liquid or solution containing mobile ions undergoing electrolysis.

  • Electrode: Conducting rods (graphite or platinum) immersed in the electrolyte.

  • Anode: Positive electrode ($+$-electrode) attracting negative ions (anions).

  • Cathode: Negative electrode ($-$-electrode) attracting positive ions (cations).

  • Mnemonic: PANIC — Positive Anode, Negative Is Cathode.

• Ion Discharge Mechanisms:

  • At Cathode ($-$-electrode): Positive metal ions (cations) gain electrons (Reduction) to form neutral metal atoms.

  • At Anode ($+$-electrode): Negative non-metal ions (anions) lose electrons (Oxidation) to form neutral non-metal molecules.

• Electrolysis of Molten Ionic Compounds Examples:

  • Molten Lithium Chloride ($LiCl_{(l)}$):

  • Cathode: $Li^+ + e^- \rightarrow Li_{(l)}$ (Lithium metal)

  • Anode: $2Cl^- \rightarrow Cl_{2(g)} + 2e^-$ (Chlorine gas)

  • Molten Lead (II) Bromide ($PbBr_{2(l)}$):

  • Cathode: $Pb^{2+} + 2e^- \rightarrow Pb_{(l)}$ (Lead metal)

  • Anode: $2Br^- \rightarrow Br_{2(g)} + 2e^-$ (Bromine gas)

  • Products Table Summary:

  • Zinc iodide ($ZnI_2$): Anode = $I_2$, Cathode = $Zn$

  • Lithium bromide ($LiBr$): Anode = $Br_2$, Cathode = $Li$

  • Iron fluoride ($FeF_2$): Anode = $F_2$, Cathode = $Fe$

  • Sodium oxide ($Na_2O$): Anode = $O_2$, Cathode = $Na$

  • Potassium chloride ($KCl$): Anode = $Cl_2$, Cathode = $K$

Part 8: Extraction of Metals and the Reactivity Series

• Occurrence of Metals:

  • Native State: Unreactive metals at the bottom of the reactivity series (gold, platinum, copper) are found as uncombined pure elements in the Earth's crust.

  • Ores: Reactive metals occur as compounds in rocks called ores (e.g., metal oxides).

• Three Stages of Metal Extraction:

  1. Mining and digging out the ore rock from the ground.

  2. Purifying the ore by removing rocky impurities.

  3. Extracting the pure metal from the compound by chemical reduction or electrolysis.

• Extraction Methods Based on Reactivity:

  • Metals More Reactive than Carbon ($K, Na, Ca, Mg, Al$): Extracted using electrolysis (requires high electrical energy consumption).

  • Metals Less Reactive than Carbon ($Zn, Fe, Sn, Pb$): Extracted by heating with carbon (reduction), where carbon displaces the metal.

  • Unreactive Metals ($Cu, Ag, Au, Pt$): Extracted by simple heating or mined directly in native form.

Part 9: Extraction of Aluminium from Bauxite

• Ore and Extraction Need:

  • Aluminium is extracted from bauxite ore ($Al_2O_3$ mixed with rocky impurities).

  • Because aluminium is more reactive than carbon, carbon cannot reduce aluminium oxide; electrolysis must be used.

• Step-by-Step Extraction Process:

  1. Purification: Bauxite ore is mined and purified to remove rock impurities, yielding pure white aluminium oxide ($Al_2O_3$).

  2. Lowering Melting Point: Pure $Al_2O_3$ has an extremely high melting point (over $2000^\circ\text{C}$). Dissolving it in molten cryolite ($Na_3AlF_6$) lowers operating temperature to around $950^\circ\text{C}$, significantly reducing energy costs.

  3. Electrolysis Cell Setup: The mixture is placed in a steel cell lined with carbon (cathode), with carbon/graphite blocks suspended into the liquid (anodes).

  4. Cathode Reaction (Reduction): Positive aluminium ions ($Al^{3+}$) move to the negative cathode, gaining electrons to form liquid aluminium:      $Al^{3+} + 3e^- \rightarrow Al_{(l)}$      The dense liquid aluminium sinks to the bottom and is tapped off.

  5. Anode Reaction (Oxidation): Negative oxide ions ($O^{2-}$) move to positive anodes, losing electrons to produce oxygen gas:      $2O^{2-} \rightarrow O_{2(g)} + 4e^-$

  6. Anode Replacement: At $950^\circ\text{C}$, oxygen gas reacts with the hot carbon anodes to produce carbon dioxide ($C_{(s)} + O_{2(g)} \rightarrow CO_{2(g)}$) or carbon monoxide ($2C_{(s)} + O_{2(g)} \rightarrow 2CO_{(g)}$). Consequently, carbon anodes burn away and must be replaced regularly.