Acids & Alkalis (AQA GCSE Combined Science: Synergy: Physical Sciences): Flashcards
Exam code: 8465
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- FrontReactions Of Acids
Define base.
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Cards in this collection (64)
Define base.
A base is a substance that can neutralise an acid. Examples include metal oxides, metal hydroxides, and metal carbonates.
What are the products when a metal reacts with an acid?
When a metal reacts with an acid, the products are a salt and hydrogen gas.
General equation: metal + acid → salt + hydrogen
True or False?
Copper reacts with dilute hydrochloric acid to produce a salt and hydrogen.
False.
Copper does not react with dilute acids because it is below hydrogen in the reactivity series. Only metals above hydrogen in the reactivity series react with dilute acids.
metal + acid → .......... + ..........
metal + acid → salt + hydrogen
Define neutralisation.
Neutralisation is a chemical reaction between an acid and a base that produces a salt and water. The pH of the solution moves closer to 7.
How do you test for hydrogen gas?
Hold a burning splint at the mouth of a test tube containing the gas. If hydrogen is present, it burns with a squeaky pop.
True or False?
Effervescence during an acid-carbonate reaction is caused by hydrogen gas.
False.
The effervescence (fizzing) is caused by carbon dioxide gas, not hydrogen. Metal carbonates react with acids to produce a salt, water, and carbon dioxide.
How do you test for carbon dioxide gas?
Bubble the gas through limewater (calcium hydroxide solution). If carbon dioxide is present, the limewater turns cloudy white.
Define redox reaction.
(Higher Tier Only)
A reaction in which oxidation and reduction occur simultaneously. One species loses electrons (oxidation) while another gains electrons (reduction).
In the reaction between zinc and hydrochloric acid, which species is oxidised and which is reduced?
(Higher Tier Only)
Zinc is oxidised: Zn → Zn2+ + 2e−
Zn loses 2 electrons.
Hydrogen ions are reduced: 2H+ + 2e− → H2
H+ ions gain electrons.
True or False?
In a metal-acid reaction, the metal atom loses electrons and is therefore oxidised.
(Higher Tier Only)
True.
Metal atoms lose electrons to form positive ions — this is oxidation. Hydrogen ions simultaneously gain electrons — this is reduction. Both happen together in a redox reaction.
In redox reactions, oxidation is the .......... of electrons, and reduction is the .......... of electrons.
(Higher Tier Only)
In redox reactions, oxidation is the loss of electrons, and reduction is the gain of electrons.
Write the two half equations for the reaction between magnesium and sulfuric acid.
(Higher Tier Only)
Oxidation: Mg → Mg2+ + 2e− (Mg loses electrons)
Reduction: 2H+ + 2e− → H2 (H+ gains electrons)
Overall: Mg + H2SO4 → MgSO4 + H2
True or False?
In the ionic equation for a metal-acid reaction, the spectator ions change during the reaction.
(Higher Tier Only)
False.
Spectator ions remain unchanged and are not included in the ionic equation. For example, SO42- is a spectator ion in the Mg + H2SO4 reaction.
Define oxidation and reduction — in terms of electrons.
(Higher Tier Only)
Oxidation: loss of electrons. Reduction: gain of electrons.
Memory aid: OIL RIG — Oxidation Is Loss, Reduction Is Gain.
Why is the reaction between iron and hydrochloric acid classified as a redox reaction?
(Higher Tier Only)
Oxidation and reduction happen simultaneously:
Fe loses electrons (oxidised): Fe → Fe2+ + 2e−
H+ gains electrons (reduced): 2H+ + 2e− → H2
Both processes occur in the same reaction.
Define base.
A chemical that neutralises acids. Examples include metal oxides, metal hydroxides and metal carbonates. Bases that dissolve in water are called alkalis.
What are the products when an acid reacts with a metal oxide?
salt + water
general equation: acid + metal oxide → salt + water
Example: 2HCl + CuO → CuCl2 + H2O
True or False?
Hydrochloric acid always produces chloride salts when it reacts with a base.
True.
Hydrochloric acid always produces chloride salts. Sulfuric acid produces sulfates; nitric acid produces nitrates.
The salt produced from a reaction with an acid is named using the .......... of the metal and the type of ..........
The salt produced from a reaction with an acid is named using the name of the metal and the type of acid.
Name the salt formed when zinc reacts with nitric acid.
Zinc nitrate (Zn(NO3)2)
Nitric acid always produces nitrate salts. The metal name gives the first part.
True or False?
When preparing a soluble salt, the insoluble base is added in excess to ensure all the acid is used up.
True.
Adding the base in excess ensures all the acid reacts. Any unreacted acid would become dangerously concentrated during evaporation. Excess base is removed by filtration.
Describe how a soluble salt is separated from excess base after the neutralisation reaction.
The mixture is filtered to remove the excess insoluble base. The filtrate (salt solution) is then heated to evaporate water until crystals form, and left to dry.
Define alkali.
A base that dissolves in water, producing an alkaline solution with pH above 7. Examples: sodium hydroxide (NaOH) and calcium hydroxide Ca(OH)2.
In RP 17, why is the insoluble base added in excess?
To ensure all the acid has reacted. Unreacted acid would become dangerously concentrated during evaporation. Excess base is harmless and removed by filtration.
Define filtration (in salt preparation).
The process used after the neutralisation in RP 17 to remove excess insoluble base. The filtrate (salt solution) passes through filter paper; the solid base remains.
True or False?
In RP 17, the acid and copper(II) oxide are mixed cold to make copper(II) sulfate.
False.
The acid is warmed using a Bunsen burner before adding the copper(II) oxide. Warming speeds up the reaction between the acid and the insoluble oxide.
In RP 17, after filtering, the salt solution is heated until .......... begin to form, then left to .......... fully.
In RP 17, after filtering, the salt solution is heated until crystals begin to form, then left to crystallise / dry fully.
What is the chemical equation for the reaction in RP 17 between copper(II) oxide and dilute sulfuric acid?
CuO (s) + H2SO4 (aq) → CuSO4 (aq) + H2O (l)
Copper(II) oxide reacts with sulfuric acid to form copper(II) sulfate and water.
True or False?
The solution is saturated when a cold glass rod dipped into it shows crystals forming on its surface.
True.
This is the standard check for saturation in RP 17. Once the solution is saturated, it is left to cool and crystallise slowly, forming larger crystals.
Define crystallisation.
The process by which salt crystals form as water evaporates from a saturated solution. Used in RP 17 to obtain a pure, dry salt sample.
State the colour and appearance of pure copper(II) sulfate crystals.
Copper(II) sulfate crystals are bright blue and regularly shaped, reflecting the ionic lattice structure of the compound.
Define exothermic reaction.
A reaction that transfers energy to the surroundings, causing the temperature of the surroundings to increase. Examples: combustion, neutralisation.
Define endothermic reaction.
A reaction that takes in energy from the surroundings, causing the temperature of the surroundings to decrease. Examples: thermal decomposition, citric acid + sodium hydrogencarbonate.
True or False?
In an exothermic reaction, the products have more energy than the reactants.
False.
In an exothermic reaction, energy is released to the surroundings. The products therefore have less energy than the reactants by the amount of energy transferred.
Give two everyday uses of exothermic reactions.
Self-heating cans — use an exothermic reaction to heat food or drink
Hand warmers — release heat via an exothermic chemical reaction
Other examples: combustion (burning fuels).
Energy is conserved in chemical reactions: it cannot be .......... or .........., only ..........
Energy is conserved in chemical reactions: it cannot be created or destroyed, only transferred.
True or False?
The reaction between citric acid and sodium hydrogencarbonate is endothermic.
True.
This reaction takes in heat from the surroundings, causing the temperature to decrease. It is a classic example of an endothermic reaction.
How can you tell from a temperature change whether a reaction is exothermic or endothermic?
Exothermic: temperature of surroundings increases — energy transferred out.
Endothermic: temperature of surroundings decreases — energy taken in.
A thermometer measures these changes.
Give two examples of endothermic reactions.
Thermal decomposition (e.g. calcium carbonate → calcium oxide + CO2)
Citric acid + sodium hydrogencarbonate
Other examples: electrolysis, early stages of photosynthesis.
In RP 18, what type of apparatus is used to measure temperature changes in reactions?
A polystyrene (Styrofoam) cup calorimeter with a lid and thermometer. The insulating cup minimises heat loss to the surroundings.
Define calorimetry.
The experimental technique used in RP 18 to measure temperature changes in chemical reactions. A calorimeter (insulated container) reduces heat loss.
True or False?
In RP 18 (HCl + NaOH), temperature rises continuously as more NaOH is added.
False.
Temperature rises as NaOH is added and neutralisation occurs, reaches a peak, then falls as excess NaOH dilutes the solution with no further reaction.
In RP 18, two straight lines of best fit are drawn on the graph. Their .......... indicates the volume at which neutralisation is complete.
In RP 18, two straight lines of best fit are drawn on the graph. Their intersection indicates the volume at which neutralisation is complete.
State four types of reaction that can be investigated in RP 18.
Acid + metal (e.g. HCl + magnesium)
Acid + carbonate (e.g. HCl + calcium carbonate)
Neutralisation (e.g. HCl + NaOH)
Displacement (e.g. zinc + copper sulfate solution)
True or False?
The neutralisation reaction between HCl and NaOH in RP 18 is exothermic.
True.
Neutralisation reactions are exothermic — they release heat, causing the temperature of the solution to rise. The polystyrene cup minimises heat loss.
Define anomalous result.
A result that does not fit the pattern of the other data. In RP 18, anomalous results are excluded from the mean calculation (e.g. a reading that is clearly too high or too low).
Why does temperature fall after the peak in the RP 18 neutralisation graph?
Once the limiting reactant is used up, adding more of the other reactant (now in excess) simply dilutes the solution. No further reaction occurs, so no more heat is released, and the solution cools.
Define pH scale.
A numerical scale from 0 to 14 measuring the acidity or alkalinity of a solution. Below 7 = acidic; 7 = neutral; above 7 = alkaline.
What ions do acids and alkalis produce in aqueous solution?
Acids produce hydrogen ions, H+ (aq).
Alkalis produce hydroxide ions, OH− (aq).
The relative concentrations of H+ and OH− determine the pH.
True or False?
A solution with pH 3 is more acidic than a solution with pH 6.
True.
The lower the pH, the more acidic the solution. A pH of 3 indicates a higher concentration of H+ ions than a pH of 6.
Universal indicator gives only an .......... value for pH. A pH probe gives a more .......... measurement.
Universal indicator gives only an approximate value for pH. A pH probe gives a more accurate measurement.
Define neutralisation.
The reaction between hydrogen ions and hydroxide ions to produce water:
H+ (aq) + OH− (aq) → H2O (l)
This is the net ionic equation for all acid-alkali neutralisations.
True or False?
Universal indicator is suitable for detecting the end-point in an acid-base titration.
False.
Universal indicator gives only an approximate pH and changes colour gradually. A titration requires a sharp colour change at the end-point, which only single-colour indicators provide.
Write the ionic equation for the neutralisation of an acid by an alkali.
H+ (aq) + OH− (aq) → H2O (l)
Hydrogen ions from the acid react with hydroxide ions from the alkali to form water, giving a neutral solution.
State the pH range for acids, neutral solutions, and alkalis.
Acids: pH below 7
Neutral: pH = 7
Alkalis: pH above 7
The pH scale runs from 0 (most acidic) to 14 (most alkaline).
Define strong acid.
(Higher Tier Only)
An acid that completely ionises (dissociates) in water, producing a high concentration of H+ ions. Examples: hydrochloric acid, nitric acid, sulfuric acid.
Define weak acid.
(Higher Tier Only)
An acid that only partially ionises (dissociates) in water, producing a low concentration of H+ ions. Examples: ethanoic acid, citric acid, carbonic acid.
True or False?
A concentrated solution of a weak acid always has a lower pH than a dilute solution of a strong acid.
(Higher Tier Only)
False.
Concentration (amount of acid per dm3) and strength (degree of ionisation) are independent. A dilute strong acid can have a lower pH than a concentrated weak acid because the strong acid ionises more completely.
What does it mean that the pH scale is logarithmic?
(Higher Tier Only)
Each change of 1 pH unit represents a change in H+ concentration by a factor of 10.
So pH 3 has 10 times the H+ concentration of pH 4, and 100 times that of pH 5.
Weak acids .......... (completely / partially) dissociate in water. The equilibrium lies to the .........., giving a low concentration of H+ ions.
(Higher Tier Only)
Weak acids partially dissociate in water. The equilibrium lies to the left, giving a low concentration of H+ ions.
True or False?
For two acids of equal concentration, the stronger acid has the lower pH.
(Higher Tier Only)
True.
The stronger acid ionises more completely, producing more H+ ions. Higher H+ concentration means lower pH.
Distinguish between the terms dilute and weak when applied to acids.
(Higher Tier Only)
Dilute: describes concentration — low number of acid molecules per dm3 of solution.
Weak: describes degree of ionisation — only a small proportion of molecules dissociate.
A dilute strong acid can be more acidic than a concentrated weak acid.
Give three examples each of strong and weak acids.
(Higher Tier Only)
Strong acids: hydrochloric acid, nitric acid, sulfuric acid
Weak acids: ethanoic acid, citric acid, carbonic acid
Strong acids completely dissociate; weak acids only partially dissociate.
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