AQA Chemistry 8462 - Chemical Changes Vocabulary Flashcards

Vocabulary flashcards covering AQA Chemistry 8462 section 4.4 Chemical Changes, including reactivity series, extraction of metals, acids and salts, titrations, strong/weak acids, electrolysis, and required practicals 1-3.

Oxidation (in terms of oxygen)

Gain of oxygen.

Reduction (in terms of oxygen)

Loss of oxygen.

Reaction of a metal with oxygen

Forms a metal oxide.

Oxidation nature of metal-oxygen reaction

It is an oxidation reaction because the metal gains oxygen.

Metal atom behavior in reactions

When metals react with other substances, their atoms form positive ions.

Reactivity of a metal

Determined by how readily its atoms form positive ions.

Reactivity series

An arrangement of elements in order of their reactivity.

AQA metal reactivity order

Potassium, sodium, lithium, calcium, magnesium, zinc, iron, copper.

Positions of carbon and hydrogen in the reactivity series

Carbon is placed below magnesium and above zinc; hydrogen is placed below iron and above copper.

Metals reacting most vigorously with water

Potassium, sodium and lithium.

Products of alkali metals with water

Potassium, sodium and lithium react with water to form a metal hydroxide and hydrogen.

General equation for metal reacting with water

$$\text{metal} + \text{water} \rightarrow \text{metal hydroxide} + \text{hydrogen}$$

Calcium reaction with water at room temperature

Reacts to form calcium hydroxide and hydrogen.

Magnesium reaction with water at room temperature

Reacts very slowly.

Reaction of zinc, iron, and copper with water at room temperature

They do not react significantly with water at room temperature.

Reaction of magnesium, zinc, or iron with dilute acid

Produces a salt and hydrogen.

Copper reaction with dilute hydrochloric or sulfuric acid

Does not react, because copper is below hydrogen in the reactivity series.

General equation for a metal reacting with an acid

$$\text{metal} + \text{acid} \rightarrow \text{salt} + \text{hydrogen}$$

Displacement reaction

A reaction where a more reactive metal displaces a less reactive metal from its compound.

Mechanism of metal displacement

A more reactive metal can displace a less reactive metal because the more reactive metal forms positive ions more readily.

Experimental ranking of metals

Metals can be ranked by comparing how readily or vigorously they react with water, acids, or compounds of other metals.

Natural occurrence of gold

Gold can occur naturally as the metal itself because it is very unreactive.

Occurrence of most metals in the Earth's crust

Most metals are found as compounds because they are reactive enough to have combined chemically with other elements.

Extraction of metals less reactive than carbon

Extracted from their oxides by reduction with carbon.

Limitation of carbon reduction for very reactive metals

Very reactive metals cannot be extracted using carbon because they are more reactive than carbon, so carbon cannot remove oxygen from their oxides.

Role of carbon in reduction of metal oxides

Carbon removes oxygen from the metal oxide.

Reduced species in metal extraction

The metal oxide is reduced because it loses oxygen.

Oxidation (in terms of electrons)

Loss of electrons.

Reduction (in terms of electrons)

Gain of electrons.

OIL RIG mnemonic

Oxidation Is Loss, Reduction Is Gain of electrons.

Oxidation in displacement reactions

The more reactive metal atoms lose electrons and are oxidised.

Reduction in displacement reactions

Ions of the less reactive metal gain electrons and are reduced.

Ionic equation

An equation showing only the particles that actually change during a reaction.

Half equation

An equation showing the electron loss or gain for one species.

Acid and metal reaction products

Acids react with some metals to produce a salt and hydrogen.

AQA required metal-acid reactions

Reactions of magnesium, zinc, and iron with hydrochloric acid and sulfuric acid.

Salts from hydrochloric acid

Produces chloride salts.

Salts from nitric acid

Produces nitrate salts.

Salts from sulfuric acid

Produces sulfate salts.

Redox nature of metal-acid reactions

It is a redox reaction because the metal loses electrons and is oxidised while hydrogen ions gain electrons and are reduced.

Neutralisation

A reaction in which an acid reacts with a base or alkali to form salt and water.

Acid + metal oxide

$$\text{acid} + \text{metal oxide} \rightarrow \text{salt} + \text{water}$$

Acid + metal hydroxide

$$\text{acid} + \text{metal hydroxide} \rightarrow \text{salt} + \text{water}$$

Acid + alkali

$$\text{acid} + \text{alkali} \rightarrow \text{salt} + \text{water}$$

Acid + metal carbonate

$$\text{acid} + \text{metal carbonate} \rightarrow \text{salt} + \text{water} + \text{carbon dioxide}$$

Positive ion in a salt

Determined by the metal or other positive ion supplied by the base, alkali, carbonate, or metal.

Deducing the formula of a salt

Combine the positive and negative ions in a ratio that gives an overall charge of zero.

Preparation of a soluble salt from an insoluble solid

Add excess solid to the acid until no more reacts, filter off the excess, then crystallise the salt solution.

Purpose of adding excess solid in salt preparation

To ensure all the acid has reacted.

Purpose of filtration in salt preparation

To remove excess unreacted insoluble solid.

Crystallisation procedure

Evaporate some water, then allow the concentrated solution to cool and crystallise.

Required practical 1 objective

To prepare a pure, dry sample of a soluble salt from an insoluble oxide or carbonate.

Warming acid in Required Practical 1

Warmed to increase the rate of reaction.

Avoiding boiling acid in Required Practical 1

Acid should not be boiled strongly to avoid excessive evaporation or splashing.

Drying crystals in Required Practical 1

Dried to remove water so a pure, dry salt sample is obtained.

Acid ions in aqueous solution

Acids produce hydrogen ions ($H^+$) in aqueous solution.

Alkali ions in aqueous solution

Aqueous alkalis contain hydroxide ions ($OH^-$).

GCSE pH scale range

Usually covers $0$ to $14$.

Neutral pH

$$\text{pH } 7$$

Acidic pH values

Below $7$.

Alkaline pH values

Above $7$.

Approximate pH measurement

Measured using universal indicator or a wide-range indicator.

Precise pH measurement

Measured using a pH probe or meter.

Ionic equation for neutralisation

$$H^+ + OH^- \rightarrow H_2O$$

Titration purpose

Used to measure accurately the volumes of acid and alkali that react completely.

Specified acids for strong acid-strong alkali titrations

Hydrochloric acid, nitric acid, and sulfuric acid.

Burette

Apparatus that accurately delivers a measured variable volume in a titration.

Pipette

Apparatus commonly used to transfer a fixed accurate volume to the conical flask in a titration.

Purpose of an indicator in titration

To show the end point by a colour change.

Use of a white tile in titration

Makes the indicator colour change easier to see.

Purpose of swirling during titration

To mix the reacting solutions thoroughly.

Dropwise addition of titrant

Done near the end point to avoid adding too much and passing the end point.

Concordant titres

Titre results that are very close to each other.

Titres used to calculate a mean

Concordant titres, excluding the rough result.

Required practical 2 objective

Determine the reacting volumes of a strong acid and strong alkali by titration.

Calculating unknown concentration from titration data

Use the known concentration and volume to calculate moles, apply the balanced equation ratio, then divide the unknown moles by its volume.

Strong acid

An acid that is completely ionised in aqueous solution.

Weak acid

An acid that is only partially ionised in aqueous solution.

Examples of strong acids

Hydrochloric acid, nitric acid, and sulfuric acid.

Examples of weak acids

Ethanoic acid, citric acid, and carbonic acid.

Difference between strong and concentrated acids

Strength describes the degree of ionisation, whereas concentration describes the amount of acid per volume.

Difference between weak and dilute acids

Weak describes partial ionisation, whereas dilute describes a low concentration.

pH comparison of strong vs weak acids at equal concentration

For equal concentrations, the stronger acid has a lower pH.

Effect on H+ concentration when pH decreases by 1

Hydrogen ion concentration increases by a factor of $10$.

Effect on H+ concentration when pH decreases by 2

Hydrogen ion concentration increases by a factor of $100$.

Electrolyte

A molten ionic compound or ionic solution that conducts electricity because its ions can move.

Electrolysis

The decomposition of an ionic compound by passing an electric current through it.

Requirement for molten or dissolved state in electrolysis

An ionic compound must be molten or dissolved so that its ions are free to move to the electrodes.

Cathode

The negative electrode in electrolysis.

Anode

The positive electrode in electrolysis.

Ions attracted to the cathode

Positive ions, or cations.

Ions attracted to the anode

Negative ions, or anions.

Discharge of ions at electrodes

Ions gain or lose electrons and form neutral products.

Cathode product in electrolysis of a molten binary ionic compound

The metal.

Anode product in electrolysis of a molten binary ionic compound

The non-metal.

Products of molten lead bromide electrolysis

Lead at the cathode and bromine at the anode.

Reason for using electrolysis in metal extraction

Used when metals are too reactive to be extracted by reduction with carbon, or when they react with carbon.

Cost factor of electrolytic metal extraction

Expensive because large amounts of energy are needed to melt compounds and supply the electric current.

Extraction method for aluminium

Extracted by electrolysis of molten aluminium oxide mixed with cryolite.

Role of cryolite in aluminium extraction

Lowers the melting point of the mixture, reducing the energy required.