Chemical changes

  • Created by: holly6901
  • Created on: 28-03-19 15:41

Acids and alkalis

Alkalis:

  • Alkalis are substances that form hydroxide (OH-) ions when they dissolve in water (aqueous solutions).
  • An alkali is a type of base.

Acids

  • Acids are substances that form hydrogen (H+) ions when they dissolve in water (aqueous solutions).
  • The acidity and alkalinity of a substance is measured using the pH scale.
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Oxidation reactions

Oxidation:

  • An oxidation reaction involves gaining oxygen.
  • E.g. when metals react with oxygen, metal oxides are produced:
    • Magnesium + oxygen → magnesium oxide
    • 2Mg + O2 → 2MgO

Reduction

  • A reduction reaction involves losing oxygen.
  • E.g. when metal oxides lose oxygen and return to their atomic form:
    • Iron oxide + carbon monoxide → iron + carbon dioxide
    • Fe2O3 + 3CO → 2Fe + 3CO2
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The reactivity series

  • When metals react with other substances, the metal atoms always form positive ions.
  • The easier it is for metal to form its positive ion, the more reactive the metal is.
  • Metals can be arranged in order of their reactivity.
  • Very unreactive metals, such as gold and platinum, are found in the Earth’s crust as pure metals. These are called native metals.
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Displacement reactions

  • A metal can only displace another metal from a compound if it is located above it in the reactivity series

Example:

  • Magnesium is more reactive than copper, so magnesium can displace copper from a copper sulfate solution to create magnesium sulfate.
    • Magnesium + copper sulfate → magnesium sulfate + copper
    • Mg(s) + CuSO4(aq) → MgSO4(aq) + Cu(s)
  • Platinum, however, is less reactive than copper and so cannot displace copper from a copper sulfate solution.
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Displacement reactions: Halogens

  • Sodium bromide + chlorine → sodium chloride + bromine.
  • 2NaBr + Cl2 → 2NaCl + Br2

Bromide ions

Bromide ions are oxidised (electrons are lost):

  • 2Br- → Br2 +2e-.

Chlorine

Chlorine is reduced (electrons are gained):

  • Cl2 + 2e- → 2Cl-.

Spectator ions

  • The sodium ions do not change in the reaction, so we call them spectator ions.
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Reactions of Reactive Metals with Water

Potassium

  • Potassium is the most reactive so reacts very quickly.
  • The hydrogen produced ignites instantly and the metal also sets alight, sparking and burning with a lilac flame.

Sodium

  • Sodium fizzes rapidly and melts to form a ball that moves around on the water surface.

Lithium

  • Lithium fizzes steadily and floats, becoming smaller until it eventually disappears.
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Reactions of Metals with Dilute Acids

Metals that do react with dilute acids

  • Metals don't have to be that reactive to react with dilute acids.
  • The metals that react with dilute acids but not cold water are magnesium, aluminium, zinc, iron and lead (iron and lead react slowly).

Metals that don't react with dilute acids

  • The only metals that aren't reactive enough are copper, silver, gold and platinum.
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Extraction of Metals

Most metals are only found as compounds because the metal has reacted with other elements in the past. Because of this, the metal has to be extracted from the ore (rock) where the metal compound is found.

  • Metals that are less reactive than carbon can be extracted from their oxides (compounds with oxygen) by reducing with carbon.
  • Carbon is used because it is cheap and abundant.
  • In the reduction, the metal oxide loses oxygen to form a pure metal.
  • Reduction with carbon normally involves heating the metal oxide in the presence of the carbon, which is often used in the form of coal.

Electrolysis is used to extract reactive metals from molten ores (melted materials containing metals). It is used to extract metals such as aluminium, which are more reactive than carbon.

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Reactivity Series - Extraction Using Carbon

  • Potassium down to aluminium are more reactive than carbon and so cannot be extracted by reduction with carbon.
  • Silver, gold and platinum are either found as natural elements or can be simply extracted by heating directly in air.
  • The remaining metals - zinc, iron, and copper - are all commonly extracted by reduction with carbon.
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Reaction of Metals with Acids

When acids react with metals, a hydrogen and a salt are always formed. The first part of the salt's name comes from the metal involved. The second part of the name comes from the acid.

iron + sulfuric acid → iron sulfate + hydrogen

zinc + sulfuric acid → zinc sulfate + hydrogen

magnesium + hydrochloric acid → magnesium chloride + hydrogen

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Neutralisation of Acids

Metal oxides

Acid + oxide → salt + water.

  • E.g. hydrochloric acid + magnesium oxide → magnesium chloride + water.

Metal carbonates

Acid + carbonate → salt + water + carbon dioxide.

  • E.g. hydrochloric acid + calcium carbonate → calcium chloride + water + carbon dioxide

Metal hydroxides

Acid + hydroxide → salt + water.

  • E.g. hydrochloric acid + sodium hydroxide → sodium chloride + water.
  • A neutralisation reaction happens when an acid reacts with an alkali. The hydrogen ions from the acid combine with hydroxide ions from the alkali to produce water.
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Soluble salts

Soluble salts are salts which dissolve in water.

Making soluble salts

  • Soluble salts can be made by reacting acids with solid, insoluble substances such as pure metals, metal oxides, metal hydroxides, or metal carbonates.

Example

Blue copper sulfate crystals are produced by adding black copper oxide to sulfuric acid:

  • CuO(s) + H2SO4(aq) → CuSO4(aq) + H2O(l)
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Required Practical - Separating Mixtures

The aim of this practical is to produce a pure, dry sample of a soluble salt from an insoluble oxide or carbonate.

Method

  1. ­­Gently warm an acid using a Bunsen burner,
  2. Add the insoluble solid (with stirring) until no more reacts.

3. Filter the solution to remove the excess insoluble solid.

  • This will leave a solution of the salt dissolved in water.
  • Heat the solution in an evaporating basin above a beaker of water.
    • The 'water bath' ensures gentle heating.
  • Leave the solution to cool and allow more water to evaporate.
    • As water evaporates, the solution will become more concentrated and the salt will begin to crystallise.
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The pH scale

The pH scale, from 0 to 14, is a measure of acidity or alkalinity:

  • pH of exactly 7 = neutral.
  • pH less than 7 = acid.
    • Strong acids have a pH close to 0.
    • Acids form hydrogen (H+) ions when they dissolve in water.
  • pH more than 7 = alkali.
    • Strong alkalis have a pH close to 14.
    • Alkalis form hydroxide (OH-) ions when they dissolve in water.]

Measuring pH

  • pH probe
  • Universal indicator
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The concentration of H+ ions

The numerical pH scale is a measure of the concentration of H+ ions in a solution. For every increase of 1 in pH, the concentration of H+ ions gets divided by 10:

  • At pH 0, the concentration of H+ ions is 1 mol/dm3.
  • At pH 1, the concentration of H+ ions is 0.1 mol/dm3.
  • At pH 2, the concentration of H+ ions is 0.01 mol/dm3.
  • Etc.
  • The numerical pH scale is a measure of the concentration of H+ ions in a solution. For every increase of 1 in pH, the concentration of H+ ions gets divided by 10.
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Electrolysis

Electrolyte

  • When ionic compounds are melted or dissolved in water, the ions (charged particles that have gained/lost electrons) are free to move around, and the liquid/solution will conduct electricity.
  • Ionic compounds that dissolve in water to make a solution that conducts electricity are called electrolytes.

Electrodes

  • When a voltage (direct current) is applied across an electrolyte, the charged ions are attracted to the electrode with the opposite charge to the ion.When an ion touches an electrode, electrons can be transferred, producing elements.
    • E.g. if copper ions (Cu2+) are in a solution and a voltage is applied, they'll move to the cathode (negative electrode). When they touch the cathode, each Cu2+ ion will gain 2 electrons and form a copper metal.
    • Positively charged ions move to the negative electrode (cathode).
    • Negatively charged ions move to the positive electrode (anode).
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Electrolysis 2

  • Elements

  • When an ion touches an electrode, electrons can be transferred, producing elements.
    • E.g. if copper ions (Cu2+) are in a solution and a voltage is applied, they'll move to the cathode (negative electrode). When they touch the cathode, each Cu2+ ion will gain 2 electrons and form a copper metal.
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Electrolysis example

Electrolysis is used to extract reactive metals from molten ores (melted materials containing metals). It is used to extract metals such as aluminium, which are more reactive than carbon.

What is formed at the anode

  • Carbon dioxide is formed at the anode.
  • The anode is usually made of carbon because it is a good conductor and is cheap.

What is formed at the cathode

  • The aluminium is formed at the cathode.

Disadvantage

  • Lots of energy is needed to:All this energy costs money, and a lot of it!
    • Melt the solid ionic compound to allow the ions to flow.
    • To produce the electrical current.
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Electrolysis of aqueous solutions

The products of the electrolysis of aqueous solutions are difficult to predict because the water molecules in the solution split up to form hydrogen (H+) and hydroxide (OH-) ions.

At the anode

  • What is formed at the anode depends on if halide ions are present:
    • If halide ions are present, the respective halogen forms.
    • If halide ions are absent, oxygen forms.

    At the cathode

  • What is formed at the cathode depends on the reactivity of the metal:

    • If the metal's more reactive than hydrogen, hydrogen is produced.
    • If the metal's less reactive than hydrogen, the metal is produced.
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