Part 4. Some important REACTIONS of important ACIDS

Acids react with a wide range of metals, oxides, hydroxides and carbonates to form salts in neutralisation reactions. The reactions of acids with metals, oxides, hydroxides, carbonates and hydrogencarbonates are described and lots of examples of word and symbol equations. Part 4 Describes and explains the reactions of common acids like hydrochloric acid, sulfuric acid and nitric acid with moderately reactive metals, metal oxides, metal hydroxides, metal carbonates and aqueous ammonia solution. What is formed in these reactions? Are the products of these reactions of any use? These revision notes on chemical reactions of acids e.g. sulfuric, hydrochloric and nitric acids, should prove useful for the new AQA chemistry, Edexcel chemistry & OCR chemistry GCSE (9–1, 9-5 & 5-1) science courses.

EQUATION NOTE: The equations are often written three times: (i) word equation, (ii) balanced symbol equation without state symbols, and, (iii) with the state symbols (g), (l), (s) or (aq) to give the complete balanced symbol equation.

•  Acids are neutralised by reaction with metals, oxides, hydroxides or carbonates to form salts and other products.

•  Apart from metals (which is an electron loss/gain redox reaction), the other reactants listed above are considered as bases, meaning they react by accepting a proton from an acid in forming the salt.

  • Water soluble bases are known as alkalis.

• The reaction between acids and bases/alkalis like oxides, hydroxides and carbonates are classified as neutralisation reactions.

• The first part of the salt name is derived from the positive ion in the base e.g. sodium, magnesium, ammonium etc.

• The second part of the salt name is derived from the acid e.g. chloride (from hydrochloric acid), sulfate (from sulfuric acid), nitrate (from nitric acid) etc.

• You need to be able to predict the products of these reactions and know how to work out the formula of a salt given common ions.

• The ionic theory of neutralisation is described and explained in section 2.

• General word equation: metal + acid ==> a salt + hydrogen

• The salt, and its name, depends on the metal and acid used in the reaction and the acid is neutralised in the process.

• e.g. the grey-silvery solid zinc dissolves in hydrochloric acid with effervescence to evolve hydrogen gas and leave a colourless solution of the salt zinc chloride.

  • zinc + hydrochloric acid ==> zinc chloride + hydrogen

  • Zn + 2HCl ==> ZnCl₂ + H₂

    • Zn(s) + 2HCl(aq) ==> ZnCl₂(aq) + H₂(g)   (equation with state symbols)

  • Its the same equation for many other Group 2 and Transition metals e.g. Mg, Ca and Fe, Co, Ni

  instead of zinc/Zn in the word/symbol equations

  • Note that hydrochloric acid gives a chloride salt.

  • Similarly

    • magnesium + hydrochloric acid ==> magnesium chloride + hydrogen

      • Mg + 2HCl ==> MgCl₂ + H₂

      • in pictures!

    • iron + hydrochloric acid ==> iron(II) chloride + hydrogen

      • Fe + 2HCl ==> FeCl₂ + H₂

    • aluminium + hydrochloric acid ==> aluminium chloride + hydrogen

      • 2Al + 6HCl ==> 2AlCl₃ + 3H₂   (testing equation to balance)

  •  (you can add the state symbols as for the Zn + HCl reaction above)

  • 

  • Illustrated above are two ways in which the zinc - hydrochloric reaction can be used to prepare a sample of hydrogen gas.

• A strip of magnesium ribbon dissolves with effervescence to evolve hydrogen gas and leave a colourless solution of the salt magnesium sulfate.

  • magnesium + sulfuric acid ==> magnesium sulfate + hydrogen

  • Mg + H₂SO₄ ==> MgSO₄ + H₂

    • Mg(s) + H₂SO₄(aq) ==> MgSO₄(aq) + H₂(g)

        (with state symbols)

  • Note that sulfuric/sulfuric gives a sulfate/sulfate salt

  • Similarly ...

  • zinc + sulfuric acid ==> zinc sulfate + hydrogen

  • Zn + H₂SO₄ ==> ZnSO₄ + H₂

    • Zn(s) + H₂SO₄(aq) ==> ZnSO₄(aq) + H₂(g)

  • Instead of Mg or Zn, you can have Fe or Al e.g.

    • iron + sulfuric acid ==> iron(II) sulfate + hydrogen

      • Fe + H₂SO₄ ==> FeSO₄ + H₂

    • aluminium + sulfuric acid ==> aluminium sulfate + hydrogen

      • 2Al + 3H₂SO₄ ==> Al₂(SO₄)₃ + 3H₂   (testing equation to balance!)

  • However some metals give little or no reaction e.g. copper, therefore to make copper salts you need to react the acid with copper oxide or copper carbonate (see below in the following sections for the details of these reactions).

• You can test the gas given off with lit splint and the sound of a squeaky pop confirms it to be hydrogen gas AND therefore the original solid was a metal.

  • 2H₂ + O₂ ==> 2H₂O + heat energy!

    • 2H₂(g) + O₂(g) ==> 2H₂O(l)

  • You might see the condensed water on side of test tube when doing this test.

• Note that nitric acid (HNO₃) doesn't usually form hydrogen with a metal, instead you get nasty brown fumes of nitrogen dioxide! but you still get the metal nitrate salt ....

  • so in general ....

  • metal + nitric acid ==> metal nitrate salt + water + nitrogen oxides (maybe some hydrogen too)

• The reaction of metals with acids is a REDOX reaction and NOT an acid-base reaction.

• See also the REACTIVITY SERIES OF METALS page for the relative rate of reaction of metals with hydrochloric acid and sulfuric acid.

• Naming salts reminder - hydrochloric acid makes chloride salts, sulfuric/sulfuric acid makes sulfate/sulfate salts and nitric acid makes nitrate salts. This is the second part of the name, and the first part of the name (in most cases) is simply the metal name from the metal compound that reacted with the acid. However, sometimes you may need to add a Roman numeral in brackets to indicate the valency of that metal in a particular compound where a metal has a variable valency e.g. copper(II) .., iron(II) ... or an iron(III) salt ... etc.

Advanced REDOX theory of the metal - acid reaction

(this theory does NOT apply to any other reaction on this page because the reactions of acids with oxides, hydroxides or carbonates does NOT involve oxidation and reduction)

Introduction to oxidation and reduction theory and application to REDOX reactions

The reaction between a metal and an acid is technically what is called a REDOX reaction.

This means the reaction takes place in two parts, an oxidation involving electron loss and a reduction involving electron gain.

The metal atoms lose electrons to form positive ions (an oxidation - electron loss - the metal atoms are oxidised).

The hydrogen ions gain electrons forming hydrogen gas molecules (reduction - electron gain - hydrogen ions are reduced).

The two simultaneous changes occur on the surface of the metal where the aqueous positive hydrogen ions hit the surface and pinch electrons from the metal atoms and then metal ions pass into solution.

These changes can be written as half equations and then combined to give the full redox ionic equation.

e.g. for the reaction of magnesium with sulfuric acid, you can write

(i) Mg(s) + H₂SO₄(aq) ==> MgSO₄(aq) + H₂(g)

(ii) Mg(s)  ===>  Mg²⁺(aq)  +  2e^-  (oxidation half equation, 2 electron loss from Mg atom)

(iii) 2H⁺(aq)  +  2e^- ===>  H₂(g)  (reduction half equation, 1 electron gain by 2 H⁺ ions)

adding (ii) + (iii) gives the full (iv) the redox ionic equation ...

(iv) Mg(s)  +  2H⁺(aq)  ===>  Mg²⁺(aq)  +  H₂(g)

... because the two electrons cancel each other out you don't see them in either the 'normal' equation or the ionic equation. The sulfate ion from the acid is the spectator ion, takes no part in the reaction and is not written into the ionic equation.

You can write exactly the same sort of equations, whatever the acid and can also substitute Mg with e.g. Zn or Fe.

• e.g. zinc + hydrochloric acid ==> zinc chloride + hydrogen
  • Zn(s) + 2HCl(aq) ==> ZnCl_2(aq) + H_2(g)
  • the chloride ion Cl^– is the spectator ion
  • zinc + hydrogen ion ==> zinc ion + hydrogen
  • The fully balanced symbol ionic equation is ...
  • Zn(s) + 2H⁺(aq) ==> Zn²⁺(aq) + H₂(g)
  • Zinc atoms are oxidised to zinc ions by electron loss, so zinc is the reducing agent (electron donor).
  • hydrogen ions are the oxidising agent (gaining the electrons) and are reduced to form hydrogen molecules by electron gain.
  • Again, you can think of it as two 'half–reactions'
    • Zn ==> Zn²⁺ + 2e^–       (the oxidation half equation, electron loss, zinc atom is oxidised to the zinc ion)
    • 2H⁺ + 2e^– ==> H₂        (the reduction half equation, electron gain, hydrogen ions reduced to hydrogen molecules)
    • again, the electron loss and gain cancel out, so you don't see the electrons in the ionic equation.
    • Here, the chloride ion is the spectator ion and is not included in the ionic equation.

• The reaction of aluminium with hydrochloric acid or sulfuric acid would be normally given as:

  • aluminium  +  hydrochloric acid  ===>  aluminium chloride  +  hydrogen

  • 2Al  +  6HCl  ===>  2AlCl₃  +  3H₂

  • aluminium  +  sulfuric acid  ===>  aluminium sulfate  +  hydrogen

  • 2Al  + 3H₂SO₄  ===> Al₂(SO₄)₃  +  3H₂

  • I've omitted the state symbols, but still tricky equations to balance!

  • The half reactions are the same because the chloride ions or sulfate ions are spectator ions.

  • The oxidation half equation is: Al  ===> Al³⁺  +  3e^-   (electron loss, aluminium atoms oxidised to aluminium ion)

  • The reduction half equation is: 2H⁺  +  2e^-  ===> H₂   (electron gain, hydrogen ions reduced to hydrogen molecules)

  • To write the overall redox ionic equation by combining the two half equations is a little bit tricky!

  • 2Al  +  6H⁺  ===>  2Al³⁺  +  3H₂

Note that the chloride ion (Cl^-) or sulfate ions (SO₄^2-) don't figure here, they are spectator ions, and don't take part in the reaction, its just the hydrogen ions that are involved from the acid - their formation is shown below.

You can think of hydrochloric acid behaving as: HCl(aq)  ===>  H⁺(aq) + Cl^-(aq)

and for sulfuric acid: H₂SO₄(aq)  ===> 2H⁺(aq)  +  SO₄^2-(aq)

because that is exactly what happens when you dissolve hydrogen chloride and sulfuric acid in water - they ionise and its the hydrogen ion that reacts with the metal surface. The chloride and sulfate ions simply become part of the salt solution formed in the reaction, but didn't actually chemically change like the metal atoms and hydrogen ions do.

See also REDOX reactions and oxidation and reduction for more examples

These may be alkalis (soluble bases) or water insoluble bases

• Metal oxides and metal hydroxides are typical bases, they may be insoluble in water, or soluble in water to give an alkaline solution (a soluble base is called an alkali).

  • An oxide formula just as an 'M' + 'O', a hydroxide formula has an 'M' + 'OH' in various simple proportions.

• The general word equation for this classic 'neutralisation' reaction is

• alkali (soluble/insoluble base) + acid ==> salt + water

• i.e. the acid is neutralised in the process.

• metal hydroxide + acid ==> a salt + water

  • The acid is neutralised in the process e.g.

  • sodium hydroxide + hydrochloric acid ==> sodium chloride + water

  • NaOH + HCl ==> NaCl + H₂O

    • NaOH(aq) + HCl(aq) ==> NaCl(aq) + H₂O(l)   (equation with state symbols)

    • in pictures!

  • Other examples of neutralising alkalis (soluble bases)

    • potassium hydroxide + hydrochloric acid ==> potassium chloride + water

      • KOH + HCl ==> KCl + H₂O

    • potassium hydroxide + hydrobromic acid ==> potassium bromide + water

      • KOH_(aq) + HBr_(aq) ==> KBr_(aq) + H₂O_(l)   (equation with state symbols)

      • Its the same for any Group 1 Alkali Metal hydroxide e.g. LiOH, NaOH, RbOH etc. and any other Group 7 halogen acids e.g. HCl hydrochloric acid.

    • calcium hydroxide + hydrochloric acid ==> calcium chloride + water

      • Ca(OH)₂ + 2HCl ==> CaCl₂ + 2H₂O

  • All solutions involved are colourless but the reaction can be monitored using universal indicator or a pH meter.

  • Soluble bases i.e. here, soluble oxides or hydroxides, are called alkalis.

  • 

  • This neutralisation reaction is used in ...

    • (i) salt preparations method (a) for more details AND

    • (ii) volumetric titrations for quantitative analysis.

    • Other neutralisation reactions of acids reacting with oxides, hydroxides and carbonates are given below.

• or  sodium hydroxide + sulfuric acid ==> sodium sulfate + water

  • 2NaOH + H₂SO₄ ==> Na₂SO₄ + 2H₂O

    • 2NaOH(aq) + H₂SO₄(aq) ==> Na₂SO₄(aq) + 2H₂O(l)   (equation with state symbols)

  • Its the same equation for any Group 1 Alkali Metal hydroxide e.g. LiOH, KOH etc.

  • e.g.

  • sodium hydroxide + nitric acid ==> sodium nitrate + water

  • NaOH + HNO₃ ==> NaNO₃ + 2H₂O

    • NaOH(aq) + HNO₃(aq) ==> NaNO₃(aq) + H₂O(l)

  • potassium hydroxide + nitric acid ==> potassium nitrate + water

  • KOH + HNO₃ ==> KNO₃ + H₂O

    • KOH(aq) + HNO₃(aq) ==> KNO₃(aq) + H₂O(l)

• Another soluble base is ammonia, and here are the equations for its neutralisation with the three most common mineral acids you come across in the school/college laboratory.

  • ammonia + nitric acid ==> ammonium nitrate

    • NH₃ + HNO₃ ==> NH₄NO₃

      • NH_3(aq) + HNO_3(aq) ==> NH₄NO_3(aq)

  • ammonia + hydrochloric acid ==> ammonium chloride

    • NH₃ + HCl ==> NH₄Cl

      • NH_3(aq) + HCl_(aq) ==> NH₄Cl_(aq)

  • ammonia + sulfuric acid ==> ammonium sulfate

    • 2NH₃ + H₂SO₄ ==> (NH₄)₂SO₄

      • 2NH_3(aq) + H₂SO_4(aq) ==> (NH₄)₂SO_4(aq)

• You should have noticed that with hydrochloric acid (HCl) and nitric acid (HNO₃), there is one hydrogen (H) that is replaced by the metal ion, and in the case of sulfuric acid (H₂SO₄), there are two protons replaced by a metal ion or ions. This is because each HCl or HNO₃ molecule produces one hydrogen ion (H⁺) in water, and each sulfuric acid molecule produces two hydrogen ions in water.

  • There is one common mineral acid that potentially can produces three protons per acid molecule.

  • This is phosphoric(V) acid, H₃PO₄, commonly known as just phosphoric acid.

  • Therefore, phosphoric acid, will react with three times as much sodium hydroxide when the solutions are mixed to form the salt trisodium phosphate(V) (or simply 'sodium phosphate').

  • phosphoric acid + sodium hydroxide ==> sodium phosphate + water

    • H₃PO₄ + 3NaOH ==> Na₃PO₄ + 3H₂O

      • H₃PO₄(aq) + 3NaOH(aq) ==> Na₃PO₄(aq) + 3H₂O(l)

• insoluble base + acid ==> salt + water

  • (note: oxides that react with acids to form salts are known as 'basic oxides')

  • The acid is neutralised in the process e.g.

  • e.g. metal oxide + acid ==> salt* + water

  • Black copper (ii) oxide dissolves in colourless sulfuric acid to give a blue solution of copper(II) sulfate from which the salt can be crystallised.

  • copper(II) oxide + sulfuric acid ==> copper(II) sulfate + water

  • CuO + H₂SO₄ ==> CuSO₄ + H₂O

    • CuO(s) + H₂SO₄(aq) ==> CuSO₄(aq) + H₂O(l)

  • Note that insoluble bases, like these insoluble oxides, although reacting with acids to form salts, are NOT alkalis.

  • Instead of copper/Cu, you can have magnesium/Mg, zinc/Zn or nickel/Ni in the word/symbol equations.

  • 

  • This neutralisation reaction is used in salt preparations method b for details.

  • Apart from copper compounds, all solutions involved here are colourless and all the salts form colourless crystal if the solution is carefully evaporated to cause crystallisation.

• calcium hydroxide + hydrochloric acid ==> calcium chloride + water

  • Ca(OH)₂ + 2HCl ==> CaCl₂ + 2H₂O

    • Ca(OH)_2(s) + 2HCl(aq) ==> CaCl₂(aq) + 2H₂O(l)

  • calcium hydroxide + sulfuric acid ==> calcium sulfate + water

    • Ca(OH)₂ + H₂SO₄ ==> CaSO₄ + 2H₂O

      • Ca(OH)₂(s) + H₂SO₄(aq) ==> CaSO₄(aq/s) + 2H₂O(l)

      • This reaction soon slows down because calcium sulfate isn't very soluble in water.

• Other examples ...

  • sulfuric acid + zinc oxide ==> zinc sulfate + water

    • H₂SO₄ + ZnO ==> ZnSO₄ + H₂O

  • magnesium oxide + nitric acid ==> magnesium nitrate + water

    • MgO + 2HNO₃ ==> Mg(NO₃)₂ + H₂O

  • calcium oxide + nitric acid ==> calcium nitrate + water

    • CaO + 2HNO₃ ==> Ca(NO₃)₂ + H₂O

  • magnesium hydroxide + hydrochloric acid ==> magnesium chloride + water

    • Mg(OH)_2(s) + 2HCl_(aq) ==> MgCl_2(aq) + 2H₂O_(l)

    • and ...

    • magnesium hydroxide  +  nitric acid  ===> magnesium nitrate  +  water

    • Mg(OH)_2(s) + 2HNO_3(aq) ==> Mg(NO₃)_2(aq) + 2H₂O_(l)

    • in pictures

    • Similar for many other Group 2 and Transition metal hydroxides  e.g. Ca, Sr, Ba and Fe, Co, Ni, Cu, Zn instead of Mg

  • zinc oxide + hydrochloric acid ==> zinc chloride + water

    • ZnO_(s) + 2HCl_(aq) ==> ZnCl_2(aq) + H₂O_(l)

    • Similar for many other Group 2 and Transition metal oxides  e.g. Mg, Ca, Ba and Co, Ni, Cu instead of Zn

  • copper(II) oxide + hydrochloric acid ==> copper(II) chloride + water

    • CuO_(s) + 2HCl_(aq) ==> CuCl_2(aq) + H₂O_(l)

  • zinc oxide + sulfuric acid ==> zinc sulfate + water

    • ZnO + H₂SO₄ ==> ZnSO₄ + H₂O

      • ZnO(s) + H₂SO₄(aq) ==> ZnSO_4(aq) + H₂O(l)

  • and finally, a bit more tricky!

  • aluminium oxide + sulfuric acid ===> aluminium sulfate + water

  • Al₂O₃ + 3H₂SO₄ ===> Al₂(SO₄)₃ + 3H₂O

  • 

• Apart from copper compounds, all solutions involved here are colourless and all the salts form colourless crystal if the solution is carefully evaporated to cause crystallisation.

• Naming salts reminder - hydrochloric acid makes chloride salts, sulfuric/sulfuric acid makes sulfate/sulfate salts and nitric acid makes nitrate salts.

Reactions of acids with soluble/insoluble carbonates and hydrogencarbonates

(also bases, most insoluble)

• Contrary to what some textbooks may say, but often do not point out, all carbonates are bases and react with acids to form salts, the difference in reaction with alkalis or insoluble bases is that carbon dioxide gas is evolved.

• On adding a solid carbonate or hydrogencarbonate to an acid you see effervescence from carbon dioxide gas and the general word equation is ...

  • metal carbonate/hydrogencarbonate + acid ==> a salt + water + carbon dioxide

  • The obvious extra observation compared to the reaction with oxides or hydroxides is the production of a gas, but the acid is still neutralised in the process to form a salt e.g.

  • The white solid calcium carbonate (limestone) dissolves in dilute hydrochloric acid to form a colourless solution of calcium chloride and colourless carbon dioxide gas.

  • This is a typical reaction between an insoluble carbonate and a dilute acid solution.

  • calcium carbonate + hydrochloric acid ==> calcium chloride + water + carbon dioxide

    • CaCO₃ + 2HCl ==> CaCl₂ + H₂O + CO₂

      • CaCO₃(s) + 2HCl(aq) ==> CaCl₂(aq) + H₂O(l) + CO₂(g)

        (equation with state symbols)

    •  

    • Illustrated above are two ways in which the limestone chips (calcium carbonate) - hydrochloric acid reaction can be used to prepare a sample of carbon dioxide gas.

  • 

  • This neutralisation reaction is used in salt preparations (see method b for details).

  • Apart from copper compounds, all solutions involved here are colourless and all the salts form colourless crystal if the solution is carefully evaporated to cause crystallisation.

• The dark turquoise-green solid copper(II) carbonate dissolves in hydrochloric acid to form a greeny-blue solution of copper(II) chloride and effervescence from the carbon dioxide formed.

  • Another reaction between an insoluble carbonate and a dilute acid solution.

  • copper(II) carbonate + hydrochloric acid ==> Copper(II) chloride + water + carbon dioxide

    • CuCO₃ + 2HCl ==> CuCl₂ + H₂O + CO₂

      • CuCO₃(s) + 2HCl(aq) ==> CuCl₂(aq) + H₂O(l) ₊ CO₂(g)

    • and with sulfuric acid a blue solution of copper(II) sulfate is formed and with nitric acid a blue solution of copper(II) nitrate - word and full symbol equations below.

  • copper(II) carbonate + sulfuric acid ==> Copper(II) sulfate + water + carbon dioxide

    • CuCO₃ + H₂SO₄ ==> CuSO₄ + H₂O + CO₂

      • CuCO₃(s) + H₂SO₄(aq) ==> CuSO₄(aq) + H₂O(l) + CO₂(g)

      • in pictures!

  • copper(II) carbonate + nitric acid ==> Copper(II) nitrate + water + carbon dioxide

    • CuCO₃ + 2HNO₃ ==> Cu(NO₃)₂ + H₂O + CO₂

      • CuCO₃(s) + 2HNO₃(aq) ==> CuSO₄(aq) + H₂O(l) + CO₂(g)

  • Similar equations for other carbonates

    • calcium carbonate CaCO₃, to make three salts - calcium chloride/nitrate (calcium sulfate is not very soluble)

    • iron(II) carbonate FeCO₃, to make three salts - iron(II) chloride/sulfate/nitrate

    • magnesium carbonate MgCO₃, to make three salts - magnesium chloride/sulfate/nitrate

    • manganese(II) carbonate MnCO₃, to make three salts - manganese(II) chloride/sulfate/nitrate

    • zinc carbonate ZnCO₃, to make three salts - zinc chloride/sulfate/nitrate

    • lead(II) carbonate PbCO₃, only nitric acid to make lead(II) nitrate

• Similarly, but forming colourless solutions from other white insoluble solid carbonates  ...

  • magnesium carbonate + hydrochloric acid ==> magnesium chloride + water + carbon dioxide

    • MgCO₃ + 2HCl ==> MgCl₂ + H₂O + CO₂

      • MgCO₃(s) + 2HCl(aq) ==> MgCl₂(aq) + H₂O(l) + CO₂(g)

  • or calcium carbonate + nitric acid ==> calcium nitrate + water + carbon dioxide

  • CaCO₃ + 2HNO₃ ==> Ca(NO₃)₂ + H₂O + CO₂

    • CaCO₃(s) + 2HNO₃(aq) ==> Ca(NO₃)₂(aq) + H₂O(l) ₊ CO₂(g)

  • Its the same equation for many other Group 2 and Transition metals e.g. Mg, Sr and Co, Ni, Cu

    • Test for carbon dioxide gas - it gives a white precipitate of calcium carbonate (cloudiness) when bubbled into limewater (calcium hydroxide solution).

    • Ca(OH)₂ + CO₂ ==> CaCO₃ + H₂O

      • Ca(OH)₂(aq) + CO₂(g) ==> CaCO₃(s) + H₂O(l)

  • Note: Using sulfuric acid and calcium carbonate you don't get much of a fizz! because the calcium sulfate salt formed, is not very soluble, and coats the remaining calcium carbonate inhibiting the reaction! This will happen with any reaction between an acid and a water insoluble reactant which forms an insoluble solid product!

• magnesium carbonate + sulfuric acid ==> magnesium sulfate + water + carbon dioxide

  • MgCO₃ + H₂SO₄ ==> MgSO₄ + H₂O + CO₂

    • MgCO₃(s) + H₂SO₄(aq) ==> MgSO₄(aq) + H₂O(l) + CO₂(g)

• AND six equations for sodium carbonate and sodium hydrogencarbonate. Both are actually soluble in water, but in each case you are likely to add the white solid directly into the acid where it dissolves to give a colourless solution of the colourless salt with the evolution of carbon dioxide gas (fizzing - effervescence) e.g....

  • sodium hydrogencarbonate + hydrochloric acid ==> sodium chloride + water + carbon dioxide

    • NaHCO₃ + HCl ==> NaCl + H₂O + CO₂

      • NaHCO₃(s) + HCl(aq) ==> NaCl(aq) + H₂O(l) + CO₂(g)

      • Note that hydrogencarbonate is commonly known as 'bicarbonate'.

      • Its the same for any Group 1 Alkali Metal hydrogencarbonate  e.g. for Li, K etc. instead of Na

  • sodium hydrogencarbonate + nitric acid ==> sodium nitrate + water + carbon dioxide

    • NaHCO₃ + HNO₃ ==> NaNO₃ + H₂O + CO₂

      • NaHCO₃(s) + HNO₃(aq) ==> NaNO₃(aq) + H₂O(l) + CO₂(g)

  • sodium carbonate + hydrochloric acid ==> sodium chloride + water + carbon dioxide

    • Na₂CO₃ + 2HCl ==> 2NaCl + H₂O + CO₂

      • Na₂CO₃(s) + 2HCl(aq) ==> 2NaCl(aq) + H₂O(l) + CO₂(g)

  • sodium hydrogencarbonate + sulfuric acid ==> sodium sulfate + water + carbon dioxide

    • 2NaHCO₃ + H₂SO₄ ==> Na₂SO₄ + 2H₂O + 2CO₂

      • 2NaHCO₃(s) + H₂SO₄(aq) ==> Na₂SO₄(aq) + 2H₂O(l) + 2CO₂(g)

  • sodium carbonate + sulfuric acid ==> sodium sulfate + water + carbon dioxide

    • Na₂CO₃ + H₂SO₄ ==> Na₂SO₄ + H₂O + CO₂

      • Na₂CO₃(s) + H₂SO₄(aq) ==> Na₂SO₄(aq) + H₂O(l) + CO₂(g)

  • sodium carbonate + nitric acid ==> sodium nitrate + water + carbon dioxide

    • Na₂CO₃ + 2HNO₃ ==> 2NaNO₃ + H₂O + CO₂

      • Na₂CO₃(s) + 2HNO₃(aq) ==> 2NaNO₃(aq) + H₂O(l) + CO₂(g)

  • phosphoric acid + sodium carbonate ==> sodium phosphate + water + carbon dioxide

    • 2H₃PO₄ + 3Na₂CO₃ ==> 2Na₃PO₄ + 3H₂O + 3CO₂

      • 2H₃PO₄(aq) + 3Na₂CO₃(aq) ==> 2Na₃PO₄(aq) + 3H₂O(l) + 3CO₂(g)

      • That's a really tricky equation to balance!!!

• You can test the gas given off with limewater and the formation of a white precipitate confirms it to be carbon dioxide AND therefore the original solid was a carbonate or hydrogencarbonate.

• For more reactions of acids and carbonates see Limestone Chemistry page

• Naming salts reminder - hydrochloric acid makes chloride salts, sulfuric/sulfuric acid makes sulfate/sulfate salts and nitric acid makes nitrate salts.

• Ammonia gas is very soluble in water to form an alkaline solution that can be neutralised by acids to form ammonium salts. All solutions involved here are colourless and all the salts form colourless crystal if the solution is carefully evaporated to cause crystallisation.

• ammonia + acid ==> ammonium salt

  • Note that no water is formed and see also note (c) below.

  • e.g. (i) ammonia + hydrochloric acid ==> ammonium chloride

    • NH₃ + HCl ==> NH₄Cl

      • NH₃(aq) + HCl(aq) ==> NH₄Cl(aq)

  • or (ii) ammonia + nitric acid ==> ammonium nitrate

    • NH₃ + HNO₃ ==> NH₄NO₃

      • NH₃(aq) + HNO₃(aq) ==> NH₄NO₃(aq)

  • or (iii) ammonia + sulfuric acid ==> ammonium sulfate

    • 2NH₃ + H₂SO₄ ==> (NH₄)₂SO₄

      • 2NH₃(aq) + H₂SO₄(aq) ==> (NH₄)₂SO₄(aq)

• NOTE that

  • (a) NH₃ is used in equations, NOT NH₄OH which doesn't exist!

  • Therefore, theoretically, no water is produced when ammonia solution is neutralised with acids, BUT it is still a neutralisation reaction because ammonia is a base/alkali.

  • (b) All these ammonium salts are colourless crystalline solids - formed if the water is carefully evaporated salt preparation method (a) for details and on ammonium salts page for details.

  • (c) Reactions (ii) and (iii) are used to make fertiliser salts - see ammonia chemistry and uses.

• Naming salts reminder - hydrochloric acid makes chloride salts, sulfuric acid makes sulfate salts and nitric acid makes nitrate salts.

NOTE (a): As already mentioned, and to summarise, the name of the particular salt formed depends on (i) the metal name, which becomes the first part of salt name, and (ii) the acid e.g. H₂SO₄ sulfuric acid on neutralisation makes a ... sulfate; HCl hydrochloric acid makes a ... chloride;  HNO₃ nitric acid makes a ... nitrate etc.

NOTE (b): There is a list of compound formulae and their solubility in section 8. The first part of the salt name is ammonium derived from ammonia (with metals or their compounds the metal retains its original name), but the second part of the salt name is always derived from the acid as in NOTE (a) above.

NOTE (c): Ammonia is an alkaline gas that is very soluble in water. It is a weak alkali or soluble base and is readily neutralised by acids in solution to form ammonium salts which can be crystallised on evaporating the resulting solution. Sometimes the equations are written with the 'fictitious' 'ammonium hydroxide'

e.g. NH₄OH(aq) + HCl(aq) ==> NH₄Cl(aq) + H₂O(l)

but this is not considered to be a correct representation of this neutralisation reaction these days!

NOTE (d): An extensive structured question on acid reaction equations.

In the formula of the salt the total positive charge must equal the negative charge i.e. the salt must be overall electrically neutral.

A list of common positive ions (cations e.g. from the base) and negative ions (e.g. from the acid) is given on the right.

In the examples below of salt formulae, the derived formula are shown in 'molecular' formula style, but they are actually ionic compounds, so the ionic formula is also shown.

For magnesium chloride: 1 of Mg²⁺ balances 2 of Cl^- because 1 x 2 = 2 x 1 gives MgCl₂ or Mg²⁺(Cl^-)₂

For iron(III) chloride: 1 of Fe³⁺ balances 3 of Cl^- because 1 x 3 = 3 x 1 gives FeCl₃ or  Fe³⁺(Cl^-)₃

For calcium nitrate: 1 of Ca²⁺ balances 2 of NO₃^- because 1 x 2 = 2 x 1 gives Ca(NO₃)₂ or Ca²⁺(NO₃^-)₂ 

For iron(III) sulfate: 2 of Fe³⁺ balances 3 of SO₄^2- because 2 x 3 = 3 x 2 gives Fe₂(SO₄)₃ or (Fe³⁺)₂(SO₄^2-)₃

For more on equations and formulae see: How to write equations, work out formula and name compounds

GCSE/IGCSE Acids & Alkalis revision notes sub–index: Index of all pH, Acids, Alkalis, Salts Notes 1. Examples of everyday acids, alkalis, salts, pH of solution, hazard warning signs : 2. pH scale, indicators, ionic theory of acids–alkali neutralisation : 4. Reactions of acids with metals/oxides/hydroxides/carbonates, neutralisation reactions : 5. Reactions of bases–alkalis like ammonia & sodium hydroxide : 6. Four methods of making salts : 7. Changes in pH in a neutralisation, choice and use of indicators : 8. Important formulae of compounds, salt solubility and water of crystallisation : 10. More on Acid–Base Theory and Weak and Strong Acids

See also Advanced Level Chemistry Students Acid-Base Revision Notes - use index

• GCSE/IGCSE foundation-easier multiple choice quiz on pH, Indicators, Acids, Bases, Neutralisation and Salts

• GCSE/IGCSE higher-harder multiple choice quiz on pH, Indicators, Acids, Bases, Neutralisation and Salts

• GCSE/IGCSE Structured question worksheet on Acid Reaction word equations and symbol equation questions

•  Word equation answers and symbol equation answers)

• GCSE/IGCSE word-fill worksheet on Acids, Bases, Neutralisation and Salts

• GCSE/IGCSE matching pair quiz on Acids, Bases, Salts and pH

• See also Advanced Level Chemistry Students Acid-Base Revision Notes - use index

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