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GCSE Chemistry Notes: The chemical reactions of common mineral acids
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4. Important Reactions of Common Acids (Part 4 includes redox half equations for the metal-acid reaction and how to work out the formula of a salt given the constituent ions)
Index of all my GCSE notes on acids, bases and salts All my GCSE Chemistry Revision notes This is a BIG website, you need to take time to explore it [ SEARCH BOX]Use your mobile phone or ipad etc. in 'landscape' mode 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. 4. Some important reactions of Acids
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) Note that sulfuric/sulfuric gives a sulfate/sulfate salt Similarly ... zinc + sulfuric acid ==> zinc sulfate + hydrogen Zn + H2SO4 ==> ZnSO4 + H2
Instead of Mg or Zn, you can have Fe or Al e.g.
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).
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
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.
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
You can write exactly the same sort of equations, whatever the acid and can also substitute Mg with e.g. Zn or Fe.
Note that the chloride ion (Cl-) or sulfate ions (SO42-) 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) + SO42-(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 REACTION OF ACIDS WITH BASES - basic oxides and hydroxides These may be alkalis (soluble bases) or water insoluble bases
Reactions of acids with soluble bases (alkalis, usually soluble metal hydroxides)
or sodium hydroxide +
sulfuric acid ==> sodium sulfate +
water 2NaOH + H2SO4
==> Na2SO4 + 2H2O 2NaOH(aq) + H2SO4(aq)
==> Na2SO4(aq) + 2H2O(l) 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
potassium hydroxide + nitric acid ==> potassium nitrate + water KOH + HNO3 ==> KNO3 + H2O
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.
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.
Reactions of acids with insoluble bases (metal oxides and metal hydroxides)
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
Other examples ...
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)
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.
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
copper(II) carbonate + nitric acid ==> Copper(II) nitrate + water + carbon dioxide
Similar equations for other carbonates
Similarly, but forming colourless solutions from other white insoluble solid carbonates ...
or calcium carbonate + nitric acid ==> calcium nitrate + water + carbon dioxide
Its the same equation for many other Group 2 and Transition metals e.g. Mg, Sr and Co, Ni, Cu
Ca(OH)2 + CO2 ==> CaCO3 + H2O
magnesium carbonate + sulfuric acid ==> magnesium sulfate + water + carbon dioxide
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 + 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
sodium hydrogencarbonate + sulfuric acid ==> sodium sulfate + water + carbon dioxide
sodium carbonate + sulfuric acid ==> sodium sulfate + water + carbon dioxide
sodium carbonate + nitric acid ==> sodium nitrate + water + carbon dioxide
phosphoric acid + sodium carbonate ==> sodium phosphate + water + carbon dioxide
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
(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'
Appendix How to work out the formula of a salt given the ions 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 potassium chloride: 1 of K+ balances 1 of Cl- because 1 x 1 = 1 x 1 gives KCl or K+Cl-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 sodium sulfate: 2 of Na+ balances 1 of SO42- because 2 x 1 = 1 x 2 gives Na2SO4 or (Na+)2SO42-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 SO42- because 2 x 3 = 3 x 2 gives Fe2(SO4)3 or (Fe3+)2(SO42-)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
Multiple choice revision quizzes and other worksheets
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