Download - C10 acids, bases and salts
LEARNING OUTCOMES
Define acid and acid anhydrideInvestigate the reactions of non-oxidising acids with metals, carbonates, hydrogen carbonates and basesDefine base and alkaliInvestigate the reaction of bases with ammonium saltsRelate acidity and alkalinity to the pH scaleDiscuss the strength of acids and alkalis on the basis of their completeness of ionisationDefine acidic, basic, amphoteric and neutral oxides
Chapter 10
Acids, Bases and Salts
LEARNING OUTCOMES
Define saltIdentify an appropriate method of salt preparation based on the solubility of the saltDistinguish between acidic and normal saltsInvestigate neutralisation reactions using indicators and temperature changes
Chapter 10
Acids, Bases and Salts
Chapter 10
Acids, Bases and Salts
What are acids? Fruits like apples, oranges and pineapples taste sour because they
contain acids.
Acids also turn blue litmus paper red.
Acids produce hydrogen ions H+ in water.
An acid is a substance which produces hydrogen ions, H+(aq) in water.
Definition of An Acid
For example, hydrochloric acid dissolves in water to form hydrogen ions and chloride ions:
HCl(aq) H+(aq) + Cl-(aq) It is the hydrogen ions which turn blue litmus to red and give acids their characteristic properties.
Chapter 10
Acids, Bases and Salts
Acids react with metals to produce hydrogen gas.E.g. Mg + H2SO4 MgSO4 + H2
Other chemical properties of acids
Acids react with carbonates to produce carbon dioxide.E.g. CaCO3 +2HCl CaCl2 + H2O + CO2
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What are acids?
( test for hydrogen gas)
(test for carbon dioxide)
Limewater turns chalky
HCl+CaCO3
pop
Other chemical properties of acids Acids react with bases to form a salt and water only. E.g. sulphuric acid reacts with copper(II) oxide to form a salt
called copper(II) sulphate and water:H2SO4 + CuO CuSO4 + H2O
This reaction is called neutralisation.
What are acids?
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A Strong Acid
A strong acid is an acid that is completely ionised in water. This means that all the acid molecules become ions in the water.
Examples of strong acids are: sulphuric acid, hydrochloric Examples of strong acids are: sulphuric acid, hydrochloric acid and nitric acid.acid and nitric acid.
Strong acid
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A Weak Acid
E.g.s. of weak acids are: ethanoic acid, citric acid and carbonic acid.
Weak acid
A weak acid is an acid that is only partially ionised in water. This means that only a few molecules of the acid become ions in water.
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Some Common Acids
Name of acid Formula
Sulphuric acid H2SO4
Hydrochloric acid HCl
Nitric acid HNO3
Citric acid C6H8O7
Ethanoic acid (vinegar) CH3COOH
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Ethanoic acid is used in vinegar for cooking and to preserve food such as vegetables.
Uses of Acids
Hydrochloric acid is used in the industry to remove rust from metals before they are painted.
Sulphuric acid is used to make fertilisers and detergents.
Citric acid is used in making fruit salts.
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Quick check 11. What ions do acids produce in water?
2. State three properties of acids.
3. Explain what is meant by a strong acid. Give one example of a strong acid.
4. Explain what is meant by a weak acid. Give one example of a weak acid.
5. Some dry citric acid crystals are placed on a dry piece of litmus paper. Will there be a colour change? Explain your answer.
Solution
Chapter 10
Acids, Bases and Salts
Solution to Quick check 1
1. Hydrogen ions
2. (a) Acids have a sour taste.(b) Acids turn blue litmus to red.(c) Acids react with metals to produce hydrogen.
3. A strong acid is an acid that is completely ionised in water. E.g. sulphuric acid.
4. A weak acid is an acid that is only partially ionised in water. E.g. ethanoic acid.
5. There will be no colour change because there is no water, so the citric acid cannot form hydrogen ions.
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Bases A base is an oxide or hydroxide of a metal. Examples of bases are:
sodium oxide, sodium hydroxide, copper(II) oxide, copper(II) hydroxide, etc.
A base reacts with an acid to form a salt and water only. E.g. CuO + H2SO4 CuSO4 + H2O
This process is called neutralisation.
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If a base is soluble in water, it is called an alkali.
Alkalis
Sodium hydroxide is an alkali because it dissolves in water to produce hydroxide ions:NaOH(aq) Na+(aq) + OH−(aq)
An alkali is a soluble base which produces hydroxide ions, OH− (aq) in water.
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Copper(II) hydroxide is a base but not an alkali. This is because it is insoluble in water and hence cannot produce hydroxide ions in water.
Difference between base and alkali
BASEALKALICuO
MgOCa(OH)2
NaOH KOH NH3(aq)
Fe2O3
Cu(OH)2
Is this true?All alkalis are bases, but not all bases are alkalis.
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Alkalis have a bitter taste and soapy feel. Alkalis turns red litmus to blue.
Chemical properties of alkalis
Alkalis react with acids to from salt and water only.E.g. 1. NaOH + HCl NaCl + H2O
E.g. 2 2KOH + H2SO4 K2SO4 + 2H2O
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Alkalis react with ammonium salts to produce ammonia gas. Ammonia gas is acidic, thus it turns red litmus paper blue. Ammonia gas is very soluble in water and gives out a pungent
smell.E.g.1: NaOH + NH4Cl NaCl + NH3 + H2O
Chemical properties of alkalis
Sodium hydroxide + ammonium chloride
E.g. 2: Ca(OH)2 + 2NH4Cl CaCl2 + 2NH3 + 2H2O
NH3 gas produced turns red litmus blue
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Sodium hydroxide and potassium hydroxide are used in making soaps.
Uses of Bases
Ammonia solution is used in window cleaners. Magnesium hydroxide is used in toothpastes to neutralise
the acid produced by bacteria. Calcium hydroxide (slaked lime) is used to neutralise
acids found in acidic soil.
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Some Common Alkalis
Name Chemical formula
Sodium hydroxide NaOH
Potassium hydroxide KOH
Calcium hydroxide Ca(OH)2
Ammonia solution (ammonium hydroxide)
NH3(aq)
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Quick check 2
1. What is a base? Give 3 examples of bases.2. Define what is an alkali. Give 3 examples of alkalis.3. State 3 properties of alkalis.4. Explain why iron(II) hydroxide is a base, but not an alkali.5. Write balanced chemical equations for the following
reactions:(a) potassium hydroxide + ammonium chloride(b) calcium hydroxide + ammonium chloride
Solution
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Solution to Quick check 2
1. A base is an oxide or hydroxide of a metal. E.g. sodium oxide, copper(II) oxide, calcium hydroxide.
2. An alkali is a soluble base which produces hydroxide ions in water. E.g. sodium hydroxide, potassium hydroxide, calcium hydroxide.
3. (i) Alkalis turn red litmus blue.(ii) Alkalis react with acids to produce a salt and water.(iii) Alkalis react with ammonium salts to produce ammonia.
4. Iron(II) hydroxide is a base, but not an alkali because it is insoluble in water, so it cannot produce hydroxide ions in water.
5. (a) KOH + NH4Cl KCl + H2O + NH3
(b) Ca(OH)2 + 2NH4Cl CaCl2 + 2H2O + 2NH3
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Indicators
Indicators are substances which show different colours in acidic and alkaline solutions.
Litmus is a common indicator. It is red in acidic solutions and blue in alkaline solutions.
Other important indicators are shown in the table on the next slide.
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Indicators
Indicator Colour in strong Acids
pH at which colour changes
Colour in strong alkalis
Methyl orange red pH 4 yellow
Litmus red pH 7 blue
Phenolphthalein colourless pH 9 pink
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The pH of a solution tells us how acidic or alkaline a solution is.
The pH is a measurement of the hydrogen ion concentration in a solution.
The pH scale ranges from 0 to 14. The pH of a solution can be measured with a pH meter.
The pH Scale
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The lower the pH, the more acidic the solution is. The higher the pH, the more alkaline the solution is. pH 7 is neutral. Distilled water, sugar solution and most salt solutions are
neutral (pH 7).
The pH Scale
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The Universal Indicator consists of a mixture of dyes which changes its colour in different pH solutions.
We can use the Universal Indicator to tell us the approximate pH of a solution.
The Universal Indicator or pH paper changes its colour according to the pH shown in the chart below.
The Universal Indicator
Box of pH paper with colour chart
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Types of Oxides
Elements burn or react with oxygen to form oxides. There are 4 types of oxides: acidic oxides, basic oxides, amphoteric
oxides and neutral oxides. An acidic oxide is an oxide of a non-metal. It dissolves in water to form an
acid. Acidic oxides react with alkalis to form salts . A basic oxide is an oxide of a metal. If soluble, it will dissolve in water to
form an alkali. Basic oxides react with acids to form salts. An amphoteric oxide is an oxide which can react with both acids and
alkalis to form salts. A neutral oxide does not react with either acids or alkalis.
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Types of Oxides
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Acidic Oxides
Basic Oxides
Amphoteric Oxides
CO2 , SO2
NO2 , NONa2O, CaO, K2O,
MgO, CuOAl2O3 , PbO ,
ZnO
React with alkalis to form
salts
React with acids to form salts
React with both acids & alkalis to
form salts
Neutral Oxides
H2O, CO , N2O
Do not react with both acids &
alkalis
4 TYPES OF OXIDES
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Quick check 31. Name 3 common indicators and their colour change in strong
acidic and strong alkaline solutions.2. What is meant by the pH of a solution? What is the pH of :
(a) hydrochloric acid, (b) citric acid, (c) sodium chloride solution, (d) sodium hydroxide solution?
3. What are the 4 types of oxides? Give one example of each type of oxide.
4. What colours would you expect to see when the following indicators are added to a solution of pH 5?(a) litmus, (b) phenolphthalein, (c) methyl orange
Solution
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Solution to Quick check 3
1. Litmus: red, blue; Phenolphthalein: colourless, pink; Universal Indicator: red, violet
2. The pH of a solution measures the acidity or alkalinity of a solution. (a) 0 – 1, (b) 3 – 4, (c) 7, (d) 13 – 14.
3. Acidic oxides, basic oxides, amphoteric oxides and neutral oxides. E.g. sulphur dioxide, sodium oxide, aluminium oxide, water.
4. (a) litmus: red, (b) phenolphthalein: colourless, (c) methyl orange: yellow
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A salt is formed when an acid is neutralised by a base.
A salt contains two parts: Metal part : cation (comes from the
base) Non-metal part : anion (comes from
the acid)
Salts
+Acid Base
Salt
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Examples of Salts
Base (alkali) Acid Salt formed
Sodium hydroxide Hydrochloric acid Sodium chloride
Potassium hydroxide Hydrochloric acid Potassium chloride
Sodium hydroxide Sulphuric acid Sodium sulphate
Potassium hydroxide Sulphuric acid Potassium sulphate
Calcium hydroxide Nitric acid Calcium nitrate
Ammonia solution Nitric acid Ammonium nitrate
Table 1
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Sodium chloride is used as table salt and to preserve meat and vegetables.
Sodium chloride is electrolysed to obtain sodium and chlorine in the industry.
Ammonium nitrate and ammonium sulphate are used as plant fertilisers.
Uses of Salts
Magnesium sulphate, commonly called Epsom salt, is used as a bath-salt.
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Methods of Preparing Salts
ACID + ALKALI SALT + WATER
1. Action of acid on alkali
This process is called neutralisation.
To carry out the neutralisation of the acid and alkali exactly, a method called titration is used.
The salts listed in Table 1 can be prepared by the titration method.
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To prepare sodium nitrate by neutralisation (titration method)
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Sodium nitrate and water (phenolphthalein as indicator)
burette
Pipette
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ACID + BASE SALT + WATER
2. Action of acid on insoluble base
This method is used for bases which are insoluble in water.
Examples of salts prepared by this method: * copper(II) sulphate from copper(II) oxide and sulphuric acid:
CuO + H2SO4 CuSO4 + H2O
* zinc chloride from zinc oxide and hydrochloric acid:ZnO + 2HCl ZnCl2 + H2O
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Methods of Preparing Salts
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Preparation of copper(II) sulphate (acid on insoluble base)
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Step 1 Place about 50 cm³ of dilute sulphuric acid in a beaker and gently warm the acid. Copper(II) oxide is added, a little at a time, to the acid, until no more can dissolve.
Equation: CuO + H2SO4 CuSO4 + H2O
Step 2 Filter off the excess copper(II) oxide using a filter paper and funnel. Collect the filtrate which contains copper(II) sulphate in an evaporating dish.
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Preparation of copper(II) sulphate (acid on insoluble base)
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Acids, Bases and Salts
Step 3 Evaporate the copper(II) sulphate solution until it is saturated. Allow the hot solution to cool to form crystals.
Step 4 Filter off the copper(II) sulphate crystals formed and dry them by pressing them between sheets of filter paper.
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Eg.1 Sulphuric acid on sodium carbonate H2SO4 + Na2CO3 Na2SO4 + H2O + CO2
Eg.2 Hydrochloric acid on calcium carbonate 2HCl + CaCO3 CaCl2 + H2O + CO2
This method is similar to the previous method; instead of the oxide, the carbonate is added in excess to the acid.
3. Action of acid on a carbonate
ACID + CARBONATE SALT + WATER + CO2
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Methods of Preparing Salts
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Eg.1 Sulphuric acid on zinc H2SO4 + Zn ZnSO4 + H2
Eg.2 Hydrochloric acid on magnesium 2HCl + Mg MgCl2 + H2
NOTE: Only metals like magnesium, zinc and iron are suitable. Metals like sodium, potassium and calcium are explosive with acids; while metals like lead and copper are unreactive with acids.
4. Action of acid on a metal
ACID + METAL SALT + HYDROGEN
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Methods of Preparing Salts
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Making zinc sulphate (acid on metal)
Chapter 10
Acids, Bases and Salts
Can you describe how zinc sulphate is prepared with the aid of the diagrams?
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5. Double Displacement (Precipitation method)
This method is used to prepare insoluble salts. Two solutions are mixed together to produce a precipitate of the insoluble salt which can then be filtered off from the mixture.
+AD (s)
AB (aq) CD (aq)
CB (aq)
E.g. Lead(II) nitrate + Sodium chloride Lead(II) chloride + Sodium nitrate Pb(NO3)2(aq) + 2NaCl(aq) PbCl2(s) + 2NaNO3(aq)
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Methods of Preparing Salts
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Silver chlorideAgNO3(aq) + HCl(aq) AgCl(s) + HNO3(aq)
Barium sulphateBa(NO3)2(aq) + H2SO4(aq) BaSO4(s) + 2HNO3(aq)
Copper(II) carbonate CuSO4(aq) + Na2CO3(aq) CuCO3(s) + Na2SO4(aq)
Other salts made by precipitation method
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Table of soluble and insoluble salts
Soluble salts Insoluble salts
All sodium, potassium and ammonium salts
All carbonates except those of sodium, potassium and ammonium
All nitrates None
All sulphates except those of calcium, lead and barium
Calcium sulphate, lead(II) sulphate and barium sulphate
All chlorides except those of silver and lead
Silver chloride and lead(II) chloride
This table will be useful to you when preparing salts
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Quick check 41. Define what is salt. Give an example of a soluble and insoluble
salt.2. State 4 methods of making salts. 3. State whether the following salts are soluble or insoluble:
(a) sodium carbonate, (b) calcium chloride, (c) barium sulphate, (d) lead(II) nitrate, (e) lead(II) chloride.
4. State the method you would choose to prepare the following salts:(a) potassium nitrate, (b) zinc nitrate, (c) magnesium sulphate, (d) copper(II) carbonate. For each method, state the chemicals you will need and write a balanced chemical equation for the reaction.
Solution
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Solution to Quick check 4
1. A salt is formed when an acid is neutralised by a base. E.g. soluble salt: sodium chloride E.g. insoluble salt: calcium sulphate
2. (a) Acid on metal, (b) acid on base, (c) acid on carbonate, (d) precipitation method
3. Soluble: sodium carbonate, calcium chloride, lead(II) nitrate; Insoluble: lead(II) chloride, barium sulphate
4. (a) potassium nitrate: titration method; potassium hydroxide and nitric acid; KOH + HNO3 KNO3 + H2O (b) zinc nitrate: acid on carbonate; nitric acid and zinc carbonate; 2HNO3 + ZnCO3 Zn(NO3)2 + H2O + CO2
(c) magnesium sulphate: acid on metal; magnesium and sulphuric acid; Mg + H2SO4 MgSO4 + H2
(d) copper(II) carbonate: precipitation method; copper(II) sulphate and sodium carbonate; CuSO4(aq) + Na2CO3(aq) CuCO3(s) + Na2SO4(aq)
Chapter 10
Acids, Bases and Salts
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The state symbols in a chemical equation tell us about the state of each reactant and product.
The following are the state symbols used: Solid (s) Liquid (l) Gas (g) Aqueous solution (aq)
Example: CaCO3(s) + 2HCl(aq) CaCl2(aq) + H2O(l) + CO2(g)
The above equation tells us that solid calcium carbonate reacts with a solution of hydrochloric acid to produce liquid water and carbon dioxide gas.
State symbols in equations
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Ionic equations are general equations which can apply to any particular reaction.
They represent ions taking part in a reaction, leaving out those ions which do not react (spectator ions).
They contain state symbols. Only solutions (aq) can form ions; gases, solids and liquids
do not ionise.
Writing ionic equations
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Steps in writing ionic equations
Step 3: Rewrite the equation with the final ions left: H+ (aq) + OH- (aq) H2O(l)
EXAMPLE 1
HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l)
Step 1: Break substances with (aq) into its ions:H+
(aq) + Cl-(aq) + Na+ (aq) + OH-
(aq) Na+ (aq) + Cl- (aq) + H2O (l)
Step 2: Remove similar ions from both sides of equation.
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Writing ionic equations
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EXAMPLE 2
2HCl(aq) + CaCO3 (s) CaCl2 (aq) + H2O (l) + CO2 (g)
Step 1: Break those with (aq) into its ions:2H+ (aq) + 2Cl-(aq) + CaCO3 (s) Ca2+ (aq) + 2Cl- (aq) + H2O (l) + CO2 (g)
Step 2: Remove similar ions on both sides.
Step 3: Rewrite the equation with the ions left: 2H+(aq) + CaCO3(s) Ca2+(aq) + H2O(l) + CO2(g)
Steps in writing ionic equations
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Writing ionic equations
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EXAMPLE 3
Pb(NO3)2(aq) + 2NaCl (aq) PbCl2 (s) + 2NaNO3 (aq)
Step 1: Break those with (aq) into its ions:Pb2+
(aq) + 2NO3-(aq) + 2Na+
(aq) + 2Cl- (aq) PbCl2(s) + 2Na+(aq) + 2NO3
- (aq)
Step 2: Remove similar ions on both sides.
Step 3: Rewrite the equation with the ions left:Pb2+(aq) + 2Cl- (aq) PbCl2(s)
Steps in writing ionic equations
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Writing ionic equations
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Quick check 5Construct (i) a balanced chemical equation and (ii) an ionic equation for each of the following reactions:
(1) Sulphuric acid + potassium hydroxide(2) Nitric acid + sodium hydroxide
(3) Silver nitrate solution + sodium chloride solution (4) Calcium carbonate + hydrochloric acid
(5) Magnesium + hydrochloric acid
Solution
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Solution to Quick check 5
1. H2SO4(aq) + 2KOH(aq) K2SO4(aq) + 2H2O(l) H+(aq) + OH-(aq) H2O(l)
2. HNO3(aq) + NaOH(aq) NaNO3(aq) + H2O(l) H+(aq) + OH-(aq) H2O(l)
3. AgNO3(aq) + NaCl(aq) AgCl(s) + NaNO3(aq) Ag+(aq) + Cl-(aq) AgCl(s)
4. CaCO3(s) + 2HCl(aq) CaCl2(aq) + H2O(l) + CO2(g)CaCO3(s) + 2H+(aq) Ca2+(aq) + H2O(l) + CO2(g)
5. Mg(s) + 2HCl(aq) MgCl2(aq) + H2(g)Mg(s) + 2H+(aq) Mg2+(aq) + H2(g)
Chapter 10
Acids, Bases and Salts
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