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


Chapter 10


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.

http://images.google.com.sg/imgres?imgurl=http://www.azfotos.com/food_meals/fruits/stockphotosalamy/many_apples_AEA24A.jpg&imgrefurl=http://www.azfotos.com/food_meals/fruits/apples_pictures.htm&h=320&w=450&sz=37&tbnid=3KnQ9mEHBTprYM:&tbnh=88&tbnw=124&hl=en&start=9&prev=/images%3Fq%3Dapples%26svnum%3D10%26hl%3Den%26lr%3D%26sa%3DG

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 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

Chapter 10


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?

Chapter 10


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

Chapter 10


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.

Chapter 10


Some Common Acids

Name of acid Formula

Sulphuric acid H2SO4

Hydrochloric acid HCl

Nitric acid HNO3

Citric acid C6H8O7

Ethanoic acid (vinegar) CH3COOH

Chapter 10


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.

Chapter 10


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


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.

Return

Chapter 10


13

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.

Chapter 10


14

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.

Chapter 10


15

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.

Chapter 10


16

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

Chapter 10


17

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

Chapter 10


18

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.

Chapter 10


19

Some Common Alkalis

Name Chemical formula

Sodium hydroxide NaOH

Potassium hydroxide KOH

Calcium hydroxide Ca(OH)2

Ammonia solution (ammonium hydroxide)

NH3(aq)

Chapter 10


20

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

Chapter 10


21


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

Return

Chapter 10


22

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.

Chapter 10


23

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

Chapter 10


24

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

Chapter 10


25

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

Chapter 10


26

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

Chapter 10


27

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.

Chapter 10


28

Types of Oxides

Chapter 10


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

29

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

Chapter 10


30


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

Return

Chapter 10


31

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

Chapter 10


32

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

Chapter 10


33

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.

Chapter 10


34

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.

Chapter 10


35

To prepare sodium nitrate by neutralisation (titration method)

Chapter 10


Sodium nitrate and water (phenolphthalein as indicator)

burette

Pipette

36

Chapter 10


To prepare sodium nitrate by neutralisation (titration method)

37

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

Chapter 10



38

Preparation of copper(II) sulphate (acid on insoluble base)

Chapter 10


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.

39

Preparation of copper(II) sulphate (acid on insoluble base)

Chapter 10


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.

40

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

Chapter 10



41

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

Chapter 10



42

Making zinc sulphate (acid on metal)

Chapter 10


Can you describe how zinc sulphate is prepared with the aid of the diagrams?

43

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)

Chapter 10



44

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

Chapter 10


45

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

Chapter 10


46

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

Chapter 10


47


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


Return

48

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

Chapter 10


49

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

Chapter 10


50

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.

Chapter 10



51

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)


Chapter 10



52

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)


Chapter 10



53

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

Chapter 10


54


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


Return

55

1. http://www.sciencebyjones.com/acids_bases_salts.htm

2. http://ull.chemistry.uakron.edu/genobc/Chapter_09/

3. http://www.chem.ubc.ca/courseware/pH/index.html

To learn more about Acids, Bases and Salts, click on the links below!

Chapter 10


References

Chemistry for CSEC Examinations by Mike Taylor and Tania Chung

Longman Chemistry for CSEC by Jim Clark and Ray Oliver

56

c10 acids, bases and salts

Science

acid molecules

weak acids

acid anhydrideinvestigate

carbonic acid

acidshydrochloric acid

strength of acids

weak acida weak acid

hydrochloric acid dissolves