ACIDS, BASES AND SALTS

ACIDS AND BASES

‘THE NATURAL SOURCES OF ACIDS AND BASES

In everyday life, we deal with many substances that chemists classify as acids. For example, ‘orange juice and grapefuit juice contain citric acid. These juices, and others of the like, contain ascorbic acid, a substance more commonly known as vitamin C. Examples of natural sources of acids and the type of acids they contain are shown in table below Some natural sources acids Source ‘Type of acid present Mineral acids (HCI

Minerals

H,$0,, HNO,, ete.)

Tobacco Salicylic acid

Tea Tannic acid

Cotte Chlorogenie acid

Sugar beet Giutarie and adipic acids

Blackberry Isocitrie acid

Spinach, tomato Onalic acid Sour (fermented milk) Lactic acid Bee, ant and nettle stings Methanoic acid (formic acid) Grapes, bananas, Tartaric acid tamarinds Citric acid ( lemons and limes have particularly high concentrations of cit Citrus fruits acid, itcan constitute as much as 8% of the dry weight of these fruits) Acids have a sour taste. Vinegar, lemon juice, grapeftuit juice and spoilt or fermented milk are all sour tasting because of the presence of acids. ‘The acids present im animal and plant materials

are known as organie acids. Salads are often flavoured with vinegar, which contains dilute acetic acid. Boric acid is a substance that is sometimes used to wash the eyes, In any chemistry laboratory, we find acids such as hydrochloric acid (HC), sulphuric acid (H:S0,), and nitric acid (HNO), These acids are called mineral acids because they can be prepared from naturally occurring compounds called minerals Mineral acids are generally

stronger and should be handled with great care, especially the concentrated acids, for they are very corrosive, They can eat away metals, skin and clothing. Nevertheless, some acids are not corrosive even when they are concentrated. They ate called weak acids. Ethanoie acid is one ‘example Its found in vinegar In general, organie acids ate weaker than natural acids You can tell fa substance is acid or not by its effect on litmus. Litmus is a purple dye. It ean be

used as a solution, or on paper, called litmus paper Litmus solution is purple, Litmus paper for testing acids is blue while that for testing bases is red in colour. Acids will turn litmus

solution red. They will aso turn blue litmus paper red.

Bases do not usually occur naturally. So they are not normally obtained from natural sources. However, they are prepared in the laboratory or in industry. Bases can be classified into oxides, hydroxides or carbonates. ‘Therefore, bases can be defined as the oxides, hydroxides or carbonates of metals. Bases taste bitter. A bitter taste isa characteristic of all bases.

Most bases are insoluble in water. The bases which dissolve in water are known as alkalis. The most common alkalis are potassium hydroxide (KOH), sodium hydroxide (NaOH), calcium hydroxide, Ca(OH)s, and ammonium hydroxide (NHOH), also known as ammonia solution Alkalis turn Titmus solution blue and red litmus paper blue A substance, such as litmus, which changes from one colour to another when mixed with an acid or base, is called an indicator.

Table 3.2 shows how acids and bases (alkalis) affect the colours of different indicators. We can use this clue of colour ehanges to tell whether an unknown substance is an acid or base (alkali, Some common indicator colour changes Indicator Colour in acid Colour in alkali (base) Methyl orange Orange yellow

Phenolphihalein Colourless Pink

Litmus Red blue

Bromothymol blue Yellow blue

THE REACTIONS OF ACIDS WITH VARIOUS MATERIALS

Acids react with different substances to produce different products. These reactions are best carried out by using dilute acid solutions. The following ate some reactions of dilute acids with various substances

• Reaction with metals

Acids react with quite reactive metals (not very reactive ones) to produce salt and liberate a hydrogen gas. metal + actd ~ salt + hydrogen This unsafe to tr this reaction with very reactive metals such as sodium or calcium. ‘The reaction with such metals is so violent, Metals less reactive than lead, such as silver and gold have no reaction with dilute acids. Even with lead, i is difficult o see any reaction in a short time.

The salt produced when a dilute acid reacts with a metal depends on the acid and a metal used: Me) + 2HNOs) + Ma(NOs} + Hay Pag + 2HChag > Z0Cla eg) + Ha

• Reaction with carbonates

Acids react with carbonates to give salt, water and carbon dioxide. In general, all carbonates give off carbon dioxide when they react with acids acid + metal carbonate + salt + water + carbon dioxide The normal methods of preparing carbon dioxide in the laboratory are based on this reaction.

Dilute hydrochloric acid 1s reacted with marble chips (calcium carbonates) DHChag + CACO 5) + CaCtjg) + H20 gy + CO

• Reaction with oxides and hydroxides (alkalis)

Hydroxides: aeids react with alkalis, forming salt and water NaOHay + HNOsay) + C2Clay + H:09 Ovides: They also react with metals oxides, forming salt and water Za) 2HClagy 7 Z0Chny + HO The bases (oxides, hydroxides) all react in the same way with acids, and in the process, salts are formed. This type of reaction is known as meutralization reaction. It can be summarized up in a general equation: acid + hase “> salt + water

Reaction with hydrogencarbonates (bicarbonates) Acids react with hycrogencarbonates, forming salt, water and liberating carbon dioxide gas: Nal1CO3pq) + HChag 7 NaCag) + HO) + COs The Reactions of Alkalis with Various Materials Determine the reactions of alkalis with various materials Alkalis react with acids to produce salt and water All alkalis, except ammonia solution, will react with ammonium compounds liberating ammonia gas. Aqueous solutions of alkatis will

precipitate the insoluble hydroxides of other metals from the solutions of metal salts. Caustic alkalis attack aluminium, zinc and lead to form salts, They react with carbon dioxide to form carbonates Characteristic reactions of bases Disvolution in water Most bases are insoluble in water. Some are soluble in water. Soluble bases are known as alkalis.

The commonest alkalis are sodium hydroxide, (NaOH), calcium hydroxide, Ca(OH)>, potassium hydroxide, (KOH), and ammonium hydroxide or ammonia solution, (NH,OH). All alkaline

solutions contain hydroxyl ions, OL”. In sodium hydroxide solution, the ions are produced like

this:

NaQ Hay? Nagi Oe

Like acids, alkalis can also be classified as strong or weak. Ammonia solution is a weak base because it ionizes just partially: ‘The rest of the bases are strong bases because they ionize fully into ions in solution Reaction with acids Bases react with acids to produce salt and water. Refer to the reactions of acids with oxides and hydroxides discussed early Reaction with ammonium compounds Alkalis tur litmus solution blue and red litmus paper blue A substance, such as litmus, which

changes from one colour to another when mixed with an acid or base, 1s called an indicator. Table 32 shows how aeids and hases (alkalis) affect the colours of different indicators, We can use this clue of colour ehanges to tell whether an unknown substance is an acid or base (alkali, Reaction with aqueous salts of metals Aqueous solutions of alkalis will precipitate the insoluble hydroxides of other metals from the

solutions of mefal salts. Only NH,OH, KOH and NaOH are soluble enough in water.

All other hydroxides are insoluble and can be precipitated from aqueous solution by these three alkalis,

• When sodium hydroxide solution is added to copper (II) sulphate solution, a pale blue precipitate

of copper (ID) hydroxide is formed. CuSO4jq) + 2NaOHag > Cu(OH)a) + Naz SO

• Another example is the reaction between potassium hydroxide and iron (II) chloride, which

precipitates iron (Il) hydroxide, FeClgq) + 3KOHiay —* Fe(OH) sq) + 3K Chay Reaction with metals Caustic alkalis attack very few metals, The metals known to be attacked by the alkalis are aluminum, zine and lead, where the aluminate, zincate and plambate (II) ae formed respectively The aluminum will react thus:2Aly,+ 6NaOHpe + 6H:0,) + 2NasAIOH a) + 3Hapsodium aluminate) Reaction with carbon dioxide When carbon dioxide gas is bubbled through aqueous solutions of the caustic alkalis, the

carbonates are formed. 2NAOH ay +COai) + NasCOsug + HzO, With excess of the gas, the hydrogencarbonates are formed NasCOsiaq) + HO, + CO) 2NaH COs) Reaction with chlorine Chlorine reacts with excess of cold dilute caustic alkalis to form the hypochlorite, (NaCIO or KCIO)2Na0Hay + Clg) > NaClag# NaClOpy+ H:0y)2KOH yy + Chiy > KChay +

KCI.) H:0

excess chlorine 1s bubbled through hot concentrated solutions of caustic alkalis, the chlorates are formed, (NaC10) oF KCIO:), GNaOH a + Chg —> SNaChag) + NaClOsiy) + 3H20u,6KOH ay + 3Cly—> SKChay + KC} say + 3H:0 0)

Applications of acid-base neutralization in everyday life

Applications of acid-base neutralization in everyday life

Acid-base neutralization has many applications in everyday life. The following are some of these

applications

Indigestion and pain relief The dilute hydrochloric acid produced in your stomach is used for digestion and killing bacteria that might have been swallowed together with food or taken with water, However, excess acid ‘causes indigestion, which can be painful, To ease the pain, we take an anti-acid treatment. Anti- acids are a broad group of compounds with no toxic effects on the body. They are used to neutralize the effects of acid indigestion.

Some of these anti-acids such as milk of magnesia [insoluble magnesium hydroxide, Me(OH)2] help to neutralize and hence counteract the excess acid in the stomach. This treatment, therefore, prevents indigestion and pains. The neutralization reaction equation is: Mg(OH) 3) * 2HChagy > MEChipq) + 2H20y) Other anti-acids such as “Alka-Seltzer” contain soluble materials, including sodium hydrogencarbonate. These tablets also contain some citric acid (a solid zeid). On adding water

the acid and some of the sodium hydrogencarbonate react, producing carbon dioxide gas. This helps to spread and dissolve the other less soluble material. When taken, more sodium hydrogenearbonate neutralizes the excess hydrochloric acid in the stomach, thus easing digestion Some anti-acid tablets also contain painkiller to relieve pain. “Soluble aspirin” tablets dissolve and work in a similar way to “Alker-Seltzer” tablets. Vitamin C (ascorbic acid) can be added to

the tablets. Note that itis important to add water to start the action of the acid, Descaling kettles The limeseale (CaCOs) is formed inside boilers, Kettles and water heaters when hard water is boiled, The limescale can be removed by treatment with an acid that is strong enough to react with CaCO;, but not strong enough to damage the metal. Vinegar can be used to discale kettles ‘Commercial “discalers” use other acid solutions such as methanoie acid

Prevention of tooth deeay

Food remnanis sticking onto teeth (plaque), after eating especially sugary food is acted upon by bacteria in your mouth, The pH of 2 sugar solution is 7. However, bacteria in your mouth break down the sugar in plaque to form acids, for example lactic acid. "These acids lower the pH. Tooth decay begins when the pH falls below 5.8. The acid attacks the tooth enamel.

To help prevent tooth decay many types of toothpaste contain basic substances to neutralize the acids produced by these bacteria in your mouth. The pH of these basic substances is alkaline {higher than 7). The pH of saliva is slightly alkaline (pH! 74), soit can also help to counteract the acid, particularly after a meal. After eating a sweat, for example, it takes about 15 minutes for saliva to raise the pH above 5.8, and stop further decay

Soil treatment Most plants row best when pH of the soil is close to 7. They prefer the pH of between 6.5 and 7.0. the soil pH is below 6 0, the soil is too acidic, Above the pl of 80, the soil is too alkaline [the soil is too acidic or too alkaline, the plants grow poorly or not at al.

Chemicals can be added to the soil to adjust its pH. Most offen, ifthe soil is too acidie, itis usually treated by liming, In this context, liming means addition of quicklime (calcium oxide), slaked lime (calcium hydroxide) or powdered chalk or limestone (calciurn carbonate) to an acidic soil, These compounds (bases) have the effect of neutralizing the acidity of the soil If the soil is too alkaline, acids such as sulphuric acid, nittie acid or hydrochloric acid may be

added to the soil to neutralize excessive alkalinity, However, these compounds are. very expensive and hence uneconomical to apply on large-scale basis. Insect stings treatment When a bee stings someone, it injects an acid liquid into the skin. The bee sting, which is acidic in nature, can be neutralized by rubbing on calamine solution, which contains zine carbonate or baking soda, which is sodium hydrogencarbonate. These compounds are hasic in nature and so

have the effect of neutralizing the acid in the sting Wasp stings are alkaline in nature, and can be neutralized with vinegar, which contains ethanoie acid, Ant and nettle stings contain methanote acid. These may be neutralized by rubbing an extract squeezed from erushed onion leaves (which contain basie compounds) on the affected skin. The acid in the sting can also be neutralized by applying weak alkalis such as ammonia

solution, ash extract, baking powder, ete

Factory wastes treatment Liquid wastes from factories often contain acid. If it reaches a river, lake or ocean, the aeid will Kill fish and other aquatic fife, This can be prevented by adding slaked lime (calcium hydroxide) to the waste, to neutralize the acid before being dumped into water bodies.

Indicators

An Indicator from Locally Available Materials

Explain an indicator from localty available materials Certain coloured substances (many extracted from plants) have been found to change colour if added to an acid or alkaline solution. The colour change is reversed if the acid or alkali is neutralized. Substances that behave like this are known as indicators.

Coloured extracts can be made from red cabbage or blackberries, but probably the most used indicator is diimus. This is extracted from lichens Litmus is purple in @ neutral solution. When added to an acid solution, it tums red. Changing this red colour of litmus needs a chemical reaction The molecules of the indicator are usually changed in the presence of the acid. Substances with the opposite chemical effect to acids are

needed to reverse the change, and these are called alkalis They tum litmus solution to blue Litmus can also be used in paper form, in which case itis called limus paper. Here it comes in the blue and red forms. Litmus isa single chemical compound It gives a single colour change Litmus is not the only single indicator that chemists find useful. Others that are used frequently are phenolphthalein and methyl orange. These indicators give different colour changes when in

acidic and alkaline solutions (see table 3.2) Another commonly used indicator is the universal indicator (or full-range indicator). This is made from a mixture of dyes Such an indicator is useful because it gives a range of colours (“spectrum”) depending on the strength of the acid or alkali added (see table 3.3) With a universal indicator, different acids produce a range of different colours. Indeed, solutions of the same acid with different concentrations (pH) give different colours.

The more acidic solutions (for example battery acid) tum the universal indicator bright red, A less acidic solution (for example vinegar) will only tum it orange-yellow. There are also colour differences produced with different alkali solutions. The most alkaline solutions give a violet colour while the less alkaline solutions give a blue colout.

We learned that many indicators are extracted from plants. Flowers and leaves of different plants have different colours. These plant organs may be used to prepare indicators locally.

Activity 1

To prepare indicators from local plant materials Procedure:

• Collect Mowers fiom different plants in your local area. You may use coloured leaves if the

coloured flowers are not available.

• Crush the flowers/leaves in a motor and pestle to make a fine paste.
• Add ethanol to the paste to wash out chlorophyll. Add about Lem’ of ethanol per gram of pestle

used.

• Grind the mixture to a very fine paste so that the ethanol can penetrate the broken plant cells

fully

• Place the mixture in the sun or heat gently to evaporate off ethanol. Make sure most of the

ethanol has evaporated

• Filter the mixture to obtain a clear but coloured fitate. To obiain as much extract as possible,

squeeze the paste in a clean piece of cloth and collect the juice in a beaker. The liguid you obtain is your indicator

• Arrange test tubes in a rack and label them A,B, CD and E
• Pour sodium hydroxide, dilute hydrochloric acid, limewater, lemon juice, vinegar and washing

soda in test tubes A. B,C, D and E respectively.

• Add two to three drops ofthe prepared indicator in each of the test tubes. Observe and record the

colour changes. ‘Questions from the experiment

• What was the colour of your indicator?
• Write down the colour changes in each of the test tubes A to E.
• Which substance showed a sharp colour change?
• Perform a similar experiment using 2 ready-made universal indicator and observe whether there

is any difference in colour changes between this commercial indicator and that one prepared from local plants The Acidity and Alkalinity of Substance Using Indicators Test the acidity and alkalinity of substance using indicators ‘The Strengths of Acids and Bases There is a big difference between the strength of an acid or base and its concentration. An acid or alkaline solution is said to be concentrated iFit contains a large amount of it ina small amount oF

water. A dilute acid or base (alkali) has a small amount of it in lot of water. The concentration fof an acid or base tells us how much of itis dissolved in a certain volume of solution The concentration is normally expressed in grams per litre (g dm) or moles per litre (mol dim”.

The sitength of an acid or alkali expresses its dissociation in water, Strong acids or alkalis will dissociate completely in water to form ions. Examples of strong acids are sulphuric acid hydrochloric acid, nitrie acid and phosphoric acid. Weak acids inchide ethanaie acid, earbonie acid and methanoic acid. Examples of strong alkalis include potassium hydroxide, sodium hydroxide, calcium hydroxide and ammonium hydroxide. Weak bases inchide ammonia solution

and sodium hydrogen carbonate. A strong acid or alkali forms many ions in water. ‘The number of hydrogen ions, H’, formed when it dissociates in water, determines the strength of an acid, The strength of an alkali depends ‘on the number of hydroxyl ions, OH, formed when it dissociates in water. Strong acids and alkalis will form many 11" and OH ions respectively. Weak acids or bases will form very few of the respective ions,

Likewise, the term weak acid or base should not be confused with the term dilute acid or base. A weak acid dissociates in water only very slightly to form very few protons, HA weak alkali also dissociates very slightly to form very few hydroxyl ions, OH

‘THE CONCEPT OF AN INDICATOR

Describe the concept ofan indicator You have seen that single indicators change their colours only once when put in different acid and alkaline solutions The single indicators most commonly used inchide litmus, phenolphthalein and methyl orange On the other hand, universal indicators show a range of colour changes depending on the strength of an acid or base Single indicators can only tell us whether a certain solution is an acid or an alkali. These types of

indicators cannot be used to compare two acids or two alkalis with different strengths. Litmus paper, for example, cannot be used to compare the strengths of sulphuric aeid and ethanoic acid.

Both acids will change the blue litmus paper to red. Likewise, you cannot compare the strengths ‘of aqueous ammonia solution (NH,OH) and sodium hydroxide by just using a litmus paper, They will both turn to red litmus paper to blue.

A universal indicator can be used to measure strengths of different acids and alkalis. This indicator is @ mixture of simple indicators. Instead of changing colour just ance, it changes colour a number of times depending on the degree of acidity or alkalinity of the substances tested The pl¥ scale is a convenient means of expressing the acidity and alkalinity in liquids. The pH seale is a numerical scale used to indicate the relative strengths of acidie or basic solutions in

terms of relative amount of hydrogen ions (protons) or hydroxyl ions in solutions. The seale ranges from 0 to 14 Acidie solutions will have pH values less than 7.0 and alkaline solutions will have pH values greater then 7.0. Ail neutral liquids c.g. pure water have pH of 7.0. Table 3.3 shows the pH and strengths of acidic and alkaline solutions and the associated indicator colour changes.

Colours of the universal indicator in different acidic and alkaline solutions pH range Colour Strength 1,2,3 Red Strongly acidic 4 Orange 5.6 Yellow Weakly acidic

7 Green Neutral

89 Blue Indigo Weakly alkaline 10,11, 12, 13,14 Purplelviotet Strongly alkaline Rememiber that there is no clear dividing line between the pH ranges as apparently shown in the above table. This means that you may have substances with, for example, pH 1.2.1.5, 3.5, 4.4 5.6, 84, ete The table just tries to simplify the concept of acidity and alkalinity of acid and alkaline solutions.

Salts

An Indicator from Locally Available Materials

Explain an indicator from locally available materials Certain coloured substances (many extracted from plants) have been found to change colour if added to an acid or alkaline solution. The colour change is reversed if the acid or alkali is neutralized. Substances that behave like this are known as indicators.

Coloured extracts ean be made from red cabbage or blackberries, but probably the most used indicator is mus. This is extracted ftom lichens. Litmus is purple ina neutral solution. When added to an acid solution, it tums red. Changing this red colour of litmus needs a chemical reaction. The molecules of the indicator are usually changed in the presence of the acid. Substances with the opposite chemical effect to acids are

needed to reverse the change, and these are called alkalis. They tum litmus solution to blue. Litmus can also be used in paper form, in which case itis called limus paper. Here it comes in the blue and red forms. Litmus is a single chemical compound. It gives a single colour change Litmus is not the only single indicator that chemists find useful. Others that are used frequently are phenolphihalein and methyl orange. These indicators give different colour changes when in

acidic and alkaline solutions (see table 3.2) Another commonly used indicator is the universal indicator (or full-range indicator). This is made from a mixture of dyes. Such an indicator is useful because it gives a range of colours (“spectrum”) depending on the strength of the acid or alkali added (see table 3.3) With a universal indicator, different acids produce a range of different colours. Indeed, solutions of the same acid with different concentrations (pH) give different colours,

The more acidic solutions (for example battery acid) tum the universal indicator bright red. A less acidic solution (for example vinegar) will only tum it orange-yellow. There are also colour differences produced with different alkali solutions. The most alkaline solutions give a violet colour while the less alkaline solutions give a blue colour.

We learned that many indicators are extracted from plants. Flowers and leaves of different plants have different colours. These plant organs may be used to prepare indicators locally.

Activity 1

To prepare indicators from local plant materials Procedure:

• Collect flowers from different plants in your local area. You may use coloured leaves if the

coloured flowers are not available

• Crush the flowers/leaves in a motor and pestle to make a fine paste,
• Add ethanol to the paste to wash out chlorophyll. Add about 10em’ of ethanol per gram of pestle

used

• Grind the mixture to a very fine paste so that the ethanol can penetrate the broken plant cells

fully.

• Place the mixture in the sun or heat gently to evaporate off ethanol. Make sure most of the

ethanol has evaporated

• Filter the mixture to obtain a clear but coloured filtrate, To obtain as much extract as possible,

squeeze the paste in clean piece of cloth and collect the juice in a beaker. The liguid you obtain is your indicator

• Arrange test tubes in a rack and label them A, B, C D and E.
• Pour sodium hydroxide, dilute hydrochloric acid, limewater, lemon juiee, vinegar and washing

soda in tes! tubes A, B, C, D and B respectively

• Add two to three drops of the prepared indicator in each of the test tubes. Observe and record the

colour changes. (Questions from the experiment

• What was the colour of your indicator?
• Write down the colour changes in each of the est tubes A to E
• Which substance showed a sharp colour change?
• Perform a similar experiment using a ready-made universal indicator and observe whether there

is any difference in colour changes between this commercial indicator and that one prepared from local plants ‘The Acidity and Alkalinity of Substance Using Indicators Test the acidity and alkalinity of substance using indicators The Strengths of Acids and Bases There is a big difference between the strength of an acid or base and its concentration. An acid or alkaline solution is said to be concentrated iF it contains a large amount of it in a small amount oF

water. A dilute acid or base (alkali) has a small amount of it in a lot of water. The concentration of an acid or base tells us how much of itis dissolved in a certain volume of solution The ‘concentration is normally expressed in grams per litre (g dm) or moles per litre (mol dm The strength of an acid or alkali expresses its dissociation in water. Strong acids or alkalis will dissociate completely in water to form ions, Examples of strong acids are sulphuric acid

hydrochloric acid, nitrie acid and phosphoric acid. Weak acids inchide ethanoie acid, carbonic acid and methanoic acid. Examples of strong alkalis include potassium hydroxide, sodium hydroxide, calcium hydroxide and ammonium hydroxide. Weak bases include ammonia solution and sodium hydrogenearbonate A strong acid or alkali forms many ions in water. The number of hydrogen ions, H’, formed when it dissociates in water, determines the strength of an acid, The strength of an alkali depends

‘on the number of hydroxyl ions, OH, formed when it dissociates in water. Strong acids and alkalis will form many 11° and OH ions respectively. Weak acids or bases will form very few of the respective ions, Likewise, the term weak acid or base should not be confused with the term dilute acid or base. A weak acid dissociates in water only very slightly to form very few protons, HA weak alkali also dissociates very slightly fo form very few hydroxyl ions, OH

The Concept of an Indicator

Describe the concept of an indicator You have seen that single indicators change their colours only onee when put in different acid and alkaline solutions, The single indicators most commonly used include litmus, phenolphthalein and methyl orange On the other hand, universal indicators show a range of colour changes depending on the strength of an acid or base Single indicators can only tell us whether a certain solution isan acid or an alkali, These types of

indicators cannot be used to compare two acids or two alkalis with different strengths. Litmus paper, for example, cannot be used to compare the strengths of sulphuric acid and ethanoic acid Both acids will change the blue litmus paper to red. Likewise, you cannot compare the strengths of aqueous ammonia solution (NH,Ol) and sodium hydroxide by just using a litmus paper. ‘They will both turn to red litmus paper to blue.

A universal indicator can be used to measure strengths of different acids and alkalis. This indicator is a mixture of simple indicators. Instead of changing colour just once, it changes colour a number of times depending on the degree of acidity or alkalinity of the substances tested.

The pH scale is a convenient means of expressing the acidity and alkalinity in liquids. The pH seale is a numerical scale used to indicate the relative strengths of acidie or basic solutions m terms of relative amount of hydrogen ions (protons) or hydroxyl ions in solutions. The seale ranges from 0 to 14.

Acidic solutions will have pH values less than 7.0 and alkaline solutions will have pH values ercater then 7.0. All neutral liquids e.g. pure water have pH of 7.0. Table 3.3 shows the pH and strengths of acidic and alkaline solutions and the associated indicator colour changes.

Colours of the universal indicator in different acidic and alkaline solutions pH range Colour Strength 1,23 Red Strongly acidic 4 Orange 5.6 Yellow Weakly acidic

7 Green Neutral

89 Blue Indigo Weakly alkaline 10,11, 12,13, 14 Purpleiviolet Strongly alkaline Remember that there is no elear dividing line between the pH ranges as apparently shown in the above table. This means that you may have substances with, for example, pH 1.2. 1.5,3.5, 44, 5.6, 84, ete The table just tries to simplify the concept of acidity and alkalinity of acid and alkaline solutions.

Salts Salts The Natural Source of Salts in Daily Life Investigate the natural source of salts in daity life A salt isa substance formed when some or all of the hydrogen atoms ofan acid are replaced by @ metal or ammonium ion. A salt, therefore, may be defined asa compound in which thereplaceable hydrogen of an acid has been wholly or partially replaced by a meta.

In sodium chloride (NaCl), for example, the hydrogen atom of hydrochloric acid (HCI) has been wholly replaced by an atom of sodium In magnesium sulphate (MgSO4) and sodium sulphate (Na:SO9), both hydrogen atoms of sulphuric acid (H1:S0s) have been replaced by one atom of magnesium and two atoms of sodium respectively. In sodium hydrogen sulphate (NaH1SO,), only ‘one out of two hydrogen atoms has been replaced by an atom of sodium. ‘This type of a salt is

called an acid salt, because it still contains a replaceable hydrogen atom. Many chemical compounds may be classified as salts. The salt most familiar to every body is table salt (sodium chloride). Baking soda is the salt, sodium bicarbonate (NaHCOs). Magnesium sulphate (also called Epsom salt) is often found in the home In general, salfs are ionic impounds that are compased of metal and non metal ions. For example, sodium chloride isis composed of metallie sodium ions (Na") and non-metallic chloride ions (Cl

) Some salts are made of metallic and non-metallic radicals € g ammonium nitrate (NINO) is composed of ammonium radical (NH,") and nitrate radical (NOS). There is a wide range of types and natural sourees of salts. Common salt is mined from underground deposits. The salt obtained from such a source contains sodium chloride mixed with rock impurities The other source of sodium chloride is seawater The silty taste of seawater is due to the

presence of salts such as sodium chloride and magnesium bromide. However, there are many different types of salts present in seawater, though in small proportions, as shown in the table below (table 35)

Salt Formula Percentage composition

Sodium ehtoride NaCl 272 Magnesium chloride Mech, 038 Magnesium sulphate MzSO, 017 Calcium sulphate C2804 013 Potassium chloride KCl 0.09 Caleium chloride CaCO 0.01 Magnesium bromide MeBrs 001 Sodium nitrate (Chile saltpetre), NaNO; and caleium carbonate, CaCO; are found in underground deposits. Calcium carbonate occurs naturally as marble, limestone or chalk in the round from which it can be mined mechanically. What other natural sources of salts do you

know? ‘Types of Salts Salts may be classified according to their mode of formation. ‘The following are types of salts srouped according to their mode of formation: Normal salt:- This is a salt formed when all of the replaceable hydrogen atoms ofan acid have been replaced by a metal atom ¢.2. sodium chloride is a normal salt because all hydrogen atoms are replaced from an acid during its formation.

2Nay) + 2HClagr> 2NAClog) + Ha Other normal salts include magnesium chloride (MgCI.), potassium chloride (KCI), copper (I) sulphate (CuSO.), sodium sulphate (Na;S03), sodium carbonate (Ne;COs), trisodium phosphate (NasPO4), ete Acid salt:- An acid salt is a salt formed when part of the replaceable hydrogen atoms of an acid are displaced by 2 metal ., sodium bisulphate (NaHSO,) is an acid salt, Hy8O ag) + NaOH ogy NaFlSO vo) + HO

Other examples of acid salts include sodium hydrogensulphate, (NaHSO.). sodium hydrogensulphide, (NaHS) and sodium hydrogencarbonate (NaHCOs). Since acid salts contain hydrogen ions, they exhibit some acidic properties. Hence, they behave like acids, for example a. they react with bases to form salts and water only. NaHSO uy +NeOHag7-+ N8:SO 4) + HO) b. they react with carbonates to yield carbon dioxide 2NSHSOtuxy + NaCOsoqy+ 2N&SOcqi*

H20q) + COs)

Basie salt:- A basic salt is formed by the action of an acid with higher proportions of the base, than is necessary for the formation ofa normal salt

Examples of basic salts are

• Basie copper carbonate, CuCOs,Cu(OH)s
• Basie lead carbonate (white lead), PbCOs.Pb(OH)s
• Basic magnesium chloride, MgCl Ma(OH
• Basie zine chloride, ZnCl; Zn(OH}z

A basic salt may also be formed by the partial replacement of the hydroxy! groups of a diacidic or triacidic base by an acid radical PLIOM)a [lead Aydrovide] + HNOsiay (nitric acid] + PHOHNOss [basic lead nitrate?+ 1,09 firater] Basie salts are usually insoluble in water. Such salts are formed by the close association of two simple salts, when erystallized from a solution of a mixture of the two.

The Solubility of Different Salts in the Laboratory Analyse the solubility of different salts in the laboratory Some salts are more soluble in water than others are. However, other salts are insoluble in water The knowledge of solubility of different salts in water is very important because it ean help us prepare different salts in the laboratory by such methods as precipitation, diteet combination (synthesis), crystallization and so forth.

As regards to solubilities, salts can be classified into two groups: salts which are soluble in water (soluble salts) and salts which do not dissolve in water (insoluble salts). Table 3.6 summarizes the solubility of different salts in water.

The patterns of solubility for various types of salts Soluble salts Insoluble salts 1 All sodium, potassium and ammonium salts, 2: All eiteates of metals silver, mercury(D) and lead chlorides bariuny lead (U1) and ealeium sulphates but oth

• All chlorides except

common carbonates are insoluble but oth AN aunpians eat common hydroxides are insoluble

• Sodium, potassium, and ammonium

cartonates, hydroxides

• Sodium, potassium and ammonium

Salts in the Laboratory Prepare salts in the laboratory Several methods are available for the preparation of salts. The solubilities of the prepared salts determine their methods of preparation. Hence, in the choice of a method of preparation of a particular salt, one has to be acquainted with its solubility properties.

Soluble salts are usually prepared by methods which involve erystallization. In this method. as the name suggests, resultant salts are inthe form of crystals Insoluble salts are usually prepared by methods which involve precipitation, These methods are sometimes referred to as double decomposition. To precipitate an insoluble salt, you must mix a solution that contains its positive ions with the one that contains its negative ions

Salts may also be prepared by direct combination (or synthesis). For example, magnesium chloride may be prepared in the laboratory by heating magnesium in a stteam of chlorine Mg) + Clag) > MgCl

Preparation of soluble salts

Soluble salts may be prepared by any of the following methods:

• Reaction between an acid aud an alkat:in this method, a dilute acid 1s added to an alkali m the

appropriate volume ratio, The reaction between an acid and an alkali is termed as neutralization For example, sodium chloride may be prepared by the following neutralization reaction: NaOHyay + HCliay7> NaClaa)* Hs, Both reactants are soluble, and no gas is given off during the reaction. So, itis difficult to know when the reaction is over. In this case, you have to use an indicator. A universal indicator or litmus could be used, but even better is

phenolphthalein, his is pink in alkaline solution, but colourless in neutral or acidic solutions. 2 Reaction of a metal with an acid:This is another general method for preparing salts. For

example, zine sulphate can be made by reacting dilute sulphurie acid with zine-Zng +

H,80 jug Z0SO uy) Hag However, this method is not suitable for all metals or all acids. is good for preparing salts of fairly reactive metals such as magnesium, aluminium, zine and iron, However, the reactions of highly reaetive metals like sodium, potassium and calcium with acids are very violent and dangerous. The reaction with lead is too slow. Copper, silver and gold do not react at al

• Reaction of a metal oxide with an acid Metal oxides, as you studied early, react with dilute

acids to produce salts. Copper oxide is an insoluble base. Although copper will not react with dilute sulphuric acid, copper (II) oxide will, The salt that forms is copper (Il) sulphate CuO + H,SOsay + CuSOsoy + H:0o

• Reaction of a metal carbonate with an acid.(he reaction between metal carbonates and dilute

acids are accompanied with evolution of carbon dioxide gas. The evolution of a gas cam be used to indicate when the reaction is over. An example of such reactions is the reaction between calcium carbonate and dilute hydrochloric acid. CaCOxy~ 2HChay > CaChiay * HO) + COs) General methods of preparing soluble salts ‘The above methods for preparing soluble salis are specific for each method mentioned Generally, soluble salts may be prepared by two broad methods

Method 1:This route is essentially the same whether starting with a solid metal, a solid base (oxide) or a solid carbonate. The route ean be divided into four stages:

• Stage I-An excess (more than enough) of the solid is added to the acid and allowed to react.

Using an excess of the solid makes sure that all the acid used up. If itis not used up at this stage, the acid would become more concentrated when the water is evaporated later (stage 3)

• Stage 2:The excess solid is filtered out after the reaction is completed.
• Stage 3:"The filtrate is gently evaporated to concentrate the solution. This can be done on

heated water bath, Do not heat so strongly or “spitting” might take place

• Stage 4The concentrated solution is eooted down to let the erystals form. Filter off the crystals

Wash them witha little distilled water. Dry the erystals carefully between the filter papers. Method 2:This method (titration method) involves the neutralization of an acid with an alkali (for example sodium hydroxide) or a soluble carbonate (for example sodium carbonate). Since both the reactants and the products are colourless, an indicator is used to find the neutralization point or end point (when all the acid has just been neutralized). Once the end point 1s reached.

the resulting salt solution is evaporated and cooled to form erystals as deseribed in method 1 General methods for preparing insoluble salts Some salts ate insoluble in water (for example silver chloride and barium sulphate ~ see table 3.6) Such salts are generally prepared by ionic precipitation Precipitation is the sudden

formation of a solid either: when two solutions are mixed; or when a gas is bubbled into @

solution.

For example, barium sulphate can be prepared by adding a solution of a soluble sulphate (for ‘example sodium sulphate) o a solution of a soluble barium salt (for example barium chloride) The insoluble barium sulphate is formed immediately. This solid falls to the bottom of the container 2s a precipitate (figure 3.2). The precipitate can be filtered off, tis then washed with distilled water and dried in a warm oven. The equation for the reaction is:

BaClayay* NazSO4ayi> BaSO ya) + 2NICly This Shows how important the state symbols can be – i is only through state symbols that we can tell this equation shows a precipitation, barium sodium chloride sulphate

solution solution

sodium — barium

solution — (EEE TEI — sulphate

precipitate Barium sulphate could also be made fom barium nitrate and sodium sulphate for example, since vill be prosipiate Ba ay + SO? aq) +BaSO4) iodides and suipates

Preparation of salts by direct combination (synthesis)

Som soluble and insoluble salts can be made diesty by reacting bwo elements together. This is bromides and iodides. For instance, if a piece of burning sodium is lowered into a gas jar of chlorine the two react violently t produce a white powder of sodium chloride 2Nayy+ Clay 2NaChy Other chlorides can also be prepared by combination, for example, iron (II1) chloride and aluminium chloride can be made by heating iron and aluminium metals in stream of chlorine:

Fey) + 3Clag) + 2FeChy 2Aley + Ih —7 2AICl The reaction between ammonia gas and hydrogen ehloride gas to produce ammonium chloride is also a synthesis reaction Nig) +HChgy > NHAC Direct combination reactions do not produce crystals ofthe salt, but only a powder The Effects of Heat on Salts Examine the effects of heat on salts When different salts are heated, they behave in different manners. The crystals of some salls

contain water of crystallization, When these hydrated salts are heated, their water of crystallization is driven off as steam. The erystals then lose their shape and become a powder The following are few examples of hydrated salts: Salt formula Chemical name CuSO..510 Copper (I) sulphate five water Na;COs.10H,0 Sodium carbonate ten water MgCl, 6H:0 Magnesium chloride six water FeCh 6120 Iron (IID) chloride six water PeS0..711,0 {ron (I) sulphate seven water

CoC. 60 Cobalt (Il) chloride six water MgS0,,7H:0 Magnesium sulphate seven water CaS0,2H,0 Caleium sulphate two water Salphates Sulphates of potassium, sodium, calcium, lithium and magnesium are stable to heat and do not decompose when heated. Other sulphates decompose to give the oxide and sulphur trioxide gas except iron (II) sulphate which decomposes to give sulphur dioxide and sulphur trioxide.

‘Copper (Il) sulphate five water erystals are blue in colour, but when heated, they are dehydrated to form a white powder: CuSO, SH,O,shvdraied (biue)—+ CuSOx,janhvatrous (white) + SH20 yy Crystals that have lost their water of crystallization are called anhydrous. If water is added back to the anhydrous copper (Il) sulphate powder. the powder tums into blue erystals again and heat is evolved, This can be used as a qualitative test for water If the white, anhydrous powder is

further heated strongly, it decomposes to black copper (I!) oxide:

CuSOg, witie—> CUuOblack= SOre)

Hydrated iron (If sulphate is green in colour, When heated, it loses all its water of erystalization and changes colour from green to white: FeSO, 7TH:O, green—+ FeSOq, white + 7H:04 When heated even more strongly, the white powder decomposes to form a black oxide: 2FeSOqihite—+ Fez0x,yblack~ SO2u) * SOx tron (IID sulphate decomposes on heating to form slightly different products: PexSOa)as) > Fes0s + 380i Chtorides ‘The chlorides of most metals are hydrated except these of potassium, lead, mercury and silver

Hydrated chlorides do not usually give the anhydrous salt when heated. Instead, a chemical change termed as hydrolysis normally occurs. The reaction is secompanied by the evolution of sieam and hydrogen chloride gas, and the formation of the basic chloride or oxide, When, for

example, hydrated magnesium chloride is heated, its basic chloride is formes

MgCl, 6H:0,) + Mg (OH)Cly + HCl) *SH:O,.) The same case applies when hydrated calcium chloride is heated. However, when hydrated aluminum chloride is heated, it does not produce the anhydrous salt, Instead, the oxide is formed thus

2AICI, 6109 + ALO) + HCl + 3H:0.)

Ammonium chloride sublimes when heated. The reaction is reversible and the products may recombine on cooling to form the salt back

NH,Cle@2NHs) + HCL)

Carbonates and hydrogencarbonates The carbonates of potassium and sodium are very stable to heat. They do not decompose even when heated to very high temperatures. All other carbonates decompose when heated to give the ‘oxide and carbon dioxide: CaCO ys CAO) + CO

CUCO%) > CUO) + COry

However, there are very few and exceptional carbonates that do not behave like this. Ammonium carbonate, for example, decomposes readily when heated to give ammonia gas, water vapour and carbon dioxide gas: (NH.)2COsq) + 2NHayg + HO) + COxy) All hydrogencarbonates decompose on heating to give the earbonates, water vapour and carbon dioxide 2NAHCO,ys) + NezCOsy + HO) C02) Nitrates When heated, potassium and sodium nitrates decompose to give the nitrite and oxygen

2KNO sq) 2KNO2)+ Ox) 2NaNOsay + 2NaNO3) + O25) The nitrates of common heavy metals (such as Pb, Al, Ca, Mg, Zn and Cu) decompose on heating to give the oxide, nitrogen dioxide and oxygen: 2PD(NOz)35) + 2PHO, + 4NOz ie) O24

2€a(NOs}34) > 2CAO + ANO2 ) + O2

The nitrates of silver and mercury are completely decomposed to the metal, nitrogen dioxide and oxygen 2ANO y+ 28g) +2NO2 + Org) Ha(NOs)as) + Ho + 2NOxy + Ox Ammonium nitrate is decomposed by heat into dinitrogen oxide and water NELNOs.) > N20 wp + 2420 Hydroxides Potassium and sodium hydroxides are very stable to heat, They do not decompose even when heated strongly. All other hydroxides decompose to give the oxide and water vapour, ¢.8

Ca(OH) 3.) Cap) + HO

DELIQUESCENCE, EFFLORESCENCE AND HYGROSCOPY

Deliquescence Deliquescence is the absorbing of moisture from the atmosphere hy a solid to form a solution I? calcium chloride (CaCh) 1s exposed to air, it absorbs water vapour from the atmosphere and ‘eventually dissolves. lis tendency to absorb water vapour explains why itis used as a drying agent for gases (not ammonia, because it combines withthe 228), Solid sodium hydroxide is also detiquescent. On exposure to air, pellets of sodium hydroxide

quickly become shiny and then sticky as they absorb water vapour from the atmosphere Eventually the sodium hydroxide pellets absorb moisture fiom the atmosphere so much that they dissolve to form 2 solution of sodium hydroxide Copper (LI) nitrate and zine chloride are the other deliquescent salts. Pure table salt (NaCl) is not deliquescent. However, if the salt 1s directly obtained from the sea, it 18 deliquescent. The salt

from the sea contains magnesium chloride as one as the impurities. It is this magnesium chloride salt that deliquesces and not sodium chloride Hygroscopy Some substances tend to absorb water vapour from the air but do not change their physical states, Copper (Il) oxide and calcium oxide are both hygroscopic solids because they can absorb moisture from the atmosphere and yet retain their solid states. Because of this behaviour, calcium

oxide is used as a drying agent, which absorbs moisture from gases prepared in the laboratory. Concentrated sulphuric acid is @ hygroscopic liquid. When exposed to ai, the acid absorbs water vapour from the stmosphere diluting itself to absorb 3 times its original volume, Therefore, Aygrescopy may be defined as the tendency of a substance to absorb water vapour from the atmosphere without changing its physical states. The word hygroscopy is a general term

applied to all substances that absorb water vapour from the air. Any substance that can take up moisture from the atmosphere is said to be hygroscopic in nature. Efflorescence Efflorescence is the tendency of a hydrated substance to lose the water of crystallization to the atmosphere, Some salt crystals give out some or all of their water of crystallization to the atmosphere when exposed to air. Such substances are said to be efflorescent and the process of

water loss is known as efflorescence. Sodium carbonate ten water (washing soda) is a good ‘example of an efflorescent substance. If washing soda crystals are exposed to open air at room temperature, they lose some of the water of erystalization. The solid loses nine of its ten molecules of water of crystallization to the air. One molecule of water, which remains fixed, can be removed only by strong heating NaxCO:.10H:Oyycrpstals + NaxCOs H2Q,opowder + 9H20\9,

The crystal lattice is broken down as the salt loses its after of crystallization. ‘Thus, transparent crystals of hydrated sodium carbonates become white and powdery on the surface.

Another efflorescent compound is Glauber’s salt, sodium sulphate ten water (NazSOx 10:0) (On exposure to air, it loses the whole of its water of erystallization to the ai. NasSO, 10H:0,.) + Na:SOya + 1020.9 Iron (11) sulphate seven water (FeSOs,7H.0) is also efflorescent.

‘The Uses of Different Types of Salts in Everyday Life Explain the uses of different types of salts in everyday life There is a wide range of salts. A great number of them play an important role in our everyday life. The following are the uses of some sats: Sodiunt chloride (common salt) Sodium chloride often ealled common salt or table salt, is essential for life and is an important raw material for industries. At home, it is used for cooking, that is, flavouring different

foods Biologically, it has a number of functions: it is involved in muscle contraction; it enables the conduction of nerve impulses in the nervous system: it regulates osmosis (the passage of solvent molecules through membranes); and it is converted into the hydrochloric acid that aids digestions in the stomach, Some industrial uses of sodium chloride include curing bacon, flavouring foods, and in the manufacture of margarine, butter and cheese. It is also used to tan leather inthe leather industry.

Rock salt is used as a fertilizer for sugar beet, and is spread on roads to melt the ice during winter, The salt is the starting point for many important chemicals, for example, the electrolysis of brine (concentrated solution of sodium chloride) gives sodium hydroxide, chlorine and hydrogen. What other uses of sodium chloride do you know’ Mention them.

Catcium carbonate (marble, limestone, chath) Caleium carbonate finds a wide range of uses:

• An important use of calcium carbonate is in the building industry. It is widely used in making

cement, lime, mortar and making steel from iron.

• Powdered limestone i used as a liming material to neutralize soil acidity. When used m this

way, itis termed as agricultural lime. When added in the soil, agricultural lime acts as a caleium source for plants as well as increasing the pH and water retaining capacity of acidic soils.

• tis also used in making paint, plastic, rubber, ceramic and glass: and in oil refining, and iron ore

purification,

• Caleium carbonate is the most preferred mineral in the paper industry. It helps in the production

of the best quality papers.

• Since calcium is essential for healthy bones and teeth, it is used as.a dietary calcium supplement
• Limestone can also be well shaped, painted, and then used as decorative stones

Ammonium salts Most ammonium salts such as (NH,)sPOs, NH\CI, NHANOs, (NH,)2S0,, CAN, urea, ete are used as nitrogenous fertilizers which are applied to the soil to improve soil fertility and hence enhance plant growth and production. Millions of tonnes of fertilizers are produced every year. Without these chemicals, world food production would probably be halved Calcium sulphate (anhydride, CaSO. gypsum, CaSOxH0) Gypsum is chiefly used for the manufacture of Plaster of Paris (P-O.P). The Plaster of Paris is

used for making casts for statuary ( the expansion during setting ensures a fine impression), in surgery to maintain joints ina fixed position and in cement and wall plasters, Among the many other uses of calcium sulphate are as a pigment in white paints, as a soil conditioner, in Portland cement, as a sizer, filler and coating agent in papers, in the manufscture of sulphuric acid and sulphur, in the metallurgy of inc ores and ax a drying agent in. many

laboratory and commercial processes, Sodiuns carbonate (washing soda) Sodium carbonate is widely used as one of the raw materials for making glass. Glass is made by heating a mixture of limestone, sand, sodium carbonate, and recycled glass in a furnace. The cosmetic industry uses it for manufacturing soap. The chemical industry uses it as a precursor to numerous sodium-containing reagents Iti also important in photography and in the textile industry. In addition to these industrial

applications, sodium carbonate is used in medicine as an anti-acid

Sodium carbonate has various environmental applications. Large quantities of the carbonate are used in sewage treatment; in water softening as washing soda erystals, NaxCO;.10H:0, and in <desulphurisation of flue gas Magnesium sulphate (Epsom salt) Magnesium sulphate is mainly employed for making health salts (laxative, mild purgative). The health salt is used as a medicine (laxative) which aids to empty the bowels following

constipation or other health problems, Copper (11) sulphate Copper (II) sulphate is used in fungicides, which are sprayed on crops, especially vines and potatoes, to kill moulds which would inure plants and hence curtail crop yield, It is also used in the manufacture of certain green pigments which are used for painting Calcium phosphate Its largely used in making phosphoric acid and fertilizers. Calcium phosphate is used in baking.

Itis also used mm cheese products.