The second chapter in the CBSE Science textbook is Acids, Bases and Salts. We all are aware about the terms Acids, Bases and Salts.
› Acids – Anything with a pH value less than 7 are considered as Acids. For e.g. HCL, H2SO4, HNO3 etc.
› Bases – Anything with a pH value more than 7 are considered as Bases. For e.g. NaOH, KOH, LiOH etc.
› Salts – Anything with a pH value of 7 are considered as salts. The most common salt is Distilled Water (H20).
The acids, bases and salts are the three most important categories of the chemical compounds. Acids have a sour taste and are found in many fruits and vegetables that are found in the grocery stores. Some common acid containing fruits and vegetables are Lemons which contain Citric Acid, Apples which contains Malic Acid, Milk which contains Lactic Acid etc. On the other hand bases are the exact opposite of acids in chemical composition. The taste of any basic substance is bitter and are soapy/oily to touch. The most common examples of bases are soaps and detergents.
Note: Both the acids and bases must be handled with care as mishandling of strong acids and strong bases can cause severe burns on naked skin.
Acids have a pH value less than 7 on the pH Scale and are sour in taste. There are two types of acids that are discussed below:
› Organic Acids – Acids that occur naturally are known as organic acids. Organic Acids can be found in plants, fruits, vegetables, animals and other naturally occurring things. Some of the most common naturally occurring Organic acids are as follows.
- Milk – Lactic Acid (pH 2.4)
- Oranges/Lemons – Citric Acid (pH 2.2)
- Tamarind (Imli) – Tartaric Acid (pH 2.2)
- Apples – Malic Acid (pH 2.2)
- Vinegar – Acetic Acid (pH 2.4)
- Ant Sting – Formic Acid (pH 2.3)
- Tomatoes – Oxalic Acid (pH 4.4)
› Inorganic/Mineral Acids – Acids that are made in laboratories with the help of naturally occurring substances are known as Inorganic or Mineral Acids. Some of the most common Inorganic/Mineral Acids are as follows.
- Acetic Acid – CH3COOH
- Hydrochloric Acid – HCl
- Sulfuric Acid – H2SO4
- Nitric Acid – HNO3
- Carbonic Acid – H2CO3
- Hydrofluoric Acid – HF
General Properties of Acids
- Has a pH value of less than 7.
- Has a sour taste.
- Acids reacts with metals (Zn, Mg etc.) to liberate hydrogen gas.
- Turns the color of litmus paper from blue to red.
- Can conduct electricity.
An acid is considered as a strong acid if the acid dissociates completely in water, in other words if an acid dissociates 100% in water then the acid is known as a strong acid. It must be noted that in these acids all the hydrogen ions (H+) combine with water molecule and exist as hydronium ions (H3O+).
Some of the commonly found strong acids are Sulfuric acid (H2SO4), Hydrochloric acid (HCl) and Nitric Acid (HNO3).
HCl(aq) → H+(aq) + Cl–(aq)
HNO3(aq) → H+(aq) + NO3–(aq)
H2SO4(aq) → 2H+(aq) + SO42-(aq)
An acid is considered as a weak acid if the acid is unable to dissociates completely in water, in other words if an acid partially dissociates in water then the acid is known as a weak acid. Also, in an aqueous solution of weak acids ions and molecules are present.
Some of the commonly found weak acids are Formic acid (HCOOH), Carbonic acid (H2CO3) and Acetic Acid (CH3COOH).
HCOOH(aq) → HCOO–(aq) + H+(aq)
H2CO3(aq) → H+(aq) + HCO3–(aq)
H2CO3(aq) → 2H+(aq) + CO32-(aq)
CH3COOH(aq) → CH3COO–(aq) + H+(aq)
Reaction of Acids with Metals
- Acids reacts with Metals to form salt and liberate hydrogen gas. The general chemical equation of the reaction between acid and metal has been shown below:
Metal + Acid → Salt + Hydrogen
i) Magnesium reacts with diluted Hydrochloric acid to give magnesium chloride and hydrogen gas.
Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)
ii) Zinc reacts with diluted hydrochloric acid to give zinc chloride and hydrogen gas.
Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)
iii) Iron reacts with diluted sulfuric acid to give iron sulfate and hydrogen gas.
Fe(s) + H2SO4(aq) → Fe2SO4(aq) + H2(g)
- The most active metals such as Potassium, Calcium and Sodium reacts with the acids in the similar manner and gives out hydrogen gas. But, the reaction is so vigorous that the chances of explosions are very high.
i) Sodium reacts with diluted hydrochloric acid to give out sodium chloride and hydrogen gas in an explosive way.
2Na(s) + 2HCL(aq) → 2NaCl(aq) + H2(g)
- Nitric acids when reacts with metals often shows oxidizing properties during the reaction. Extremely diluted nitric acid reacts with magnesium to give out hydrogen gas.
- Acids reacts with carbonates and bicarbonates to form salt, water and carbon dioxide gas. The chemical equation for the same reaction is as follows.
Acid + Carbonate/Bi-Carbonate → Salt + Water + Carbon Dioxide (CO2)
i) Zinc carbonate reacts with sulfuric acid to give out zinc sulfate and carbon dioxide gas.
ZnCO3(s) + H2SO4(aq) → ZnSO4(aq) + H2O(l) + CO2(g)
ii) Sodium carbonate reacts with hydrochloric acid to give out sodium chloride with water and carbon dioxide gas.
Na2CO3(aq) + 2HCl(aq) → 2NaCl(aq) + H20(l) + CO2(g)
iii) Sodium hydrogen carbonate reacts with sulfuric acid to give out sodium sulfate with water and carbon dioxide gas.
2NaHCO3(s) + H2SO4(aq) → Na2SO4(aq) + 2H2O(l) + 2CO2(g)
The reaction between the hydrogen ions of an acid and the hydroxyl ions of a base is called neutralization. In a neutralization reaction the acid reacts with base to form salt and water. It is known as a neutralization reaction as both the reactants that the acids and bases are neutralized.
Some examples of neutralization reaction are as follows:
i) Sodium hydroxide reacts with hydrochloric acid to give out sodium chloride and water.
NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l)
ii) Ammonium hydroxide reacts with phosphoric acid to give out ammonium phosphate and water.
3NH4OH(aq) + H3PO4(aq) → (NH4)3PO4(aq) + 3H20(l)
iii) Lead oxide reacts with nitric acid to give out lead nitrate and water.
PbO(s) + 2HNO3(aq) → Pb(NO3)2(aq) + H20(l)
The presence of Hydrogen ions (H+) is what contributes to the acidic property of an acid, while the hydroxyl ions (OH–) contributes to the basic property of a base. When both the acid and base reacts together then the positive ion of the acid combines with the negative ion of base and gets neutralized resulting in the formation of water molecule which does not contain any charge.
H+(aq) + OH–(aq) → H20(l)
Reaction of Metallic Oxides with Acids
- Basic oxides reacts with acids to form salt and water which can be seen in the chemical equation given below.
Basic Oxide + Acid → Salt + Water
i) Lead oxide reacts with nitric acid to form lead nitrate and water.
PbO(s) + 2HNO3(aq) → Pb(NO3)2(aq) + H2O(l)
ii) Copper oxide reacts with sulfuric acid to give out copper sulfate and water.
CuO(s) + H2SO4(aq) → CUSO4(aq) + H2O(l)
iii) Sodium oxide reacts with hydrochloric acid to form sodium chloride and water.
Na2O(s) + 2HCL(aq) → 2NaCl(aq) + H2O(l)
- Basic Hydroxides reacts with acids to form salt and water which can be seen in the chemical equation given below.
Basic Hydroxide + Acid → Salt + Water
i) Potassium hydroxide reacts with carbonic acid to form potassium carbonate and water.
2KOH(aq) + H2CO3(aq) → K2CO3(aq) + 2H2O(l)
ii) Sodium hydroxide reacts with phosphoric acid to form tri-sodium phosphate and water.
3NaOH(aq) + H3PO4(aq) → Na3PO4(aq) + 3H2O(l)
iii) Ammonium hydroxide reacts with sulfuric acid to form ammonium sulfate and water.
NH4OH(aq) + H2SO4(aq) → (NH4)2SO4(aq) + 2H2O(l)
Reaction of Non Metallic Salts with Bases
The non-metallic oxides are acidic in nature as when a base such as calcium hydroxide reacts with a non-metallic oxide such as carbon dioxide then the products formed are salt and water. Thus, this is similar to the reaction between acid and base which upon reaction gives out salt and water.
Ca(OH)2(aq) + CO2(g) → CaCO3↓ + H2O(l)
Bases have a pH value greater than 7 on the pH Scale and are bitter in taste. They have a soapy or oily texture upon touch such as soaps and detergents. Strong bases are powerful enough to burn the skin so it is always advisable to be extra careful while handling the bases in the laboratory. Sometimes bases can be dissolved in water and those bases are known as “Alkali”.
Example of chemical reaction of a base (Sodium Hydroxide) in water to produce hydroxide ions.
NaOH(aq) → Na+(aq) + OH–(aq)
Some common examples of Alkalis are as follows:
- Sodium Hydroxide
NaOH(aq) → Na+(aq) + OH–(aq)
- Potassium Hydroxide
KOH(aq) → K+(aq) + OH–(aq)
- Calcium Hydroxide
Ca(OH)2(aq) → Ca2+(aq) + 2OH–(aq)
- Ammonium Hydroxide
NH4OH(aq) → NH4+(aq) + OH–(aq)
General Properties of Bases
- Has a pH value of more than 7.
- Has a bitter taste.
- Strong bases can cause burns on the skin.
- Commonly found bases in daily life are soaps and detergents.
- Caustic soda, Milk of Magnesia, Tooth paste, KOH etc. are the bases that are commonly found in laboratories.
A base is considered as a strong base if there is a high concentration of hydroxyl ions upon dissociation in water. In other words, a base that is dissolved completely in water to give a high concentration of hydroxyl ions (OH–) is known as a strong base. The number of hydroxyl ions is directly proportional to the nature of a base. KOH, NaOH and LiOH are some commonly found strong bases.
NaOH → Na+ + OH–
A base is considered as a weak base if the base is unable to dissociates completely in water, in other words if a base partially dissociates in water and gives a low concentration of hydroxyl (OH–) ions then the base is known as a weak base. Calcium hydroxide Ca(OH)2 and Ammonium hydroxide NH4OH are the two examples of weak bases.
NH4OH(aq) → NH+4(aq) + OH–(aq)
Reaction of Bases/Alkali with Ammonium Salts
Alkalis reacts with ammonium salts to liberate ammonia gas. The reaction is as follows:
Alkali + Ammonium Salt → Salt + Water + Ammonia
Other examples include:
i) Calcium hydroxide reacts with ammonium chloride to give calcium chloride with water and ammonia.
Ca(OH)2(aq) + 2NH4Cl(s) → CaCl2(s) + 2H20(l) + 2NH3(g)
ii) Sodium hydroxide reacts with ammonium sulfate to give out sodium sulfate with water and ammonia.
2NaOH(aq) + (NH4)2SO4(aq) → Na2SO4(aq) + 2H2O(l) + 2NH3(g)
The pH of a solution is defined as the negative logarithm of hydrogen ion concentration in moles per liter. The formula for pH is given below.
pH = – log [H+(aq)]
The pH scale is a continuous scale and the value of pH varies between 0 to 14. The pH of pure or neutral water is 7. Solutions having pH less than 7 are acidic in nature and the solutions with pH more than 7 are basic in nature.
- pH of Acids > 7
- pH of Bases < 7
- pH of Salts = 7
pH Value table of some common substances:
|Pure Water||7.0 (Neutral)|
Indicators are the substances that tells us about the nature of a solution. These indicators plays an important role in distinguishing between acids and bases by the unique property of changing color. The contact of any acid or base solution changes the color of indicator due to which one gets the idea if the solution is acidic or basic by observing the change in color.
Some indicators are as follows:
- Litmus – It is a purple dye that is extracted from a plant named as lichen. The purple dye is often coated on small strips of paper known as litmus paper that is quite common in laboratories. Upon contact with acid the litmus solution turns from purple to red, while with base the litmus solution turns from purple to blue.
- Turmeric – Turmeric is a common household spice that is used in cooking. This also acts as an indicator and tells us about the nature of a solution. A stain of turmeric will turn to reddish brown if soap is rubbed on it, indicating the basic nature of the soap.
- Red Cabbage Extract – The red cabbage extract can also tell us whether a solution is acidic or basic. It turns to red color when comes in contact with acid while gives yellow color in case of base.
- Onion – Onions have a characteristic smell. In basic solutions like NaOH, there is no smell. Acids however, do not destroy the smell of onions.
- Vanilla Extract – It has a pleasant smell in acidic solutions, whereas in basic solutions there is no smell.
The list of common indicators used in the lab along with the color change is as follows:
The universal indicator is basically a mixture of commonly used indicators helpful in identifying acids and bases over a wide range of pH scale accurately.
Importance of pH in Our Daily Life
pH plays a very important role in our daily life. Many companies use the pH scale to make their products. Some other importance of pH in our daily life are given below.
i) pH and Plants: Proper pH of soil is required for healthy growth of plants. It should not be too acidic or too basic.
ii) pH in the Digestive System: Human body secrets hydrochloric acid which aids in digestion. Hyper-acidity: The Condition of excess acid in the stomach. Hyper-acidity can be cured by taking ant-acid tablets or suspensions.
iii) pH and Tooth Decay: Tooth enamel which is the hardest substance in our body is corroded when the pH of the mouth is below 5.5. Cleaning of teeth using toothpaste helps in preventing tooth decay. Toothpastes are basic in nature, therefore neutralize the excess acid in the mouth and thus prevent tooth decay.
Salts and pH of Salts
Salts are the neutral products that are obtained when an acid reacts with a base. The general chemical formula to obtain salt is given below.
Acid + Base → Salt + Water
Unlike acids which only consists of the positive ions or cations, and base which only contain the negative ions or anions, salts contains both the anions and cations. In salts, the anions are called acidic radicals as these are obtained from acids whereas, the cations are called basic radicals as they are obtained from bases.
The reaction of a strong base (Sodium Hydroxide) with a strong acid (Hydrochloric Acid) is shown below:
Na+OH– + H+Cl–(aq) → NaCl(aq) + H2O(l)
The dissociation of salt in water is shown below. It gives positive ions and negative ions.
NaCl(aq) → Na+(aq) + Cl–(aq)
Family of Salts
There are three different types of salts in chemistry, acidic sat, basic salt and neutral salt. Though, here we will study about the normal salts or neutral salts.
A neutral salt is formed when there is a complete replacement of the hydrogen (of acid) ions by a metal ion. When an acid reacts with a base resulting in a neutralization reaction, then the products formed are neutral salt and water. These salts have a pH value of 7 on the pH scale. NaCl, NH4Cl, Na3PO4, Na2SO4, K2CO3 etc. are the common examples of natural/neutral salts.
Sodium Chloride (NaCl) Introduction, Properties and Uses
Sodium Chloride also know as common salt is the most commonly available salt. Seawater is the main source of sodium chloride or common salt. Seawater contains about 3.5% of soluble salts, the most common of which is sodium chloride (2.7 to 2.9%). Saline water of inland lakes is also a good source of this salt. Sodium chloride is also found as rock salt.
Common salt is generally obtained by evaporation of seawater. Crude sodium chloride is obtained by crystallization of ‘brine‘ that contains sodium sulfate, calcium sulfate, calcium chloride and magnesium chloride as impurities. Pure sodium chloride is obtained from the crude salt by dissolving it in minimum amount of water and filtering it to remove insoluble impurities. The solution is then saturated with hydrogen chloride gas, when crystals of pure sodium chloride separate out. Calcium and magnesium chlorides, being more soluble than sodium chloride, remain in solution.
Properties of Sodium Chloride
- The density of sodium chloride is 2.17 g/ml present as a white crystalline solid.
- The melting point of sodium chloride is 1080K or 807°C.
- The boiling point of sodium chloride is 1713K or 1440°C.
- It is soluble in water with a solubility of 36g per 100g of water at 273K or 0°C.
- In solid form sodium chloride does not conduct electricity at room temperature, but the molten state is capale and a very good conductor of electricity.
Uses of Sodium Chloride
- Sodium chloride is used in food as common salt.
- Na2CO3, NaOH, Cl2 etc. are manufactured with the help of sodium chloride.
- Useful in freezing of mixtures.
- In textile industries and for tanning purpose.
- Widely used as a preservative to keep fish, meat, butter and other food items preserved for long time.
Sodium Carbonate (Na2CO3) Introduction, Properties and Uses
The sodium carbonate is another commonly used salt and is used as washing soda and soda ash. Sodium carbonate exists as both anhydrous as well as hydrated form. The anhydrous form is called soda ash and the hydrated form is called washing soda (Na2CO3.10H2O).
Properties of Sodium Carbonate
- Sodium Carbonate has a very high solubility in water.
- It exists as a white crystalline solid.
- It has four forms viz. Anhydrous Salt (Na2CO3), Monohydrate Salt (Na2CO3.H2O), Heptahydrate Salt (Na2CO3.7H2o) and Decahydrate Salt (Na2CO3.10H2o).
- Upon heating it gradually loses water to form Soda Ash (Na2CO3).
Na2CO3.10H2O ——373K—→ Na2CO3.H2O ——Above 373K—→ Na2Co3 (Soda Ash)
Uses of Sodium Carbonate
- Use in the refining of petroleum and textile industry.
- Glass manufacturing is done using sodium carbonate.
- In laundries for washing purposes.
- Other sodium compounds such as Sodium silicates, sodium hydroxide, borax etc. are manufactured using sodium carbonate.
- Can be used as a cleaning agents in house.
- Used for the softening of water.
Manufacture of Sodium Carbonate
The manufacturing of Sodium Carbonate is usually done with the Ammonia-Soda process also know as the Solvay Process. Common Salt, Ammonia and Limestone are the raw materials used in Solvay process for the manufacturing of Sodium Carbonate.
The concentrated solution of brine which is saturated with ammonia is passed through carbon dioxide which produces ammonium bicarbonate.
NH4OH + H2CO3 → NH4HCO3 + H2O
Then, the ammonium bicarbonate is made to react with common salt resulting in the formation of Sodium Bicarbonate.
NH4HCO3 + NaCl → NaHCO3 + NH4Cl
Sodium bicarbonate gets precipitated and is removed by filtration process to obtain sodium carbonate by heating.
2NaHCO3 ——Δ—→ a2CO3 + H2O + CO2
The Solvay process as discussed above is used in the manufacturing of Sodium Carbonate with the help of Common Salt, Ammonia and Limestone. The Solvay Process has been shown in detail below.
› Step 1 – Ammoniacal Brine (NaCl, NH3, H2o) is made to react wit carbon dioxide to obtain hydrogen carbonate.
NaCl + NH3 + H2O +CO2 → NH4Cl +NaHCO3
› Step 2 – The obtained Sodium Hydrogen Carbonate is then heated to get Sodium Carbonate.
2NaHCO3 ——Δ—→ Na2CO3 +H2O +CO2
› Step 3 – The recrystallization of sodium carbonate is done by dissolving it into wate to obtain washing soda.
Na2CO3 + 10H2O → Na2CO3.10H2O
Also, Limestone is heated to obtain Carbon dioxide.
CaCO3 → CaO +CO2
Sodium Hydrogen Carbonate (NaHCO3) Introduction, Properties and Uses
Sodium Hydrogen Carbonate is a common household item goes by the name Baking Soda that is used in baking of cakes, pastries to make them light and fluffy. Upon heating, the baking soda generates produces carbon dioxide that makes the cake porous resulting in the softness. Other name of the Sodium Hydrogen Carbonate is Sodium Bicarbonate.
Preparation of Sodium Hydrogen Carbonate
Sodium carbonate (Na2CO3) is saturated with carbon dioxide resulting in the formation of Sodium Hydrogen Carbonate. As the NaHCO3 is less soluble in water, gets separated.
Na2CO3 + H2O + CO2 → NaHCO3
On large scale the Sodium Hydrogen Carbonate is obtained as an intermediate product in the Solvay Process (discussed above) that is used for the manufacturing of Sodium Carbonate.
Properties of Sodium Hydrogen Carbonate
- The Sodium Hydrogen Carbonate is present as white crystalline solid.
- It has a density of 2.2 g/ml.
- The taste of Sodium Hydrogen Carbonate is alkaline.
- It is not very soluble in water, though the solubility increases with the rise in temprature.
Uses of Sodium Hydrogen Carbonate
- Used as a household product for baking of cakes.
- It is used in fire extinguishers.
- Used an an antacid to neutralize the acidity of stomach.
- Can also be used as an antiseptic for skin disease.
- It is used as a chemical agent in laboratories.
Sodium Hydroxide (NaOH) Introduction, Properties and Uses
Sodium Hydroxide is a strong base and alkali that is often called Caustic Soda. It is used for a number of industrial purposes from decomposition of plant and animal tissues to making manufacturing of soap and paper. The sodium hydroxide in large quantities is made by a process known as “Chlor-alkali Process”. In the manufacturing process, chlorine gas is give at anode and hydrogen gas at the cathode and the Sodium Hydroxide solution is formed.
Properties of Sodium Hydroxide
- Sodium Hydroxide is present as a white runny solid.
- It has a melting point of 591K or 318°C.
- It has a good heat tolerance is remains stable when heated.
- It is highly soluble in water and the process is exothermic as it releases a considerable amount of heat due to the formation of various hydrates such as NaOH.H2O, NaOH.2H2O.
- It is also soluble in alcohol.
- The solution of sodium hydroxide is soapy to touch and has a bitter taste.
- It is advisable not to touch the sodium hydroxide with bare skin as it can cause severe burn due to the break down property of skin and flesh.
- Sodium Hydroxide in water is strongly alkaline and completely dissociates into Na+ and OH–.
NaOH + H2O → Na+(aq) + OH–(aq)
Uses of Sodium Hydroxide
- It is used in industries to produce paper, soap, artificial silk (Viscose Rayon) and other chemicals.
- It is also used in the refining of petroleum and other oils.
- It is also used in the extraction of Aluminum by the purification of bauxite.
- It is an ingredient in the washing powder used in washing machines.
- In laboratories, it is used as a chemical agent.
- Soda Lime is also manufactured by Sodium Hydroxide.
Plaster of Paris (CaCO4.1/2H20) Introduction, Properties and Uses
Plaster of Paris is a common construction material that is often used during construction of houses to make beautiful designs and textures on the walls or roof. The other name for the Plaster of Paris is Calcium Sulfate. When calcium sulfate is mixed with half a molecule of water per molecule of sale then Plaster of Paris is formed.
Preparation of Plaster of Paris
The basic element in the manufacturing of Plaster of Paris is Gypsum (CaSO4.2H2O). When gypsum is added at 120°C in the rotary kilns, then partial dehydration takes place.
2(CaSO4.2H2O) ——120°C—→ 2(CaSO4).H20 + 2H2O
In this process, the temperature must strictly be kept under 140°C otherwise the dehydration will continue and the property of the plaster will be reduced partially.
Properties of Plaster of Paris
- It is present in a white powder form.
- Mixing Plaster of Paris with water generates heat.
- It dries within 5 to 15 minutes after mixing with water and takes the perfect shape of the mold.
- There are two stages which is undergone by Plaster of Paris after mixing with water viz. Setting stage (First Stage) and Hardening Stage (Last Stage).
- The process of setting is shown below:
CaSO4.1/2H2o ——Setting—→ CaSO4.2H2O ——Hardening—→ CaSO4.2H20
Uses of Plaster of Paris
- It is used by doctors to fix broken or fractured bones.
- It is also used by dentists to fill the tooth cavities.
- Plaster of Paris is used widely in making statues, decorating house walls and rooftops.
- Chalks used in classrooms are made using Plaster of Paris.
- Widely used in construction industry to build structures.