Acids, Bases and Salts Class 10 Science Notes

Acids and Bases

Acids and bases are fundamental chemical substances with distinct properties. In everyday life, we encounter them in food (like sour curd or bitter neem leaves) and in household products (like cleaning agents). Understanding their chemical nature helps us explain and predict their reactions.

Indicators

Indicators are substances that change colour in the presence of an acid or a base, helping us to identify them.

Indicator	Colour in Acidic Solution	Colour in Basic Solution
Blue Litmus	Red	No change
Red Litmus	No change	Blue
Phenolphthalein	Colourless	Pink
Methyl Orange	Red	Yellow

Litmus, a purple dye extracted from lichens, is a natural indicator. In a neutral solution, it is purple.

Some substances have their odour change in acidic or basic media; these are called olfactory indicators. Examples include:

Onion: Loses odour in basic solutions, retains odour in acidic solutions.
Vanilla Essence: Retains odour in acidic solutions, loses odour in basic solutions.
Clove Oil: Retains odour in acidic solutions, loses odour in basic solutions.

How do Acids and Bases React with Metals?

Acids react with reactive metals to produce a salt and hydrogen gas ( $H_{2}$ ). This is because the metal displaces the hydrogen from the acid.

The general reaction is:

Acid + Metal \to Salt + Hydrogen gas

Example: Reaction of Zinc with dilute Hydrochloric Acid:

Z n (s) + 2 H C l (a q) \to Z n C l_{2} (a q) + H_{2} (g)

Testing for Hydrogen Gas: Hydrogen gas burns with a pop sound when a burning candle is brought near it. This is often demonstrated by passing the evolved gas through soap solution; the bubbles of hydrogen gas are then tested.

Reaction of Bases with Certain Metals: Some bases, like Sodium Hydroxide ( $N a O H$ ), also react with certain metals (like Zinc) to produce a salt and hydrogen gas.

Example: Reaction of Sodium Hydroxide with Zinc:

2 N a O H (a q) + Z n (s) \to N a_{2} Z n O_{2} (a q) + H_{2} (g)

(Sodium zincate)

Remember
Not all metals react with acids in this manner. Only metals that are more reactive than hydrogen will displace it from acids.

Exam-smart
You might be given a metal and an acid, and asked to identify the products and the type of gas evolved. You would need to recall the general equation and the test for hydrogen gas.

How do Metal Carbonates and Metal Hydrogencarbonates React with Acids?

Metal carbonates and metal hydrogencarbonates react with acids to produce a salt, carbon dioxide gas ( $C O_{2}$ ), and water ( $H_{2} O$ ).

The general reactions are:

Metal Carbonate + Acid \to Salt + Carbon Dioxide + Water

Metal Hydrogencarbonate + Acid \to Salt + Carbon Dioxide + Water

Examples:

Sodium Carbonate with Hydrochloric Acid:

N a_{2} C O_{3} (s) + 2 H C l (a q) \to 2 N a C l (a q) + H_{2} O (l) + C O_{2} (g)

Sodium Hydrogencarbonate with Hydrochloric Acid:

N a H C O_{3} (s) + H C l (a q) \to N a C l (a q) + H_{2} O (l) + C O_{2} (g)

Testing for Carbon Dioxide Gas: Carbon dioxide gas turns lime water (calcium hydroxide solution, $C a (O H)_{2}$ ) milky. This is because it forms a white precipitate of calcium carbonate ( $C a C O_{3}$ ).

The reaction with lime water is:

C a (O H)_{2} (a q) + C O_{2} (g) \to C a C O_{3} (s) + H_{2} O (l)

If excess carbon dioxide is passed, the calcium carbonate dissolves to form soluble calcium hydrogencarbonate:

C a C O_{3} (s) + H_{2} O (l) + C O_{2} (g) \to C a (H C O_{3})_{2} (a q)

Note: Limestone, chalk, and marble are different natural forms of calcium carbonate ( $C a C O_{3}$ ).

Remember
The effervescence (bubbling) observed when acids react with carbonates or hydrogencarbonates is due to the release of carbon dioxide gas.

Exam-smart
A question might describe a reaction producing effervescence that extinguishes a burning candle. You should identify the gas as $C O_{2}$ and deduce that the reactant must be a metal carbonate or hydrogencarbonate reacting with an acid.

How do Acids and Bases React with each other?

Acids and bases react with each other in a process called neutralisation. In this reaction, the acid neutralises the base, and the base neutralises the acid, forming a salt and water.

The general reaction is:

Acid + Base \to Salt + Water

Example: Reaction of Sodium Hydroxide (a base) with Hydrochloric Acid (an acid):

H C l (a q) + N a O H (a q) \to N a C l (a q) + H_{2} O (l)

(Hydrochloric Acid) + (Sodium Hydroxide) $\to$ (Sodium Chloride) + (Water)

Demonstration: If phenolphthalein indicator is added to a base (like $N a O H$ ), it turns pink. When an acid (like $H C l$ ) is added drop by drop, the pink colour disappears as neutralisation occurs. If more base is added, the pink colour reappears, showing that the solution has become basic again.

Remember
This reaction is exothermic, meaning it releases heat.

Exam-smart
You might be asked to explain why antacids are used to relieve indigestion (acidity). The reasoning would involve the neutralisation of excess stomach acid by the basic antacid.

Reaction of Metallic Oxides with Acids

Metallic oxides react with acids to produce a salt and water. This is similar to the reaction between a base and an acid. Therefore, metallic oxides are generally basic in nature.

The general reaction is:

Metallic Oxide + Acid \to Salt + Water

Example: Reaction of Copper(II) Oxide with Hydrochloric Acid:

C u O (s) + 2 H C l (a q) \to C u C l_{2} (a q) + H_{2} O (l)

(Copper(II) Oxide) + (Hydrochloric Acid) $\to$ (Copper(II) Chloride) + (Water)

The solution turns blue-green due to the formation of copper(II) chloride ( $C u C l_{2}$ ).

Remember
If a metallic oxide dissolves in an acid to form a salt and water, it behaves like a base.

Exam-smart
You could be given a reaction where a solid dissolves in acid to form a coloured solution and water, and asked to identify the nature of the solid. You should conclude it's a metallic oxide acting as a base.

Reaction of a Non-metallic Oxide with Base

Non-metallic oxides react with bases to produce a salt and water. This is similar to the reaction between an acid and a base. Therefore, non-metallic oxides are generally acidic in nature.

Example: Reaction of Carbon Dioxide (a non-metallic oxide) with Calcium Hydroxide (a base):

C O_{2} (g) + C a (O H)_{2} (a q) \to C a C O_{3} (s) + H_{2} O (l)

(Carbon Dioxide) + (Calcium Hydroxide) $\to$ (Calcium Carbonate) + (Water)

This is the same reaction that causes lime water to turn milky.

Remember
When a non-metal oxide reacts with a base, it forms a salt and water, indicating its acidic character.

Exam-smart
If you observe that a gas evolved from burning a non-metal causes a base solution to form a precipitate, you can infer that the non-metal oxide is acidic.

What do all Acids and all Bases have in Common?

All acids share common chemical properties because they all produce hydrogen ions ( $H^{+}$ ) when dissolved in water. Similarly, all bases share common chemical properties because they produce hydroxide ions ( $O H^{-}$ ) when dissolved in water.

Acids in Water Solution

When an acid is dissolved in water, it dissociates to release hydrogen ions ( $H^{+}$ ). These ions are responsible for the acidic nature of the substance. For example, when hydrochloric acid ( $H C l$ ) is dissolved in water:

H C l (a q) \to H^{+} (a q) + C l^{-} (a q)

However, hydrogen ions ( $H^{+}$ ) are highly reactive and cannot exist independently in water. They immediately combine with water molecules to form hydronium ions ( $H_{3} O^{+}$ ):

H^{+} + H_{2} O \to H_{3} O^{+}

Therefore, it is more accurate to represent the hydrogen ion in aqueous solution as $H_{3} O^{+}$ . The presence of these hydronium ions in the solution makes it acidic.

Why do some hydrogen-containing compounds not show acidic properties?

Compounds like glucose and alcohol contain hydrogen atoms, but they do not dissociate to produce $H^{+}$ ions in water. When dissolved in water, they do not conduct electricity, unlike acids. This is because they do not form ions.

Activity 2.8: Electrical Conductivity

Observation: Solutions of acids conduct electricity, causing a bulb to glow, while solutions of glucose and alcohol do not.
Reason: Acids produce mobile ions ( $H^{+}$ and anions) in water, which carry electric current. Glucose and alcohol do not produce ions in water, hence they do not conduct electricity.

Activity 2.9: Dry HCl Gas

Observation: Dry HCl gas does not change the colour of dry blue litmus paper. However, HCl gas dissolved in water (aqueous solution) turns blue litmus red.
Inference: Acids exhibit acidic character only in the presence of water. Dry HCl gas is molecular and does not produce $H^{+}$ ions. In water, HCl dissociates to form $H^{+}$ and $C l^{-}$ ions, which then form hydronium ions.

Bases in Water Solution

When a base is dissolved in water, it dissociates to produce hydroxide ions ( $O H^{-}$ ). For example:

N a O H (a q) \to N a^{+} (a q) + O H^{-} (a q)

K O H (a q) \to K^{+} (a q) + O H^{-} (a q)

M g (O H)_{2} (a q) \to M g^{2 +} (a q) + 2 O H^{-} (a q)

Alkali: Bases that are soluble in water are called alkalis. They are often described as feeling soapy to touch, having a bitter taste, and being corrosive.

Remember
Not all bases are soluble in water. Only the soluble ones are called alkalis.

What Happens When Water is Mixed with an Acid or a Base?

Mixing concentrated acids or bases with water is a highly exothermic process, meaning it releases a significant amount of heat.

Activity 2.10: Dissolving Acid/Base in Water

Observation: When concentrated sulphuric acid or sodium hydroxide pellets are added to water, the beaker becomes hot, indicating a rise in temperature.
Procedure: It is crucial to always add acid slowly to water with constant stirring. Never add water to concentrated acid. Adding water to concentrated acid can cause the heat generated to splash the corrosive mixture out of the container, potentially causing severe burns and even breaking the glass.
Dilution: The process of decreasing the concentration of ions (like $H_{3} O^{+}$ or $O H^{-}$ ) per unit volume by adding water is called dilution. The resulting solution is then called diluted acid or diluted base.

Exam-smart
A question might describe a scenario where a student is diluting an acid. You might be asked to explain the correct procedure and the reasons behind it, focusing on the exothermic nature of the reaction and the potential dangers of adding water to acid. You could also be asked to predict the effect of dilution on the concentration of ions in the solution.

How Strong are Acid or Base Solutions?

The strength of an acid or base solution refers to the concentration of hydronium ions ( $H_{3} O^{+}$ ) in an acidic solution or hydroxide ions ( $O H^{-}$ ) in a basic solution. This concentration determines how effectively the substance can donate or accept protons, and thus its reactivity.

The pH Scale

To quantitatively measure the hydrogen ion concentration ( $[H^{+}]$ ), a pH scale is used. The pH scale ranges from 0 to 14.

pH is defined as the negative logarithm (base 10) of the hydrogen ion concentration:

p H = - \log_{10} [H^{+}]

The 'p' in pH comes from the German word 'potenz', meaning 'power'.
Higher hydronium ion concentration means lower pH value.
Lower hydronium ion concentration means higher pH value.

Nature of Solutions Based on pH

pH Value	Nature of Solution	$[H^{+}]$ Concentration vs $[O H^{-}]$ Concentration
< 7	Acidic	$[H^{+}]$ > $[O H^{-}]$
= 7	Neutral	$[H^{+}]$ = $[O H^{-}]$
> 7	Basic (Alkaline)	$[H^{+}]$ < $[O H^{-}]$

A neutral solution has a pH of 7.
As the pH increases from 7 to 14, the hydroxide ion concentration ( $[O H^{-}]$ ) increases, indicating a stronger base.
As the pH decreases from 7 to 0, the hydrogen ion concentration ( $[H^{+}]$ ) increases, indicating a stronger acid.

Universal Indicator

A universal indicator is a mixture of several indicators that shows different colors at different pH values. It is often used in the form of pH paper, which is impregnated with the universal indicator.

Strength of Acids and Bases

The strength of an acid or base is determined by the extent to which it ionizes in water to produce $H^{+}$ (or $H_{3} O^{+}$ ) ions or $O H^{-}$ ions, respectively.

Strong Acids: Acids that produce a large amount of hydrogen ions in solution. For example, hydrochloric acid (HCl).
Weak Acids: Acids that produce a relatively small amount of hydrogen ions in solution. For example, acetic acid ( $C H_{3} C O O H$ ).
Strong Bases: Bases that produce a large amount of hydroxide ions in solution. For example, sodium hydroxide (NaOH).
Weak Bases: Bases that produce a relatively small amount of hydroxide ions in solution. For example, ammonium hydroxide ( $N H_{4} O H$ ).

Remember
When an acid is diluted, the concentration of $H_{3} O^{+}$ ions decreases, leading to an increase in pH. Similarly, when a base is diluted, the concentration of $O H^{-}$ ions decreases, leading to a decrease in pH.

Exam-smart
You might be given the pH of two solutions and asked to determine which is more acidic or basic, or which has a higher concentration of $H^{+}$ or $O H^{-}$ ions. Remember that lower pH means higher $[H^{+}]$ and higher pH means higher $[O H^{-}]$ .

Importance of pH in Everyday Life

The pH of a solution plays a crucial role in many biological and chemical processes.

Living Organisms: Most living organisms can survive only within a narrow range of pH, typically between 7.0 and 7.8.
Acid Rain: Rainwater with a pH less than 5.6 is considered acid rain. It lowers the pH of rivers and soil, harming aquatic life and affecting plant growth.
Soil pH: Plants require a specific pH range for healthy growth. Farmers often adjust soil pH using substances like quicklime ( $C a O$ ), slaked lime ( $C a (O H)_{2}$ ), or chalk ( $C a C O_{3}$ ) to neutralize excess acidity or alkalinity.
Digestive System: Our stomach produces hydrochloric acid (HCl) for digestion. In cases of indigestion, excess acid causes pain. Antacids, like magnesium hydroxide ( $M g (O H)_{2}$ ) (Milk of Magnesia), are mild bases used to neutralize this excess acid.
Tooth Decay: Tooth decay begins when the pH in the mouth drops below 5.5. Bacteria produce acids from food particles, which corrode tooth enamel. Basic toothpastes help neutralize these acids.
Animal and Plant Defence: Some animals and plants use chemical defence mechanisms involving acids. For instance, bee stings inject an acid, and applying a mild base like baking soda ( $N a H C O_{3}$ ) provides relief. Nettle leaves secrete methanoic acid; the juice of the dock plant, which is alkaline, is a traditional remedy.

Natural Source	Acid Present
Vinegar	Acetic acid
Curd	Lactic acid
Lemon/Orange	Citric acid
Tamarind	Tartaric acid
Ant sting	Methanoic acid
Tomato	Oxalic acid
Nettle sting	Methanoic acid

Remember
Neutralization is a key process in nature's defence and in our body's functions.

Exam-smart
Case studies might describe scenarios like acid rain affecting crops, or the use of antacids for indigestion, requiring you to apply your knowledge of pH and neutralization.

Family of Salts

Salts can be grouped into families based on the common positive ion (cation) or negative ion (anion) they contain.

Sodium family: Salts containing the sodium ion ( $N a^{+}$ ), such as sodium chloride ( $N a C l$ ), sodium sulphate ( $N a_{2} S O_{4}$ ), sodium nitrate ( $N a N O_{3}$ ), sodium carbonate ( $N a_{2} C O_{3}$ ).
Chloride family: Salts containing the chloride ion ( $C l^{-}$ ), such as sodium chloride ( $N a C l$ ), potassium chloride ( $K C l$ ), ammonium chloride ( $N H_{4} C l$ ).

pH of Salts

The pH of a salt solution depends on the strength of the acid and base from which it is formed.

Type of Salt	pH of Solution	Nature of Solution	Example
Salt of a strong acid and strong base	7	Neutral	Sodium chloride ( $N a C l$ )
Salt of a strong acid and weak base	< 7	Acidic	Ammonium chloride ( $N H_{4} C l$ )
Salt of a weak acid and strong base	> 7	Basic	Sodium carbonate ( $N a_{2} C O_{3}$ )

Acidic salts are formed from the reaction of a strong acid with a weak base. The weak base ionizes less, leaving an excess of $H^{+}$ ions in solution.
Basic salts are formed from the reaction of a weak acid with a strong base. The weak acid ionizes less, leaving an excess of $O H^{-}$ ions in solution.
Neutral salts are formed from the reaction of a strong acid and a strong base; both ions are relatively stable and do not significantly alter the pH of water.

Remember
Hydrolysis of salts is the key concept here – the reaction of the salt's ions with water to affect the pH.

Exam-smart
You might be given a list of salts and asked to identify their nature (acidic, basic, neutral) or the type of acid and base used to form them, based on their pH.

Chemicals from Common Salt

Common salt (Sodium chloride, $N a C l$ ) is a vital raw material for producing several important chemicals.

Sodium Hydroxide ( $N a O H$ )

Produced by the chlor-alkali process, which involves the electrolysis of an aqueous solution of sodium chloride (brine).
Reaction:

2 N a C l (a q) + 2 H_{2} O (l) \overset{E l e c t r o l y s i s}{\to} 2 N a O H (a q) + C l_{2} (g) + H_{2} (g)

Products:
Chlorine gas ( $C l_{2}$ ) is produced at the anode.
Hydrogen gas ( $H_{2}$ ) is produced at the cathode.
Sodium hydroxide ( $N a O H$ ) solution is formed near the cathode.
Uses: $N a O H$ is used in making soaps, detergents, paper, and textiles. $C l_{2}$ is used in water purification, bleaching, and making PVC. $H_{2}$ is used as fuel and in making ammonia.

Bleaching Powder ( $C a O C l_{2}$ )

Manufactured by passing chlorine gas over dry slaked lime ( $C a (O H)_{2}$ ).
Reaction:

C a (O H)_{2} (s) + C l_{2} (g) \to C a O C l_{2} (s) + H_{2} O (l)

(Note: The actual composition is complex, often represented as $C a (C l O)_{2} \cdot C a C l_{2} \cdot C a (O H)_{2}$ )

Uses: Bleaching textiles and paper pulp, as an oxidizing agent, and for disinfecting drinking water.

Baking Soda (Sodium Hydrogencarbonate, $N a H C O_{3}$ )

Produced using sodium chloride as a raw material.
Reaction (Solvay process):

N a C l (a q) + H_{2} O (l) + C O_{2} (g) + N H_{3} (g) \to N H_{4} C l (a q) + N a H C O_{3} (s)

Properties: It is a mild, non-corrosive basic salt.
Heating reaction: When heated, it decomposes to form sodium carbonate, water, and carbon dioxide.

2 N a H C O_{3} (s) \overset{H e a t}{\to} N a_{2} C O_{3} (s) + H_{2} O (g) + C O_{2} (g)

Uses:
Baking powder: A mixture of baking soda and a mild edible acid (like tartaric acid). The $C O_{2}$ produced makes bread and cakes soft and spongy.
Antacids: Neutralizes excess acid in the stomach.
Fire extinguishers: Reacts with acids to produce $C O_{2}$ gas, which smothers flames.

Washing Soda (Sodium Carbonate Decahydrate, $N a_{2} C O_{3} \cdot 10 H_{2} O$ )

Obtained by recrystallisation of sodium carbonate (which is produced by heating baking soda).
Reaction:

N a_{2} C O_{3} (s) + 10 H_{2} O (l) \to N a_{2} C O_{3} \cdot 10 H_{2} O (s)

Properties: It is a basic salt.
Uses:
Used in glass, soap, and paper industries.
Manufacture of other sodium compounds like borax.
Cleaning agent for domestic purposes.
Removing permanent hardness of water.

Remember
The chlor-alkali process is a cornerstone for producing essential chemicals from common salt.

Exam-smart
Questions may involve describing these processes, writing their equations, and listing the uses of the products. Understanding the role of each reactant and product is key.

Are the Crystals of Salts Really Dry?

Many crystalline salts contain a fixed number of water molecules within their crystal structure. This water is called water of crystallisation.

Hydrated Salts: Salts containing water of crystallisation.
Anhydrous Salts: Salts that have lost their water of crystallisation, usually by heating.

Copper Sulphate ( $C u S O_{4}$ )

Hydrated copper sulphate is blue ( $C u S O_{4} \cdot 5 H_{2} O$ ).
When heated, it loses water and becomes anhydrous copper sulphate, which is white ( $C u S O_{4}$ ).

C u S O_{4} \cdot 5 H_{2} O (s) \overset{H e a t}{\to} C u S O_{4} (s) + 5 H_{2} O (g)

If water is added to anhydrous copper sulphate, its blue colour is restored.

Gypsum ( $C a S O_{4} \cdot 2 H_{2} O$ )

Gypsum is a common salt containing two water molecules of crystallisation.

Plaster of Paris ( $C a S O_{4} \cdot \frac{1}{2} H_{2} O$ )

Produced by heating gypsum at 373 K (100°C).

C a S O_{4} \cdot 2 H_{2} O (s) \overset{373 K}{\to} C a S O_{4} \cdot \frac{1}{2} H_{2} O (s) + 1 \frac{1}{2} H_{2} O (g)

It is a white powder.
When mixed with water, it sets into a hard solid mass (gypsum) within a few minutes. This property is used by doctors to immobilize fractured bones.

C a S O_{4} \cdot \frac{1}{2} H_{2} O (s) + 1 \frac{1}{2} H_{2} O (l) \to C a S O_{4} \cdot 2 H_{2} O (s)

Uses: In making toys, materials for decoration, and for making surfaces smooth.

Remember
Water of crystallisation is a fixed part of the crystal structure, not just adsorbed moisture.

Exam-smart
Questions might describe heating a hydrated salt and observing a colour change, or the setting of Plaster of Paris, asking for explanations or chemical equations.

Acids, Bases and Salts — Notes

Read the notes

Acids and Bases

Indicators

How do Acids and Bases React with Metals?

How do Metal Carbonates and Metal Hydrogencarbonates React with Acids?

How do Acids and Bases React with each other?

Reaction of Metallic Oxides with Acids

Reaction of a Non-metallic Oxide with Base

What do all Acids and all Bases have in Common?

Acids in Water Solution

Bases in Water Solution

What Happens When Water is Mixed with an Acid or a Base?

How Strong are Acid or Base Solutions?

The pH Scale

Nature of Solutions Based on pH

Universal Indicator

Strength of Acids and Bases

Importance of pH in Everyday Life

Family of Salts

pH of Salts

Chemicals from Common Salt

Sodium Hydroxide ( $N a O H$ )

Bleaching Powder ( $C a O C l_{2}$ )

Baking Soda (Sodium Hydrogencarbonate, $N a H C O_{3}$ )

Washing Soda (Sodium Carbonate Decahydrate, $N a_{2} C O_{3} \cdot 10 H_{2} O$ )

Are the Crystals of Salts Really Dry?

Copper Sulphate ( $C u S O_{4}$ )

Gypsum ( $C a S O_{4} \cdot 2 H_{2} O$ )

Plaster of Paris ( $C a S O_{4} \cdot \frac{1}{2} H_{2} O$ )

All Class 10 Science chapters

Acids, Bases and Salts — Notes

Read the notes

Acids and Bases

Indicators

How do Acids and Bases React with Metals?

How do Metal Carbonates and Metal Hydrogencarbonates React with Acids?

How do Acids and Bases React with each other?

Reaction of Metallic Oxides with Acids

Reaction of a Non-metallic Oxide with Base

What do all Acids and all Bases have in Common?

Acids in Water Solution

Bases in Water Solution

What Happens When Water is Mixed with an Acid or a Base?

How Strong are Acid or Base Solutions?

The pH Scale

Nature of Solutions Based on pH

Universal Indicator

Strength of Acids and Bases

Importance of pH in Everyday Life

Family of Salts

pH of Salts

Chemicals from Common Salt

Sodium Hydroxide (NaOH)

Bleaching Powder (CaOCl2)

Baking Soda (Sodium Hydrogencarbonate, NaHCO3)

Washing Soda (Sodium Carbonate Decahydrate, Na2CO3⋅10H2O)

Are the Crystals of Salts Really Dry?

Copper Sulphate (CuSO4)

Gypsum (CaSO4⋅2H2O)

Plaster of Paris (CaSO4⋅12H2O)

All Class 10 Science chapters

Sodium Hydroxide ( $N a O H$ )

Bleaching Powder ( $C a O C l_{2}$ )

Baking Soda (Sodium Hydrogencarbonate, $N a H C O_{3}$ )

Washing Soda (Sodium Carbonate Decahydrate, $N a_{2} C O_{3} \cdot 10 H_{2} O$ )

Copper Sulphate ( $C u S O_{4}$ )

Gypsum ( $C a S O_{4} \cdot 2 H_{2} O$ )

Plaster of Paris ( $C a S O_{4} \cdot \frac{1}{2} H_{2} O$ )