Acids and Bases for Class 10

Acids

Acids are substances that can release positively charged hydrogen ions (H⁺) when they are dissolved in water. These hydrogen ions are what give acids their unique properties.

Properties of Acids

1. Taste: Acids have a sour or tart taste. However, it's essential to note that tasting acids can be dangerous, so it's not recommended.
2. Odour: Some acids may have distinctive odours. For example, acetic acid (found in vinegar) has a characteristic odour.
3. State of Matter: Acids can exist in various physical states, including liquids (e.g., hydrochloric acid, sulphuric acid), solids (e.g., citric acid in citrus fruits), and gases (e.g., carbonic acid in carbonated beverages).
4. Solubility and Dissociation in Water: Acids readily dissolve in water, which means they can mix with water to form a solution. This property is essential because many chemical reactions involving acids occur in aqueous (water-based) solutions. When acids dissolve in water, they dissociate (or ionise) to form hydrogen ions along with other ions specific to the acid.
   For example, for HCl: HCl (aq) → H⁺ (aq) + Cl^- (aq)
   In this reaction, the HCl molecule dissociates into two ions: a positively charged hydrogen ion (H⁺) and a negatively charged chloride ion (Cl^-).
   The positively charged hydrogen ions (H⁺) are key to the acidic properties of the solution. These hydrogen ions are responsible for the sour taste, corrosive nature, and other characteristics of acids. In the absence of water or an aqueous environment, a substance will not form hydrogen ions (H⁺) or hydronium ions (H₃O⁺), and therefore, it will not exhibit its acidic behaviour.

Learn more about Chemical Reactions and Equations

5. Formation of Hydronium Ions (H₃O⁺): It's important to note that in aqueous solutions, hydrogen ions (H⁺) don't exist freely as H+ ions. Instead, they tend to attach themselves to water molecules (H₂O) to form hydronium ions (H₃O⁺). This happens because water is a polar molecule, and the positive hydrogen ion is attracted to the partially negative oxygen atom in water.
   The reaction can be represented as H⁺ + H₂O → H₃O⁺
   So, you can consider H⁺ (aq) and H₃O⁺(aq) to be equivalent; they represent the same thing in the solution.
6. Electrical Conductivity: When acids are dissolved in water, they can conduct electricity. This is because they release those hydrogen ions (H⁺) into the solution, creating a flow of electric charge.
7. Corrosiveness: Many acids have a corrosive nature, meaning they can chemically attack and deteriorate certain materials, including metals, organic substances, and even skin. This property is often harnessed for industrial applications, such as in the pickling of metals.
8. pH Level: Acids have pH values below 7 on the pH scale, with lower pH values indicating stronger acidity. The pH scale is a measure of the concentration of hydrogen ions (H⁺) in a solution, and acids have a higher concentration of these ions.

Strong and Weak Acids

Strong and weak acids are categories used to describe how effectively an acid ionises or dissociates in water, leading to the production of hydrogen ions (H⁺ or hydronium ions, H₃O⁺). The key difference between strong and weak acids lies in the extent to which they ionise in an aqueous solution:

Strong Acids

1. Complete Ionisation: Strong acids completely ionise or dissociate when dissolved in water, meaning nearly all of their molecules break apart into ions.
2. Large Concentration of Hydrogen Ions: As a result of complete ionisation, strong acids produce a high concentration of hydrogen ions (H⁺ or H₃O⁺) in the solution.
3. Conduct Electricity Well: Due to the high concentration of ions, solutions of strong acids are good conductors of electricity.
4. Low pH: Strong acids have a low pH (usually less than 2), indicating their strong acidity.
5. Examples: Common examples of strong acids include hydrochloric acid (HCl), sulphuric acid (H₂SO₄), and nitric acid (HNO₃).

Weak Acids

1. Partial Ionisation: Weak acids only partially ionise or dissociate when dissolved in water, which means only a small fraction of their molecules break apart into ions.
2. Lower Concentration of Hydrogen Ions: Because of partial ionisation, weak acids produce a relatively low concentration of hydrogen ions (H⁺ or H₃O⁺) in the solution.
3. Conduct Electricity Less Effectively: Due to the lower concentration of ions, solutions of weak acids are poorer conductors of electricity compared to strong acids.
4. Slightly Acidic pH: Weak acids have a pH that is slightly acidic but closer to neutral (typically between 3 and 6).
5. Examples: Common examples of weak acids include acetic acid (found in vinegar, CH₃COOH) and carbonic acid (H₂CO₃).

Concentrated and Dilute Acid

Concentrated Acid

1. A concentrated acid is an acid solution that contains a high proportion of the acid itself and a minimal amount of water.
2. These solutions are typically very strong and can be highly corrosive and dangerous.
3. Examples include concentrated sulphuric acid (H₂SO₄) and concentrated hydrochloric acid (HCl).

Dilute Acid

1. A dilute acid is an acid solution that has been mixed with a significant amount of water, reducing the concentration of the acid.
2. Dilute acids are less concentrated and are generally safer to handle than concentrated acids.

Dilutions of Acids

The dilution of concentrated acids is an exothermic process and can release intense heat. The heat produced during the process arises from the strong attraction between the acid molecules. As the concentrated acid molecules mix with water, the attraction between them is weakened, and energy is released in the form of heat. The dilution should always be done with care and following specific safety guidelines.

1. Gradual Addition: Concentrated acid should be added gradually to water while stirring. This process is exothermic, meaning it releases heat, and adding acid to water helps dissipate this heat more safely.
2. Never Add Water to Acid: It is crucial to avoid adding water to concentrated acid. Adding water to concentrated acid can cause a sudden release of heat, potentially leading to a violent reaction, splattering of the acid, or even a container rupture.
3. Safety Precautions: When diluting concentrated acids, safety precautions, such as wearing appropriate personal protective equipment (PPE), including safety goggles, lab coats, and gloves, should be taken.
4. Ventilation: Dilution should be performed in a well-ventilated area or in a fume hood to prevent inhaling any potentially harmful fumes.

Reactions of Acids

1. Reaction with Metals

Acids react with certain metals to produce hydrogen gas and a metal salt. For example, when hydrochloric acid (HCl) reacts with zinc (Zn), it forms zinc chloride (ZnCl₂) and releases hydrogen gas (H₂):

2HCl + Zn → ZnCl₂ + H₂

2. Reaction with Metal Carbonates

Acids react with metal carbonates to produce salt, carbon dioxide gas, and water. For instance, when sulphuric acid (H₂SO₄) reacts with calcium carbonate (CaCO₃), it forms calcium sulphate (CaSO₄), carbon dioxide (CO₂), and water (H₂O):

H₂SO₄ + CaCO₃ → CaSO₄ + CO₂ + H₂O

Test for Carbon Dioxide (CO₂): When carbon dioxide gas is passed through lime water (calcium hydroxide, Ca(OH)₂), it turns milky due to the formation of calcium carbonate (CaCO₃) as a white precipitate:

Ca(OH)₂ (aq) + CO₂ (g) → CaCO₃ (s) + H₂O (l)

Excess carbon dioxide causes the milkiness to disappear, forming soluble calcium hydrogen carbonate.

3. Reaction with Metal Hydrogen Carbonates (Bicarbonates)

Acids also react with metal hydrogen carbonates to produce salt, carbon dioxide gas, and water. An example is the reaction between hydrochloric acid (HCl) and sodium bicarbonate (NaHCO₃) to form sodium chloride (NaCl), carbon dioxide (CO₂), and water (H₂O):

2NaHCO₃ + 2HCl → 2NaCl + 2CO₂ + 2H₂O

4. Reaction with Metal Oxides

Acids can react with certain metal oxides to form salts and water. For example, when copper oxide (CuO) reacts with hydrochloric acid (HCl), it produces copper chloride (CuCl₂) and water (H₂O):

CuO + 2HCl → CuCl₂ + H₂O

5. Reaction with Bases

The reaction of acids with bases is known as a neutralisation reaction. This chemical reaction involves the combination of an acid and a base to produce water and salt. The key components of this reaction are hydrogen ions (H⁺) from the acid and hydroxide ions (OH^-) from the base. Here's the general equation for a neutralisation reaction:

Acid + Base → Water + Salt

For example, if hydrochloric acid (HCl) reacts with sodium hydroxide (NaOH), the resulting salt is sodium chloride (NaCl):

HCl (aq) + NaOH (aq) → NaCl (aq) + H₂O (l)

In this reaction, the hydrogen ion (H⁺) from HCl combines with the hydroxide ion (OH^-) from NaOH to form water (H₂O), and the sodium ion (Na⁺) from NaOH combines with the chloride ion (Cl^-) from HCl to form sodium chloride (NaCl), which is a salt.

Uses of Acids

Mineral acids, such as sulphuric acid, nitric acid, and hydrochloric acid, find various industrial applications:

1. Manufacture of fertilisers, including ammonium sulphate.
2. Production of paints, dyes, and chemicals.
3. Used in the plastics industry.
4. Essential in the production of synthetic fibres.
5. An ingredient in detergents.
6. A key component in the production of explosives.
7. Used in the manufacturing of car batteries.

1. Used in the production of fertilisers, including ammonium nitrate.
2. Essential in the manufacturing of explosives like TNT (Trinitrotoluene).
3. A crucial component in the dye industry.
4. Used in the production of plastics.

1. Employed for removing oxide films from steel objects, particularly before galvanisation.
2. Used to eliminate scale deposits from inside boilers.
3. Applied in various industries, including dyeing, textiles, food, and leather processing.

Bases

Bases, also known as alkaline substances, have distinct properties that help distinguish them from other types of chemicals.

Properties of Bases

1. Bitter Taste: Bases typically have a bitter taste, which is quite different from the sour taste of acids. However, it's important to note that many bases are not safe to taste due to their corrosive nature.
2. Slippery or Soapy Feel: When you touch a base or its solution, it feels slippery or soapy. This tactile sensation is often used as a simple test to identify bases. It's important to handle strong bases with caution as they can be caustic and cause skin irritation.
3. Turn Red Litmus Blue: Bases have the remarkable ability to change the colour of red litmus paper to blue. Litmus paper is a widely used indicator to determine whether a substance is acidic or basic, and this colour change is a characteristic property of bases.
4. State at Room Temperature: Some common bases exist as solids at room temperature, such as sodium hydroxide and potassium hydroxide. Others, like ammonia (NH₃) and sodium hydroxide solution (liquid form), are in liquid states at room temperature.
5. Production of Hydroxide Ions (OH^-): One of the defining characteristics of bases is their capacity to produce hydroxide ions (OH^-) when they are dissolved in water. This property distinguishes them from acids, which produce hydrogen ions (H⁺) instead.
   For example: Sodium hydroxide (NaOH) dissolves in water to form sodium ions (Na⁺) and hydroxide ions (OH^-):
   NaOH → Na⁺ + OH^-
   The presence of hydroxide ions gives the solution its alkaline nature.
6. Electrical Conductivity: Bases, when dissolved in water, conduct electricity because they release ions, including hydroxide ions (OH^-), which can carry an electrical charge. This property makes them electrolytes in aqueous solutions.
7. Solubility in Water: Bases can vary in their solubility in water. While some bases are highly soluble in water, forming clear and homogeneous solutions, others may have limited solubility, resulting in cloudy or partially dissolved mixtures. Bases that readily dissolve in water are often referred to as alkalis, and they form alkaline solutions.
8. pH Level: Aqueous solutions of bases have pH values greater than 7 on the pH scale.

Strong and Weak Bases

Strong bases and weak bases are categories used to describe how completely a base ionises in water and, consequently, the concentration of hydroxide ions (OH^-) it produces in solution.

Strong Bases

1. Complete Ionisation: A strong base is a base that completely ionises (dissociates) in water to yield a high concentration of hydroxide ions (OH^-) in solution. In other words, strong bases fully break apart into ions when mixed with water.
2. Examples: Sodium hydroxide (NaOH) and potassium hydroxide (KOH) are classic examples of strong bases.
   When strong bases like NaOH or KOH are dissolved in water, they dissociate entirely into their constituent ions, producing a large number of hydroxide ions:
   NaOH → Na⁺+ OH^-
   KOH → K⁺ + OH^-
3. Conductivity: Solutions of strong bases are excellent conductors of electricity due to the high concentration of hydroxide ions.

Weak Bases

1. Partial Ionisation: A weak base is a base that partially ionises (dissociates) in water, resulting in a relatively low concentration of hydroxide ions (OH^-) in the solution. Weak bases do not completely break apart into ions when mixed with water.
2. Examples: Ammonium hydroxide (NH₄OH), calcium hydroxide [Ca(OH)₂], and magnesium hydroxide [Mg(OH)₂] are examples of weak bases.
   They only partially ionise, yielding fewer hydroxide ions:
   NH₄OH ⇌ NH₄⁺ + OH^- (partial ionisation)
   Ca(OH)₂ ⇌ Ca²⁺ + 2OH^- (partial ionisation)
   Mg(OH)₂ ⇌ Mg²⁺ + 2OH^- (partial ionisation)
3. Conductivity: Solutions of weak bases have lower electrical conductivity than strong bases because of the lower concentration of hydroxide ions.

Reaction of Bases

Bases, also known as alkalis when they are water-soluble, participate in various chemical reactions. These reactions involve the ability of bases to accept protons (H⁺ ions) or donate hydroxide ions (OH^- ions).

1. Neutralisation Reactions with Acids

When a base reacts with an acid, a neutralisation reaction occurs. In this reaction, the base's hydroxide ions (OH^-) and the acid's hydrogen ions (H⁺) combine to form water (H₂O) and salt. The resulting solution is neutral, with a pH of 7.

2. Reaction with Metals

Bases can react with certain metals, particularly active metals like aluminium and zinc, to produce hydrogen gas (H₂) and a metal salt. The metal displaces hydrogen ions from the base.

Example: Sodium hydroxide reacts with aluminium to produce aluminium hydroxide [Al(OH)₃] and hydrogen gas (H₂):

2NaOH + 2Al → 2NaAlO₂ + 3H₂

This reaction can be used to test for the presence of a base. When a metal reacts with a base, you will observe the evolution of bubbles of hydrogen gas.

3. Reaction with Ammonium Salts

Bases can react with ammonium salts to produce ammonia gas (NH₃), water (H₂O), and the corresponding salt. This reaction is often used to identify the presence of ammonium ions (NH₄⁺) in a solution.

Example: Ammonium hydroxide (NH₄OH) reacts with ammonium chloride (NH₄Cl) to produce ammonia gas (NH₃), water (H₂O), and ammonium chloride (NH₄Cl):

NH₄OH + NH₄Cl → NH₃ + H₂O + NH₄Cl

The release of ammonia gas is characterised by its pungent odour.

4. Reaction with Non-metal Oxides

Bases react with non-metal oxides to form metal salts and water. These reactions are similar to neutralisation reactions with acids, but instead of hydrogen ions, they involve the hydroxide ions in the base reacting with non-metallic oxides.

Example: Calcium hydroxide (Ca(OH)₂) reacts with carbon dioxide (CO₂) to produce calcium carbonate (CaCO₃) and water (H₂O).

Ca(OH)₂ + CO₂ → CaCO₃ + H₂O

This reaction is used in various applications, including the removal of acidic gases from industrial processes.

Uses of Some Common Bases

1. Used in soap production.
2. Essential in paper manufacturing and rayon production.

1. Used in making bleaching powder.

1. Acts as an antacid to relieve indigestion.

1. Used as washing soda and for water softening.

1. Used in cooking and baking.
2. An ingredient in baking powder.
3. Can be used as an antacid.
4. Found in some fire extinguishers.

pH Scale

The pH scale is a numerical scale that measures the acidity or basicity (alkalinity) of a solution. It ranges from 0 to 14, with 7 being neutral.

1. Neutral (pH 7): Pure water has a pH of 7 and is considered neutral because it contains an equal concentration of hydrogen ions (H⁺) and hydroxide ions (OH^-). Substances with a pH of 7 are neither acidic nor basic.
2. Acidic (pH < 7): Solutions with a pH less than 7 are acidic. The lower the pH, the stronger the acidity. For example, a solution with a pH of 2 is more acidic than a solution with a pH of 5. Acidic solutions have a higher concentration of hydrogen ions H⁺) than hydroxide ions (OH^-). They turn blue litmus paper red and may contain hydrogen ions.
3. Basic (pH > 7): Solutions with a pH greater than 7 are basic (or alkaline). The higher the pH, the stronger the basicity. For example, a solution with a pH of 12 is more basic than a solution with a pH of 9. Basic solutions have a higher concentration of hydroxide ions (OH^-) than hydrogen ions (H⁺). They turn red litmus paper blue and may contain hydroxide ions.

Calculation of pH

The pH scale is logarithmic, meaning each whole number change on the scale represents a tenfold difference in the concentration of hydrogen ions (H⁺) or hydroxide ions (OH^-). For example, a solution with a pH of 4 is ten times more acidic than a solution with a pH of 5.

The pH of a solution is calculated using the formula:

pH = - log₁₀[H⁺]

In pure water, the concentration of hydrogen ions ([H⁺]) is equal to the concentration of hydroxide ions ([OH^-]), and both are approximately 10^-7 mol/L. Therefore, the pH of pure water is 7, which is considered neutral.

Importance of pH in Everyday Life

The pH scale, which measures the acidity or alkalinity of a substance, plays a significant role in various aspects of everyday life.

1. Digestive Health: The human digestive system relies on maintaining specific pH levels for optimal function. The stomach's acidic environment, with a pH of around 1.4 to 3.5, helps break down food and kill harmful bacteria. Proper pH balance is crucial for preventing indigestion and ensuring effective digestion.
2. Oral Health: Maintaining the right pH in the mouth is vital for oral health. When we consume sugary foods, bacteria produce acids, which can lead to tooth decay if the mouth's pH drops below 5.5. Toothpaste and mouthwash with a slightly basic pH (around 8.0) can help neutralise mouth acids, protecting teeth from decay.
3. Agriculture: Soil pH is a critical factor in agriculture. Most plants thrive in soil with a pH close to 7 (neutral). Soil that is too acidic or alkaline can affect nutrient availability to plants, leading to stunted growth or crop failure. Farmers adjust soil pH using additives like lime (to raise pH) or organic matter (to lower pH) to optimise plant growth.
4. Aquatic Ecosystems: pH levels in aquatic ecosystems, such as lakes and rivers, impact aquatic life. Acid rain, which results from air pollution, can lower the pH of water bodies, harming fish and other aquatic organisms. Monitoring and controlling pH in aquatic environments are essential for protecting aquatic ecosystems.
5. Body Function: The human body functions optimally within a narrow pH range, typically around 7.0 to 7.8. Disturbances in pH can lead to health issues. Blood, for example, has a tightly regulated pH of about 7.4. Deviations from this pH can affect enzyme activity, cellular function, and overall health.
6. Cosmetics and Personal Care Products: Many cosmetic and personal care products, like shampoos and skincare items, are formulated to have a specific pH to match the body's natural pH. This helps maintain healthy skin and hair.
7. Chemical Reactions: pH influences chemical reactions in various industries, including water treatment, manufacturing, and food production. Proper pH control ensures the desired outcomes in these processes.
8. Environmental Conservation: Monitoring pH levels in natural water bodies, soils, and ecosystems is crucial for environmental conservation efforts. Changes in pH can signal pollution or imbalances in these systems, allowing for timely intervention.

Indicators

An indicator is a substance that undergoes a noticeable colour change when it comes into contact with an acidic or basic solution. It serves as a visual or sensory signal to determine whether a substance is acidic or basic. If the indicator shows no colour change, the substance is considered neutral.

Indicators can be classified into two main categories:

1. Natural Indicators: These are indicators derived from natural sources, such as plants or foods. Examples include litmus (from lichen), turmeric (from the turmeric plant), and red cabbage juice (from red cabbage leaves). They change colour in response to pH changes.
2. Chemical (Synthetic) Indicators: These are artificially synthesised compounds designed to exhibit specific colour changes at different pH levels. Common examples are phenolphthalein and methyl orange. These chemical indicators are often used in laboratory settings for precise pH measurements.

Indicators play a crucial role in qualitative chemical analysis, allowing scientists and chemists to quickly identify whether a solution is acidic, basic or neutral. They are valuable tools for understanding the properties of substances and their pH characteristics.

Universal Indicator

A universal indicator is a special type of indicator used to measure the pH (acidity or basicity) of a solution and determine its relative strength as a strong acid, weak acid, strong base, or weak base. It is a mixture of several different indicators or dyes, each of which changes colour at a different pH value. Universal indicators can provide a more detailed and accurate measurement of the pH compared to common indicators like litmus.

For example, if the universal indicator paper turns dark red, the solution likely has a pH of around 0, indicating a strong acid. If the paper turns orange, the pH is approximately 4, suggesting a weak acid. Blue colours indicate basic solutions, with stronger bases producing darker shades of blue or violet.

Frequently Asked Questions

1. Why do acids conduct electricity in an aqueous solution but not in a dry state?

In an aqueous solution, acids dissociate into ions which are responsible for conducting electricity. In the dry state, acids do not have free ions and hence cannot conduct electricity.

2. How can you demonstrate that hydrochloric acid is a stronger acid than acetic acid using electrical conductivity?

Since hydrochloric acid dissociates completely in water (strong acid) and acetic acid only partially dissociates (weak acid), a solution of hydrochloric acid will have more ions, making it a better conductor of electricity than acetic acid. You can measure the conductivity of both solutions using a conductivity meter, and hydrochloric acid should show higher conductivity.

3. Why do some metals like copper not react with dilute hydrochloric acid, while metals like zinc react vigorously?

Metals like copper are less reactive than hydrogen and cannot displace hydrogen from acids like hydrochloric acid. However, zinc is more reactive and can displace hydrogen from the acid, resulting in a vigorous reaction that produces hydrogen gas.

4. Why does sodium hydroxide feel slippery when touched, but concentrated solutions can cause burns?

Sodium hydroxide feels slippery because it reacts with the oils and fats on your skin to form soap-like compounds (saponification). However, concentrated solutions are highly corrosive and can damage skin tissue, causing chemical burns due to their strong alkaline nature.

5. How does a universal indicator differ from phenolphthalein or litmus in determining the pH of a solution?

A universal indicator provides a full range of colours that correspond to different pH values across the entire pH scale, from red (acidic) to violet (basic). Phenolphthalein and litmus only show colour changes over a narrow pH range, making universal indicators more versatile for determining exact pH values.