Acids and Bases: Properties, Reactions, and Biological Importance

Acids and Bases

Introduction

Acids and bases are fundamental chemical substances that play crucial roles in chemistry and biology. Their properties, reactions, and effects on living organisms are essential topics in general biology and chemistry.

Acids

Definition and Characteristics

• Acid: A solution with an excess of hydrogen ions (H+).

• The term "acid" comes from the Latin word acidus, meaning "sharp" or "sour".

• The more H+ ions present, the more acidic the solution.

Properties of Acids

• Taste: Sour (e.g., lemon juice, vinegar).

• Electrical Conductivity: Acids conduct electricity in solution.

• Corrosiveness: Acids can break down certain substances, including metals, fabrics, and paper.

• Reaction with Metals: Some acids react strongly with metals, often producing hydrogen gas.

• Litmus Test: Acids turn blue litmus paper red.

Common Uses of Acids

• Acetic Acid: Found in vinegar.

• Citric Acid: Present in lemons, limes, oranges, and some candies.

• Ascorbic Acid (Vitamin C): Essential for human health.

• Sulfuric Acid: Used in fertilizers, steel production, paints, plastics, and car batteries.

Strength of Acids

• Strong Acids: Completely dissociate in water, releasing all their H+ ions. Example: Hydrochloric acid (HCl).

• Weak Acids: Do not completely dissociate in water. Example: Acetic acid (CH3COOH).

Equation for strong acid dissociation:

$\mathrm{HCl} \rightarrow \mathrm{H}^+ + \mathrm{Cl}^-$

Equation for weak acid partial dissociation:

$\mathrm{CH}_3\mathrm{COOH} \rightleftharpoons \mathrm{H}^+ + \mathrm{CH}_3\mathrm{COO}^-$

Bases

Definition and Characteristics

• Base: A solution with an excess of hydroxide ions (OH-).

• Another term for base is alkali.

• Bases are substances that can accept hydrogen ions.

Properties of Bases

• Taste: Bitter.

• Texture: Slippery or soapy feel.

• Corrosiveness: Strong bases are caustic and can damage organic tissue.

• Electrical Conductivity: Bases conduct electricity in solution.

• Litmus Test: Bases turn red litmus paper blue.

• Reaction with Metals: Bases generally do not react with metals.

Common Uses of Bases

• Sodium Hydroxide (NaOH): Used in soaps and cleaning products.

• Ammonia (NH3): Found in household cleaners.

• Chalk and Oven Cleaner: Contain basic substances.

• Blood: Human blood is slightly basic.

Strength of Bases

• Strong Bases: Completely dissociate in water, releasing all their OH- ions. Example: Sodium hydroxide (NaOH).

• Weak Bases: Do not completely dissociate in water. Example: Ammonia (NH3).

pH Scale

Measuring Acidity and Alkalinity

• pH: A measure of how acidic or basic a solution is.

• The pH scale ranges from 0 (very acidic) to 14 (very basic), with 7 being neutral.

• Acidic solutions: pH below 7.

• Neutral solutions: pH of 7 (e.g., pure water).

• Basic solutions: pH above 7.

• Each change of 1 pH unit represents a tenfold change in H+ concentration.

pH Equation:

$\mathrm{pH} = -\log_{10}[\mathrm{H}^+]$

Examples of pH Values

Substance	Approximate pH
Gastric acid	1-2
Lemon juice	2
Tomato juice	4
Pure water	7
Sea water	8
Baking soda	9
Ammonia solution	11
Bleach	13

Acid-Base Reactions

Neutralization

• A reaction between an acid and a base is called neutralization.

• The products of neutralization are typically a salt and water.

• The resulting solution is less acidic or basic than the original reactants.

General Neutralization Equation:

$\mathrm{Acid} + \mathrm{Base} \rightarrow \mathrm{Salt} + \mathrm{Water}$

Example:

$\mathrm{HCl} + \mathrm{NaOH} \rightarrow \mathrm{NaCl} + \mathrm{H}_2\mathrm{O}$

Common Salts Formed by Acid-Base Reactions

Salt	Uses
Sodium chloride (NaCl)	Table salt, food flavoring, preservative
Potassium iodide (KI)	Added to "iodized" salt to prevent iodine deficiency
Calcium chloride (CaCl2)	De-icer for roads, food additive
Potassium chloride (KCl)	Salt substitute in foods
Calcium carbonate (CaCO3)	Found in limestone, antacids
Ammonium nitrate (NH4NO3)	Fertilizer, instant cold packs

Buffers

Definition and Biological Importance

• Buffer: A solution that resists changes in pH when small amounts of acid or base are added.

• Buffers are essential for maintaining homeostasis in living organisms.

• They act like chemical sponges, soaking up excess H+ or OH- ions to keep pH constant.

• Once a buffer is "used up," it can no longer regulate pH.

Example: Antacids are buffers made of calcium carbonate (CaCO3).

Acid/Base Balance in Biology

Homeostasis and pH Regulation

• Maintaining a stable internal pH is vital for biological processes.

• Human blood must stay near pH 7.35–7.45 for proper function.

• The stomach requires an acidic environment (pH 2–3) for digestion.

• Disruption of acid-base balance can lead to health problems.

Dissociation of Water

Formation of Acids and Bases

• Water can dissociate into hydrogen ions and hydroxide ions:

$\mathrm{H}_2\mathrm{O} \rightleftharpoons \mathrm{H}^+ + \mathrm{OH}^-$

• Neutral water has equal concentrations of H+ and OH-.

• Acids increase the concentration of H+ in solution.

• Bases increase the concentration of OH- in solution.

• Acids and bases neutralize each other when mixed.

Salts

Definition and Examples

• Salt: An ionic compound formed from the reaction of an acid and a base, producing cations other than H+ and anions other than OH-.

• The most familiar salt is sodium chloride (NaCl), or table salt.

• Other examples include baking soda (NaHCO3) and Epsom salts (MgSO4).

The Role of Buffers in Salts

• Certain salts act as buffers, combining with excess H+ or OH- to reduce acidity or alkalinity.

• Buffers are vital for maintaining pH in organisms and can be "used up" over time.

Example: Calcium carbonate (CaCO3) acts as a buffer in antacids and natural systems.