Acids, bases & their properties

Acids, bases & their properties

{{TABLE: title=Acids vs. Bases: A Quick Comparison

Welcome to your complete guide on acids, bases, and their properties. You've likely encountered these substances in your daily life—from the sharp taste of a lemon to the soapy feel of a detergent. But what makes an acid an acid, and a base a base? Let's dive deep into the chemistry that defines them, their reactions, and their importance in our world.

At its core, the study of acids and bases is the study of chemical reactivity, often revolving around the movement of tiny particles: protons (hydrogen ions, H⁺) and electrons. Over time, chemists have developed several theories to define and explain their behaviour, each one building upon the last to give us a more complete picture.

Defining Acids and Bases: Three Key Theories

Understanding acids and bases starts with learning how scientists have defined them over the years. There isn't just one definition; there are three main theories, each useful in different contexts.

1. The Arrhenius Theory (The Water-Based Definition)

This is the oldest and simplest theory, proposed by Svante Arrhenius in 1884. It defines acids and bases based on their behaviour in water.

• An Arrhenius acid is a substance that ionises in water to produce hydrogen ions (H⁺).
  • Example: Hydrochloric acid (HCl) dissolves in water: HCl(aq) → H⁺(aq) + Cl⁻(aq)
• An Arrhenius base is a substance that ionises in water to produce hydroxide ions (OH⁻).
  • Example: Sodium hydroxide (NaOH) dissolves in water: NaOH(aq) → Na⁺(aq) + OH⁻(aq)

This theory is excellent for many common situations but has a key limitation: it only applies to aqueous (water-based) solutions and doesn't explain the basicity of substances like ammonia (NH₃), which don't contain OH⁻ ions.

2. The Brønsted-Lowry Theory (The Proton-Transfer Definition)

In 1923, Johannes Brønsted and Thomas Lowry independently proposed a more general definition. This theory focuses on the transfer of protons (H⁺ ions).

• A Brønsted-Lowry acid is a proton (H⁺) donor.
• A Brønsted-Lowry base is a proton (H⁺) acceptor.

Consider the reaction between hydrochloric acid (HCl) and ammonia (NH₃): HCl + NH₃ → NH₄⁺ + Cl⁻ In this reaction, HCl donates a proton to NH₃. Therefore, HCl is the acid, and NH₃ is the base. This theory cleverly explains why ammonia is a base without needing it to be in water or to produce OH⁻ ions directly. It simply accepts a proton. A key concept that emerges from this theory is the idea of conjugate pairs.

{{KEY: type=concept | title=Conjugate Acid-Base Pairs | text=When a Brønsted-Lowry acid donates a proton, the species that remains is called its conjugate base. When a Brønsted-Lowry base accepts a proton, the species that is formed is called its conjugate acid. In the reaction HCl + NH₃ ↔ NH₄⁺ + Cl⁻, Cl⁻ is the conjugate base of the acid HCl, and NH₄⁺ is the conjugate acid of the base NH₃. They always exist in pairs.}}

3. The Lewis Theory (The Electron-Pair Definition)

The most general and comprehensive theory was proposed by G.N. Lewis. It shifts the focus from protons to electron pairs.

• A Lewis acid is an electron-pair acceptor. These species are typically electron-deficient.
• A Lewis base is an electron-pair donor. These species have at least one lone pair of electrons to donate.

Let's look at the reaction between boron trifluoride (BF₃) and ammonia (NH₃): BF₃ + :NH₃ → F₃B←:NH₃ Here, the nitrogen atom in ammonia has a lone pair of electrons, which it donates to the electron-deficient boron atom in BF₃. Thus, NH₃ is the Lewis base, and BF₃ is the Lewis acid. This theory is broad enough to cover reactions that don't even involve hydrogen!

Properties of Acids

Acids exhibit a distinct set of physical and chemical properties that allow us to identify and classify them.

Physical Properties of Acids

• Taste: They generally have a sour taste. Think of citric acid in lemons or acetic acid in vinegar. (Warning: Never taste chemicals in the lab!)
• Corrosive Nature: Strong acids are highly corrosive. They can cause severe burns on skin and can eat through materials like metal and cloth.
• Electrical Conductivity: Acidic solutions are good conductors of electricity because they release mobile ions (H⁺ and the anion) when dissolved in water.

Chemical Properties of Acids

This is where the action happens! Acids participate in several important types of reactions.

1. Reaction with Metals Acids react with most metals (those above hydrogen in the reactivity series) to produce a salt and hydrogen gas.

• General Equation: Acid + Metal → Salt + Hydrogen gas (H₂)
• Example: Hydrochloric acid reacts with magnesium: 2HCl(aq) + Mg(s) → MgCl₂(aq) + H₂(g)
• The 'Pop' Test: The presence of hydrogen gas can be confirmed by bringing a burning splint near the mouth of the test tube. The gas extinguishes the splint with a characteristic 'pop' sound.

{{VISUAL: diagram: A labelled laboratory setup showing zinc granules being dropped into a test tube of dilute sulphuric acid. Bubbles of hydrogen gas are shown rising and being collected over water in an inverted trough. A burning splint is being brought to the gas, showing a 'pop' sound icon.}}

2. Reaction with Metal Carbonates and Metal Hydrogen Carbonates Acids react with metal carbonates and hydrogen carbonates to produce a salt, water, and carbon dioxide gas.

• General Equation:
  • Acid + Metal Carbonate → Salt + Water + Carbon Dioxide (CO₂)
  • Acid + Metal Hydrogen Carbonate → Salt + Water + Carbon Dioxide (CO₂)
• Examples:
  • With sodium carbonate: 2HCl(aq) + Na₂CO₃(s) → 2NaCl(aq) + H₂O(l) + CO₂(g)
  • With sodium bicarbonate (baking soda): HCl(aq) + NaHCO₃(s) → NaCl(aq) + H₂O(l) + CO₂(g)
• The Limewater Test: The evolved carbon dioxide gas can be tested by passing it through limewater (a solution of calcium hydroxide, Ca(OH)₂). The limewater turns milky or cloudy due to the formation of a white precipitate of calcium carbonate (CaCO₃).

3. Reaction with Bases (Metal Oxides and Hydroxides) This is one of the most fundamental reactions in chemistry: neutralisation. An acid reacts with a base to form a salt and water.

• General Equation: Acid + Base → Salt + Water
• Example: Hydrochloric acid reacts with sodium hydroxide: HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l) In this reaction, the acidic properties of HCl and the basic properties of NaOH are destroyed or 'neutralised'. The resulting solution of sodium chloride (common salt) is neutral.

Properties of Bases

Bases have their own unique set of properties that often contrast with those of acids.

Physical Properties of Bases

• Taste: They have a bitter taste.
• Feel: Aqueous solutions of bases feel soapy or slippery to the touch. This is because they react with the fatty acids and oils on your skin to form soap-like substances.
• Electrical Conductivity: Like acids, basic solutions conduct electricity due to the presence of mobile ions (OH⁻ and the cation).

Chemical Properties of Bases

1. Reaction with Acids (Neutralisation) As we've seen, this is a defining reaction. Bases neutralise acids to form a salt and water.

• Example: Sulphuric acid reacts with potassium hydroxide: H₂SO₄(aq) + 2KOH(aq) → K₂SO₄(aq) + 2H₂O(l)

2. Reaction with Non-metallic Oxides Bases react with non-metallic oxides (which are acidic in nature) to form a salt and water. This is similar to a neutralisation reaction.

• General Equation: Base + Non-metallic Oxide → Salt + Water
• Example: Carbon dioxide (a non-metallic oxide) reacts with sodium hydroxide: 2NaOH(aq) + CO₂(g) → Na₂CO₃(aq) + H₂O(l)

{{KEY: type=exam | title=Acidic and Basic Oxides | text=A common exam question involves identifying oxides. Remember this simple rule: Oxides of metals are generally basic in nature (e.g., Na₂O, CaO), while oxides of non-metals are generally acidic in nature (e.g., CO₂, SO₂).}}

Indicators: The Chemical Detectives

How do we know if a substance is acidic or basic without tasting or touching it? We use indicators. An indicator is a substance that changes colour when it comes into contact with an acid or a base.

{{TABLE: title=Common Laboratory Indicators

• Olfactory Indicators: Some substances have a different smell in acidic and basic media. These are called olfactory indicators. For example, onion has a characteristic smell, but it loses its smell in a basic solution like NaOH. Clove oil and vanilla extract behave similarly.

Strength of Acids & Bases: The pH Scale

Not all acids and bases are created equal. Some are dangerously corrosive, while others are safe enough to eat. Their strength depends on how well they dissociate or ionise in water.

• Strong Acids/Bases: These dissociate completely in water, releasing a high concentration of H⁺ or OH⁻ ions.
  • Strong Acids: HCl, H₂SO₄, HNO₃
  • Strong Bases: NaOH, KOH, Ca(OH)₂
• Weak Acids/Bases: These dissociate only partially in water, releasing a low concentration of H⁺ or OH⁻ ions. An equilibrium is established between the ions and the undissociated molecules.
  • Weak Acids: CH₃COOH (acetic acid), H₂CO₃ (carbonic acid)
  • Weak Bases: NH₄OH (ammonium hydroxide), Mg(OH)₂

To quantify the acidity or basicity of a solution, we use the pH scale.

{{VISUAL: diagram: The pH scale shown as a horizontal coloured bar from 0 to 14. 0 is coloured red (highly acidic) with examples like battery acid. 7 is green (neutral) with an example of pure water. 14 is dark blue/purple (highly alkaline) with examples like drain cleaner. Other common items like lemon juice, milk, and soap are placed at their approximate pH values along the scale.}}

{{KEY: type=definition | title=The pH Scale | text=The pH scale is a logarithmic scale ranging from 0 to 14 that measures the concentration of hydrogen ions (H⁺) in a solution. The 'p' in pH stands for 'potenz', the German word for power. A lower pH indicates higher acidity, a higher pH indicates higher basicity, and a pH of 7 is neutral.}}

• pH < 7: Acidic solution
• pH = 7: Neutral solution
• pH > 7: Basic (or alkaline) solution

It's important to remember that the pH scale is logarithmic. This means that a change of one pH unit represents a tenfold change in the concentration of H⁺ ions. For example, a solution with pH 3 is ten times more acidic than a solution with pH 4, and one hundred times more acidic than a solution with pH 5.

Salts: More Than Just Table Salt

When an acid and a base neutralise each other, the ionic compound that is formed is called a salt. While we commonly think of table salt (NaCl), there are countless types of salts.

NaCl is formed from a strong acid (HCl) and a strong base (NaOH), so its solution is neutral (pH = 7). However, the nature of the resulting salt solution depends on the strengths of the acid and base that created it.

{{KEY: type=points | title=Types of Salt Solutions | text=

• Neutral Salts (pH ≈ 7): Formed from a strong acid and a strong base. Example: NaCl, KNO₃.
• Acidic Salts (pH < 7): Formed from a strong acid and a weak base. Example: NH₄Cl, Al₂(SO₄)₃.
• Basic Salts (pH > 7): Formed from a weak acid and a strong base. Example: CH₃COONa, Na₂CO₃. }}

Salts have a vast range of uses. Sodium chloride is essential for life and food preservation. Sodium bicarbonate (NaHCO₃) is used in baking and as an antacid. Calcium carbonate (CaCO₃) is a major component of limestone, chalk, and marble.

Importance in Everyday Life

Acids, bases, and pH are crucial in many biological and industrial processes.

• Human Body: Our stomach uses hydrochloric acid (pH 1.5-3.5) to digest food. Our blood is slightly basic (pH 7.35-7.45) and is kept in this narrow range by a buffer system.
• Tooth Decay: Bacteria in our mouth produce acids that can lower the pH below 5.5. This acidic environment starts to dissolve the enamel on our teeth, leading to cavities. Using toothpaste, which is basic, helps neutralise these acids.
• Agriculture: Plants require a specific pH range in the soil to absorb nutrients effectively. Farmers often test soil pH and add substances like lime (a base) or sulphur (which forms an acid) to adjust it.
• Acid Rain: Pollutants like sulphur dioxide and nitrogen oxides from burning fossil fuels can dissolve in rainwater to form sulphuric acid and nitric acid, creating acid rain. This can lower the pH of lakes and soil, harming aquatic life and forests.

{{FLASHCARD: q=What are the three products of the reaction between an acid and a metal carbonate? | a=A salt, water (H₂O), and carbon dioxide (CO₂).}}