Undergraduate
  1. General chemistry  1.1. Introduction to Chemistry  1.2. Atomic Structure  1.2.1. Subatomic particles  1.2.2. Atomic Model  1.2.3. Quantum numbers  1.2.4. Electron Configuration  1.2.5. Periodic trends  1.2.6. Isotopes and their applications  1.3. Chemical bond  1.3.1. Ionic bond  1.3.2. Covalent bonds  1.3.3. Metal Bond  1.3.4. Hybridization  1.3.5. Molecular geometry and VSEPR theory  1.3.6. Bond polarity and dipole moment  1.4. Stoichiometry  1.4.1. Mole concept  1.4.2. Empirical and molecular formula  1.4.3. Limiting reactant and excess reactant  1.4.4. Percent Yield and Purity  1.4.5. Dilution calculation  1.5. States of matter  1.5.1. Gas Laws  1.5.2. Matter and Intermolecular Forces  1.5.3. Solid and Crystal Structures  1.5.4. Phase diagram  1.5.5. Plasma and supercritical fluid  1.6. Chemical reactions  1.6.1. Types of Chemical Reactions  1.6.2. Balancing Chemical Equations  1.6.3. Thermochemical reactions  1.6.4. Reaction energy profiles  1.7. Solutions and Mixtures  1.7.1. Concentration units  1.7.2. Syndrome properties  1.7.3. Solubility and Solubility Rules  1.7.4. Henry's Law and Raoult's Law  1.7.5. Partition coefficient  1.8. Chemical equilibrium  1.8.1. Law of mass action  1.8.2. Le Chatelier's principle  1.8.3. Reaction quotient  1.8.4. To use this online calculator for Equilibrium Constant, enter Equilibrium Constant (Eq.) and hit the calculate button. Here is how the Equilibrium Constant calculation can be explained with given input values -> 0.05 = 0.05/0.  1.8.5. Acid-base balance  1.9. Acids and bases  1.9.1. Arrhenius, Brønsted–Lowry, and Lewis definitions  1.9.2. pH and pOH  1.9.3. Acid-base titration  1.9.4. Buffer Solution  1.9.5. Acid-base indicator  1.9.6. Polyprotic Acids and Bases  1.10. Electrochemistry  1.10.1. redox reactions  1.10.2. Galvanic and Electrolytic Cells  1.10.3. Nernst equation  1.10.4. Electrolysis  1.10.5. Faraday's laws of electrolysis  1.11. Thermodynamics  1.11.1. Laws of Thermodynamics  1.11.2. Enthalpy, entropy and Gibbs free energy  1.11.3. Spontaneity of reactions  1.11.4. Thermodynamic cycles (Hess's law, Carnot cycle)  1.12. Kinetics  1.12.1. Rate law and reaction order  1.12.2. Activation Energy and Catalyst  1.12.3. Collision theory  1.12.4. Reaction mechanism  1.12.5. Transition state theory
  2. Organic chemistry  2.1. Structure and relationships  2.1.1. Hybridisation in Carbon Compounds  2.1.2. Resonance and aromatization  2.1.3. Molecular orbitals  2.1.4. Inductive and mesomeric effects  2.2. Hydrocarbons  2.2.1. Hydrocarbons  2.2.2. Alkene  2.2.3. Alkynes  2.2.4. Aromatic compounds  2.2.5. Cycloalkanes and Conformation  2.3. Functional Group  2.3.1. Alcohols and phenols  2.3.2. Ethers and epoxides  2.3.3. Aldehyde and ketone  2.3.4. Carboxylic acids and derivatives  2.3.5. Amin  2.3.6. Esters and amides  2.3.7. Thiols and sulfides  2.4. Stereoscopic  2.4.1. Isomerism in Stereochemistry  2.4.2. Chirality and optical activity  2.4.3. Structural analysis  2.4.4. Diastereomers and Enantiomers  2.5.1. Substitution reactions  2.5.2. Elimination reactions  2.5.3. Addition reactions  2.5.4. Radical reactions  2.5.5. Rearrangement reactions  2.5.6. Pericyclic Reactions  2.6. Spectroscopy and structural analysis  2.6.1. Infrared Spectroscopy  2.6.2. Nuclear magnetic resonance spectroscopy  2.6.3. Mass Spectrometry  2.6.4. UV-visible spectroscopy  2.6.5. X-ray crystallography  2.7. Polymer chemistry  2.7.1. Addition and Condensation Polymers  2.7.2. Polymerization mechanism  2.7.3. Biodegradable Polymer  2.7.4. Conducting and smart polymers
  3. Inorganic chemistry  3.1. Coordination chemistry  3.1.1. Ligands and coordination compounds  3.1.2. Crystal field theory  3.1.3. Molecular Orbital Theory in Coordination Compounds  3.1.4. Jahn–Taylor distortion  3.2. Main Group Chemistry  3.2.1. Alkali and alkaline earth metals  3.2.2. Halogens and Noble Gases  3.2.3. Transition metals and their complexes  3.2.4. Lanthanides and Actinides  3.3. Solid state chemistry  3.3.1. Crystal lattices and unit cells  3.3.2. Defects in solids  3.3.3. Electrical and magnetic properties  3.3.4. Superconductivity  3.4. Bio-inorganic chemistry  3.4.1. Metal ions in biology  3.4.2. Enzymatic reactions with metals  3.4.3. Metal poisoning and detoxification  3.4.4. Metalloproteins and Metalloenzymes
  4. Physical chemistry  4.1. Quantum Chemistry  4.1.1. Wave–particle duality  4.1.2. Schrödinger Equation  4.1.3. Quantum Numbers and Orbitals  4.1.4. Atomic and molecular spectroscopy  4.2. Statistical mechanics  4.2.1. Maxwell–Boltzmann distribution  4.2.2. Partitioning functions  4.2.3. Bose–Einstein and Fermi–Dirac statistics  4.3. Chemical thermodynamics  4.3.1. Gibbs Free Energy  4.3.2. Phase transition  4.3.3. Thermodynamic efficiency  4.4. Surface chemistry  4.4.1. Absorption and Catalysis  4.4.2. Colloids and Emulsions
  5. Analytical Chemistry  5.1. Classical methods  5.1.1. Gravimetric Analysis  5.1.2. titrimeter  5.2. Instrumental methods  5.2.1. Chromatography  5.2.2. Spectroscopy  5.2.3. Electroanalytical methods  5.2.4. X-ray diffraction
  6. Industrial Chemistry  6.1. Petrochemistry  6.2. Chemical engineering principles  6.3. Green Chemistry  6.4. Pharmaceuticals and Drug Synthesis
  7. Biochemistry  7.1. Carbohydrates  7.2. Proteins and enzymes  7.3. Lipids and membranes  7.4. Nucleic acids  7.5. Metabolism and Bioenergetics  7.6. Enzyme kinetics and mechanism

Undergraduate → General chemistry → Acids and bases ↓


Acid-base indicator


Acid-base indicators are substances that change color in response to changes in pH. They are extremely useful in chemistry for determining the acidity or alkalinity of a solution. These indicators, usually weak acids or bases themselves, display different colors when they lose or gain protons (H + ions) in a solution over a specific pH range.

Understanding pH and the pH scale

The pH scale is a numerical scale used to specify the acidity or basicity (alkalinity) of an aqueous solution. It runs from 0 to 14, with lower values indicating high acidity, higher values indicating high alkalinity, and a pH of 7 indicating a neutral solution.

pH example

        Water (neutral): pH 7
        Lemon juice (acidic): pH 2
        Soap solution (basic): pH 10

Mechanism of acid-base indicators

Indicators change color depending on the dissociation of the weak acid or base they contain. As the pH of the solution changes, the equilibrium between the protonated and deprotonated form of the indicator shifts, causing a change in color.

Equilibrium reaction

The normal equilibrium for the acid (HIn) indicator is shown below:

        HIn ↔ H + + In – (acid form ↔ base form)

Where HIn is the protonated form of the indicator, and In - is the deprotonated form.

Visual example of indicator change

Acidic (pH 3) Neutral (pH 7) Alkaline (pH 10)

Common acid-base indicators and their transitions

1. Litmus

The most well-known indicator, litmus, turns red in acidic solutions (pH < 7) and blue in alkaline solutions (pH > 7).

Litmus in acid: red
Litmus in base: blue

2. Phenolphthalein

Phenolphthalein is colorless in acidic solution but turns pink in alkaline solution (pH > 8.2).

Phenolphthalein in acid: colourless
Phenolphthalein in base: pink

3. Methyl orange

Methyl orange is red in acidic solutions (pH < 3.1) and turns yellow at pH above 4.4, which is typical of less acidic solutions.

Methyl orange in acid: red
Methyl orange in base: yellow

4. Bromothymol blue

Bromothymol blue is yellow in acidic solution (pH < 6.0), and turns blue in alkaline solution (pH > 7.6).

Bromothymol blue in acid: yellow
Bromothymol blue in base: blue

Applications of various indicators

Titration

Acid-base indicators are widely used in titration procedures to determine the end point of a chemical reaction. During a titration, the indicator helps to visually show the change in pH, which indicates completion. For example, phenolphthalein is popularly used for titrations involving strong bases and weak acids.

pH papers

pH papers are strips soaked with indicator and are used to quickly test the pH of a solution. When it is dipped into a solution, the paper changes color, making it easy to identify the pH by comparing the color to a standard chart.

Domestic and industrial use

Outside the laboratories, indicators have practical uses in a variety of fields, such as testing soil pH in agriculture, testing the acidity of foods in the culinary industry, and determining suitable conditions for aquariums.

Preparation of acid-base indicator

Some natural substances can be used as acid-base indicators. For example, red cabbage contains a pigment called anthocyanin, which changes color at different pH levels. It can be used as a natural indicator.

Making a cabbage indicator

        1. Chop red cabbage leaves and boil them in water.
        2. Strain the liquid, which will turn purple in color.
        3. This liquid can be used as a pH indicator - it will appear red in acidic solutions and greenish-yellow in alkaline solutions.

Limitations of indicators

Although indicators are useful for quick reference, they also have their limitations. Each indicator has a specific range of pH where a color change is apparent. Outside this range, the indicator may not work properly. Indicators can also be less accurate than pH meters, which provide exact numerical values.

Example of limits in precision

Consider bromothymol blue, which has a transition range of approximately pH 6.0 to 7.6. When used in a solution with a pH of 5.5, the result will still appear yellow, indicating no significant difference in pH from a solution with a pH of 6.0.

Alternatives to visual indicators

Electronic pH meters provide numerical pH measurements and are useful when accuracy is important. They work by measuring the voltage difference between two electrodes placed in a solution, which translates into a change in pH level.

Use of pH meter

        1. Calibrate the pH meter using a standard buffer solution.
        2. Dip the meter electrode into the solution being tested.
        3. Read the digital pH value displayed on the meter.

Conclusion

Acid-base indicators are important tools in chemistry, providing a simple visual means to determine the acidity or alkalinity of a solution. From common laboratory titrations to practical everyday uses, indicators help identify pH levels quickly and effectively. Although they have limitations in accuracy, the combination of traditional indicators and electronic options such as pH meters provides a comprehensive approach to pH measurement.


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Acid-base indicator

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