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 titration


Acid-base titration is a method in chemistry that allows chemists to determine the concentration of an acid or base in a solution. In this process a solution of known concentration, called the titrant, is slowly added to a solution of the analyte whose concentration is unknown. By accurately measuring the amount of titrant needed to achieve the neutralization reaction, the concentration of the unknown solution can be calculated.

Fundamentals of titration

The basic principle behind an acid-base titration is the neutralization reaction, in which an acid and a base react to form water and, usually, a salt:

Acid + Base → Water + Salt

This reaction continues until the solution reaches its equivalence point, the point where the number of moles of hydrogen ions is equal to the number of moles of hydroxide ions.

Understanding acids and bases

Before proceeding with the titration process, it is important to understand what acids and bases are.

Acid

  • An acid is a substance that donates hydrogen ions (H +) in solution.
  • Common examples include hydrochloric acid (HCl), sulfuric acid (H 2 SO 4), and acetic acid (CH 3 COOH).

Bases

  • A base is a substance that donates hydroxide ions (OH -) in solution.
  • Examples include sodium hydroxide (NaOH), potassium hydroxide (KOH), and ammonia (NH 3).

Titration process

The titration process involves several main steps. Below is a detailed description of these steps, including visual examples to aid understanding.

To install

The usual setup for a titration includes a burette, a pipette, and a flask. The burette is filled with the titrant, a solution of known concentration, and the flask contains the analyte, or a solution of unknown concentration.

Process

  1. Prepare the analyte by using a pipette to measure an accurate volume in the flask.
  2. Add a few drops of a suitable indicator to the analyzer. The indicator will help determine the end point of the titration, the point at which neutralization occurs.
  3. Slowly pour the titrant from the burette into the analyzer, swirling the flask constantly to ensure thorough mixing.
  4. Note the color change in the solution, which indicates that the end point is near.
  5. At a slight and stable colour change, known as the end point, stop adding the titrant. This indicates that neutrality has been achieved.

Example titration calculations

To illustrate a common titration calculation, assume you have an unknown solution of hydrochloric acid (HCl) and you are using sodium hydroxide (NaOH) as the titrant with a known concentration of 0.1 M. Let's say you added 25 mL of NaOH to reach the end point.

The reaction can be represented as follows:

HCl + NaOH → NaCl + H 2 O

Since the mole ratio between HCl and NaOH is 1:1, you can use the formula:

M 1 V 1 = M 2 V 2

where M 1 is the molarity of NaOH, V 1 is the volume of NaOH, M 2 is the molarity of HCl (unknown), and V 2 is the volume of HCl.

For example, suppose you used 50 ml of HCl:

0.1 M * 25 ml = M 2 * 50 ml

Solve for M 2:

M 2 = (0.1 M * 25 ml) / 50 ml = 0.05 M

Thus, the concentration of the unknown HCl solution is 0.05 M.

Indicators in titration

Indicators are important in titrations because they visually indicate the end point. The choice of indicator depends on the specific acid and base being titrated. Here are some examples:

Common indicators used

  • Phenolphthalein: Colourless in acidic solution and pink in alkaline solution. Ideal for strong acid-strong base titrations.
  • Methyl orange: Red in acidic solution and yellow in alkaline solution. Useful for strong acid-weak base titration.
  • Bromothymol Blue: Yellow in acidic solution and blue in alkaline solution. Suitable for weak acid-strong base titrations.

Types of titration curve

The titration curve is a graph that shows the changes in pH when the titrant is added to the analyte. The curve diagram helps in understanding the progress of the titration and determining the equivalence point.

Visual example of a titration curve

Here's a basic example of a titration curve:

Volume of titrant (mL) pH

To understand the curve better:

  • A sharp increase in pH indicates approaching the equivalence point.
  • The midpoint of a steep slope is usually the equivalence point.
  • The initial flat portion represents the pH of the analyte before any titrant is added.

Conclusion

Acid-base titrations are essential analytical techniques used in chemistry to determine the concentrations of unexpected solutions. Understanding the basics of acids and bases, the role of indicators, and solution stoichiometry forms the basis for performing accurate and meaningful titrations. Skill in handling instruments such as burettes and pipettes, combined with an understanding of titration curves, enables chemists to perform these exciting procedures with precision. As you become familiar with titration concepts, you will find that these methods play an important role not only in laboratories but also in a variety of industries and real-world applications.


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