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 → Chemical equilibrium ↓


Acid-base balance


Acid-base chemistry is an important part of chemistry, especially when analyzing chemical equilibria in solution. This concept revolves around the idea that an acid donates protons (H + ions), while a base accepts protons. Understanding these interactions and equilibrium conditions is essential for predicting how acids and bases behave in various contexts. This lesson explores acid-base equilibrium by explaining fundamental concepts and providing visual and text examples to illustrate the topic.

1. Basic concepts of acids and bases

According to the Bronsted-Lowry theory, an acid is a substance that can donate a proton (H + ), and a base is a substance that can accept a proton. This theory extends the earlier Arrhenius definition by including non-aqueous solvents and reactions that do not directly involve hydroxide ions.

Depending on the strength of their ability to donate or accept protons, acids and bases are classified as strong or weak. A strong acid such as hydrochloric acid (HCl) dissociates completely in water:

HCl + H 2 O → H 3 O + + Cl -

In contrast, a weak acid such as acetic acid ( CH3COOH ) is only partially dissociated:

CH 3 COOH ⇌ H 3 O + + CH 3 COO -

2. The concept of chemical equilibrium

When a reaction occurs in a closed system, it eventually reaches a state where the concentrations of reactants and products remain constant over time. This is called chemical equilibrium. For acid-base reactions, the equilibrium position determines whether the acid or the base will predominate in the solution.

Reactants Products balance K A

For any weak acid dissociation in water, the equilibrium can be described by an equilibrium constant, more specifically called the acid dissociation constant (K a). The larger the value of K a, the stronger the acid, since it reflects a greater degree of ionization.

3. Calculation of pH

The pH scale is used to measure the acidity or alkalinity of a solution. It is defined as the negative logarithm of the hydrogen ion concentration:

pH = -log[H + ]

A pH less than 7 is considered acidic, while a pH greater than 7 is considered alkaline. A solution with a pH of 7 is considered neutral, as is the case with pure water.

Example calculation

Consider a solution with a hydrogen ion concentration 1.0 × 10 -3 M Use the formula:

pH = -log(1.0 × 10 -3 ) = 3.0

This solution is acidic because its pH value is less than 7.

0 7 14 PH scale

4. Buffer solution

Buffer solutions resist large changes in pH when small amounts of acid or base are added. These solutions are usually composed of a weak acid and its conjugate base or a weak base and its conjugate acid.

Common examples include acetic acid and sodium acetate for acidic buffers or ammonia and ammonium chloride for basic buffers. Buffers are important in biological systems where a stable pH is necessary for enzymes and other biochemical processes.

Henderson–Hasselbalch equation

The Henderson–Hasselbalch equation provides a way to calculate the pH of buffer solutions:

pH = pK a + log 10 ([A - ]/[HA])

where [A - ] is the concentration of the conjugate base, [HA] is the concentration of the acid, and pK a is the negative logarithm of the acid dissociation constant.

Example calculation

For a buffer system with acetic acid and sodium acetate, with concentrations [CH 3 COOH] = 0.1 M and [CH 3 COO - ] = 0.1 M , and pK a = 4.76, the pH can be calculated as:

pH = 4.76 + log(0.1/0.1) = 4.76

This buffer system will maintain a nearly constant pH of 4.76 even with small amounts of acids or bases.

5. Le Chatelier's principle in acid-base equilibria

Le Chatelier's principle predicts how changes in conditions (such as concentration, pressure, or temperature) will affect the equilibrium of a reaction. In the context of acid-base equilibrium, changing the concentrations of reactants or products will shift the equilibrium to oppose the change.

Example scenario

Consider the balance:

HA + H 2 O ⇌ H 3 O + + A -

If more H 3 O + is added, the equilibrium will shift to the left, resulting in increased formation of HA to reduce the stress on the system. Conversely, removal of H 3 O + will shift the equilibrium to the right, leading to increased dissociation of HA.

6. Practical application

Biological systems

Acid-base balance is essential in biology. Enzymes function in specific pH ranges, where the active sites maintain their structural integrity. For example, human blood is maintained at a nearly constant pH of about 7.4 by bicarbonate and phosphate buffer systems.

Industrial applications

In industry, controlling pH is important for processes such as pharmaceuticals, food and beverage manufacturing, and water treatment. Acid-base balance regulates these processes by ensuring that reactions occur under optimal conditions.

7. Conclusions

Understanding acid-base equilibrium is important for chemists and related professionals, helping them to predict reaction outcomes and control conditions effectively. From calculating pH and buffer capacities to using principles such as Le Chatelier's, these concepts form the backbone of much practical and theoretical chemistry.


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