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


Introduction to acids and bases

Acids and bases are fundamental concepts in chemistry that are important for understanding chemical reactions both in the laboratory and in nature. These substances have different properties and are involved in a wide variety of chemical processes.

Basic definitions

Acid

An acid is generally a substance that can donate a proton (H + ion) or accept an electron pair in reactions. The traditional definition of an acid comes from the Arrhenius concept, which states that an acid is a substance that increases the concentration of hydrogen ions (H +) when dissolved in water. A more generalized definition given by the Bronsted-Lowry theory describes an acid as a substance that can donate a proton to another substance.

Bases

A base is a substance that can accept a proton or donate an electron pair in reactions. According to the Arrhenius definition, a base is a substance that increases the concentration of hydroxide ions (OH -) when dissolved in water. The Bronsted-Lowry theory similarly extends this by defining a base as a substance that can accept a proton.

Properties of acids and bases

Properties of acids

  • Acids usually taste sour.
  • They turn blue litmus paper red.
  • They react with metals such as zinc and magnesium to release hydrogen gas (H 2).
  • Acids react with bases to form salt and water, this process is called neutralization.
  • In aqueous solution, acids conduct electricity by releasing ions.

Properties of bases

  • Bases taste bitter and feel slippery to touch.
  • They turn red litmus paper blue.
  • Bases react with acids to form salt and water.
  • In aqueous solutions, they conduct electricity due to the presence of ions.

Principles of acids and bases

Arrhenius theory

The Swedish scientist Arrhenius first defined acids and bases in terms of their dissociation in water. According to Arrhenius:

  • Acids: Substances that dissociate in water to form hydrogen ions (H +). 
     HCl (aq) → H + (aq) + Cl - (aq)
  • Bases: Substances that dissociate in water to form hydroxide ions (OH -). 
     NaOH (aq) → Na + (aq) + OH - (aq)

Bronsted–Lowry theory

Extending the work of Arrhenius, the Brønsted–Lowry theory defines:

  • Acid: Proton (H +) donor.
  • Base: Proton acceptor.

For example, in the reaction between ammonia and water:

 NH 3 (aq) + H 2 O (l) ⇌ NH 4 + (aq) + OH - (aq)

Here, water acts as an acid by donating protons to ammonia, which acts as a base by accepting protons.

Lewis theory

Further development in understanding acids and bases comes from Lewis theory, according to which:

  • Acid: Electron pair acceptor.
  • Base: Electron pair donor.

This definition is broad and includes many more substances such as acids and bases. For example, in the formation of the ammonia-boron trifluoride complex:

 NH 3 + BF 3 → NH 3 BF 3

Ammonia donates an electron pair to boron trifluoride, making ammonia a Lewis base and boron trifluoride a Lewis acid.

PH scale

The pH scale is a measure of the acidity or alkalinity of a solution. The scale ranges from 0 to 14, where:

  • pH 7: Neutral (like pure water).
  • pH less than 7: Acidic.
  • pH greater than 7: Alkaline (or alkaline).

The pH scale is logarithmic, meaning that each whole number change represents a tenfold change in hydrogen ion concentration. The formula for calculating pH is:

 pH = -log[H + ]
0 7 14 Acidic Neutral Basic

Acid-base reactions

Acid-base reactions are fundamental to chemistry and involve the exchange of protons. Such reactions usually result in the formation of salt and water. A classic example of this is the reaction between hydrochloric acid and sodium hydroxide:

 HCl (aq) + NaOH (aq) → NaCl (aq) + H 2 O (l)

Here, H + from hydrochloric acid combines with OH– from sodium hydroxide to form water, and Na + combines with Cl– to form sodium chloride.

Buffer

Buffers are special solutions that resist changes in their pH when acids or bases are added. They are important in maintaining the pH of biological systems. For example, human blood is a buffer system where the carbonic acid-bicarbonate balance is:

 H 2 CO 3 ⇌ HCO 3 - + H +

Visualization of concepts

Acid Base

The above illustration shows how an acid can donate a proton (shown moving to the right) which is accepted by a base.

Conclusion

Understanding the concepts of acids and bases is important in chemistry, as they play a vital role in a wide range of chemical reactions. From the simple process of neutralization to maintaining biological pH levels, acids and bases significantly affect chemical behavior. Common theories - Arrhenius, Bronsted-Lowry, and Lewis - provide several lenses through which these substances can be understood, allowing for a wide range of investigations of their behavior and characteristics.


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Acids and bases

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