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 ↓


Arrhenius, Brønsted–Lowry, and Lewis definitions


The study of acids and bases is fundamental to chemistry. Over time, scientists have developed various models to explain the behavior of acids and bases. The three main models are the Arrhenius definition, the Bronsted-Lowry definition, and the Lewis definition. Each of these theories has its own approach and application. In this article, we will take a deeper look at each of these definitions, explore how they characterize acids and bases, as well as provide several examples and explanations.

1. Arrhenius definition

One of the earliest frameworks for understanding acids and bases is the Arrhenius definition, named after Swedish scientist Svante Arrhenius. It is a straightforward concept that deals with the production of hydrogen ions (H +) and hydroxide ions (OH -) in water.

Arrhenius Acid: An Arrhenius acid is a substance that increases the concentration of H + ions when dissolved in water.

Arrhenius Base: An Arrhenius base is a substance that increases the concentration of OH - ions when dissolved in water.

Example:

Hydrochloric acid (HCl)

HCl (aq) → H⁺ (aq) + Cl⁻ (aq)

Here, hydrochloric acid dissociates in water to release hydrogen ions, which conforms to the definition of an Arrhenius acid.

Sodium hydroxide (NaOH)

NaOH (aq) → Na⁺ (aq) + OH⁻ (aq)

Sodium hydroxide dissociates to release hydroxide ions, which causes it to be classified as an Arrhenius base.

2. Bronsted-Lowry definition

Arrhenius' definition works well, but it is limited to aqueous solutions. To overcome this limitation, Johannes Nicolaus Bronsted of Denmark and Thomas Martin Lowry of England independently proposed a more general theory in 1923, known as the Bronsted-Lowry definition.

Bronsted-Lowry Acid: It is a substance that can donate a proton (H +) to another substance.

Bronsted-Lowry Base: It is a substance that can accept a proton (H +) from another substance.

Example:

Ammonia (NH3) as a base

NH₃ + H₂O ⇌ NH₄⁺ + OH⁻

In this equilibrium, ammonia accepts a proton from a water molecule, forming an ammonium ion and releasing a hydroxide ion. Therefore, ammonia is a Bronsted-Lowry base.

The Bronsted–Lowry theory introduces the concept of conjugate acid–base pairs, where every acid has a conjugate base formed after donating a proton, and every base has a conjugate acid formed after accepting a proton.

Conjugate acid-base pair:

When hydrochloric acid (HCl) dissociates, it forms hydrogen ions and chloride ions:

HCl + H₂O ⇌ H₃O⁺ + Cl⁻

In this reaction, HCl is the acid, and Cl⁻ is its conjugate base. Similarly, H₂O acts as a base, and H₃O⁺ is its conjugate acid.

3. Lewis definition

The Lewis definition, proposed by Gilbert N. Lewis in 1923, provides a broader view of acids and bases that is not limited to the presence of hydrogen ions. This definition is based on electron pairs.

Lewis Acid: Lewis acid is an electron pair acceptor.

Lewis Base: Lewis base is an electron pair donor.

Example:

Boron trifluoride (BF3)

BF₃ + NH₃ → BF₃NH₃

In this example, boron trifluoride acts as a Lewis acid by accepting an electron pair from the nitrogen atom present in ammonia, which acts as a Lewis base.

The Lewis model is particularly useful for understanding reactions that do not involve hydrogen ions, and it highlights the role of electron pair interactions in acid–base chemistry.

Comparison of models

Each definition - Arrhenius, Bronsted-Lowry and Lewis - adds layers to our understanding of acids and bases. While the Arrhenius model is limited to aqueous solutions and direct ion release, the Bronsted-Lowry and Lewis theories provide more general descriptions.

The Bronsted–Lowry theory emphasizes proton transfer, and extends the concept to non-aqueous environments, while the Lewis theory highlights electron pair interactions, which apply to a wider range of chemical reactions.

In short:

  • Arrhenius: Focuses on H⁺ and OH⁻ ions in water.
  • Bronsted-Lowry: extends this idea to proton donors and acceptors.
  • Lewis: Focuses on electron pair acceptance and donation.

Textual examples of each model

Each model describes common acid–base reactions as follows:

Arrhenius example:

Reaction between hydrochloric acid (HCl) and sodium hydroxide (NaOH):

HCl + NaOH → NaCl + H₂O

In Arrhenius terms, HCl provides H⁺ ions, and NaOH provides OH⁻ ions, resulting in the formation of water.

Bronsted-Lowry example:

Equilibrium reaction of ammonia (NH₃) and water:

NH₃ + H₂O ⇌ NH₄⁺ + OH⁻

Ammonia acts as a base, accepting protons from water (acid), and forming ammonium ions and hydroxide ions.

Lewis example:

Reaction of ammonia with boron trifluoride:

BF₃ + NH₃ → F₃B:NH₃

The ammonia donates an electron pair to boron trifluoride, exposing the Lewis interaction.

Conclusion

A journey through the models of acids and bases - from Arrhenius to Bronsted-Lowry to Lewis - reveals the evolving understanding of chemical interactions. Each definition enhances our ability to predict and explain the behavior of acids and bases in different contexts.

These models are indispensable tools in chemistry, providing perspectives that help students and professionals appreciate the subtle and varied nature of chemical reactions. By adopting these definitions, we gain a robust framework for exploring the complex dance of atoms, ions, and molecules in the world of chemistry.


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Arrhenius, Brønsted–Lowry, and Lewis definitions

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