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 bond ↓


Ionic bond


Ionic bonds are a type of chemical bond that forms when electrons are transferred from one atom to another. This transfer results in the formation of ions: positively charged ions (cations) and negatively charged ions (anions), which are attracted to each other because of their opposite charges. Ionic bonds are fundamental to understanding the structure and properties of many types of compounds, especially salts.

What are ions?

To fully understand ionic bonding, we must first understand ions. An ion is an atom or molecule that has a net electrical charge due to the loss or gain of one or more electrons. When an atom loses one or more electrons, it becomes positively charged and is called a cation. Conversely, when an atom gains electrons, it becomes negatively charged and is called an anion.

Example of cation and anion formation:

Consider the example of sodium (Na) and chlorine (Cl):

na → na⁺ + e⁻
Cl + e⁻ → Cl⁻

In this reaction, the sodium atom loses an electron to become a sodium ion (Na⁺), a cation, while the chlorine atom gains an electron to become a chloride ion (Cl⁻), an anion.

Formation of ionic bonds

Ionic bonds are formed through electrostatic attraction between cations and anions. This process can be broken down into a few basic steps:

  1. Electron transfer: An atom (usually a metal) donates one or more of its electrons to another atom (usually a nonmetal).
  2. Ion formation: A cation is formed as a result of the loss of electrons by a metal atom, and an anion is formed as a result of the gain of electrons by a nonmetal.
  3. Attraction: Oppositely charged ions attract each other and form ionic bonds.

Visual example:

Na Cl e⁻

In the above visual example, we can see the visual representation of the transfer of electron from sodium atom to chlorine atom, resulting in the formation of ionic bonds between Na⁺ and Cl⁻.

Properties of ionic compounds

Ionic compounds have unique properties that distinguish them from other types of compounds. Because of the strong electrostatic forces between the ions, ionic compounds typically have these properties:

  • High melting and boiling points: Considerable energy is required to break the strong attraction between ions, leading to high melting and boiling points.
  • Solubility in water: Many ionic compounds are soluble in water because the polar nature of water molecules can help separate the positive and negative ions in the compound.
  • Electrical conductivity: Although solid ionic compounds do not conduct electricity, they do conduct electricity when melted or dissolved in water (where the ions are free to move and carry charge).
  • Crystal lattice structure: Ionic compounds often form crystalline solids where the ions are arranged in a regular, repeating pattern, called a crystal lattice.

Example: Sodium chloride (NaCl)

The most common example of an ionic compound is sodium chloride (table salt). In sodium chloride:

Na⁺ + Cl⁻ → NaCl

Here, sodium ions (Na⁺) and chloride ions (Cl⁻) attract each other to form the compound NaCl. In its solid form, NaCl forms a crystal lattice structure.

Energy considerations in ionic bonding

The formation of ionic bonds is affected by energy changes. Several energy terms are involved in this process:

  • Ionization energy: The energy required to remove an electron from an atom to form a cation.
  • Electron affinity: The change in energy that occurs when an electron is added to an atom to form an anion.
  • Lattice energy: The energy released when ions come together to form a crystalline lattice.

Overall, the formation of an ionic compound is typically exothermic, meaning it releases energy. This release of energy is due to the strong attraction between the ions in the crystal lattice, which makes the ionic compound more stable than the individual ions.

Visual representation - energy diagram

Ionization energy Electron affinity Lattice energy Total released energy

Factors affecting ionic bond strength

Several factors can affect the strength and stability of an ionic bond:

  • Charge of ions: Ions with higher charge will form stronger ionic bonds because of greater electrostatic force between them.
  • Size of the ions: Smaller ions can bind more closely with each other, increasing the strength of the electrostatic interactions and thus the bond strength.
  • Presence of other forces: Additional forces such as polarization (distortion of electron clouds) can also affect the stability of an ionic bond.

Common misconceptions about ionic bonds

Despite the relatively simple nature of ionic bonds, there are several common misconceptions:

  • Not covalent: Ionic bonds involve complete transfer of electrons, while covalent bonds involve shared electrons.
  • Ionic compounds are not molecules: Ionic compounds are often mistakenly referred to as single molecules, but they actually form large lattice structures rather than isolated molecules.

Comparing ionic bonds to covalent and metallic bonds

Understanding the distinctive features of ionic bonds helps to distinguish them from other types of chemical bonds, such as covalent bonds (where electrons are shared between atoms) and metallic bonds (which involve a ‘sea of electrons’ shared between a lattice of metal atoms).

Table comparing different types of bonds:

Type of bond Primary feature Example
Ionic Ions are formed by electron transfer Sodium chloride
Covalent Electron sharing H2O
Metallic Delocalized electron ‘sea’ Fe (Iron)

Conclusion

Ionic bonds are an integral part of chemical bonding and are essential for forming compounds, especially salts. By exploring the mechanics of ionic bonds - electron transfer, ion formation, and electrostatic attraction - we can better understand the properties and behaviors of ionic compounds. Considering the energy changes during bond formation and the factors that affect bond strength helps to further understand this influential interaction.

To understand these bonds, it is important to consider not only the theoretical aspects but also practical examples in real-life applications and experiments in the laboratory.


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Ionic bond

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