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 ↓


Metal Bond


Metallic bonding is a fundamental concept in chemistry, which explains the properties and behaviors of metals. Metals have unique properties such as electrical conductivity, malleability, ductility, and lustrous appearance. Understanding metallic bonds helps us know how metals interact at the atomic level to give them these properties. In metallic bonding, atoms in a metal contribute their valence electrons to form a "sea" of delocalized electrons. These electrons move freely around a lattice of metal cations, which facilitates the unique properties of metals.

Electron sea model

The electron sea model is a common way to describe metallic bonding. In this model, the valence electrons of metal atoms are not bound to any specific atom. Instead, they are free to move throughout the metal lattice. This "sea" of electrons at the atomic level facilitates the conductivity of metals because the electrons can move in response to an electric field.

To illustrate the electron sea model, consider a simple metal such as sodium (Na). Each sodium atom donates its outermost electron to the electron sea. This gives the sodium atom its positively charged core (nucleus and inner electrons), known as a sodium cation (Na +). The free electrons provide a "glue" that holds these cations together.

Na ↔ Na + + e -

Here is a basic visual depiction of a row of metal atoms and their group of shared electrons:

Na + Na + Na + Na + Na + Sea of Electrons

Properties of metallic bonds

Electrical conductivity

The delocalized electrons in metallic bonds allow metals to conduct electricity. When a potential difference is applied to a piece of metal, the electrons can move from one side to the other, causing an electric current to flow. This is why metals like copper and aluminum are used in electrical wires.

Thermal conductivity

Metals are also good conductors of heat. Free electrons can carry thermal energy through the metal lattice, allowing heat to transfer quickly from hotter areas to cooler areas. This property makes metals useful for cookware and heat exchangers.

Malleability and ductility

Because metallic bonds involve a flexible sea of electrons, metal atoms can slide past one another without breaking the bond. This gives metals the ability to be hammered into sheets (malleability) or drawn into wires (ductility). These physical properties are used in a variety of manufacturing processes.

Aura

The lustre of metals, their shiny appearance, is due to the interaction of light with delocalised electrons. When light falls on the surface of a metal, the electrons can absorb and reemit light, giving metals their characteristic lustre.

Metallic bond strength

The strength of a metallic bond is affected by several factors, including the number of delocalized electrons, the charge of the metal cations, and the size of the cations. For example, magnesium (Mg) forms a stronger metallic bond than sodium because it has two valence electrons that are delocalized rather than one.

Mg ↔ Mg 2+ + 2e -

The higher charge on the magnesium cation increases the electrostatic attraction between the cations and the electron sea, forming stronger metallic bonds than with monovalent metals such as sodium.

Comparison with other types of bonds

Metallic bonds are one of the three primary types of chemical bonds, along with ionic bonds and covalent bonds.

Ionic bond

Ionic bonds are formed between metals and nonmetals. Metals lose electrons to become cations, while nonmetals gain those electrons to become anions. The electrostatic attraction between these oppositely charged ions forms ionic bonds. Unlike metallic bonds, ionic compounds are usually brittle and nonconductors in solid form, but can conduct electricity when dissolved in water.

Covalent bonds

Covalent bonding involves the sharing of electron pairs between atoms, usually between non-metals. This can lead to the formation of molecules. For example, two hydrogen atoms share electrons to form a molecule of hydrogen gas (H 2).

H: + :H ↔ H::H

Covalent substances can be gases, liquids or solids, with conductivity varying depending on their structure. Metals, with their metallic bonds, have different properties, as described earlier.

Models of metallic bonding

Band theory

Band theory is an advanced concept used to describe metallic bonding at the quantum level. It arises from ideas involving the overlap of atomic orbitals in the metal lattice and the formation of continuous bands of energy states. Metals have partially filled conduction bands that allow free movement of electrons, which contributes to their conducting properties.

Conduction band Fermi Level Valence band

Free electron model

The free electron model simplifies the concept of the electron sea by treating the valence electrons as a gas of free electrons. This model takes into account conduction properties and is useful in explaining basic metallic bonding behavior, but does not consider electron-electron interactions.

In summary, metallic bonding is a versatile and essential concept that explains why metals exhibit certain properties. Understanding the nature of the electron sea and its effect on the properties of metals provides fundamental insights into metallurgy, materials science and solid state physics. Whether considering a practical application such as electrical wiring or the development of new alloys, the principles of metallic bonding remain invaluable in both research and industry.


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Metal Bond

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