Metallic Bond
A metallic bond is the type of chemical bonding that holds metal atoms together in a solid, arising from the electrostatic attraction between positively charged metal ions (cations) and a delocalized "sea" of electrons that move freely throughout the structure. This electron sea model explains the characteristic properties of metals such as electrical conductivity, thermal conductivity, malleability, and lustre. The strength of metallic bonds varies with the number of valence electrons and the size of the metal ion, which accounts for differences in hardness and melting point among metals.
Worked Example
Problem
Copper has a melting point of 1085°C while sodium melts at just 98°C. Using the metallic bond model, explain this difference.
Solution
Step 1: Sodium (Na) has 1 valence electron per atom contributing to the electron sea, and Na⁺ ions have a large ionic radius (102 pm). Step 2: Copper (Cu) has more electrons contributing to bonding (including d-electrons), and Cu²⁺/Cu⁺ ions have a smaller ionic radius (73 pm for Cu²⁺). Step 3: More delocalized electrons in copper create stronger electrostatic attraction between the cations and the electron sea. Step 4: Greater charge density and more electrons per atom mean more energy is needed to overcome metallic bonding in copper. Step 5: Therefore, copper requires a much higher temperature (1085°C) to melt than sodium (98°C).
Answer
Copper's higher charge density and more delocalized electrons result in a stronger metallic bond and a much higher melting point (1085°C vs 98°C for sodium).
Metallic Bond Strength and Metal Properties
| Metal | Valence Electrons | Melting Point (°C) | Electrical Conductivity | Hardness (Mohs) |
|---|---|---|---|---|
| Sodium (Na) | 1 | 98 | Good | 0.5 |
| Aluminium (Al) | 3 | 660 | Excellent | 2.75 |
| Iron (Fe) | 2–3 | 1538 | Good | 4 |
| Copper (Cu) | 1–2 | 1085 | Excellent | 3 |
| Tungsten (W) | 6 | 3422 | Good | 7.5 |
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Related Terms
Ionic Bond
An ionic bond is a type of chemical bond formed through the complete transfer of one or more electrons from a metal atom to a non-metal atom, creating oppositely charged ions that attract each other electrostatically. This electrostatic attraction between cations (positively charged) and anions (negatively charged) holds the compound together in a crystal lattice structure. Ionic bonds are responsible for the properties of salts such as high melting points, brittleness, and electrical conductivity when dissolved in water.
Covalent Bond
A covalent bond is a type of chemical bond formed when two atoms share one or more pairs of electrons, resulting in a stable arrangement for both atoms. This sharing occurs most commonly between non-metal atoms that have similar electronegativities, allowing each atom to achieve a full valence shell without complete electron transfer. Covalent bonds are the foundation of organic chemistry and molecular biology, governing the structure of molecules ranging from water (H₂O) to complex proteins.
Electronegativity
Electronegativity is a measure of the tendency of an atom to attract a shared pair of electrons toward itself within a covalent bond, expressed on a dimensionless scale. The most widely used scale is the Pauling scale, where fluorine is assigned the highest value of 3.98, making it the most electronegative element, and caesium the lowest at 0.79. Electronegativity determines bond polarity, the character of chemical bonds (ionic vs. covalent), and influences molecular properties such as reactivity, acid strength, and solubility.
From Latin "metallum" (metal, mine) and Old French "bond" (binding). The electron-sea model was developed in the early 20th century by Paul Drude (1900) and later refined by Arnold Sommerfeld (1928).