Chemistry · Chemical bonding · Grade 8

Types of chemical bonding

Why is table salt a hard crystal, water a liquid and copper wire a conductor of current?

The situation

On the tablecloth there are three things: white table salt in a salt cellar, water in a glass, and copper strands at the end of a lamp wire. The salt crystals are hard and brittle and melt only at 801 °C; the water is liquid at room temperature, while copper bends and conducts current well. All three are made of atoms — the difference lies in the “glue” that holds the atoms together. How do atoms play “tug of war” for shared electrons, and why does a covalent bond form in one place and an ionic or metallic bond in another?

Topic

Chemical bonding: covalent, ionic, metallic and hydrogen bonds. Table salt, water and copper wire: atoms play “tug of war” for the electron pair — the electronegativity difference decides the type of bond

Key ideas

Electronegativity
Covalent and ionic bonds
Metallic and hydrogen bonds

Formulas

Δ χ = | χ_{A} - χ_{B} |

Electronegativity difference (Pauling scale): < 0.4 — non-polar covalent, 0.4–1.7 — polar covalent, > 1.7 — ionic (conventional limits).

ionic character \approx 1 - e^{- Δ χ^{2} / 4}

Pauling's approximate formula: at Δχ = 1.7 the bond is about 50 % ionic.

H \cdot + \cdot H ⟶ H : H

A covalent bond: a shared electron pair. H–H: 74 pm, and 436 kJ/mol is needed to break it.

Na - e^{-} \overset{}{\to} Na^{+} Cl + e^{-} \overset{}{\to} Cl^{-}

An ionic bond: an electron is transferred and the ions attract each other.

NH_{3} + H^{+} \overset{}{\to} NH_{4}^{+}

A donor–acceptor bond: the lone pair of nitrogen is donated into the empty orbital of H⁺.

Theory

Relative electronegativity (§ 14)

Electronegativity χ is the ability of an atom to attract the shared electrons in a chemical bond. On the Pauling scale the greatest is that of fluorine (3.98), followed by O (3.44), N and Cl (3.04 and 3.16); the smallest are those of the alkali metals (Na 0.93, K 0.82). Along a period from left to right χ increases (the nuclear charge increases and the radius decreases), and down a group it decreases. The type of bond is determined by the difference in electronegativity Δχ: the shared pair shifts towards the atom that attracts more strongly, where a partial negative charge δ⁻ appears, and δ⁺ on the other atom. The limits are conventional: 1.7 is the difference at which, by Pauling's formula, the bond is 50 % ionic.

The covalent bond (§ 15)

A covalent bond is a shared electron pair between two atoms. When atoms approach, their electron clouds overlap and the energy falls; the distance of lowest energy is the bond length (H₂: 74 pm), and the energy needed to break the bond is the bond energy (436 kJ/mol; the depth of the curve is 458 kJ/mol — the bond energy is slightly smaller because the molecule vibrates even in its lowest state). Between identical atoms the bond is non-polar (H₂, Cl₂, N₂); between different atoms it is polar (HCl, H₂O): the molecule becomes a dipole (H₂O 1.85 D). The number of pairs is the bond order: N≡N is a triple bond, very strong (945 kJ/mol). In the donor–acceptor mechanism one atom provides a ready-made pair (NH₃ is the donor) and the other provides an empty orbital (H⁺ is the acceptor); in the NH₄⁺ formed, all four N–H bonds are identical.

The ionic bond (§ 16)

If Δχ is large, the electron is transferred completely: sodium gives up its 3s electron (ionisation energy 496 kJ/mol) and becomes a cation Na⁺ with the Ne configuration, and chlorine accepts it (349 kJ/mol is released) and becomes an anion Cl⁻ with the Ar configuration. Oppositely charged ions attract each other electrostatically — this is the ionic bond. An ionic substance contains no molecules: in a crystal of NaCl each Na⁺ is surrounded by 6 Cl⁻ and vice versa, the Na–Cl distance is 282 pm and the edge of the cubic unit cell is 564 pm. When the crystal lattice forms, 787 kJ/mol is released — which is why ionic substances are hard and have high melting points (NaCl 801 °C), and their melt and solution conduct current.

Crystal lattices (§ 17)

Depending on what stands at the lattice points: ionic (NaCl, KCl — ions, ionic bond: hard, brittle, high melting point), atomic (diamond, silicon, quartz — a network of covalent bonds between atoms: very hard, very high melting point), molecular (iodine, sulfur, ice, CO₂ — molecules held by weak intermolecular forces: soft, melt or sublime easily; iodine 114 °C), metallic (Cu, Fe — ions and free electrons: malleable, conduct current). In lab work 2 models of the KCl, sulfur and iodine lattices are built; the geometry of lattices is covered in detail in the “Crystalline and amorphous solids” experiment in the physics section.

Metallic and hydrogen bonds

The metallic bond: metal atoms readily share their outer electrons — in the lattice there are positive ions, and between them move free electrons that belong to the whole crystal (the “electron gas”). When a voltage is applied the electrons drift in one direction — a current; layers can slide without breaking bonds — the metal bends and stretches; free electrons reflect light — lustre. The hydrogen bond: hydrogen attached to a strongly electronegative atom (F, O, N) is additionally attracted to the lone pair of such an atom in a neighbouring molecule (in water ≈ 21 kJ/mol — 20 times weaker than a covalent bond). That is why water boils at 100 °C, while H₂S boils at −60 °C; without hydrogen bonds water would boil at about −80 °C (extrapolated from the series H₂Te, H₂Se, H₂S). The lattice of ice is open — ice is lighter than water.

Hybridisation and Gillespie's theory (University)

Valence bond theory: a covalent bond is the overlap of atomic orbitals; a σ bond lies along the axis, a π bond sideways (p orbitals). In forming bonds the s, p (and d) orbitals of the central atom mix to give identical hybrid orbitals: sp (180°, BeCl₂, C₂H₂), sp² (120°, BF₃, C₂H₄), sp³ (109.5°, CH₄), sp³d (trigonal bipyramid, PCl₅), sp³d² (octahedron, SF₆). Gillespie (VSEPR) theory: the electron pairs around the central atom (AXₙEₘ) move as far from one another as possible; lone pairs take up more space, so the angle is 106.7° in NH₃ and 104.5° in H₂O. As the bond order increases, the length decreases and the energy increases: C–C 154 pm / 348 kJ/mol, C=C 134 / 614, C≡C 120 / 839. About the model: the sizes of atoms and ions are schematic (in crystals the spheres are shrunk), the electron cloud and the “rope” are illustrations; the vibration of the H₂ molecule is greatly slowed down (the real period is ≈ 8 fs); the values are averages from the CRC, NIST and textbook tables.

Did you know?

Linus Pauling proposed the electronegativity scale in 1932; in 1954 he received the Nobel Prize in Chemistry “for his research into the nature of the chemical bond”.
Ice is lighter than water (density ≈ 0.92 g/cm³): hydrogen bonds hold the molecules in an open, loose lattice. That is why lakes and ponds freeze from the top, and water remains beneath the ice.

Inside the experiment

The experiment starts from a real-life situation: you predict the result, change the controls, run the experiment and draw a conclusion. Inside are missions, theory, a quiz and a guided tour in which the camera moves and explains the key points.

More experiments in this section

The molecule: orbitals and symmetry — Year 1