Table of Contents
Chemical bonds are the “glue” that holds atoms together in molecules and larger structures. In this chapter, the focus is on the main types of chemical bonds that matter in biology and on how these bonds relate to carbon-based life.
We will look at:
- Ionic bonds
- Covalent bonds (including polar vs. nonpolar, single/double/triple, and resonance)
- Coordinate (dative) bonds
- Metallic bonds (only briefly, for contrast)
- Intermolecular forces (especially hydrogen bonds and Van der Waals forces)
The periodic table background is assumed; here we concentrate on how atoms combine.
Ionic Bonds
Formation of Ionic Bonds
An ionic bond is an electrostatic attraction between oppositely charged ions. It typically forms between atoms with very different tendencies to lose or gain electrons (large difference in electronegativity).
- A metal atom (e.g. Na) tends to lose one or more electrons:
- It becomes a positively charged ion (cation), e.g.$\text{Na} \rightarrow \text{Na}^+ + e^-$
- A nonmetal atom (e.g. Cl) tends to gain electrons:
- It becomes a negatively charged ion (anion), e.g. $\text{Cl} + e^- \rightarrow \text{Cl}^-$
The resulting ions attract each other:
$$\text{Na}^+ + \text{Cl}^- \rightarrow \text{NaCl}$$
The bond strength comes from the Coulomb force between charges:
$$F \propto \frac{q_1 q_2}{r^2}$$
(You do not need to calculate it here; the key point is: stronger charges and closer ions → stronger attraction.)
Properties Relevant for Biology
Ionic compounds:
- Form crystalline solids (e.g. NaCl).
- Often dissociate in water into free ions (e.g. Na⁺, Cl⁻), because water stabilizes the ions.
- Conduct electricity when dissolved or molten (free ions can move).
In biology:
- Ionic bonds are crucial for salts and electrolytes (e.g. Na⁺, K⁺, Ca²⁺, Cl⁻).
- In solid form they are strong, but in aqueous solution they can be easily separated by water molecules; thus in cells ionic interactions are often reversible and dynamic.
Ionic bonds usually involve metals like Na, K, Ca, Mg and nonmetals such as O, Cl, not directly carbon–carbon bonding. Carbon-based biomolecules are therefore rarely held together internally by ionic bonds, but they can interact ionically with ions (e.g. negatively charged phosphate groups with Mg²⁺).
Covalent Bonds
Basic Idea
In a covalent bond, atoms share electron pairs. This usually occurs between nonmetal atoms (e.g. C, H, O, N, S, P).
Each shared pair is a bonding pair of electrons. Atoms aim to reach a filled outer shell (often an “octet” of 8 valence electrons for main-group elements, or a duet – 2 electrons – for hydrogen).
Example: molecular hydrogen, H₂
Each H has 1 electron. They share a pair:
- Each H “feels” 2 electrons in its valence shell → stable duet.
Covalent bonds are the dominant bond type in carbon chemistry and in biological macromolecules (proteins, nucleic acids, lipids, carbohydrates).
Single, Double, and Triple Bonds
The number of shared electron pairs determines bond type:
- Single bond: 1 shared pair (e.g. C–H, C–C in ethane, H–O in water)
- Double bond: 2 shared pairs (e.g. C=C in ethene, C=O in carbonyl groups)
- Triple bond: 3 shared pairs (e.g. C≡C in ethyne)
Consequences:
- Bond length:
- Single > double > triple (single bonds are longest, triple are shortest).
- Bond strength:
- Triple > double > single (triple bonds are strongest).
- Flexibility:
- Single bonds allow rotation (important for 3D shapes of biomolecules).
- Double and triple bonds restrict rotation, making parts of molecules more rigid.
This flexibility vs. rigidity is crucial in biological structure (e.g. the double bonds in unsaturated fatty acids introduce kinks into the chains).
Polar vs. Nonpolar Covalent Bonds
Covalent bonds are not all the same. The electronegativity difference between atoms decides whether a bond is polar or nonpolar.
- Nonpolar covalent bond:
- Electrons shared approximately equally.
- Atoms have similar electronegativities (e.g. C–C, C–H approximately).
- No significant partial charges.
- Polar covalent bond:
- Electrons shared unequally.
- One atom attracts electrons more strongly (higher electronegativity).
- Leads to partial charges:
- More electronegative atom: partial negative ($\delta^-$)
- Less electronegative atom: partial positive ($\delta^+$)
Examples important in biology:
- O–H in water: O is more electronegative → O$\delta^-$, H$\delta^+$
- N–H in amino groups: N$\delta^-$, H$\delta^+$
- C=O in carbonyl groups: O$\delta^-$, C$\delta^+$
These partial charges:
- Make molecules polar → affect solubility in water.
- Enable hydrogen bonding and other specific interactions (critical in DNA base pairing and protein folding).
Transition Between Ionic and Covalent
Bond type is not always purely ionic or purely covalent. With very large electronegativity differences (e.g. Na–Cl), bonds are mostly ionic; with small differences (e.g. C–H) they are mostly covalent and nonpolar. Many biologically relevant bonds (e.g. O–H, N–H) are polar covalent, giving molecules both covalent structure and electrostatic interaction potential.
Multiple Bonds and Resonance
Some structures cannot be described fully by a single arrangement of single/double bonds. Resonance is used to represent these.
- In resonance, several equivalent Lewis structures can be drawn.
- The real structure is a hybrid, not flipping between them, but intermediate.
- Electron density is delocalized over more than two atoms.
Biologically relevant examples:
- Carboxylate groups (–COO⁻):
- The negative charge is delocalized over two oxygens → both C–O bonds are between single and double bond character.
- Aromatic rings, e.g. benzene-like structures in amino acids (phenylalanine, tyrosine) and nucleotides:
- Alternating single and double bonds are best described with resonance; electrons are delocalized over the ring.
Consequences of resonance:
- Increased stability of the molecule or group.
- Characteristic reactivity important in enzyme reactions and energy transfer.
- Contributes to color, absorption of light, and electron transport in biological systems.
Coordinate (Dative) Covalent Bonds
In a coordinate (dative) bond, both electrons in a shared pair come from one atom.
- Donor atom: has a lone pair of electrons.
- Acceptor atom: has an empty orbital to accept the pair.
Once formed, the bond is still a covalent bond (shared pair), but its origin is special.
Biological importance:
- Metal ion complexes: e.g. coordination of Mg²⁺, Fe²⁺/Fe³⁺, Zn²⁺ by proteins, nucleotides or small molecules.
- Lone pairs on O, N, or S atoms donate to metal ions.
- Ammonium ion, NH₄⁺:
- Formed when NH₃ (with a lone pair on N) donates this pair to H⁺.
These interactions are often central to enzyme active sites, electron transport, and oxygen binding (e.g. iron–porphyrin complex in hemoglobin involves coordinate bonds).
Metallic Bonds (for Contrast)
While not central in organic and biological molecules, metallic bonding is useful as a contrast.
In metallic bonds:
- Metal atoms share a “sea” of delocalized valence electrons.
- Positive metal ions are surrounded by these mobile electrons.
- This explains:
- Good electrical conductivity
- Malleability and ductility
- Typical metallic luster
Biology rarely uses bulk metallic bonding, but some organisms incorporate metals (e.g. iron, copper) into structures or enzymes. In these cases, however, the metals are usually bound via ionic or coordinate covalent interactions, not metallic bonding.
Intermolecular Forces (Secondary Bonds)
So far, the focus was on primary (intramolecular) bonds that hold atoms together within molecules. In biological systems, forces between molecules (intermolecular forces) are just as important, especially in determining:
- 3D shapes of proteins, DNA, and membranes.
- How molecules recognize and bind to each other (e.g. enzyme–substrate, hormone–receptor).
These are generally weaker than covalent bonds, but collectively they can be very strong.
Hydrogen Bonds
A hydrogen bond is a special type of interaction with three partners:
- A donor: an electronegative atom (usually O or N) covalently bonded to H
- The H bears a partial positive charge ($\delta^+$), e.g. O–H$\delta^+$.
- A hydrogen atom with this partial positive charge.
- An acceptor: another electronegative atom with a lone pair (often O or N).
Depiction:
- Often written as
O–H···OorN–H···Oetc. - The
···symbolizes the hydrogen bond.
Features:
- Weaker than covalent bonds, stronger than most other intermolecular forces.
- Highly directional (strongest when nearly linear: donor–H–acceptor).
- Easily formed and broken → crucial for dynamic biological systems.
Biological roles:
- Water structure: hydrogen bonds between water molecules give water its unique properties (covered elsewhere).
- DNA: base pairs (A–T, G–C) are held together by hydrogen bonds.
- Proteins: hydrogen bonds stabilize α-helices and β-sheets and help maintain overall folding.
- Enzyme–substrate interaction: often involves multiple hydrogen bonds for specificity.
Hydrogen bonds require polar covalent bonds (mostly N–H and O–H) and therefore directly reflect basic bonding properties of atoms in the periodic table.
Dipole–Dipole Interactions
Polar molecules have permanent dipoles (separation of charge). They can attract each other via dipole–dipole forces:
- Positive end ($\delta^+$) of one molecule is attracted to negative end ($\delta^-$) of another.
- Weaker than hydrogen bonds, but still important for:
- Interactions between polar organic molecules.
- Orientation and solubility of molecules in polar environments.
Van der Waals (London Dispersion) Forces
Van der Waals forces (especially London dispersion forces) arise from temporary fluctuations in electron distribution:
- Even nonpolar molecules can have instantaneous dipoles.
- These induce dipoles in neighboring molecules.
- Resulting weak attractions operate at very short distances.
Individually very weak, but:
- They add up when many atoms are in close contact (e.g. in long hydrocarbon chains, lipid bilayers).
- Help stabilize:
- Packing of fatty acid chains in membranes.
- Folding of nonpolar regions inside proteins.
These forces do not require charge or polarity; they occur between all atoms and molecules.
Comparing Bond Types in a Biological Context
A concise overview:
- Covalent bonds:
- Strong, define the primary structure of biomolecules.
- Carbon-based skeletons are built almost entirely from them.
- Ionic bonds:
- Strong in dry crystals, but often weaker in water.
- Important for salts, ion gradients, and some protein–ligand interactions.
- Coordinate bonds:
- Variation of covalent bonds; essential for metal-containing biomolecules (enzymes, transporters).
- Hydrogen bonds and dipole–dipole:
- Moderate strength, high specificity.
- Critical for secondary and tertiary structures of macromolecules and for molecular recognition.
- Van der Waals forces:
- Very weak individually, but numerous.
- Important in hydrophobic interactions and molecular packing.
Carbon’s special position in the periodic table (multiple stable covalent bonds, moderate electronegativity) allows it to use these bond types in countless combinations, making it especially suited as the backbone of living matter.