1.2 Atoms and Chemical Bonds

Atoms and the Key Elements of Life

Atoms are a ubiquitous subject of study across the natural sciences. For good reason, too, as they make up the vast cosmos: from the tips of your fingers to the centers of stars (most stars, technically). The atom is comprised of a nucleus with proton(s) and neutron(s). Electrons orbit the nucleus in a probability cloud. The protons have a slightly positive charge and the electrons have a slightly negative charge. The number of protons in any given element is unique and determines that element: this is known as the element's atomic number. The mass number is the total number of protons and neutrons that a given atom has. Many elements have different isotopes (some with special names), i.e. they can have varying numbers of neutrons in their nucleus. Generally isotopes are formed when atoms combine or decay (when they generally release subatomic particles). Most isotopes are stable, however, some are not—we call these radioisotopes. Radioisotopes give off , , and radiation from its nucleus: this is known as radioactive decay. The location of a given electron is impossible to determine (the Heisenberg uncertainty principle); we can only describe a region of space with a probabilistic inference on where the electron(s) may be. The area with the most probabilistic concentration of electrons (90%) is known as the electron's orbital (outer shell).

In life, there are certain key elements that come up over and over again. The six most important are:

These are sometimes remembered by the acronym CHONPS. Carbon is particularly important because of its ability to form four bonds, allowing for long chains and complex structures.

Chemical Bonds and Valency

A bond holds molecules together because bonding is due to attraction between atoms. When atoms share electrons, they form a powerful bond called a covalent bond. A stable number of electrons is achieved in the outer shell during covalent bondage. Each atom contributes one electron pair member for each electron pair. Consider : this is a compound, i.e. a pure substance made up of two or more elements as a fixed ratio. Every compound has a molecular weight determined by the sum of all atomic weights within the molecule.

The valencies of the key elements of life are:

For a carbon atom, the outer shell has four valence electrons; thus, carbon has a valency of four—it can bond to four other atoms, forming four covalent bonds. This is the reason carbon is a pivotal atom in life: because it can make many bonds, you can get long chains of carbons added together, building complex molecules.

Electronegativity and Molecular Polarity

Whether a molecule is polar or nonpolar depends primarily on whether it has polar bonds (where electrons are unequally distributed) and also on the geometry of the molecule. We distinguish two types:

• Polar molecules have unequal electronegativity—one part of the molecule is more negative than another. These molecules tend to be hydrophilic (water-loving).
• Nonpolar molecules have equal electronegativity throughout—electrons are evenly distributed. These molecules tend to be hydrophobic (water-fearing).

Consider water. Oxygen is highly electronegative—it pulls electrons towards itself, depleting them somewhat from the hydrogens it is bonded to. This creates what is called a dipole: the oxygen end is partially negative () and the hydrogen ends are partially positive ():

The geometry of water is bent (not linear), so these dipoles do not cancel out, making water a polar molecule overall.

Determining polarity. To determine if a molecule is polar:

1. Check for polar bonds: If the electronegativity difference between bonded atoms is , the bond is polar. If , the bond is nonpolar.
2. Check for symmetry: Even with polar bonds, a symmetric molecule can be nonpolar if the dipoles cancel.

Examples:

• (linear, ): Each bond is polar, but the molecule is symmetric—the dipoles point in opposite directions and cancel. Nonpolar.
• (tetrahedral): Each bond is polar, but the four bonds are symmetrically arranged, so dipoles cancel. Nonpolar.
• (tetrahedral but asymmetric substituents): The bond is polar and there is no symmetry to cancel it. Polar.
• (bent): Polar bonds + asymmetric geometry = dipoles don't cancel. Polar.

Let us compare two molecules—propane and isopropanol:

Propane consists only of carbons and hydrogens, neither of which are particularly electronegative. There is an equal distribution of electrons throughout the molecule, so it is nonpolar and hydrophobic.

Isopropanol has a hydroxyl group () attached. The oxygen is highly electronegative and pulls electrons away from the rest of the molecule, creating a partial negative character at the hydroxyl end. This makes isopropanol polar and hydrophilic.

Types of Chemical Bonds

Bonds can be classified by their strength:

Covalent bonds are strong bonds formed by the sharing of electrons between atoms. During covalent bondage of two different atoms, one atom's nucleus may exert a greater attracting force on electrons, resulting in an unequal electron distribution. If two atoms have similar electronegativity, they form a nonpolar covalent bond. If they have differing electronegativity, they form a polar covalent bond.

Non-covalent bonds do not involve sharing electrons. In order of decreasing strength:

• Ionic bonds (strong): Electrostatic attraction between oppositely charged ions. The chloride ion () has a charge because it has one more electron than protons—such negatively charged ions are called anions. Positively charged ions are cations. Ionic bonds form stable solid compounds called salts, but covalent bonds dominate ionic bonds, so salts dissolve in polar water.
• Hydrogen bonds (medium): Attraction between a hydrogen atom (bonded to an electronegative atom like O or N, giving it ) and another electronegative atom (). These are crucial in biology—they hold DNA strands together and give water its special properties.
• Van der Waals forces (weak): Temporary attractions between atoms due to momentary dipoles. The greater the number of electrons an atom has, the greater its polarizability. Also called London dispersion forces or non-polar interactions.

Although non-covalent bonds are individually weak, many of them together can provide significant stability (as in protein folding or DNA base pairing).

The bond length of a molecule is the energy minimum at which the molecule is stable.

Representing Organic Molecules

Organic molecules—those built on carbon skeletons—can be represented in several ways, each useful for different purposes. Understanding these representations is essential for reading biochemical structures. Let us use 1-butanol (), an alcohol, as our example.

Molecular formula shows only the number of each type of atom: . This tells us the composition but nothing about how atoms are connected. The is often written separately () to indicate that there is a hydroxyl group.

Full structural formula shows every atom and every bond explicitly. This is the most complete but also the most cumbersome representation:

Condensed formula groups atoms together for compactness: or . Each carbon is understood to have enough hydrogens to complete its four bonds.

Bond-line (skeletal) formula is the most common representation in biochemistry. The conventions are:

• Carbon atoms are implied at each vertex and endpoint
• Hydrogen atoms bonded to carbon are omitted (inferred to satisfy carbon's valency of four)
• Heteroatoms (O, N, S, P, etc.) and their hydrogens are shown explicitly
• Single bonds are lines, double bonds are two parallel lines

The bond-line structure of 1-butanol is simply:

The carbons are at every corner and at the left endpoint. You can figure out where the hydrogens are because of valency: carbon has a valency of four, so each carbon must have four bonds total. A carbon at a corner (two lines) has two hydrogens; a carbon at an endpoint (one line) has three hydrogens.

Summary of 1-butanol in all representations:

Rings are drawn as polygons. For instance, cyclohexane () and benzene ():

The alternating double bonds in benzene (sometimes drawn as a circle inside) indicate the delocalized electrons of the aromatic ring.

Disubstituted benzene nomenclature. When two substituents are attached to a benzene ring, their relative positions are named:

• Ortho (o-): Substituents on adjacent carbons (1,2-positions)
• Meta (m-): Substituents separated by one carbon (1,3-positions)
• Para (p-): Substituents on opposite carbons (1,4-positions)

For example, the three isomers of dimethylbenzene (xylene, ) are ortho-xylene, meta-xylene, and para-xylene.