Atoms and Bonding

Atoms of Life

The key biologically relevant elements are hydrogen (H), carbon (C), nitrogen (N), oxygen (O), phosphorous (P) and sulfur (S). These elements represent more than 95 percent of the mass of a cell. Carbon is a major component of nearly all biological molecules.

Elements are characterized by their atomic structure. While the subatomic structure of the atom is a major topic of interest in chemistry, physics and biophysics, we will only discuss the basic structure that will provide sufficient information for the construction of molecules in the context of this course. Atoms have a central nucleus with positively charged protons and neutral neutrons; negatively charged electrons circle the nucleus. The electrons that are involved in chemical bonding are those electrons in the outermost orbit, referred to as valence electrons. On the periodic table below, you can view each of the atoms while hiding all but the outermost electrons.

Atomic mass, the sum of the number of protons and neutrons in the atomic structure, is a particularly useful measure of each element. By summing the atomic mass of all the atoms in a molecule, one can estimate the molecular mass of the molecule, which is then expressed in atomic mass units, or Daltons. This table shows the masses of the six atoms of the elements listed above, which can also be found in the upper right-hand corner of the box for each element in the periodic table.

To calculate the mass of a molecule, we find the mass of each individual atom in the molecule and add them together. For example, a water molecule (H2O) contains one oxygen atom that has a mass of 16 amu (atomic mass units) and two hydrogen atoms that each have a mass of one amu. Therefore, the mass of a water molecule is 16 amu + 2 x 1 amu = 18 amu.

The electronegativity of an element is the degree to which an atom will attract electrons in a chemical bond. Elements with higher electronegativities, such as N, O, and F (fluorine), have a strong attraction for electrons in a chemical bond and will therefore “pull” electrons away from less electronegative atoms. Elements with low electronegativity, such as metals, tend to “give away” electrons easily.

Some atoms with important biological relevance are shown.
orbits

 

Chemical Bonding and Molecules

Chemical Bonding and Molecules

Chemical bonds result when atoms of the same element (e.g., C-C) or different elements (e.g., C-O, C-N, O-H) combine into relatively strong, commonly neutral, structures. There are two major types of chemical bonds: ionic and covalent. Covalent bonds can further be divided into polar covalent and nonpolar covalent bonds. A polar covalent bond is a type of covalent bond that results in unique interaction between molecules

A molecule is a group of at least two atoms in a specified arrangement held together by covalent chemical bonds.

These polar bonds will interact with other polar bonds through an intermolecular attraction known as hydrogen bonding, such as that found between water molecules. Both the strong ionic and covalent chemical bonds and the weaker intermolecular forces are important in the functioning of the cell.

Ionic Bonds

Ions

Recall, that an ion is an atom with a gain or loss of electrons, always valence electrons. The number of protons is not equal to the number of electrons. This occurs through addition or loss of electrons.

There are many important ions in physiology including sodium (Na+), calcium (Ca2+) and chloride (Cl).

form when an atom or group of atoms gains or loses one or more electrons. When an atom gains electrons, it becomes a negatively charged ion, called an anion. When an atom loses electrons, it becomes a positively charged ion called a cation. Atoms with higher electronegativities tend to gain electrons and become anions, whereas those with lower electronegativities tend to lose electrons and become cations. The electrostatic attraction between a positively charged ion and a negatively charged ion is the basis of an ionic bond.

An ionic bond generally forms between an atom of low electronegativity and an atom of high electronegativity. In many cases this will be between a metal and a nonmetal. In this situation, one or more electrons is transferred from the atom with low electronegativity, which readily “gives away” its electrons, to the atom with high electronegativity, which strongly attracts those electrons.

For example, as illustrated in the animation below, a sodium atom will transfer its one valence electron to a chlorine atom, resulting in the formation of a sodium cation, Na+, and a chloride anion, Cl. Because these are oppositely charged particles, they are attracted to each other and form table salt which is stable in air.

When an ionic compound, like table salt, is put into water, it dissolves. This happens because the polar water molecule pulls these oppositely charged ions apart, as will be discussed further in the next module.

Covalent Bonds

After single elemental atoms, we can think of small molecules as the next level in chemistry’s hierarchy. Molecules result from the covalent bonding of two or more elements’ atoms.

Covalent bonds are strong bonds in which electrons circling the atomic nucleus are shared. The nature of the covalent bond is determined by the number of electrons shared and the nature of the two elements sharing the bond. Two or more atoms held together by covalent bonds in a specified arrangement is called a molecule. The diagram below illustrates the covalent bond that forms between two hydrogen atoms to form a molecule of hydrogen. Nonpolar covalent bonds form between atoms of the same or similar electronegativities, most often two nonmetals.

Each atom typically forms a specific number of covalent bonds when in a molecule with other atoms. The number of bonds that a particular atom will form is based on the atom’s valence electrons. Carbon for instance, which has four valence electrons, will form four bonds when it is in a molecule, as you can see from the diagram of methane below. Nitrogen, which has five valence electrons, will form three bonds, as seen in the ammonia molecule. The number of covalent bonds that a nonmetal will typically form is provided in this table for the biologically important elements.

The number of bonds between two atoms, such as single bonds, double bonds or triple bonds, helps determine the stability of the atomic interactions. Double bonds, which share two pairs of valence electrons between two atoms, are very strong. The strong bond of carbon double bonded to oxygen is found in amino acids (these will be discussed later). The number of valence electrons shared also controls the “shape” of the atomic interactions. Carbon double bonded to oxygen forms a “flat” (planar) bond that does not rotate. This limits the shapes that the larger macromolecule, with repetitive double bonds, can form.

diagram of 2 hydrogen atoms forming a hydrogen moleculeFormation of Hydrogen Molecule
A covalent bond is formed between two hydrogen atoms.

diagram of carbon and hydrogen forming a methane molecule and nitrogen and hydrogen forming an ammonia molecule

Formation of Methane and Ammonia Molecules

Polar Covalent Bonds

In a molecule such as hydrogen, the electrons are shared equally because each atom has the same electronegativity. However, in some molecules one atom is more electronegative than another, in which case the electrons are not shared equally. For example, in a water molecule, one oxygen atom is covalently bonded to two hydrogen atoms. Because an oxygen atom is more electronegative than a hydrogen atom, the oxygen atom draws the electrons being shared toward itself and away from the less electronegative hydrogen. When electrons in a covalent bond are shared in an unequal manner it is termed a polar, or polar covalent, bond. This unequal sharing of electrons results in the more electronegative element, in this example the oxygen atom, having a slightly negative charge and the less electronegative element, in this example the hydrogen atom, having a slightly positive charge. Molecules with polar bonds have characteristics of both ionic and covalent bonds. Whether or not a molecule is polar has significant implications on how that molecule interacts with other molecules and ions in biological systems.