Premed · Premed · General Biology 1
Lecture 2: Chemistry of Life — Atoms, Bonds, and Water
General Biology I — Molecular & Cellular
Learning Objectives
By the end of this lecture, students will be able to:
- Describe the structure of an atom and define atomic number, mass number, and isotopes
- Explain how electrons are arranged in energy levels and why valence electrons determine chemical behavior
- Distinguish between ionic, covalent, and hydrogen bonds
- Explain the unique properties of water and their biological significance
- Define pH, acids, bases, and buffers and explain their importance in biological systems
Lecture Content
I. Atomic Structure
All matter is composed of elements--pure substances that cannot be broken down further by ordinary chemical reactions. Of the ninety-two naturally occurring elements, about twenty-five are essential for life. Six of these--carbon, hydrogen, nitrogen, oxygen, phosphorus, and sulfur, collectively remembered by the acronym CHNOPS--are by far the most abundant in living organisms, accounting for roughly 98% of body mass. A handful of trace elements such as iron, zinc, copper, manganese, and iodine are required in tiny amounts but are nonetheless indispensable, often serving as cofactors that enable enzymes to function.
The atom is the smallest unit of an element that retains that element's chemical properties. Every atom consists of three subatomic particles. Protons, which carry a positive charge, reside in the dense central nucleus and define the element: an atom's atomic number equals its number of protons. Neutrons occupy the nucleus alongside protons, contributing to mass but carrying no charge. The mass number of an atom equals the sum of its protons and neutrons, while the atomic mass reported on the periodic table represents a weighted average across all naturally occurring isotopes. Electrons, bearing a negative charge, occupy regions of space called energy levels or electron shells around the nucleus.
Isotopes are atoms of the same element that differ in their number of neutrons. Most isotopes are stable, but some--called radioisotopes--possess unstable nuclei that spontaneously emit radiation as they decay. Carbon-14 and iodine-131 are well-known radioisotopes with important applications in biology and medicine: carbon-14 is used in radiocarbon dating and as a tracer in metabolic research, while iodine-131 is employed in medical imaging and the treatment of thyroid cancer.
<image>A labeled diagram of an atom showing the nucleus containing protons (red, labeled +) and neutrons (gray, no charge) at the center. Surrounding the nucleus are concentric electron shells (K, L, M) with electrons (blue, labeled -) distributed in each shell. Side panel shows three isotopes of carbon: carbon-12 (6p, 6n), carbon-13 (6p, 7n), and carbon-14 (6p, 8n), with carbon-14 marked as radioactive.</image>
II. Electron Configuration and Chemical Behavior
Electrons are not scattered randomly around the nucleus; they occupy discrete energy levels (shells). The first shell can hold a maximum of two electrons, the second up to eight, and the third up to eight in the biologically relevant atoms. The electrons in the outermost shell are called valence electrons, and they are the key determinants of an atom's chemical reactivity and bonding behavior. Atoms are most stable when their valence shell is full--a principle known as the octet rule (or the duet rule for hydrogen, which needs only two). Atoms with incomplete valence shells tend to form chemical bonds, either by sharing, donating, or accepting electrons.
Electronegativity is the tendency of an atom to attract shared electrons toward itself when it participates in a chemical bond. Oxygen and nitrogen are highly electronegative, which is why bonds involving these atoms are frequently polar--the electrons spend more time near the more electronegative partner, creating an unequal distribution of charge. Differences in electronegativity between bonded atoms determine whether a bond is nonpolar covalent, polar covalent, or ionic.
III. Chemical Bonds
Ionic bonds form when one atom completely transfers one or more electrons to another atom. The atom that loses electrons becomes a positively charged ion called a cation, while the atom that gains electrons becomes a negatively charged anion. Sodium chloride (NaCl) is a classic example: sodium donates one electron to chlorine, producing Na+ and Cl- ions that are held together by electrostatic attraction. In water, ionic compounds dissociate into their constituent ions, forming electrolytes that are essential for nerve impulse transmission, muscle contraction, and osmotic balance.
Covalent bonds form when two atoms share one or more pairs of electrons. When the atoms have similar electronegativities, the electrons are shared equally, producing a nonpolar covalent bond--as seen in the O-O bond of molecular oxygen or the C-H bonds of methane. When electronegativities differ, the electrons are pulled toward the more electronegative atom, creating a polar covalent bond with partial positive (delta+) and partial negative (delta-) charges. The O-H bonds in water are a prime example: oxygen's greater electronegativity pulls the shared electrons closer, giving oxygen a partial negative charge and each hydrogen a partial positive charge. Atoms can share one, two, or three pairs of electrons, forming single, double, or triple bonds respectively. Double bonds are shorter and stronger than single bonds, and triple bonds (such as the one in molecular nitrogen, N2) are shorter and stronger still.
Hydrogen bonds are weak attractive forces that form between a hydrogen atom covalently bonded to a highly electronegative atom (typically nitrogen, oxygen, or fluorine) and another electronegative atom nearby. Individually, each hydrogen bond is only about 5% as strong as a typical covalent bond, yet collectively they are enormously significant--they give water its extraordinary properties, hold the two strands of DNA together, and stabilize the three-dimensional shapes of proteins.
Van der Waals forces are the weakest intermolecular interactions, arising from transient, temporary asymmetries in electron distribution around molecules. Though individually negligible, they become significant when large molecules with extensive surface areas come into contact. The remarkable adhesive ability of gecko feet, for instance, depends on the cumulative effect of billions of van der Waals interactions between the microscopic setae on their toes and the surface they cling to.
<image>A comparative panel figure. Panel A: Ionic bond formation between sodium and chlorine, showing electron transfer and resulting Na+ and Cl- ions arranged in a crystal lattice. Panel B: Covalent bond types — a nonpolar covalent bond in H2 with equal electron sharing, a polar covalent bond in H2O with partial charges shown (delta+ on H, delta- on O), and a double covalent bond in O2. Panel C: Hydrogen bonding between water molecules, with dashed lines representing hydrogen bonds between the delta+ hydrogen of one molecule and the delta- oxygen of another.</image>
IV. Properties of Water
Water is the most abundant molecule in living organisms, typically comprising 60-70% of body mass, and its unique properties--all arising from its polarity and capacity to form hydrogen bonds--make it indispensable for life.
Cohesion and adhesion are two related properties with profound biological consequences. Cohesion refers to the attraction of water molecules to one another through hydrogen bonds, while adhesion is the attraction of water to other polar surfaces. Together, these forces produce capillary action--the ability of water to move up narrow tubes against gravity--which is how water ascends from roots to leaves through the xylem of plants. Cohesion at the air-water interface also creates surface tension, a "skin" on the water surface strong enough to support the weight of small insects like water striders.
Water has a remarkably high specific heat capacity (1 calorie per gram per degree Celsius), meaning it resists changes in temperature. This resistance exists because a great deal of energy must be absorbed to break the hydrogen bonds that restrain molecular motion. The practical consequences are enormous: water stabilizes body temperature in organisms and moderates the climate of coastal regions. Water also has a high heat of vaporization, meaning substantial energy is required to convert liquid water to gas. This property enables evaporative cooling--sweating in humans and transpiration in plants both dissipate heat effectively because each departing water molecule carries a large amount of thermal energy with it.
As a solvent, water is sometimes called the "universal solvent," though this is an overstatement. Polar and ionic substances dissolve readily in water and are termed hydrophilic ("water-loving"). Water molecules form hydration shells around dissolved ions, surrounding each ion and preventing them from reassociating into a crystal. Nonpolar substances, by contrast, do not dissolve in water and are termed hydrophobic ("water-fearing"). Virtually all biological reactions take place in aqueous solution, making water's solvent properties essential for life.
One of water's most unusual properties is that its solid form--ice--is less dense than its liquid form. In most substances, the solid phase is denser. But when water freezes, hydrogen bonds lock the molecules into a rigid, open hexagonal lattice that spaces them farther apart than in liquid water. Ice therefore floats, and this simple fact has profound ecological importance: a layer of ice on the surface of a lake insulates the liquid water below, allowing aquatic organisms to survive through winter.
<image>A four-panel figure illustrating properties of water. Panel A: Cohesion and surface tension — a water strider walking on water, with a molecular-level inset showing hydrogen bonds at the surface. Panel B: Water as a solvent — NaCl dissolving in water with hydration shells of water molecules surrounding Na+ and Cl- ions. Panel C: Ice vs. liquid water — side-by-side molecular arrangements showing the open hexagonal lattice of ice (less dense) vs. the more compact arrangement of liquid water. Panel D: Capillary action — water rising in a narrow glass tube, with adhesion to the glass walls and cohesion between water molecules labeled.</image>
V. Acids, Bases, and pH
Water undergoes a process called autoionization, in which a small fraction of water molecules spontaneously dissociate: H2O reversibly yields H+ and OH-. In pure water at 25 degrees Celsius, the concentrations of hydrogen ions and hydroxide ions are equal, each at 10^-7 molar. The pH scale provides a convenient way to express hydrogen ion concentration: pH equals the negative logarithm of [H+]. The scale ranges from 0 (extremely acidic) to 14 (extremely basic), with pH 7 representing neutrality. Because the scale is logarithmic, each unit change represents a tenfold change in hydrogen ion concentration--a solution at pH 3 has ten times more H+ than one at pH 4.
Acids are substances that donate hydrogen ions (protons) to a solution, thereby lowering the pH. Hydrochloric acid in the stomach (pH approximately 2), carbonic acid, and lactic acid are all biologically important acids. Bases accept hydrogen ions or release hydroxide ions, raising the pH. Sodium hydroxide, ammonia, and the bicarbonate ion are common examples. Living systems are exquisitely sensitive to pH changes, which is why buffers--substances that resist shifts in pH by absorbing or releasing H+ as needed--are essential. A buffer typically consists of a weak acid paired with its conjugate base. The bicarbonate buffer system is the most important buffer in human blood: carbon dioxide reacts with water to form carbonic acid, which dissociates into H+ and bicarbonate (HCO3-). This equilibrium can shift in either direction to absorb excess acid or base. The phosphate buffer system plays a similar role inside cells. Human blood pH is maintained within the narrow range of 7.35 to 7.45; deviation beyond this range--acidosis (too low) or alkalosis (too high)--is life-threatening.
VI. Functional Groups in Organic Chemistry (Preview)
Carbon occupies a central place in biological chemistry because of its remarkable ability to form four covalent bonds, enabling it to serve as the backbone of an enormous diversity of molecules. The chemical behavior of organic molecules is largely determined by their functional groups--specific clusters of atoms that confer characteristic properties. The hydroxyl group (-OH) is polar and is found in alcohols and sugars. The carbonyl group (C=O) appears in aldehydes (when terminal) and ketones (when internal). The carboxyl group (-COOH) is acidic, readily donating a proton, and is found in amino acids and fatty acids. The amino group (-NH2) is basic, tending to accept a proton, and is also a hallmark of amino acids. The sulfhydryl group (-SH), found in the amino acid cysteine, can form covalent disulfide bonds that stabilize protein structure. The phosphate group (-PO4) carries a negative charge and is a critical component of ATP, DNA, and phospholipids. Finally, the methyl group (-CH3) is nonpolar and plays an important regulatory role: the addition of methyl groups to DNA (methylation) is a key mechanism of gene regulation.
