An atom is the smallest identifiable unit of matter. Only four types of atoms—hydrogen (H), carbon (C), nitrogen (N), and oxygen (O)—make up about 96 percent of all matter in living organisms today.
Many cellular molecules contain thousands or millions of atoms bonded together; early Earth’s simple substances (like water and carbon dioxide) contained just a few atoms per molecule.
Central biological theme: Structure affects function. To understand how a molecule affects the body or chemical evolution, you must understand how it is put together.
Fundamental questions guiding study of chemical evolution:
What is the physical structure of H, C, N, O atoms in living cells?
What is the structure of simple molecules (e.g., H2O, CO2) that served as building blocks for chemical evolution?
Atomic scale depiction (Fig. 2.1): extremely small electrons orbit a nucleus made of protons and neutrons.
Most of an atom’s volume is empty space; the nucleus is tiny relative to the whole atom (illustrative scale: if an atom occupied the same volume as a stadium, the nucleus would be about the size of a pea).
Protons have a positive charge (+1), neutrons have no charge (0), electrons have a negative charge (-1).
Electrical neutrality of an atom occurs when number of protons equals number of electrons.
In the periodic table, each atom has a characteristic atomic number Z (number of protons).
The sum of protons and neutrons gives the mass number A.
Atomic notation: A is written as a superscript to the left of the symbol; Z as a subscript to the left (e.g., ZAextX).
Most of the mass of an atom comes from protons and neutrons; electrons contribute negligible mass.
Early landmark concept: isotopes are forms of an element with different neutron counts but the same number of protons; they have different masses.
Example isotopes of carbon: 12extC,ext13extC,ext14extC with masses 12, 13, and 14 Da, respectively.
The atomic weight (atomic mass) is a weighted average of naturally occurring isotopes’ masses based on their abundances (e.g., carbon’s atomic weight ≈ 12.01extDa).
Radioactive isotopes (radioisotopes) are unstable and decay over time (e.g., ^{14}{6} ext{C}
ightarrow ^{14}{7} ext{N} + e^{-} + ar{
u}_e); decay timing helps date key Earth events.
Most abundant elements in living cells are C, H, N, O, P, S; together they comprise over 99% of atoms in the body.
Understanding electron arrangement around the nucleus is key to predicting chemical behavior and bonding.
Subatomic Particles and Charges
Electrons: negatively charged particles orbiting the nucleus; occupy orbitals grouped into electron shells (levels).
Shells are numbered 1, 2, 3, …; inner shells fill first; smaller numbers are closer to the nucleus.
Each orbital can hold up to two electrons; filling rule: orbitals in a shell are filled with one electron each before any orbital receives a second (paired) electron.
The region around the nucleus where electrons reside is composed of orbitals; these are the regions that determine chemical behavior.
Important consequence: valence shell (outermost shell) contains valence electrons, which largely determine bonding properties and reactivity.
For many highlighted elements in biology, the valence shell is not full, hence unpaired valence electrons exist (e.g., carbon has 4 unpaired valence electrons; oxygen has 6 valence electrons with 4 paired and 2 unpaired).
Valence (v): the number of unpaired electrons in the valence shell; key determinant of how many bonds an atom can form. Examples:
Carbon: v = 4
Oxygen: v = 2
Atoms prefer to achieve a full valence shell for stability, which drives bond formation (e.g., covalent bonds).
Ionization, Atomic Numbers, and Isotopes
Atomic number Z (protons) determines the identity of the element; changing Z changes the element itself.
Neutrons can vary within an element, producing isotopes with the same Z but different A and masses.
Mass number A = number of protons + number of neutrons: A=Z+Nextneutrons
In neutral atoms, the number of protons equals the number of electrons.
The Dalton (Da) is a special unit for atomic masses; 1 Da ≈ mass of a proton or neutron (~1.0 Da), electrons contribute negligibly to atomic mass; practically, atom mass ≈ A Da.
Example: the carbon atomic weight (≈ 12.01 Da) reflects the natural isotope abundances (C-12 is most abundant, with mass 12 Da).
Electron Configuration and Valence
The outermost shell is the valence shell; electrons there are called valence electrons.
In highlighted biologically relevant elements, the valence shell is not full (there is at least one unpaired electron in some orbital).
The number of unpaired valence electrons is the atom’s valence: e.g., C has four unpaired valence electrons (v = 4), O has two unpaired valence electrons (v = 2).
A stable atom tends to fill its valence shell; one primary mechanism is forming chemical bonds that share or transfer electrons.
Covalent Bonding and Bond Polarity
Covalent bonding occurs when atoms share electrons to achieve filled valence shells; example: H2 forms when two hydrogen atoms share their single electrons.
Covalent bond: a bond formed by sharing one or more pairs of electrons.
Covalent bonds can be:
Nonpolar covalent: electrons shared approximately equally (e.g., H–H)
Polar covalent: electrons shared unequally due to electronegativity differences (e.g., O–H in H2O)
Polar covalent bonds create partial charges (δ+ and δ−) on atoms; in H2O, O bears δ− and H bears δ+ due to greater electronegativity of O.
Electron sharing in covalent bonds is a continuum from equal sharing (nonpolar) to unequal sharing (polar) to complete electron transfer (ionic bonds).
The concept of electronegativity explains why electrons are shared unevenly in bonds: it depends on the balance of:
Number of protons in the nucleus (pull strength) and
Distance between the nucleus and the valence shell (how strongly electrons are held).
Ranking of electronegativity among common biologically relevant elements (approximate): ext{O} > ext{N} > ext{S}
oughly ext{C}
oughly ext{H}
oughly ext{P}.
In C–H bonds, electronegativities are close enough that electrons are shared nearly equally, yielding nonpolar covalent bonds.
Ionic Bonding and the Electron-Sharing Continuum
Ionic bonds arise from complete transfer of electrons from one atom to another, resulting in full valence shells for both atoms.
Example: Sodium (Na) tends to lose one electron to form Na+, which has a full second shell and a net +1 charge.
Chlorine (Cl) tends to gain an electron to form Cl−, which has a full valence shell and a net −1 charge.
In table salt (NaCl), Na+ and Cl− ions pack into a crystal lattice held together by strong electrostatic attraction between opposite charges.
The degree of electron sharing across bonds forms a continuum: Nonpolar covalent (equal sharing) → Polar covalent (unequal sharing) → Ionic (complete transfer).
Simple Molecules Formed from C, H, N, and O
Count unpaired electrons in valence shells of carbon, nitrogen, oxygen, and hydrogen:
Each unpaired electron can form one covalent bond.
Carbon can form up to 4 covalent bonds (tetravalent); nitrogen up to 3; oxygen up to 2; hydrogen up to 1.
Common molecules formed:
Methane: extCH4 (C forms 4 C–H single bonds)
Ammonia: extNH3 (N forms 3 N–H bonds)
Water: extH2extO (O forms 2 bonds to H)
Hydrogen gas: extH2 (two H atoms share one bond)
Double bonds and triple bonds occur when atoms share more than one or three pairs of electrons, respectively.
Examples:
Carbon dioxide: extCO2 (double bonds with two O atoms, linear molecule)
Molecular nitrogen: extN2 (triple bond; very strong)
Geometry of Simple Molecules and Bonding
The overall shape of a molecule is governed by bond geometry and electron-pair repulsion (VSEPR-like reasoning): bonds arrange to minimize repulsions.
Linear structures occur when there are two regions of electron density around a central atom (e.g., extN<em>2 and extCO</em>2).
Methane (CH4) has a tetrahedral geometry: bond angles ≈ 109.5∘.
Water (H2O) has a bent/angular geometry with bond angle ≈ 104.5∘ due to repulsion from lone pairs on O.
Representing Molecules: Formats and Pros/Cons
Molecular formulas: indicate the numbers and types of atoms only (e.g., CH4, H2O, CO2).
Structural formulas: show which atoms are bonded and indicate bond types with single, double, triple bonds; provide two-dimensional geometry.
Ball-and-stick models: show three-dimensional shape and relative sizes; atoms colored to denote types (e.g., carbon often represented in black).
Space-filling models: emphasize actual relative atom sizes and spatial relationships; interpretation can be more challenging.
Each representation has advantages depending on the information you want to convey.
Practical Relevance: From Molecules to Life and Earth’s History
Some small molecules occur in volcanic gases, planetary atmospheres, and deep-sea hydrothermal vents, suggesting they were important components of Earth’s ancient atmosphere and oceans.
These building blocks could have combined in aqueous (water-based) environments to drive chemical evolution and the origin of life.
Water's properties—and its bent geometry and polar bonds—play a central role in chemical reactions and the stability of many biomolecules.
Check Your Understanding 2.1 (Key Concepts to Revisit)
Question 1: What changes when the radioactive isotope tritium (³H) decays and one neutron becomes a proton?
Question 2: Draw the structural formula for formaldehyde (CH₂O) using dots to indicate electron sharing in each covalent bond; indicate any partial charges based on electronegativity differences among C, H, and O.
Question 3: Will the bonds that chlorine forms with sodium (to form NaCl) and with carbon (to form CCl) be the same in both compounds? Consider differences in bond type and polarity.
Note: Answers are in Appendix AB.
Key Definitions and Notation recap (quick glossary)
Atomic number (Z): number of protons in the nucleus; defines element identity.
Mass number (A): total number of protons and neutrons in the nucleus; A=Z+Nextneutrons
Isotopes: same Z, different N; different masses; e.g., 12extC,ext13extC,ext14extC
Atomic weight: weighted average of isotope masses based on abundance; often not an integer (e.g., carbon ≈ 12.01extDa).