Simple Chemical Reactions and Bonds: 1.1.1
2.1 Atoms, Isotopes, Ions, and Molecules: The Building Blocks
Matter and elements
Life is built from matter; matter = substance that occupies space and has mass.
Elements are unique forms of matter with specific chemical and physical properties that cannot be broken down by ordinary chemical reactions.
There are 118 elements; only 92 occur naturally; remaining elements are lab-synthesized and unstable.
Each element has a chemical symbol: a single capital letter or a pair when the first letter is already taken by another element. Some symbols follow English names (e.g., C for carbon, Ca for calcium); others come from Latin (e.g., Na for natrium).
The four elements common to all living organisms are oxygen (O), carbon (C), hydrogen (H), and nitrogen (N).
In the non-living world, element abundance varies by reservoir (atmosphere, Earth’s crust, living organisms). Example: atmosphere rich in nitrogen and oxygen but little carbon and hydrogen; crust contains oxygen but less nitrogen and carbon. Despite abundance differences, all elements obey the same chemical/physical laws in living and non-living realms.
Approximate elemental composition in living humans vs the non-living world:
Life (Humans): Oxygen ~65%, Carbon ~18%, Hydrogen ~10%, Nitrogen ~3%.
Atmosphere: Oxygen ~21%, Nitrogen ~78%, others in trace amounts.
Earth’s Crust: Oxygen ~46%, Hydrogen ~0.1%, others in trace amounts.
All elements and their reactions follow universal chemical/physical laws regardless of being part of living or non-living systems.
The Structure of the Atom
An atom is the smallest unit of matter that retains all chemical properties of an element (e.g., one gold atom has all properties of gold). A coin is many gold atoms arranged as a coin, but still gold atoms.
An atom has two main regions:
Nucleus (center): contains protons and neutrons.
Electron cloud/orbitals around the nucleus where electrons are found.
Hydrogen is the exception: it consists of 1 proton and 1 electron with no neutrons.
Subatomic particles and their basic properties:
Proton: charge +1; mass ≈ (≈ 1 amu or 1 Dalton).
Neutron: charge 0; mass ≈ (≈ 1 amu).
Electron: charge −1; mass ≈ (≈ 1/1800 amu).
Because electrons are so light, their mass contributes negligibly to atomic mass; the mass of an atom is conventionally calculated from protons and neutrons only (electrons are ignored for mass number calculations).
In neutral atoms the number of electrons equals the number of protons, giving zero net charge (the positive and negative charges cancel).
Most of the atom’s volume is empty space; the apparent solidity of matter arises from the repulsive forces between negatively charged electrons and the negative charges of nearby atoms.
Protons, neutrons, and electrons have the following charges and locations:
Proton: charge +1; mass ≈ 1 amu; location = nucleus.
Neutron: charge 0; mass ≈ 1 amu; location = nucleus.
Electron: charge −1; mass ≈ 0 amu (negligible); location = orbitals around the nucleus.
Atomic number Z = number of protons; distinguishes one element from another. The number of neutrons N varies, giving isotopes of the same element. The mass number A = Z + N. For mass number calculations, electron mass is neglected.
Isotopes add mass without changing chemical identity; isotopes have slightly different mass numbers but the same atomic number.
Isotopes
Isotopes are different forms of the same element that have the same number of protons but different numbers of neutrons.
Natural isotopes exist for some elements (e.g., carbon, potassium, uranium).
Examples:
Carbon-12: protons = 6, neutrons = 6, electrons = 6 → mass number A = 12.
Carbon-13: protons = 6, neutrons = 7, electrons = 6 → mass number A = 13.
Carbon-14: protons = 6, neutrons = 8, electrons = 6 → mass number A = 14; this isotope is radioactive.
Radioactive isotopes (radioisotopes) emit particles and energy as they decay to more stable configurations (lower potential energy).
Carbon dating relies on the radioactive decay of Carbon-14 (14C) to determine the age of formerly living organisms.
Carbon Dating and Radiometric Dating
14C is produced in the atmosphere from 14N by cosmic ray interactions; it is incorporated into living organisms via CO2 during photosynthesis and metabolism.
Upon death, 14C decay begins; the ratio of 14C to 12C declines over time via beta decay, releasing electrons/positrons and energy.
The half-life of 14C is roughly . After this time, half of the original 14C remains.
Carbon dating is effective for materials younger than roughly (50,000 years).
Other isotopes and their half-lives used in radiometric dating:
Potassium-40:
Uranium-235:
Radiometric dating uses the known half-lives to date fossils and remains, helping trace evolutionary history.
The Periodic Table
The periodic table groups elements that share chemical properties; it was devised by Dmitri Mendeleev (1834–1907) in 1869.
Elements are arranged by atomic number (number of protons); rows (periods) and columns (groups) reflect recurring chemical and physical properties.
Each element entry shows symbol, name, atomic number (Z), and atomic mass (approximately the average mass of the natural isotopes, A-weighted):
In the text example, carbon is shown with symbol C, atomic number 6, and atomic mass ~ 12.11.
The differences in chemical reactivity arise from the number and spatial distribution of electrons in the outer (valence) shell.
When atoms bond to form molecules, their outer electrons are involved in bonding.
Grouping of elements reflects similar bonding behavior; the columns represent common valence configurations that drive shared chemical characteristics.
Electron Shells and the Bohr Model
There is a correlation between the number of protons (atomic number) and the number of electrons in neutral atoms (electrons = protons).
Bohr model (1913) depicts a central nucleus with protons and neutrons and electrons in circular orbits at fixed distances (energy levels) from the nucleus, called shells with principal quantum number n.
Electron transitions: an electron in a lower-energy shell can absorb a photon and move to a higher-energy shell; the excited state is unstable and the electron returns to the ground state, emitting a photon.
Electron filling order: electrons occupy the lowest available energy shell first; if multiple subshells have the same energy, electrons are singly occupied before pairing (Hund’s rule concept integrated into the discussion).
The outermost electron shell determines an atom’s chemical reactivity and its tendency to form bonds.
Shell capacity and stability under standard conditions:
Innermost shell (1st): max 2 electrons.
Next shells (2nd and 3rd, in the simplified description): up to 8 electrons each.
In more detailed descriptions, the 3rd shell can hold up to 18 electrons and the 4th shell up to 32 electrons.
Example electron configurations and implications:
Helium: fully occupied 1s shell (1s2).
Neon: fully occupied 1s, 2s, and 2p subshells (1s2 2s2 2p6); inert gas with high stability and low reactivity.
Lithium: 1s2 2s1.
The periodic table’s far-right group (Group 18) contains noble gases (He, Ne, Ar) with filled outer shells; they are highly inert and non-reactive.
Group 1 elements (H, Li, Na) have one electron in their outer shell and can achieve a stable configuration by losing that electron (forming +1 cations).
Group 17 elements (e.g., F, Cl) have seven electrons in their outer shell and tend to gain one electron to achieve a full outer shell (forming −1 anions).
Group 14 elements (e.g., carbon) have four electrons in their outer shell and can form several covalent bonds.
The periodic table’s column structure reflects the potential shared state of outer-shell electrons, which governs chemical characteristics.
Electron Orbitals
The Bohr model is limited; electrons are not in fixed circular orbits but occupy orbitals defined by wave functions (quantum mechanics).
An orbital is the region where an electron is most likely to be found; orbitals are classified as subshells s, p, d, f.
Subshell shapes and capacities:
s subshell: spherical; 1 orbital; holds up to 2 electrons in each principal shell with an s orbital (e.g., 1s, 2s, 3s, …).
p subshell: dumbbell-shaped; 3 orbitals; in the second shell (2n) there is a 2s and a 2p subshell (2p has 3 orbitals, total 6 electrons).
d and f subshells: more complex shapes (not all shown) with 5 and 7 orbitals respectively.
The closest orbital to the nucleus is the 1s orbital, which is filled first (e.g., H: 1s1; He: 1s2).
The second shell (2n) contains 2s and 2p orbitals; together they can hold up to 8 electrons (Ne: 1s2 2s2 2p6).
The third and fourth shells introduce additional subshells (s, p, d, and f) and larger capacities, enabling more complex electron configurations for heavier elements.
Orbitals provide a more accurate depiction of electron distribution than the simplified Bohr shells, including the shapes and spatial orientations of the regions where electrons are likely to be found.
Chemical Reactions and Molecules
Atoms are most stable when their outermost (valence) shell is filled according to the octet rule (with the exception of the innermost shell).
To achieve stability, atoms form chemical bonds by donating, accepting, or sharing electrons with other atoms to fill their outer shells.
A molecule is formed when two or more atoms bond together.
Example: Water (H2O) consists of two hydrogen atoms covalently bonded to one oxygen atom, sharing electrons to fill their outer shells.
Chemical reactions involve breaking and forming bonds; reactants (left side) are transformed into products (right side).
Representing reactions: the arrow indicates the direction of the reaction (not always one-way).
Balanced chemical equation exemplified by hydrogen peroxide decomposition:
Reactants and products balance to conserve atoms: the equation is typically written as
A compound contains atoms of more than one element (e.g., H2O, NaCl); a molecule can be a single element (e.g., O2 is a diatomic molecule, not a compound in this context).
Some reactions are irreversible (unidirectional arrows) while others are reversible (double-headed arrows) and may reach equilibrium.
In biological systems, true chemical equilibrium is seldom achieved because concentrations of reactants and products are continually changing within the body. The concept of the law of mass action governs these processes and helps explain homeostatic mechanisms (e.g., the bicarbonate buffer system in blood):
Changes in carbonic acid and CO2 elimination via respiration can shift the equilibrium and maintain acid–base balance in blood.
Ions and Ionic Bonds
Atoms can gain or lose electrons to reach a more stable outer shell, forming ions with net charges.
Cations: positively charged ions formed by losing electrons.
Anions: negatively charged ions formed by gaining electrons; their names often end with -ide (e.g., chloride, sulfide).
Electron transfer forms ionic bonds between oppositely charged ions; these bonds create ionic compounds (e.g., NaCl).
Ions are essential in physiology (electrolytes) for nerve impulse conduction, muscle contraction, and water balance; common electrolytes include sodium, potassium, and calcium.
Example of ion formation and charge:
Sodium (Na) has 1 electron in its outer shell and tends to lose it to form Na+ (11 protons, 11 neutrons, 10 electrons).
Chlorine (Cl) has 7 electrons in its outer shell and tends to gain 1 electron to form Cl− (17 protons, 17 neutrons, 18 electrons).
In ionic bonding, the donor atom loses electrons and the acceptor atom gains electrons, resulting in complete outer shells for both and formation of a crystalline lattice (e.g., NaCl).
Covalent Bonds and Other Bonds and Interactions
Covalent bonds form when atoms share electrons to satisfy the octet rule; these bonds are strong and common in life (e.g., carbon-based organic molecules, DNA, proteins) and in inorganic molecules (e.g., H2O, CO2, O2).
Bond order reflects how many electron pairs are shared: single, double, and triple bonds (1, 2, or 3 pairs of electrons shared).
The strength of covalent bonds increases with bond order; triple bonds are the strongest.
Nitrogen gas (N2) has a very strong triple bond, which makes it relatively inert and difficult for living systems to access nitrogen in this form.
Examples:
H–O bonds in water: covalent sharing of electrons between H and O to fill outer shells; water’s formula is H2O because left together, two H atoms share electrons with O to complete their shells.
O2 and CO2 are covalently bonded molecules with varying bond types and polarities.
Polar Covalent Bonds, Nonpolar Covalent Bonds, and Molecular Polarity
Polar covalent bonds involve unequal sharing of electrons, creating partial charges: δ+ on the less electronegative atom and δ− on the more electronegative atom.
Water (H2O) is a polar molecule because oxygen is more electronegative than hydrogen; this induces a partial negative charge on oxygen and partial positive charges on hydrogens, enabling hydrogen bonding.
Nonpolar covalent bonds occur when electrons are shared more evenly (or equally), typically between same elements or similarly electronegative elements.
Examples:
O2: nonpolar covalent bond due to equal sharing of electrons between two identical oxygen atoms.
CH4 (methane): nonpolar covalent bonds between carbon and hydrogen due to similar electronegativities.
The overall polarity of a molecule also depends on its geometry; for example, CO2 has polar bonds but is linear, causing dipole moments to cancel and making the molecule nonpolar.
Hydrogen Bonds and Van der Waals Interactions
Not all bonds are ionic or covalent; weaker interactions also contribute to biology:
Hydrogen bonds: occur when a hydrogen atom covalently bonded to a highly electronegative atom (e.g., O or N) experiences attraction to another electronegative atom in the same or another molecule. They are crucial for water’s properties and for stabilizing protein and DNA structures (e.g., zipping of the DNA double helix).
Van der Waals interactions: weak attractions between molecules due to transient fluctuations in electron density; require close proximity and contribute to the three-dimensional structure of proteins and other macromolecules.
Pharmaceutical Chemist
Pharmaceutical chemists develop new drugs and study drug mechanisms; involved in the entire drug development process.
Drugs can be discovered in nature or synthesized; natural compounds may be chemically modified to improve safety and efficacy.
FDA approval involves extensive testing and trials; the process can take years and requires collaboration among physicians, scientists, and chemists.
Examples of drugs with natural origins or origins in observed biology:
Paclitaxel (Taxol): an anti-cancer drug originally discovered in the bark of the Pacific yew tree.
Aspirin (acetylsalicylic acid): originally isolated from willow tree bark.
Drug discovery sometimes benefits from traditional medicine clues (e.g., willow bark lore leading to aspirin).
Minoxidil (Rogaine) is another example: originally developed for high blood pressure but observed to promote hair growth, leading to its use for baldness.
2.2 Water
By the end of this section, you will be able to:
Describe the properties of water critical to life.
Explain why water is an excellent solvent.
Provide examples of water’s cohesive and adhesive properties.
Discuss the role of acids, bases, and buffers in homeostasis.
Connections and overarching concepts
The building blocks of life (atoms, isotopes, ions, and molecules) are governed by universal chemical rules and physical laws, regardless of whether they are in living systems.
The chemical properties of elements, dictated by electron arrangements, drive why atoms bond and how biological macromolecules such as DNA and proteins attain their structure and function.
The interplay of covalent, ionic, hydrogen, and van der Waals interactions underpins the physical chemistry of cells, tissues, and organ systems.
Real-world relevance: understanding isotopes and dating informs reports on evolution and paleontology; knowledge of electrolytes and pH balance is essential in physiology and medicine; water’s properties underpin virtually all biological processes.
Key formulas and numbers to remember
Atomic masses and masses of subatomic particles:
Proton mass ≈ Neutron mass ≈
Electron mass ≈ (≈ 1/1800 amu)
Atomic mass unit (amu) = Dalton (Da)
Mass of electrons is negligible for atomic mass calculations; mass number A ≈ Z + N.
Isotopes: carbon-12, carbon-13, carbon-14 (A = 12, 13, 14 respectively)
14C half-life:
Radiometric dating uses isotopes such as (half-life ≈ ) and (half-life ≈ )
Water’s polar covalent bonds lead to hydrogen bonding and water’s unique properties (cooling, solvent capabilities, cohesion/adhesion).
Note: 2.2 Water (upcoming content)
The following section will cover the properties of water in more detail, including solvent behavior, cohesion and adhesion, and acid–base chemistry.