Chapter 2 Notes: The Chemical Basis of Life I (2.1–2.5)

2.1 Atoms

Atom: the smallest functional unit of matter that forms all chemical substances; cannot be broken down into simpler substances by ordinary chemical or physical means.
Element: a pure substance made up of only 1 kind of atom.
Molecule: formed when 2 or more atoms are bonded together.
Subatomic particles inside an atom:
- Protons: positive charge (+); located in the atomic nucleus.
- Neutrons: electrically neutral; located in the atomic nucleus.
- Electrons: negative charge (-); located in orbitals around the nucleus.
In neutral atoms, protons and electrons are typically present in equal numbers, giving the atom no net charge. The number of neutrons can vary for a given atom (forming isotopes).
Electrons occupy orbitals around the nucleus:
- An atom is not a literal miniature solar system; electrons occupy 3D volumes called orbitals.
- Each orbital can hold a maximum of 2 electrons.
- Orbitals are organized into electron shells; shells have characteristic energy levels.
- Shells are numbered from the nucleus outward (shell 1 is closest).
- A shell can contain 1 or more orbitals; e.g., the first shell contains 1 orbital (max 2 e−); the second shell contains 4 orbitals (max 8 e−).
- Electrons fill orbitals in order from the nucleus outward.
- Valence electrons are those in the outermost shell and participate in chemical bonding.
Atomic number and the periodic table:
- The number of protons in an atom is its atomic number (Z).
- The atomic number uniquely identifies an element; it also equals the number of electrons in a neutral atom (net charge = 0).
- The periodic table is arranged by atomic number and electron shell structure.
- Rows correspond to the number of electron shells; columns indicate the number of electrons in the outer shell (valence electrons).
- Elements in the same column have similar properties due to the same valence electron configuration.
Atomic mass and particles:
- Protons and neutrons have masses that are nearly equal and together are much heavier than electrons. Electron mass is negligible by comparison.
- Atomic mass is measured in daltons (Da); 1 Da = 1/12 the mass of a carbon-12 atom.
- Some sources use unified atomic mass units (amu); 1 Da = 1 amu.
- Practical masses: proton ≈ 1 Da, neutron ≈ 1 Da, electron ≈ 0 Da.
- Example exercise: an atom with 4 protons, 5 neutrons, and 4 electrons has an approximate mass of $4 + 5 = 9~\text{Da}$ (electrons contribute negligibly).
Isotopes: variants of an element with the same atomic number but different numbers of neutrons.
- Example: Carbon isotopes
- $^{12}C$ : 6 protons, 6 neutrons
- $^{14}C$ : 6 protons, 8 neutrons
- Atomic masses reported in the periodic table are averages of naturally occurring isotopes.
- Some isotopes are unstable radioisotopes that decay by emitting particles or radiation.
Major elements in living organisms:
- The elements oxygen (O), carbon (C), hydrogen (H), and nitrogen (N) typically make up about 95% of an organism’s mass.
- Minerals and trace elements are also required for growth and normal function.

2.2 Chemical Bonds and Molecules

A molecule contains 2 or more atoms bonded together; a compound is a molecule that contains different kinds of atoms.
Molecular formulas are used to represent molecules; chemical symbols denote elements; subscripts indicate the number of each atom (e.g., $\text{C}6\text{H}{12}\text{O}_6$ ).
Emergent properties: compounds can have chemical characteristics different from the individual elements that compose them.
The atoms in a molecule are held together by chemical bonds: covalent bonds (polar & nonpolar), hydrogen bonds, and ionic bonds are important in biology.
Covalent bonds: form when atoms share electrons to fill their outer shells.
- Many atoms seek to fill their outer shell; the typical stable configuration is the octet rule (8 electrons in the outer shell).
- Hydrogen is an exception, filling its outer shell with 2 electrons.
- Covalent bonds are strong and stable.
- Bond multiplicity: a single bond (1 pair of electrons), a double bond (2 pairs), a triple bond (3 pairs).
- Each atom forms a characteristic number of covalent bonds depending on how many electrons are needed to fill its outer shell; e.g., carbon can form up to 4 covalent bonds, enabling the formation of diverse biological macromolecules (carbohydrates, lipids, proteins, nucleic acids).
Electronegativity and bond polarity:
- Electronegativity is the ability of an atom to attract electrons in a bond.
- The relative electronegativities of a few biologically important atoms:
- $\mathrm{H} = 2.20, \ \mathrm{C} = 2.55, \ \mathrm{N} = 3.04, \ \mathrm{O} = 3.44, \ \mathrm{Na} = 0.93, \ \mathrm{Cl} = 3.16.$
- Nonpolar covalent bonds form between atoms with similar electronegativities; electrons are shared roughly equally; the molecule is electrically neutral overall.
- Polar covalent bonds form between atoms with different electronegativities; electrons are shared unequally, creating partial charges (polarity) across the molecule.
- Example: in water, the O—H bonds are polar due to O's higher electronegativity relative to H.
Polar vs nonpolar bonds and molecules:
- Nonpolar molecules/bonds: primarily consist of bonds like \mathrm{C}!-
  m{C} or \mathrm{C}!-
  m{H} where electrons are shared evenly.
- Polar molecules: contain polar bonds such as \mathrm{O}!-
  m{H}, \mathrm{N}!-
  m{H}, or \mathrm{O}!-
  m{C}, leading to overall molecular dipoles.
- Water (H₂O) is a classic polar molecule due to differential electronegativity between O and H.
Hydrogen bonds and van der Waals forces:
- Hydrogen bonds (H bonds): form when a hydrogen atom with a partial positive charge (δ+) is attracted to a highly electronegative atom (like O or N) on another molecule; represented as dashed lines; individually weak, but collectively can be strong; crucial for protein and DNA structure.
- van der Waals dispersion forces: very weak transient attractions arising from momentary uneven electron distribution in nonpolar molecules; can sum to significant attraction when many such interactions occur.
Ionic bonds:
- Occur between positively charged ions (cations) and negatively charged ions (anions).
- Formation of ionic compounds (salts): examples include $\text{NaCl}, \text{KCl}, \text{CaCl}_2.$
Molecular shapes:
- Molecules can adopt different shapes based on rotation around single bonds; shapes influence interactions with other molecules and biological activity.

2.3 Chemical Reactions

A chemical reaction occurs when one or more substances are changed into other substances by making and/or breaking chemical bonds.
Key features:
- Reactions typically require energy input (e.g., heat) for molecules to move and collide.
- In living systems, enzymes often act as catalysts to speed up reactions.
- Reactions tend to proceed toward a state of equilibrium where forward and reverse reaction rates become equal.
- In biological systems, many reactions do not reach true equilibrium because products are rapidly consumed or transformed by subsequent reactions, or because water provides the reaction medium.
Reactants and products:
- Starting materials are called reactants; end materials are called products.
- Example (balanced form as commonly shown):
- $\mathrm{CH4 + 2\,O2 \rightleftharpoons CO2 + 2\,H2O}$
- The double-headed arrow indicates the reaction can proceed in both directions.
Chemical equilibrium in biology:
- Biological systems often do not remain at equilibrium because products are continually utilized in other reactions, maintaining a steady state.

2.4 Properties of Water

The majority of organisms are water-based; water constitutes up to ~95% of some plants and ~60–70% of human body weight.
Solutes, solvents, and solutions:
- A substance dissolved in a liquid is a solute; the liquid is the solvent.
- A solution is a mixture of solutes in a solvent.
- In an aqueous solution, water is the solvent.
Hydrophilic vs hydrophobic:
- Hydrophilic substances readily dissolve in water and can interact with water's partial charges via electrical attractions.
- Hydrophobic substances do not dissolve in water; nonpolar molecules (mostly C and H) lack partial charges and are not attracted to water.
- Example of a hydrophobic substance: oil.
Amphipathic molecules:
- Molecules with both hydrophilic and hydrophobic regions.
- In water, amphipathic molecules may form micelles or bilayers; nonpolar (hydrophobic) regions orient toward the center while polar (hydrophilic) regions face the water.
- This arrangement promotes stable interactions and enables formation of cellular membranes.
Concentration and solution terminology:
- Concentration is the amount of solute per unit volume of solution.
- Common units include grams per liter (g/L) and moles per liter (M).
- Example 1: Dissolving 1 g NaCl in enough water to make 1 L of solution gives a concentration of $1\ ext{g/L}\,\text{NaCl}.$
- Example 2: Glucose has higher molar mass than NaCl, so a 1 g/L glucose solution contains far fewer glucose molecules than a 1 g/L NaCl solution.
- To quantify precisely, use molarity:
- $M = \dfrac{n}{V}$ where n is moles of solute and V is volume in liters.
- The mole is defined as: $1\ ext{mol} = 6.022 \times 10^{23} \text{ particles}$
- For glucose, the molecular mass is $\text{C}6\text{H}{12}\text{O}_6 = 180\ \text{Da}$ ; thus, 180 g of glucose per mole gives a 1 M solution in 1 L.
- For NaCl, the molecular mass is $58.4\ \text{g/mol}$ ; thus, 58.4 g NaCl per liter yields a 1 M NaCl solution.
Water's three states and energy changes:
- Water exists as a solid (ice), liquid (water), and gas (water vapor).
- Phase changes involve energy transfer:
- Heat of vaporization (energy to boil) is high.
- Heat of fusion (energy to melt) is high.
- Specific heat is high, meaning water can absorb/release substantial energy with little temperature change.
- Hydrogen bonds between water molecules contribute to these properties.
Water’s roles in living organisms:
- Acts as a solvent for biochemical reactions.
- Participates in reactions, provides mechanical support, aids in waste removal, facilitates evaporative cooling, and supports cohesion, adhesion, surface tension, and lubrication.

2.5 pH and Buffers

Hydrogen ion concentrations are changed by acids and bases:
- Water autoionizes: $\mathrm{H_2O \rightleftharpoons H^+ + OH^-}$
- The dissociation rate is small (about 2 dissociations per 10^9 water molecules).
- In pure water at 25°C, $[\mathrm{H^+}] = [\mathrm{OH^-}] = 10^{-7}~\text{M}$ , and the product is $[\mathrm{H^+}][\mathrm{OH^-}] = 10^{-14}~\text{M}^2!$ (the ion product constant, $K_w$).
- When acids or bases dissolve, they alter the relative concentrations of H+ and OH−.
Acids and bases:
- An acid donates H+ to solution, increasing $[\mathrm{H^+}]$ ; acids can be strong (fully dissociate) or weak (partial dissociation).
- A base accepts H+ in solution and decreases $[\mathrm{H^+}]$ ; bases can release OH− or bind H+; bases can be strong or weak.
The pH scale:
- $\text{pH} = -\log_{10} [\mathrm{H^+}]$
- Values range from 0 to 14; pH and [H+] are inversely related.
- Pure water is neutral with pH ≈ 7; acidic solutions have pH < 7; basic solutions have pH > 7.
- The pH scale is logarithmic; moving by 1 unit changes the [H+] by a factor of 10.
- In biological contexts, pH affects molecular shapes, reaction rates, binding interactions, and solubility; organisms regulate pH to stay within narrow ranges (e.g., human blood pH 7.35–7.45).
Buffers:
- A buffer is a substance (or system of substances) that resists pH changes when excess acid or base is added.
- Buffers work by accepting H+ from excess acid or donating H+ to neutralize excess base.
- A key biological buffer system is the carbonic acid/bicarbonate buffer:
- $\mathrm{CO2 + H2O \rightleftharpoons H2CO3 \rightleftharpoons H^+ + HCO_3^-}$
- This system helps maintain stable pH in body fluids.

Chapter 2 Summary

2.1 Atoms
- Atoms are composed of subatomic particles (protons, neutrons, electrons).
- Electrons occupy orbitals around a nucleus; each orbital holds up to 2 electrons.
- Each element has a unique number of protons (the atomic number); neutral atoms have equal numbers of protons and electrons.
- Atoms have a small but measurable mass; mass is typically given in daltons (Da).
- Isotopes vary in neutron number; atomic masses in the periodic table are averages of isotopes; some isotopes are radioisotopes.
- Four elements (O, C, H, N) dominate living matter (≈95% of mass); minerals/trace elements are also required.
2.2 Chemical bonds and molecules
- Covalent bonds form when atoms share electrons to fill outer shells; the octet rule generally applies (H = 2 electrons).
- Covalent bonds can be polar or nonpolar depending on electronegativity differences; electronegativity values influence bond type.
- Hydrogen bonds and van der Waals dispersion forces enable interactions within and between molecules.
- Ionic bonds involve attraction between oppositely charged ions; salts like NaCl, KCl, CaCl2 are common examples.
- Molecules can adopt different shapes via rotation around bonds, affecting biological interactions.
2.3 Chemical reactions
- Reactions create new compounds by making/breaking bonds; often require energy input and/or enzymes; they tend toward equilibrium, but biological systems may not reach equilibrium due to subsequent reactions.
- Reactants and products are interconverted; example: $\mathrm{CH4 + 2\,O2 \rightleftharpoons CO2 + 2\,H2O}$
2.4 Properties of water
- Water is a major component of organisms; bulk properties and roles include solvent activity, participation in reactions, support, waste removal, cooling, cohesion/adhesion, surface tension, and lubrication.
- Substances can be hydrophilic (water-loving), hydrophobic (water-fearing), or amphipathic (both regions).
- Solutions are described by solute concentration; molarity ( $M = \dfrac{n}{V}$ ) expresses moles per liter; 1 mole = $6.022 \times 10^{23}$ particles.
- Water exists in three states (solid, liquid, gas); phase changes involve energy changes (high heat of vaporization, fusion, and specific heat) due to hydrogen bonding.
2.5 pH and buffers
- Hydrogen ion concentrations are modulated by acids, bases, and the autoionization of water; at 25°C, $[\mathrm{H^+}] [\mathrm{OH^-}] = 10^{-14}~\text{M}^2$ .
- The pH measures the H+ concentration: $\text{pH} = -\log_{10} [\mathrm{H^+}]$ ; pH range is 0–14 with 7 neutral.
- Pure water is neutral (pH ≈ 7); acids have pH < 7; bases have pH > 7; the scale is logarithmic, so each unit represents a tenfold change in [H+].
- Buffers resist pH changes by stabilizing [H+], notably the carbonic acid/bicarbonate system in body fluids.

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