Comprehensive Biology Chemistry Notes

Why Chemistry in Biology

• Biologists study the chemical nature of substances and the properties of water.
• Matter interacts with other matter; chemical interactions in aqueous environments underpin biology.
• Atoms are the fundamental constituents of matter; chemistry is essential to understanding biology.

ATOMS

• An element is made up of only one kind of atom.
• The number of protons identifies the element (atomic number).
• Isotopes differ in the number of neutrons; atomic mass is the average mass of all isotopes of an element.
• The chemical symbol and atomic data are used to display elements in tables (example data shown for several elements on the slides).

The Periodic Table and Elementary Data (Mendeleev era to modern table)

• Elements are organized by atomic number; also by recurring chemical properties.
• Periodic table illustrates groups (columns) with similar properties and periods (rows) of increasing atomic number.
• The table contains specialized series: Lanthanide series and Actinide series (shown on the slides).
• Atomic number vs. atomic mass: atomic number identifies the element; atomic mass is the average mass of isotopes.
• Common illustrative data shown include:
  • H: chemical symbol H; atomic number 1; atomic mass ~1.0079
  • He: atomic number 2; atomic mass ~4.003
  • Na, Mg, Al, Si, Cl, Ar, K, Ca, etc., with their atomic masses and numbers (examples from the slide set).
• The visual data include a mix of elements’ symbols, atomic numbers, and approximate atomic masses or atomic weights, along with some notes about isotopes and lanthanide/actinide series.

Quarks and Nucleons

• The two light flavors of quarks that make up ordinary matter are the 'up' (u) and 'down' (d) quarks.
• Up quark charge: + frac{2}{3} of the electron charge.
• Down quark charge: - frac{1}{3} of the electron charge.
• A proton is composed of three quarks: two ups and one down.
• A neutron is composed of two downs and one up.
• Quark-charge combinations yield the correct overall charged particle:
  • Proton: two ups and one down → $\big(+ frac{2}{3}\big) + \big(+ frac{2}{3}\big) + \big(- frac{1}{3}\big) = +1.$
  • Neutron: two downs and one up → $\big(- frac{1}{3}\big) + \big(- frac{1}{3}\big) + \big(+ frac{2}{3}\big) = 0.$
• This underscores that protons are +1 charge and neutrons are neutral.

Atomic Mass and Atomic Weight

• Atomic mass number (A): the sum of protons and neutrons in an atom's nucleus.
• Atomic weight (atomic mass): the average of the atomic mass numbers of a sample of atoms of an element, weighted by isotopic abundances.
• Isotopes vary in neutron number; common hydrogen isotopes: protium (1H), deuterium (2H), tritium (3H).
• Common carbon isotopes shown: 12C (stable) and 14C (radioactive).

Isotopes (examples)

• Hydrogen isotopes:
  • 1H: 1 proton, 0 neutrons.
  • 2H (deuterium): 1 proton, 1 neutron.
  • 3H (tritium): 1 proton, 2 neutrons.
• Carbon isotopes:
  • 12C: 6 protons, 6 neutrons.
  • 14C: 6 protons, 8 neutrons.
• These illustrate how isotopes have identical chemical behavior (same Z) but different masses due to neutron count.

Electron Behavior and Chemical Bonding

• Electron behavior determines chemical bonding; electrons occupy shells (K, L, M, …).
• Outer-shell (valence) electrons govern chemical reactivity: atoms with incomplete outer shells are chemically reactive; filled outer shells are chemically inert.
• The octet rule arises from the tendency to form eight electrons in the valence shell for stability (often described in terms of s and p orbitals: s^2 p^6).
• Hydrogen and light elements fill shells differently, but the general principle of achieving a stable valence shell drives bonding.

Shell structure and reactivity

• Examples show which elements tend to form bonds due to incomplete valence shells (e.g., H, C, N, O, Cl, etc.).
• The slides note that outer-shell filling correlates with chemical reactivity; fully filled shells imply inertness.

MOLECULES AND CHEMICAL BONDS

• Chemical bonds are attractive forces linking atoms to form molecules.
• Main bond types:
  • Covalent bond
  • Polar covalent bond
  • Ionic bond
  • Hydrogen bond
  • van der Waals attractions

Covalent bond

• A covalent bond occurs when two atoms share one or more pairs of electrons in their outer shells.
• The shared electrons are distributed around the bonded atoms; the bond strength depends on how electrons are shared.
• Example: Methane (CH4). Carbon completes its outer shell by sharing electrons with four hydrogen atoms, forming a tetrahedral arrangement.
• Visual notes from slides show that the carbon atom forms four bonds in a tetrahedral geometry around a central carbon.
• Textual representation (conceptual): each line/pair of electrons represents a shared electron pair.

Hydrogen atoms and covalent bonding (H2)

• Two hydrogen atoms covalently bond to form a diatomic hydrogen molecule.
• The shared electrons are attracted by both protons.

Other covalent bonding examples

• Methane (CH4) as a tetrahedral molecule formed by carbon sharing electrons with four hydrogens.
• Carbon can form multiple covalent bonds (single, double, triple) depending on how many electrons are shared.

Polar covalent bond

• A polar covalent bond occurs when electrons are shared unequally in a covalent bond, giving partial charges on atoms within the molecule.
• Example: Water (H2O) has polar covalent bonds; electron density is greater around the oxygen atom, producing partial negative charge on O and partial positive charges on H.
• Dipole: water has a dipole moment due to uneven electron distribution.

Water as a polar molecule

• Water's polar covalent bonds cause a dipole with partial charges: δ- on O, δ+ on H.
• This polarity facilitates hydrogen bonding and solvent properties.

Ionic bonds

• Ionic bonds form by electrical attraction between oppositely charged ions (cations and anions).
• Ions arise when atoms gain or lose electrons, forming charged species:
  • Anions: negative charge (e.g., Cl^-)
  • Cations: positive charge (e.g., Na^+)
• Example: Na and Cl exchange electrons to form Na^+ and Cl^- ions, which attract each other in a lattice (sodium chloride).

Hydrogen bonds

• Hydrogen bonds are electrostatic attractions between partially positive hydrogen atoms and electronegative atoms (e.g., O in water, N in amines).
• They are weaker than covalent or ionic bonds but crucial for the structure of water, DNA, proteins, and other biomolecules.

van der Waals attractions

• Occur between nonpolar molecules (or nonpolar regions) when induced dipoles cause transient attractions.
• These forces are weaker and require close contact; they contribute to the behavior of nonpolar substances and molecular interactions.
• Nonpolar molecules have little to no attraction with polar molecules at a distance, but can exhibit van der Waals interactions with other nonpolar molecules when very close.

MOLECULES: Water and Solutions

• Water is polar and drives many biological processes.
• Water properties include:
  • Ice floats on liquid water due to lower density of the solid phase (ice lattice).
  • Water is an excellent solvent for polar substances and ionic compounds.
  • Water is cohesive (sticking to itself) and adhesive (sticking to other substances).
  • Water moderates temperature changes (high heat capacity and high heat of vaporization).

pH, acids, bases, and buffers

• Water self-ionizes to hydronium (H3O+) and hydroxide (OH-) ions:
  • Hydrogen ion behavior can be represented as:
    $\mathrm{H_2O} \rightleftharpoons \mathrm{H^+} + \mathrm{OH^-}$
• In solution, free protons exist primarily as H3O+ (hydronium).
• pH is the measure of H+ ion concentration: $\mathrm{pH} = -\log_{10}[\mathrm{H^+}]$
  • Pure water at 25°C has [H+] = 1.0 × 10^-7 M, so pH = 7 (neutral).
• The strength of acids and bases is given by their tendency to donate/accept protons; stronger acids have weaker conjugate bases, and stronger bases have weaker conjugate acids.
• Acid-base conjugate pairs: removing a proton from an acid forms its conjugate base; adding a proton to a base forms its conjugate acid.
• Equilibrium nature of weak acids: not all acids completely dissociate in water; the equilibrium constant is Ka, and pKa = -log Ka.
• Henderson-Hasselbalch equation (buffering context): $\mathrm{pH} = \mathrm{p}K<em>a + \log</em>{10}\left(\frac{[A^-]}{[HA]}\right)$
  • When [A^-] = [HA], pH = pKa; this implies half-dissociation of the weak acid.
• Buffers: systems that resist pH changes by absorbing or releasing H+; typically consist of a weak acid and its conjugate base; buffers work best near the pKa of the weak acid.
• Le Chatelier’s principle applied to buffers:
  • Adding H+ shifts toward the weak acid;
  • Adding OH- shifts toward the conjugate base and water;
  • Net H+ concentration remains relatively constant.

Functional groups important to living organisms

• Functional group: a specific group of atoms that imparts characteristic properties to molecules.
• Common functional groups (with class, structural formula, example):
  • Hydroxyl group (-OH): class – Alcohols; example – Ethanol; structural representation: R-OH; H-C-C-OH.
  • Carbonyl group (-CHO) for aldehydes (Acetaldehyde).
  • Carboxyl group (-COOH) for carboxylic acids (Acetic acid).
  • Amino group (-NH2) for amines (Methylamine).
  • Phosphate group (-OPO3^2-) for organic phosphates (varied representations show phosphate esters).
  • Sulfhydryl group (-SH) for thiols (Mercaptoethanol).
• 3D depictions and ionization at cellular pH are noted for these groups; the slide shows conventional and 3D representations of amino and carboxyl groups in amino acids.

Isomers

• Isomers are compounds with the same molecular formula but different arrangements of atoms.
• Structural isomers differ in connectivity (how atoms are bonded); examples include butane and isobutane.
• Optical (stereoisomers) are mirror images that cannot be superimposed; many biomolecules exhibit chirality.

Chiral Centers and Amino Acids

• An asymmetric carbon (chiral center) is bonded to four different groups; common in amino acids.
• In biology, most amino acids in proteins are L-enantiomers (L-amino acids).

Why Life is Carbon-Based

• Carbon atoms can link to form stable, long chains (catenation) forming backbones for large, complex molecules.
• Carbon has valency IV: each atom can form four covalent bonds, allowing diverse bonding patterns (single, double, or triple bonds).
• Branched carbon chains enable an almost limitless variety of compounds, supporting biological complexity.
• Silicon (valency IV) is less suitable for long, stable chains under biological conditions; Si-Si bonds are weaker and less conducive to forming biochemically relevant macromolecules.
• Hence, life is based on carbon rather than silicon for forming complex molecules.

Elements and Biochemical Bonding Capacities (Biologically Important Elements)

• Hydrogen, Oxygen, Nitrogen, Carbon, Phosphorus, Sulfur are central to biology.
• Covalent bonding capabilities differ by element (for example, H tends to form 1 bond; C tends to form 4 bonds; O tends to form 2 bonds; N tends to form ~3 bonds; P and S can form multiple bonds depending on oxidation state).
• The slides illustrate that different elements have characteristic bonding capacities, which underpin macromolecule formation (proteins, nucleic acids, lipids, carbohydrates).

Reactions in Biology

• Molecules undergo various chemical reactions important to biology, including:
  • Nucleophilic substitution: replacement of one atom/group by a nucleophilic atom/group (common nucleophiles are O, N, S in biomolecules).
  • Elimination: removal of atoms to form a double bond; dehydration is a common type (removal of water to form a double bond).
  • Addition: two molecules combine across a double bond; hydration is an example (addition of water across a double bond).
  • Isomerization: rearrangement within a molecule to form an isomer.
  • Hydrolysis: water adds to break a covalent bond (e.g., hydrolysis of ATP).
  • Oxidation–Reduction (Redox): transfer of electrons between donor and acceptor; oxidation is loss of electrons, reduction is gain of electrons; these processes are coupled.

A Note on Chemistry in Biological Contexts

• Water’s polarity underpins most properties of life-supporting chemistry: solvent capacity, ion solvation, acid-base chemistry, and temperature regulation.
• The concept of pH and buffers is central to maintaining stable intracellular environments, enabling enzymes and macromolecules to function optimally.
• The carbon-based backbone concept explains the diversity and complexity of biomolecules; while other elements contribute functional groups and properties, carbon-based structures enable long, branched, and cyclic architectures.

Summary of Key Formulas and Concepts (LaTeX)

• Proton and neutron sum in protons and neutrons (quark composition):
  • Proton: $\big(+ frac{2}{3}\big) + \big(+ frac{2}{3}\big) + \big(- frac{1}{3}\big) = +1.$
  • Neutron: $\big(- frac{1}{3}\big) + \big(- frac{1}{3}\big) + \big(+ frac{2}{3}\big) = 0.$
• Atomic structure concepts:
  • Atomic mass number: A = Z + N (protons + neutrons).
  • Atomic weight: weighted average of isotopic masses based on abundance.
• pH and [H+]:
  • $ext{pH} = -\log_{10}[\mathrm{H^+}]$
• Acids, bases, conjugates: Bronsted concept (donate/accept protons) and conjugate pairs.
• Henderson–Hasselbalch equation (buffers):
  • $\text{pH} = \text{p}K<em>a + \log</em>{10}\left(\frac{[A^-]}{[HA]}\right)$
  • When [A^-] = [HA], $\text{pH} = \text{p}K_a$.
• Avogadro’s number and mole concept:
  • $N_A = 6.023 \times 10^{23} \text{ molecules per mole}$
• Water autoionization (simplified):
  • $\mathrm{H_2O} \rightleftharpoons \mathrm{H^+} + \mathrm{OH^-}$
  • In aqueous solutions, H^+ is mainly present as H_3O^+ (hydronium).
• CH4 molecular weight:
  • $M<em>r(\mathrm{CH</em>4}) = 16.00$
• Example reactions (not balanced in the source):
  • $\mathrm{CH<em>3Br} + \mathrm{Cl} \rightarrow \mathrm{CH</em>3Cl} + \mathrm{Br}$
  • General hydrolysis, hydration, dehydration and redox reactions as described in the slides.

Connections to Previous and Real-World Contexts

• The chemical basis of life emerges from simple principles: atomic structure, bonding, and polarity shape macromolecules and cellular processes.
• Water’s properties enable solvent behavior, transport, and metabolic reactions in cells.
• The concept of pH and buffering explains how organisms maintain stable internal environments (homeostasis) despite external changes.
• Carbon’s versatility explains the diversity of biomolecules (carbohydrates, lipids, proteins, nucleic acids) and the vast number of possible structures.
• Isomerism and chirality influence biomolecule function (e.g., L-amino acids in proteins).

Practical and Ethical/Philosophical Implications

• Understanding molecular bonding informs drug design, metabolic engineering, and environmental chemistry.
• The reliance on carbon chemistry highlights why life as we know it is conditional on chemistry that supports complex, yet stable macromolecular architectures.
• The global importance of pH balance and buffering underscores the need for careful management of biological systems in medicine and industry.

Quick Reference Tables (conceptual)

• Bond types: covalent, polar covalent, ionic, hydrogen, van der Waals.
• Major properties of water: polar solvent, high heat capacity, cohesion, adhesion, ice less dense than liquid water.
• Functional groups: hydroxyl, carbonyl (aldehyde), carboxyl, amino, phosphate, sulfhydryl.
• Common amino acid chirality: L-amino acids predominate in proteins.
• Common biochem reactions: nucleophilic substitution, elimination/dehydration, addition/hydration, isomerization, hydrolysis, redox.
• pH and buffers: Henderson-Hasselbalch context; buffering near pKa.
• Isotopes: mass differences do not usually alter chemistry for light elements; isotopes differ in neutron count.
• Avogadro’s number: a mole contains approximately $6.023\times 10^{23}$ molecules.