Lecture 2: Chemistry of Life (BIO 111)

Atomic Structure and Basic Concepts

• Matter basics: Matter is anything that takes up space and has mass.

• The building blocks of matter are atoms and molecules; compounds are made of two or more different elements.

• An atom is the smallest functional unit of matter; an element is defined by the number of protons in its nucleus (the atomic number).

• Molecules are two or more atoms bound together; if the molecule contains two or more different elements, it is a compound. If the molecule contains only one element, it is an element molecule (e.g., O₂).

• Macromolecules are large, complex molecules formed when smaller units bond (e.g., carbohydrates, proteins, lipids, nucleic acids).

Subatomic Particles and Atomic Structure

• Subatomic particles: protons (+), neutrons (no charge), and electrons (−).

• Protons and neutrons are located in the atomic nucleus; electrons occupy orbitals around the nucleus.

• The nucleus contains protons and neutrons; electrons are,held near the nucleus by electrostatic attraction between the

  . Ïn negative electrons and the positive protons in the nucleues

• Atomic number Z = number of protons in the nucleus.

• Mass number A = total number of protons and neutrons in the nucleus; mass number ≈ atomic mass unit count.

• The sum of protons and neutrons in an atom is the mass number: A = Z + N where N is the number of neutrons.

• The mass of protons and neutrons is about 1 dalton each (1 Da ≈ 1 amu).

• Isotopes are forms of an element with different numbers of neutrons (same Z, different A). Example

• - Hydrogen I (Protium): mass ~ 1, abundance ≈

  99.9%

  • - ‘ Hydrogen II (Deuterium): mass ~ 2, abundance ≈

    0.011%%

  • Hydrogen III (Tritium): mass ~ 3, abundance ≈ 0.003%

• Atomic mass is the average mass of an element’s naturally occurring isotopes weighted by abundance: ext{Atomic mass} = ext{weighted average of isotope masses}.

Electron Configuration and Energy Levels

• Electrons possess energy and occupy regions called orbitals; each orbital can hold up to 2 electrons.

• Orbitals are grouped into energy levels, or electron shells.

• Innermost shell (closest to the nucleus) has 1 orbital; the second shell has 4 orbitals; and so on.

• Electrons fill innermost shells first following the Aufbau principle.

• The energy (potential energy) of electrons increases with distance from the nucleus: more distant electrons have higher energy.

• Electron configuration example (second shell): 1s^2\;2s^2\;2p^6

• Valence electrons are those in the outermost shell; they determine bonding behavior.

• An atom is most stable when its valence shell is full (octet rule).

The Octet Rule and Bonding Propensity

• Octet rule: Atoms tend to gain, lose, or share electrons to achieve a full valence shell of 8 electrons (except for some elements like H which follow a duet).

• Consequences:

  • Reactivity is driven by the number of electrons in the valence shell.

  • Valence electrons allow atoms to form chemical bonds with other atoms.

  • Elements with the same number of valence electrons have similar bonding properties.

• Representative valences (illustrative): many main-group elements seek to complete 2 or 8 electrons in their valence shell depending on the period and orbital structure.

Periodic Table Basics and Trends

• The primary determinant of an element’s chemical properties is the electron configuration, especially the outermost (valence) electrons.

• Period (horizontal row) corresponds to the number of electron shells; it determines the period number.

• Group (vertical column) corresponds to the number of electrons in the outermost shell; it determines the group number.

• The highlighted elements commonly found in organisms

  include H, C, N, O, P, S, and others; periodic diagrams

  illustrate their positions and valence relationships..

• The periodic table is organized into blocks (alkali metals, alkaline earth metals, transition metals, metalloids, nonmetals, Halogens, Noble Gases) and shows trends in atomic radius and electronegativity.

Element Abundance in the Human Body

• Table (approximate body mass percentages including water):

  • Oxygen (O) — 65.0%

  • Carbon (C) — 18.5%

  • Hydrogen (H) — 9.5%

  • Nitrogen (N) — 3.3%

  • Calcium (Ca) — 1.5%

  • Phosphorus (P) — 1.0%

  • Potassium (K) — 0.4%

  • Sulfur (S) — 0.3%

  • Sodium (Na) — 0.2%

  • Chlorine (Cl) — 0.2%

  • Magnesium (Mg) — 0.1%

  • Trace elements (less than 0.01%): B, Cr, Co, Cu, F, I, Fe, Mn, Mo, Se, Si, Sn, V, Zn

• Source: 2014 Pearson Education, Inc. (Table of Elements in the Human Body)

Atomic Structure Details

• Subatomic particles: protons (positive), neutrons (neutral), electrons (negative).

• Protons and neutrons reside in the nucleus; electrons orbit in regions called orbitals.

• The nucleus exerts an attractive force on electrons; electrons are held in orbit by this attraction.

• Atomic number Z equals the number of protons in the nucleus.

• Mass number A equals the sum of protons and neutrons: A = Z + N.

• Isotopes differ in neutron number; identical Z but different A.

Bonding and Bond Types

• Bond types vary by how electrons are distributed between atoms and by their electronegativity differences.

Covalent Bonds

• Occur when atoms share electrons; electrons are effectively associated with both atoms.

• Strength: generally very strong; electrons contribute to full valence shells for both atoms.

• Types:

  • Single covalent bond: one shared pair of electrons.

  • Double covalent bond: two shared pairs of electrons.

  • Triple covalent bond: three shared pairs of electrons.

• Examples:

  • Carbon dioxide: CO₂ (double bonds between C and O)

  • Molecular nitrogen: N₂ (triple bond between two N atoms)

  • Oxygen gas: O₂ (double bond between two O atoms)

• Covalent bonds can be either polar or nonpolar depending on electronegativity difference (see electronegativity section).

Ionic Bonds

• Occur when one atom transfers one or more electrons to another, creating oppositely charged ions that attract each other.

• Cation: positively charged ion; Anion: negatively charged ion.

• Example question from lecture: why Na and Cl do not just share electrons? Because the difference in electronegativity is large enough to favor electron transfer and formation of ions rather than equal sharing.

Hydrogen Bonds

• Weak bonds that form between a hydrogen atom attached to a strongly electronegative atom (like O or N) and another electronegative atom with lone pairs.

• In cells, often involve N or O as the hydrogen bond acceptor/donor.

• Many hydrogen bonds together can provide significant stabilization of 3D macromolecule structures.

Van der Waals Interactions

• Very weak, arises from transient partial charges due to movement of electrons.

• Occur when molecules are in very close proximity; can sum to a meaningful force if many interactions occur simultaneously.

Electronegativity and Bond /.Electronegativity: the ability of an atom to attract bonding

electrons toward itself when forming a bond.

• The more electronegative an atom, the more it pulls electrons toward itself in a covalent bond.

• Shared electrons are drawn closer to the more electronegative atom, creating partial charges and polar covalent bonds.

• Nonpolar covalent bonds form when electronegativities are equal or nearly equal; electrons are shared roughly equally.

• Examples:

  • Polar covalent bond: H–O in water (oxygen is more electronegative than hydrogen)

  • Nonpolar covalent bond: C–H bonds (relative electronegativities lead to near-equal sharing)

• Electronegativity trend (general): increases across a period and decreases down a group (illustrated by the trend in the periodic table). The likelihood of ionic bonding increases with larger electronegativity differences between bonded atoms.

• Electronegativity scale (illustrative values):

  • H ≈ 2.20, Li ≈ 0.98, O ≈ 3.44, F ≈ 3.98 (values approximate; see lecture table)

• Bond type from electronegativity difference ΔEN:

  • Small ΔEN (< 0.5): nonpolar covalent

  • Moderate ΔEN (≈ 0.5–≈ 1.7): polar covalent

  • Large ΔEN (> ≈ 1.7): ionic

• Numerical example from lecture:

  • Electronegativity values shown include H (2.20) and O (3.44) leading to polar covalent character in water bonds.

Water: Structure, Bonding, and Solvent Properties

• Water is a polar molecule formed by two hydrogen atoms and one oxygen atom linked by polar covalent bonds.

• Bonding in water: polar covalent due to unequal sharing of electrons (O more electronegative than H).

• Molecular geometry: bent shape due to lone pair repulsion on oxygen; partial negative charge on oxygen, partial positive charges on hydrogens.

• Water forms hydrogen bonds: each water molecule can hydrogen-bond with multiple partners; bonds constantly break and reform in liquid water.

• Water can hydrogen-bond with other polar molecules and ions, facilitating dissolution.

• Hydrophilic vs hydrophobic:

  • Hydrophilic substances interact with water due to polar or charged nature or ability to form hydrogen bonds.

  • Hydrophobic substances (nonpolar, uncharged) do not interact well with water and tend to aggregate with each other.

• Water as solvent: water dissolves more substances than any other molecule known due to its polarity and hydrogen-bonding capacity.

• Water’s 3D structure and solvent properties underpin many biological processes and real-life questions.

Four Remarkable Properties of Water (and What They Mean)

• 1) Universal (versatile) solvent: dissolves many substances via polar interactions and hydrogen bonds.

• 2) Cohesion, adhesion, and surface tension: cohesion = attraction between like molecules; adhesion = attraction between different susbstances; /surface tension reflects theresistance to external force at the surface.

• 3) Density anomaly: liquid water is denser than ice; ice floats on liquid water.

• 4) High capacity to absorb energy: high specific heat, high heat of vaporization; water requires substantial energy to raise temperature or to convert to gas.

• Practical implications: water moderates climate, supports biochemical reactions, and enables life-sustaining solvent environments.

Practical and Real-World Relevance

• Dissolution of ionic compounds in water (e.g., NaCl) occurs because water stabilizes ions through hydration shells.

• Polar and charged molecules are hydrophilic and dissolve in water; nonpolar molecules are hydrophobic and resist dissolution.

• Water’s solvent properties influence chemical reactions, nutrient transport, and metabolic pathways in living organisms.

• The geometry of water and hydrogen bonding plays a critical role in shaping macromolecular structures (proteins, nucleic acids, polysaccharides).

Isotopes, Atomic Mass, and Electron Shells (Additional Details)

• Mass and isotopes:

  • Atomic mass reflects the weighted average of naturally occurring isotopes (more abundant isotopes contribute more to the average).

  • Isotopes can differ in neutron count but share the same atomic number (Z).

• Electron shells and capacity:

  • Each shell can hold up to 2n^2 electrons, where n is the shell number.

  • Innermost shell (n=1) holds up to 2 electrons; second shell (n=2) up to 8 electrons; etc.

• Electron configuration influences chemical behavior and reactivity via valence electrons.

Quick Reference: Key Definitions and Concepts

• Atom: the basic unit of an element; defines element properties by its number of protons (Z).

• Element: substance consisting of atoms with the same number of protons.

• Molecule: two or more atoms bonded together.

• Compound: molecule composed of two or more different elements.

• Bond types: covalent (shared electrons), ionic (electrons transferred), hydrogen (special dipole-based interactions), van der Waals (weak intermolecular forces).

• Electronegativity: tendency of an atom to attract bonding electrons.

• Polar covalent bond: unequal sharing of electrons due to electronegativity differences; partial charges exist.

• Nonpolar covalent bond: approximately equal sharing of electrons; no permanent dipole.

• Ionic bond: transfer of electrons forming cations and anions with electrostatic attraction.

• Valence electrons: electrons in the outermost shell that participate in bonding.

• Octet rule: atoms prefer full valence shells (often 8 electrons for main-group elements).

• Hydrophilic: interacts with water; dissolves in water.

• Hydrophobic: repels water; tends to cluster with other nonpolar molecules.

• Hydrogen bond: a weak bond between a partially positive hydrogen atom and a lone pair on a highly electronegative atom (N, O, or F).

• Van der Waals forces: weak, transient attractions between molecules due to momentary dipoles.

• Periodic trends: periodic table structure reflects electron configurations; periods correspond to energy shells; groups reflect valence electron count.

• Water structure: bent due to lone-pair repulsion; polarity arises from asymmetric electron distribution.

• Real-life implications: water’s properties enable life, biochemical reactions, nutrient transport, and climate regulation.