The Chemical Foundation of Life
Atomic Theory and the Foundations of Chemistry
Chemistry is founded on four fundamental assumptions about atoms and matter (modern atomic theory):
1) All matter is composed of atoms.
2) Atoms of a given element differ from atoms of all other elements.
3) Chemical compounds consist of atoms combined in specific ratios.
4) Chemical reactions change only the way atoms are combined in compounds; the atoms themselves are unchanged.
Major Elements in the Human Body
Percent of total body weight by element (rounded values):
Oxygen (O): 65.0%
Carbon (C): 18.6%
Hydrogen (H): 9.7%
Nitrogen (N): 3.2%
Calcium (Ca): 1.8%
Phosphorus (P): 1.0%
Potassium (K): 0.4%
Sodium (Na): 0.21%
Chlorine (Cl): 0.21%
Magnesium (Mg): 0.11%
Sulfur (S): 0.05%
Iron (Fe): 0.03%
Iodine (I): 0.03%
These four elements (O, C, H, N) plus a few others make up about 96.5% of body weight.
Compounds and Elements
Sodium chloride (NaCl) is a compound composed of sodium ions (Na+) and chloride ions (Cl-).
A compound has characteristics different from its constituent elements.
Example concept: Sodium disposal and chemical combination form a compound with properties distinct from Na or Cl alone.
Periodic Table (Overview)
The periodic table organizes elements by increasing atomic number and groups them into main groups and transition metals.
Main groups are labeled 1A, 2A, 3A, 4A, 5A, 6A, 7A, and 8A, with noble gases typically in group 8A.
Elements in the same column (group) tend to have similar valence electron configurations and similar chemical properties.
The table includes alkali metals (Group 1), alkaline earth metals (Group 2), halogens (Group 7A), noble gases (Group 8A), and transition metals (d-block).
Atoms: Building Blocks of Elements
Atom: smallest unit of matter that retains all chemical properties of an element.
Two regions in an atom:
Nucleus: center containing protons (positive) and neutrons (neutral).
Electron cloud: outer region where electrons (negative) are found in orbit around the nucleus.
Subatomic particles:
Protons and neutrons: mass ~ 1 amu each (approximately).
Electrons: mass ~ 1/1836 of a proton; electrons are ~1836 times lighter than protons and neutrons.
Charges: protons +, electrons -, charges are equal in magnitude but opposite in sign.
Atomic Structure: Z, A, and Isotopes
Atomic Number Z: number of protons in each atom of an element. All atoms of a given element have the same Z.
Mass Number A: total number of protons and neutrons in an atom (A = Z + N).
Atoms are electrically neutral overall because the number of protons equals the number of electrons.
Isotopes: atoms of the same element that differ in the number of neutrons (hence different A but same Z).
Isotopes and Atomic Mass
Isotopes differ in neutron number; some isotopes are radioactive and decay over time.
Applications in biology: dating fossils, tracing atoms in metabolic processes, diagnostic imaging.
Atomic weight (atomic mass) is a weighted average of all naturally occurring isotopes:
ext{Atomic weight} =
\sumi [\text{abundance}i \times \text{mass}_i]
The Greek letter summation symbol is often used: \Sigma (\text{abundance} \times \text{mass})
Example: Chlorine exists as two main isotopes with non-equal abundances:
Cl-35: mass ≈ 34.97 amu, abundance ≈ 75.77%
Cl-37: mass ≈ 36.97 amu, abundance ≈ 24.23%
Weighted atomic weight ≈ 35.45 amu (illustrative calculation shown in class example)
Isotopes and their masses are depicted in isotope simulations and tables for learning.
Nuclear Model vs Quantum Mechanical Model
Early model: nucleus with protons and neutrons; electrons in a surrounding space model.
Quantum mechanical model: electrons occupy shells (energy levels) around the nucleus; shells are labeled n = 1, 2, 3, …
As the distance from the nucleus increases, shells are larger, can hold more electrons, and have higher energy.
The closest shell is Shell 1, next is Shell 2, etc.
Electron Configuration and Subshells
Within a shell, electrons occupy subshells: s, p, d, f in order of increasing energy.
Each shell contains a number of subshells equal to the shell number:
1st shell: 1s
2nd shell: 2s, 2p
3rd shell: 3s, 3p, 3d
etc.
Within each subshell, electrons occupy orbitals:
Each orbital can hold up to 2 electrons.
s subshell: 1 orbital (2e-)
p subshell: 3 orbitals (6e-)
d subshell: 5 orbitals (10e-)
f subshell: 7 orbitals (14e-)
The maximum electrons in a shell: 2n^2
1st shell: 2 e-
2nd shell: 8 e-
3rd shell: 18 e-
and so on
Electron configurations follow rules:
Aufbau principle: electrons fill the lowest-energy orbitals first.
Pauli exclusion principle: each orbital holds up to two electrons with opposite spins.
Hund’s rule: orbitals of the same energy are singly occupied before any orbital is doubly occupied.
Visualizing Orbitals and Shapes
s orbitals: spherical in shape.
p orbitals: dumbbell-shaped.
Electrons fill as per the energy ordering; some overlap occurs for shells 3 and 4.
Valence Electrons and the Periodic Table
Valence shell: outermost, highest-energy shell of an atom.
Valence electrons: electrons in the outermost shell; these electrons largely determine an element’s chemical properties.
Periodic table organization supports valence electron patterns: elements in the same group have the same number of valence electrons and similar chemistry.
Bonding Concepts: Covalent vs Ionic
Covalent bonds: formed when atoms share electrons; a group of atoms held together by covalent bonds is a molecule.
Ionic bonds: arise from electrostatic attraction between oppositely charged ions; typically form between metals and nonmetals.
Octet rule (main-group elements): seek to have eight electrons in their valence shell (two for hydrogen) through covalent bonding.
Bond types and counts:
Single bond: 1 shared electron pair.
Double bond: 2 shared electron pairs.
Triple bond: 3 shared electron pairs.
Represented by 1, 2, or 3 lines between atoms in diagrams.
Some atoms can form multiple bonds to satisfy octet (e.g., CO2, N2).
Polar Covalent Bonds and Electronegativity
In a covalent bond, electrons are shared in the region between two bonded atoms.
If atoms are identical, electrons are shared equally (nonpolar covalent bonds).
If atoms differ in electronegativity, electrons are pulled more toward the more electronegative atom, creating polar covalent bonds.
Dipole moment: separation of charge within a molecule; polar molecules have a net dipole moment.
Electronegativity trend:
Fluorine is the most electronegative element (value ≈ 4).
Electronegativity generally increases from left to right across a period and decreases from top to bottom in a group.
Noble gases are not assigned typical electronegativity values.
Bond polarity scale (rule of thumb):
Electronegativity difference < 0.5 → nonpolar covalent bond.
Difference up to 1.9 → increasingly polar covalent bonds.
Difference ≥ 2.0 → predominantly ionic bonds.
Polar molecules: sum of bond polarities and lone-pair contributions; polarity impacts melting/boiling points and solubility.
Representations: dipoles shown with an arrow toward the more negative end; δ+ (positive) on the less electronegative atom, δ− (negative) on the more electronegative atom.
Polar vs Nonpolar Molecules and Solutions
Polar molecules dissolve well in water; water is a versatile solvent.
Nonpolar compounds (e.g., oils) tend not to mix with water; oil and water form droplets due to immiscibility.
Hydration shells form around ions in water, stabilizing dissolved ions (e.g., NaCl in water).
Hydrophobic vs hydrophilic: hydrophilic substances have affinity for water; hydrophobic substances repel water (often nonpolar).
Ions: Cations and Anions; Ion Formation
Atoms can gain or lose electrons to form ions:
Cation: positively charged ion formed by loss of one or more electrons (e.g., Na+).
Anion: negatively charged ion formed by gain of one or more electrons (e.g., Cl-).
Ionization energy (IE): energy required to remove one electron from a gaseous atom.
Low IE favors cation formation.
Electron affinity (EA): energy released when adding an electron to a gaseous atom.
High EA favors anion formation.
Across the periodic table (left to right): IE and EA tend to increase.
Noble gases have high IE and low EA; they are generally nonreactive and do not form ions easily.
Ionic bonds form when opposite charges attract (e.g., Na+ and Cl- in NaCl).
Ionic Bonding: NaCl Crystal
In solid NaCl, Na+ and Cl- ions are arranged in a crystal lattice.
Each Na+ is surrounded by six Cl- ions, and each Cl- is surrounded by six Na+ ions, held together by ionic bonds.
Ionic compounds form crystalline lattices with properties distinct from the parent elements.
Acid-Base Chemistry and Aqueous Solutions
Brønsted concept preview: acids donate H+; bases accept H+.
Water is amphiprotic; it can act as both an acid and a base.
Acids dissociate to release H+ in solution (Bronsted definition): e.g., HCl → H+ + Cl-.
Common acids and their ions (examples):
Acetic acid: CH3COOH → CH3COO− + H+ (acetate ion)
Carbonic acid: H2CO3 → HCO3− + H+ (bicarbonate) or CO3^{2−} (carbonate) depending on proton loss
Hydrochloric acid: HCl → H+ + Cl−
Nitric acid: HNO3 → NO3− + H+
Nitrous acid: HNO2 → NO2− + H+
Phosphoric acid: H3PO4 → H2PO4− + H+ (and further dissociation to HPO4^{2-} and PO4^{3-})
Sulfuric acid: H2SO4 → HSO4− + H+ (and HSO4− → SO4^{2−} + H+)
Water autoionization: 2 H2O ⇌ H3O+ + OH− (often written as H2O ⇌ H+ + OH− in simplified form)
Water: The Solvent of Life
Water is essential to life due to its polarity and ability to form hydrogen bonds.
Composition: 60–70% of human body mass is water.
Hydrogen bonding: The polarity of the O–H bonds makes water a versatile solvent and enables hydrogen bonding between water molecules.
States of water: liquid, gas, and solid (ice) formed via hydrogen bonding; bonds are constantly formed/broken in liquid water; in ice, a more ordered lattice forms, making ice less dense.
Density anomaly: ice is less dense than liquid water, allowing ice to float.
Water’s heat-related properties:
High specific heat capacity: resists temperature changes, stabilizing climates and organisms.
High heat of vaporization: requires substantial energy to convert liquid water to gas.
Cohesion and surface tension: cohesive forces between water molecules create surface tension; allows objects to float on water surface or move across it (e.g., water striders).
Adhesion and capillary action: water climbs against gravity in narrow tubes due to adhesion to surfaces and cohesive forces.
Surfactants: amphipathic molecules that reduce surface tension (e.g., soaps, detergents).
Temperature moderation and environmental relevance: oceanic and climatic temperature regulation via water’s properties.
pH and acidity considerations in aqueous solutions: pH measures [H+]; pure water has equal [H+] and [OH−].
Acid precipitation and environmental impact: rain/snow with pH < 5.6 caused by atmospheric pollutants; effects on ecosystems.
Greenhouse effect: CO2 and other gases trap heat in the atmosphere, warming the planet.
Ocean acidification: increased CO2 lowers ocean pH, affecting calcifying organisms (e.g., corals).
Carbon: The Backbone of Life
Carbon’s specialty: tetravalence enables formation of large, diverse, complex molecules: proteins, DNA, carbohydrates, lipids.
Carbon bonding patterns:
Single bonds, double bonds, triple bonds adjust saturation and geometry.
Carbon chains form skeletons of organic molecules; branching and ring structures are common.
Hydrocarbons: molecules composed only of carbon and hydrogen; many biological molecules contain hydrocarbon components; hydrocarbons release large amounts of energy when broken down.
Isomerism and Molecular Diversity
Isomers: compounds with the same molecular formula but different structures/properties.
Structural (constitutional) isomers: different covalent arrangements.
Geometric (cis/trans) isomers: same covalent arrangement but different spatial arrangement around a double bond or ring.
Enantiomers: non-superimposable mirror images of each other; important in biology and pharmacology.
Example: Pentane (C5H12) shows structural isomers; cycloalkanes (e.g., cyclohexane) and branched isomers (e.g., 2-methylpropane) illustrate structural diversity.
Enantiomer significance in drugs: different enantiomers can have different biological activities (e.g., ibuprofen, albuterol).
Functional Groups: Key Reactive Groups in Biological Molecules
Functional groups are specific groupings of atoms that impart characteristic chemical properties to molecules.
Seven major functional groups important in biology:
Hydroxyl group (-OH)
Carbonyl group (C=O)
Carboxyl group (-COOH)
Amino group (-NH2)
Sulfhydryl group (-SH)
Phosphate group (-OPO3^{2-})
Methyl group (-CH3)
Hydroxyl group: polar; can form hydrogen bonds with water; contributes to solubility of alcohols and sugars.
Carboxyl group (-COOH): acidic; when deprotonated, exists as carboxylate (-COO−); found in amino acids and fatty acids; participates in acid-base chemistry and can form hydrogen bonds.
Carbonyl group: C=O; when within a carbon skeleton, forms ketones (R-CO-R') or aldehydes (R-CHO); polar; increases solubility of sugars.
Carboxyl vs carbonyl discussion shows isomeric relationships (e.g., acetone vs propanal).
Role of carbonyls in sugars leading to aldoses (contain aldehyde) and ketoses (contain ketone).
Methyl group: nonpolar; influences hydrophobicity and gene expression when attached to DNA or proteins; methylation can regulate function and structure.
Phosphate group: phosphorus atom bonded to four oxygens; one oxygen attached to carbon skeleton; two oxygens typically carry negative charges; forms organic phosphates and backbone structures in nucleic acids and energy carriers (e.g., ATP).
Sulfhydryl group (-SH): sulfur-containing analog of hydroxyl; forms cross-links in proteins (cysteine residues) stabilizing structure; can participate in redox reactions.
Amino group (-NH2): acts as a base; can accept a proton to form -NH3+ under physiological conditions; fundamental in amino acids.
Examples and mentions:
5-Methyl cytidine in DNA modification; glycerol phosphate as backbone for phospholipids; cysteine in proteins; glycine as an amino acid; methylation in DNA affects gene expression.
Phosphate groups contribute negative charge to molecules (e.g., ATP, DNA backbone).
Practical and Environmental Contexts
Water’s role in biology intersects with environmental processes:
Water acts as solvent for ionic and polar substances; ions form hydration shells in solution.
Hydrogen bonding and polarity underpin water’s unique properties that support life.
Acid rain and CO2-induced ocean acidification alter the chemical environment of organisms and ecosystems.
Summary: Core Concepts in a Nutshell
Atomic theory frames matter as composed of atoms; chemical reactions rearrange atoms, not change them.
The human body is built from a small set of major elements; these form compounds with distinctive properties.
Atoms have a nucleus with protons and neutrons, and electrons in orbitals organized into shells, subshells, and orbitals.
Electron configurations follow principled rules (aufbau, Hund’s, Pauli), determining element behavior and valence.
Bonding types (covalent, ionic) and electronegativity drive molecular structure and properties; polarity influences interactions and solubility.
Water’s properties (polarity, hydrogen bonding, high specific heat, high heat of vaporization, solvent capabilities) are critical for life and climate.
Carbon’s tetravalence enables diverse, complex organic molecules; isomerism expands chemical diversity (structural, geometric, enantiomeric).
Functional groups (hydroxyl, carbonyl, carboxyl, amino, sulfhydryl, phosphate, methyl) determine reactivity and function of biomolecules.
Acids and bases in water modulate pH and biological processes; water participates in many chemical equilibria in living systems.
The environment (greenhouse effect, acid rain, ocean acidification) intersects chemistry with biology and ecology.
ext{Atomic weight}=
\sumi [\text{abundance}i \times \text{mass}_i]
ext{Shell capacity} = 2n^2, ext{ with subshell capacities: } s:2,
p:6,
d:10,
f:14.
ext{Electron configuration rules: Aufbau, Pauli, Hund.}
ext{pH} = -\log_{10}([\text{H}^+])
In pure water at 25°C, [H+] = [OH−] = 1.0×10^{-7} M, giving pH = 7 (neutral).
Note
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