chapter 2: The Chemical Foundation of Life
Elements, Atoms, and Compounds
Organisms are composed of elements, in combinations called compounds.
Matter: anything that occupies space and has mass; is composed of chemical elements.
Element: a substance that cannot be broken down further by chemical means; there are 92 natural elements, only a few exist in pure state.
Major and trace elements in the human body
Table 2.1 (elements in the human body and approximate body-weight percentages):
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% of body weight): Boron (B), chromium (Cr), cobalt (Co), copper (Cu), fluoride (F), iodine (I), iron (Fe), manganese (Mn), molybdenum (Mo), selenium (Se), silicon (Si), tin (Sn), vanadium (V), zinc (Zn).
Four elements make up about 96% of the weight of most living organisms: Oxygen, Carbon, Hydrogen, and Nitrogen.
Trace elements are essential but needed only in minute quantities.
Compounds and their properties
Compound: two or more different elements in a fixed ratio; compounds are more common than pure elements.
Example: Sodium chloride (NaCl) is a common compound made of equal parts sodium (Na) and chlorine (Cl).
Emergent properties: compounds can have properties not predictable from their constituent elements (e.g., NaCl’s edible properties).
Sodium chloride (NaCl) is formed from Sodium (Na) and Chlorine (Cl) and dissolves to form Na⁺ and Cl⁻ in water.
Food additives and trace elements
Trace elements are common additives to food and water to preserve, enhance nutrition, or improve appearance.
Iron helps transport oxygen; Iodine helps prevent goiter; Fluoride helps reduce tooth decay.
Goiter as a symptom of iodine deficiency.
Fluoride is common in mouthwash and toothpaste.
Fortified foods (e.g., fortified cereal) include added vitamins and minerals; nutrition labels show daily values based on a 2,000-calorie diet.
Nutrition facts (fortified cereal example)
Serving size: 3/4 cup (30 g); about 17 servings per container.
Per serving (cereal vs skim milk):
Calories: 100 (cereal) / 140 (skim milk)
Calories from fat: 5 (cereal) / 10 (milk)
Total Fat: 0.5 g (1% DV) cereal; 1% DV milk
Saturated Fat: 0 g (0%) for both
Trans Fat: 0 g
Polyunsaturated Fat: 0 g
Monounsaturated Fat: 0 g
Sodium: 135 mg (6%) cereal; 9% DV in the combined label section
Potassium: 125 mg (4%) cereal; 10% DV combined
Total Carbohydrate: 23 g (8%) cereal; 10% DV combined
Dietary Fiber: 3 g (10%) cereal; 10% DV combined
Sugars: 5 g (11 g sugars noted in some sections)
Protein: 2 g
Vitamin A: 10% / 15% DV
Vitamin C: 100% / 100% DV
Calcium: 100% / 110% DV
Iron: 100% / 100% DV
Vitamin D: 10% / 25% DV
Other vitamins/minerals listed include Niacin, Vitamin B6, Vitamin E, Thiamin, Riboflavin, Folic Acid, Vitamin B12, Pantothenic Acid, Phosphorus, Magnesium, Zinc, Copper, etc. (DV percentages shown for cereal vs milk in the packaging).
Summary note: A serving of cereal with skim milk provides approximately:
Total fat: ~1 g
Cholesterol: <5 mg
Sodium: ~260 mg
Potassium: ~290 mg
Total carbohydrate: ~29 g (about 11 g sugars)
Protein: ~7 g
Atoms, subatomic particles, and isotopes
Atom: the smallest unit of matter that retains the properties of an element.
Subatomic particles: protons (positive), neutrons (neutral), electrons (negative).
Nucleus: protons and neutrons; Electron cloud/orbit around the nucleus.
Atomic number (Z): number of protons in the nucleus.
Mass number (A): total number of protons and neutrons in the nucleus; Mass number = protons + neutrons.
Atomic mass: approximately equal to the mass number for most atoms.
Isotopes: same number of protons, different numbers of neutrons; isotopes have the same chemical behavior but different masses and sometimes stability.
Radioactive isotopes: isotopes that decay spontaneously, emitting particles and energy.
Table 2.3 (isotopes of carbon):
Carbon-12: Protons 6, Neutrons 6, Electrons 6, Mass number 12
Carbon-13: Protons 6, Neutrons 7, Electrons 6, Mass number 13
Carbon-14: Protons 6, Neutrons 8, Electrons 6, Mass number 14
Applications and risks:
Radioactive isotopes can act as tracers in metabolic processes and are used in medical diagnosis (e.g., PET imaging).
Dangers of radioactive substances include damage to molecules in living cells and the potential to form abnormal chemical bonds.
Models of the atom and electron distribution
Two models of a helium atom are used to illustrate electron distribution: nucleus with protons and neutrons; electron cloud surrounding the nucleus.
Electron distribution determines an atom’s chemical properties; only electrons participate in chemical activity.
Electron shells and valence shell (outermost electrons) determine reactivity; atoms react to complete outer shells by sharing, donating, or receiving electrons.
The periodic table encodes typical electron distributions and chemical properties.
Chemical bonds: covalent, ionic, and hydrogen bonds
Covalent bonds: atoms share a pair of electrons; molecules are formed when covalent bonds connect atoms.
Electronegativity: the strength with which an atom pulls shared electrons toward itself.
Polar covalent bonds: unequal sharing of electrons due to differences in electronegativity; slight partial charges develop (e.g., in H2O).
Nonpolar covalent bonds: equal sharing of electrons; no significant charge separation.
Examples of covalent bonds:
H2: two hydrogen atoms share a pair of electrons (single covalent bond).
O2: two oxygen atoms share two pairs of electrons (double covalent bond).
CH4 and H2O: compounds made of two or more different elements.
Ionic bonds: attractions between ions of opposite charge; transfer of electrons creates cations and anions that attract each other.
Example: NaCl (sodium chloride).
Hydrogen bonds: weak bonds important in biology, formed between polar molecules (e.g., water). Each water molecule can hydrogen-bond to as many as four partners.
Visual representations include molecular formula, electron distribution diagrams, structural formulas, and space-filling models for common molecules.
Water: life’s solvent and its life-supporting properties
Water’s properties arise from hydrogen bonding between water molecules:
Cohesion: attraction between like molecules; water has unusually strong cohesion.
Adhesion: attraction between different substances.
Surface tension: high surface tension due to hydrogen bonding; enables water striders to walk on water.
Water moderates temperature:
Thermal energy relates to the random movement of molecules; heat is energy transfer between bodies; temperature measures heat intensity.
Hydrogen bonds must be broken to absorb heat; more heat is required to raise water’s temperature, stabilizing climates.
Evaporative cooling: as water evaporates, the surface left behind cools because the most energetic molecules leave.
Ice vs liquid water:
Ice is less dense than liquid water due to hydrogen-bonded lattice in which molecules are spaced apart; this causes ice to float.
Water as solvent:
Solution: homogeneous mixture of solute dissolved in solvent.
Solvent: dissolving agent; Solute: substance dissolved.
Aqueous solution: water is the solvent.
Salt (NaCl) in water: dissolves to form Na⁺ and Cl⁻; water molecules orient their polar ends (negative ends attract Na⁺, positive ends attract Cl⁻).
Acids, bases, pH, and buffers
In water, a small fraction dissociates into ions: H⁺ and OH⁻.
Acid: donates hydrogen ions to solutions; base: reduces hydrogen ion concentration.
pH scale: describes how acidic or basic a solution is; ranges from 0 to 14; each unit represents a 10-fold change in [H⁺].
The pH scale is commonly summarized by pH = - ext{log}_{10} [ ext{H}^+].
Buffers: minimize changes in pH by accepting H⁺ when in excess or donating H⁺ when depleted.
Common pH examples (illustrative): battery acid, lemon juice, vinegar (acidic); pure water, human blood, seawater, milk of magnesia, household bleach (basic).
The CO2 problem: oceans, pH, and coral reefs
Rising atmospheric CO₂, a product of fossil fuel combustion, is increasing in the atmosphere and dissolves in oceans, lowering seawater pH and causing ocean acidification.
Experimental and observational studies with naturally varying pH show dire implications for coral reef health and the diversity of organisms they support.
Carbonate chemistry in seawater: increasing CO₂ bubbles lower pH and reduce carbonate ion availability; carbonate saturation state (e.g., CO₃²⁻) affects calcification rates in corals.
The figure (illustrative) shows calcification rate vs carbonate concentration; rising CO₂ reduces carbonate and lowers calcification rates in coral reefs.
Visual: a “champagne” reef with bubbles of CO₂ rising from a volcanic seep illustrates acidification effects.
Evolution connection: water and life beyond Earth
The search for extraterrestrial life centers on the search for water.
NASA has found evidence that water was once abundant on Mars, highlighting water’s central role in the potential for life elsewhere in the universe.
Quick recap of key concepts
Elements and compounds form the basis of all matter; most of life depends on a small set of elements (O, C, H, N) and trace elements.
Atoms comprise protons, neutrons, and electrons; chemical properties are governed by electron arrangement and valence shells.
Isotopes differ in neutron number; radioactive isotopes can be used as tracers but pose biological risks.
Covalent, ionic, and hydrogen bonds drive molecular interactions; water’s polarity enables exceptional solvent properties and diverse chemistry.
Water’s cohesive and thermal properties stabilize environments and enable life processes; pH, acids, bases, and buffers regulate biochemical reactions.
Human activities are altering global chemistry (e.g., ocean acidification), with cascading effects on ecosystems like coral reefs.
The study of these processes connects chemistry to biology, ecology, health, and planetary science.
Key formulas and diagrams to remember
Atomic number, mass number, and atomic mass:
Atomic number Z = number of protons
Mass number A = number of protons + neutrons = Z + neutrons
Atomic mass ≈ A (in atomic mass units)
Isotopes example (Carbon):
Carbon-12: Z = 6, neutrons = 6, A = 12
Carbon-13: Z = 6, neutrons = 7, A = 13
Carbon-14: Z = 6, neutrons = 8, A = 14
Covalent bond (example):
ext{H}_2: two H atoms share one pair of electrons
ext{O}_2: two O atoms share two pairs of electrons (double bond)
ext{CH}4 and ext{H}2 ext{O}: examples of covalently bonded compounds
Polar covalent vs nonpolar covalent bonds:
Polar: electrons shared unequally due to electronegativity differences; partial charges develop (e.g., in H₂O)
Nonpolar: electrons shared roughly equally
Ionic bond example: formation of NaCl from Na⁺ and Cl⁻ ions
pH concept: ext{pH} = - ext{log}_{10} [ ext{H}^+]
Dissolution of salts in water: NaCl → Na⁺ + Cl⁻
Reaction example (photosynthesis, simplified):
Carbon dioxide and water react to form glucose and oxygen under light energy: notated as a general chemical equation; a common representative is summarized by photosynthesis stoichiometry, though the exact balanced form depends on the organism.
Connections to real-world relevance
Nutritional labeling and fortified foods impact daily nutrient intake and health outcomes.
Food additives (trace elements) prevent disease and influence public health policies.
Ocean acidification is a global-change issue with local and global ecological consequences, including coral reef calcification and biodiversity.
Understanding chemical bonds and water chemistry is foundational for biochemistry, physiology, and environmental science.