Ch 2 Notes: The Chemical Basis of Life
The Chemical Basis of Life
Ch. 2 The Chemical Basis of Life (Chemistry puns and themes in the slides)
Learning Outcomes
Distinguish between elements and atoms.
State the four most abundant elements in the body.
Name the subatomic particles of an atom, their charge, and their location in an atom.
Compare and contrast types of chemical bonds.
Differentiate molecules and compounds.
Describe solubility and the process of dissociation.
Explain how enzymes work and describe the factors that can affect the rate of chemical reactions.
Describe inorganic and organic compounds.
Describe the pH scale and its relationship to acidic, basic, and neutral solutions.
Describe the structural organization and major functions of carbohydrates, lipids, proteins, and nucleic acids.
Basic Chemistry – Matter, Elements, and Atoms
Matter: anything that occupies space and has mass.
Element: the simplest type of matter with unique chemical properties; composed of atoms of only one kind.
Atom: smallest particle of an element that has chemical characteristics of that element.
Quick pun: "Never trust an atom. They make up everything." (mnemonic for atoms making up all matter)
Common Elements in the Body
Major elements by mass:
Oxygen ~65%
Carbon ~18.5%
Hydrogen ~9.5%
Nitrogen ~3.5%
Other chemical elements make up the remaining ~3.5%
Atomic Structure
Subatomic particles:
Protons: positive charge +
Neutrons: no electrical charge 0
Electrons: negative charge −
Nucleus: formed by protons and neutrons.
Most of the volume of an atom is occupied by electrons (electron cloud).
Atomic Number and Mass Number
Atomic Number (Z): equal to the number of protons in an atom, which equals the number of electrons in a neutral atom.
So Z = number of protons = number of electrons in a neutral atom.
Mass Number (A): A = number of protons + number of neutrons = p + n.
Molecules and Compounds
Molecules: two or more atoms chemically combined to form an independent unit.
Examples: H n2 (H2), H2O (water)
Compounds: substances composed of two or more different types of atoms chemically combined.
Examples: H2O (water), NaCl (sodium chloride)
Electrons and Chemical Bonding
Chemical bonds form when electrons in the outermost energy level (valence shell) are shared with or transferred to another atom.
Ionic Bonding: electrons are transferred from one atom to another.
Covalent Bonding: two or more atoms share electron pairs.
The number of electrons in the valence shell determines the chemical nature of the atom.
The octet rule: atoms will give up, accept, or share electrons in order to have 8 electrons in their outer shell.
Bond Type and Electronegativity
The type of chemical bond between two atoms is determined by their difference in electronegativity.
Bond types (in increasing difference):
Nonpolar covalent: equal sharing of electrons.
Polar covalent: unequal sharing of electrons.
Ionic: transfer of electrons, leading to charged ions.
Ionic Bonds
Ions: atoms that have gained or lost 1 or more electrons.
Cations: positively charged because they lost electrons.
Anions: negatively charged because they gained electrons.
In an ionic bond, electrons are transferred between atoms, creating oppositely charged ions that are attracted to each other.
Important Ions in the Human Body
Calcium (Ca^{2+}) – part of bones and teeth; functions in blood clotting, muscle contraction, release of neurotransmitters.
Sodium (Na^{+}) – membrane potentials, water balance.
Potassium (K^{+}) – membrane potentials.
Hydrogen (H^{+}) – acid-base balance.
Hydroxide (OH^{-}) – acid-base balance.
Chloride (Cl^{-}) – water balance.
Bicarbonate (HCO_{3}^{-}) – acid-base balance.
Ammonium (NH_{4}^{+}) – acid-base balance in some contexts.
Phosphate (PO_4^{3-}) – acid-base balance; part of bone matrix; energy exchange.
Iron (Fe^{2+}) – red blood cell formation.
Magnesium (Mg^{2+}) – necessary for enzymes.
Iodide (I^{-}) – present in thyroid hormones.
Covalent Bonds
Atoms share one or more pairs of electrons because the atoms have similar electronegativities.
Single covalent bond: two atoms share one pair of electrons.
Double covalent bond: two atoms share four electrons.
Nonpolar covalent: electrons shared equally because nuclei attract electrons equally.
Polar covalent: electrons not shared equally because one nucleus attracts the electrons more than the other.
Nonpolar Covalent Bond Example
Illustration: hydrogen atoms share electrons equally, resulting in a nonpolar covalent bond.
Polar Covalent Bond Example
Example illustrating polar covalent bond (partial charges on atoms, unequal electron sharing).
Intermolecular Forces
Forces between molecules; result from weak electrostatic attractions between oppositely charged parts of molecules, or between ions and molecules.
Weaker than bonds that form chemical bonds.
Determine properties such as solubility and dissociation.
Hydrogen Bonds
Type of intermolecular force: an H atom covalently bonded to O, N, or F in one molecule is attracted to O, N, or F in another molecule.
Intramolecular hydrogen bonds influence the shape of complex molecules (e.g., proteins, nucleic acids).
Water: Adhesion and Cohesion
Adhesion: attraction of water molecules to polar or charged surfaces.
Cohesion: attraction between water molecules.
Together, they enable water transport in vessels and contribute to capillary action.
Solubility and Dissociation
Solubility: the ability of one substance to dissolve in another.
Examples: sugar and salt dissolve in water; oils do not dissolve well.
Dissociation (in ionic compounds): separation into ions in solution.
Electrolytes and Nonelectrolytes
Electrolytes: solutions formed by the dissociation of cations (+) and anions (−) in water; conduct electrical current; essential for physiological processes.
Nonelectrolytes: substances that do not dissociate in solution; do not conduct electricity.
Chemical Reactions and Energy
Chemical reaction: atoms, ions, molecules, or compounds interact to form or break chemical bonds.
Bonds are formed (synthesis; anabolic) and broken (decomposition; catabolism) during reactions.
Metabolism: sum of all anabolic and catabolic reactions in the body.
Energy
Energy: the capacity to do work (e.g., move matter).
Conservation of energy: total energy in the universe remains constant; energy changes form but is not created or destroyed.
Energy and Chemical Reactions (ADP/ATP)
(a) ATP production from ADP + Pi + Energy (endergonic process):
(b) ATP hydrolysis to release energy:
In (a) the reaction stores energy in the high-energy phosphate bond; in (b) energy is released to power cellular work.
Fun fact – Heat Energy
When a chemical bond breaks and energy is released, not all energy is captured as ATP; some energy is released as heat.
Heat from this energy supports body temperature regulation in animals, including humans.
Solubility Quiz Question
Question: Solubility refers to the ability of one substance to in another.
Correct answer: C) dissolve
Inorganic Chemistry vs Organic Chemistry
Inorganic Chemistry: substances that do not contain carbon-hydrogen bonds.
Examples: water, oxygen, carbon dioxide, calcium phosphate, metal ions.
Organic Chemistry: study of carbon-containing substances.
Exceptions: CO, CO2, and HCO3^{-} lack C–H bonds but are discussed in this context.
The pH Scale
pH measures hydrogen ion concentration in solution.
Neutral solution: pH 7 (equal amounts of H^+ and OH^−).
Acidic solution: pH < 7 (higher H^+ concentration).
Alkaline/basic solution: pH > 7 (higher OH^− concentration).
Visual reference points on the pH scale (examples from the slide):
14: sodium hydroxide (NaOH)
~13–12: strong bases like oven cleaner, bleach
7: distilled water, blood ~7.35–7.45 (blood pH within the narrow physiologic range)
~6.5–6.0: saliva around 6.5
0: stomach acid (HCl)
Other common reference values include tomato (~4), vinegar (~2–3), lemon juice (~2)
Normal vs Abnormal Body pH
Normal blood pH: approximately 7.35 to 7.45.
Acidosis: pH drops below 7.35; nervous system becomes depressed; possible disorientation or coma.
Alkalosis: pH rises above 7.45; nervous system becomes overexcitable; possible nervousness or seizures.
Reversible Reactions and Equilibrium
In a reversible reaction, when the rate of product formation equals the rate of reactant formation, the system is at equilibrium.
Answer: B) at equilibrium.
Organic Chemistry – Functional Groups
Functional groups impart distinct properties to molecules.
Major organics and their functional groups include:
Hydroxyl group: -OH (Alcohols) — polar, hydrophilic; increases solubility in water.
Sulfhydryl group: -SH (Thiols) — polar; can form disulfide bonds (e.g., cysteine) that stabilize protein structure.
Carbonyl group: C=O (Ketones and Aldehydes) — polar, hydrophilic; ketones within carbon chain; aldehydes at the end of chains.
Carboxyl group: -COOH (Carboxyl) — acidic; donates H^+; at physiological pH often negatively charged; common in amino acids; contributes to energy-related reactions.
Ester group: -COOR (Esters) — less hydrophilic; found in fats (triglycerides) and other lipids; important for energy storage.
Phosphate group: -PO_4^{2-} (Phosphates) — highly hydrophilic; used in ATP, phospholipids, and signaling via phosphorylation.
Amino group: -NH_2 (Amines) — base; can accept H^+; common in amino acids; at physiological pH often positively charged.
Major Organic Molecules
Four major groups essential to living organisms:
Carbohydrates: composed of C, H, O; includes monosaccharides, disaccharides, polysaccharides; functions as energy source, structural support, and bulk for elimination; generally water-soluble.
Lipids: composed mainly of C, H, O; relatively insoluble in water; functions include protection, insulation, regulation, component of cell membranes, and energy storage.
Proteins: composed of C, H, O, N (and sometimes S); functions to regulate processes, aid transport, provide protection, enable muscle contraction, provide structure, and supply energy.
Nucleic Acids: composed of C, H, O, N, P; functions as the code for genes; examples include ATP, DNA, and RNA.
Carbohydrates – Structure and Examples
Examples and relationships:
Monosaccharides: glucose (C6H{12}O6), fructose (C6H{12}O6), galactose (C6H{12}O_6)
Disaccharide: sucrose
Structural isomers and stereoisomers exist among monosaccharides.
Glycogen represents a main storage polysaccharide in animals.
Structural diagrams show that carbohydrate molecules can form various isomeric forms (e.g., glucose, fructose, galactose) and branching structures (glycogen).
Lipids
Fats (Triglycerides): formed by glycerol + three fatty acids (via ester bonds and dehydration synthesis).
Functions: protection, insulation, energy storage.
Saturated vs Unsaturated Fatty Acids:
Saturated: all single bonds in carbon chain; tends to be solid at room temperature.
Unsaturated: one or more double bonds; tends to be liquid at room temperature.
Trans fats: artificially saturated-like unsaturated fats; highest cardiovascular risk.
Phospholipids: one fatty acid replaced with a phosphate group; amphipathic (polar head, nonpolar tail); major component of cell membranes.
Proteins
Amino acids: building blocks of proteins.
Peptide bonds: covalent bonds that link amino acids via dehydration synthesis.
Protein Structure:
Primary structure: amino acid sequence.
Secondary structure: hydrogen bonding between amino acids leading to alpha-helices and beta-pleated sheets.
Denaturation: loss of structure due to heat or pH changes, disrupting function.
Tertiary structure: large-scale folding driven by interactions with the surrounding aqueous environment; hydrophobic regions fold inward, hydrophilic regions interact with water; determines protein shape and function.
Quaternary structure: association of two or more polypeptide subunits.
Proteins as Enzymes:
Enzymes are proteins that lower the activation energy of reactions.
Active site is where substrates bind to enzymes and reactions occur.
Nucleic Acids – DNA and RNA
Nucleic acids (DNA and RNA) are polymers made of nucleotides.
Nucleotides consist of a five-carbon sugar (deoxyribose in DNA, ribose in RNA), a nitrogenous base, and a phosphate group.
DNA:
Genetic material copied from generation to generation.
Composed of two antiparallel strands of DNA nucleotides.
Sugar-phosphate backbones oriented in opposite directions.
Gene: a sequence of bases that codes for synthesis of RNA or protein.
Base sequence determines the primary structure of a protein.
DNA structure details:
Bases: Cytosine (C) pairs with Guanine (G); Adenine (A) pairs with Thymine (T).
Antiparallel strands with a 5' to 3' directionality.
ATP – Adenosine Triphosphate
ATP is the energy currency of the body; it stores energy and provides energy for cellular processes.
All energy-requiring chemical reactions stop when there is inadequate ATP.
Quick Recap of Key Concepts (Integrated Points)
Matter, elements, and atoms form the basis of all biological substances.
The four most abundant body elements are O, C, H, and N by mass, with others making up the remainder.
Valence electrons and the octet rule govern bond formation and compound stability.
Ionic vs covalent bonds influence molecule polarity, structure, and biological function.
Intermolecular forces (e.g., hydrogen bonds) influence solubility, structure, and biological interactions.
Water’s adhesive and cohesive properties enable transport and physiology in living organisms.
Solubility and dissociation underlie electrolyte balance and physiological processes.
Chemical reactions underpin metabolism, with energy captured in ATP and released via ATP hydrolysis.
The pH scale tracks hydrogen ion concentration; cells and blood require tight pH ranges for proper function.
Organic molecules are organized into carbohydrates, lipids, proteins, and nucleic acids, each with distinctive structures and functions.
Functional groups determine the chemical behavior and properties of organic molecules, affecting solubility, reactivity, and role in metabolism.
DNA structure, replication, and the genetic code hinge on base pairing and antiparallel strand arrangement.
Connections to Foundational Principles and Real-World Relevance
Structure determines function: the arrangement of atoms, bonds, and functional groups dictates molecular behavior in biology (e.g., enzyme active sites, protein folding, nucleic acid structure).
Energy flow and metabolism rely on ATP generation and utilization, with tight regulation by enzymes and cellular compartments.
Homeostasis depends on maintaining ionic balance, pH, and solubility of electrolytes; disruptions can impair nervous system function and overall physiology.
The distinction between inorganic and organic chemistry underpins understanding of physiology, pharmacology, and biochemistry (e.g., nutrient metabolism, lipid membranes, carbohydrate utilization).
Ethical, Philosophical, or Practical Implications
Understanding pH and electrolyte balance is critical in clinical settings (e.g., monitoring blood pH, treating dehydration, electrolyte disorders).
Nutritional choices (saturated vs unsaturated fats, trans fats) have real-world health implications for cardiovascular risk and metabolic health.
Enzymes as drug targets: many medications work by modulating enzyme activity or altering substrate access to active sites, highlighting the importance of biochemistry in medicine.
Key Equations and Notation (LaTeX)
Octet rule: atoms seek 8 electrons in the outer shell: outer-shell stability is achieved by achieving 8 valence electrons.
Atomic number and mass number:
where Z = proton count, N = neutron count, A = mass number.
ATP/ADP energetic relationships:
Base pairing in DNA:
C pairs with G; A pairs with T.
pH concepts (conceptual, not a derived equation here): neutral pH = 7; acidic < 7; basic > 7.
Note: This note consolidates all major and minor points from the transcript across the sections on basic chemistry, bonding, solubility, energy, acids and bases, organic vs inorganic chemistry, major biomolecules, and DNA/ATP physiology to provide a comprehensive study guide for exam preparation.