The Chemistry of Life

Chapter 2: The Chemistry of Life

Module 2.1: Atoms and Elements

A. Definition of Matter

  • Matter: Defined as anything that has mass and occupies space.

  • States of Matter:

    • Solid

    • Liquid

    • Gas

  • Chemistry: The study of matter and its interactions.

B. Atoms and Atomic Structure

  1. Definition of Atom

    • An atom is defined as the smallest unit of matter that retains its original properties, composed of smaller structures known as subatomic particles.

  2. Subatomic Particles

    • Three basic forms:

      • Protons (p+): Found in the atomic nucleus.

      • Charge: Positively charged.

      • Neutrons (n0): Also in the atomic nucleus, slightly larger than protons.

      • Charge: No charge.

      • Electrons (e-): Located outside the atomic nucleus.

      • Charge: Negatively charged.

    • Atoms are electrically neutral; the number of protons equals the number of electrons, cancelling their charges. Neutrons may vary in number among atoms.

  3. Electron Shells

    • Regions surrounding the atomic nucleus indicating the probability of finding electrons.

      • First shell holds 2 electrons.

      • Second shell holds 8 electrons.

      • Third shell can hold 18 electrons but is satisfied with 8.

      • Some atoms may possess more than three shells.

C. Elements in the Periodic Table and the Human Body

  1. The atomic number indicates the number of protons in an atom’s nucleus.

  2. Element Definition: A substance that cannot be broken down into simpler substances by chemical means.

    • Each element consists of atoms with the same number of protons.

  3. Periodic Table:

    • Elements are organized by increasing atomic numbers.

    • Each element has a chemical symbol.

  4. Major Elements in the Human Body:

    • Hydrogen, Oxygen, Carbon, Nitrogen.

    • Seven mineral elements and thirteen trace elements also contribute to bodily composition.

D. Isotopes and Radioactivity

  1. Mass Number: Equal to the sum of protons and neutrons in the nucleus.

  2. Isotope: An atom with the same atomic number but different mass numbers due to varying neutrons.

  3. Radioisotopes: Unstable isotopes that release energy/radiation through radioactive decay, resulting in stable forms.

Module 2.2: Matter Combined: Mixtures and Chemical Bonds

A. Mixtures

  • Definition: A mixture is a physical combination of two or more elements without altering their chemical nature.

B. Types of Mixtures

  1. Suspension: Contains large, unevenly distributed particles that settle out when undisturbed.

  2. Colloids: Components have small, evenly distributed particles that do not settle out.

  3. Solutions: Composed of very small, evenly distributed particles that do not settle, containing a solute dissolved in a solvent.

    • Solute: The substance dissolved.

    • Solvent: The substance that dissolves the solute.

C. Chemical Bonds

  • Definition: Matter combines chemically to form a molecule through chemical bonding, which is an energy relationship between atoms.

  1. Definition of Molecule: Formed by chemical bonding of two or more atoms of the same element.

  2. Definition of Compound: Formed when two or more different element atoms bond chemically.

  3. Macromolecules: Very large molecules composed of numerous atoms.

  4. Molecular Formulas: Represent molecules using letters and numbers indicating the kinds and numbers of atoms.

D. Valence Electrons and Chemical Bonds

  1. Valence Shell: The outermost electron shell where interactions occur.

  2. Valence Electrons: Determine an atom's interactions and bonding tendencies.

  3. Octet Rule: An atom is most stable with eight electrons in its valence shell.

  4. Duet Rule: Applicable to atoms with five or fewer electrons; stability occurs with two electrons in the first shell.

E. Ions and Ionic Bonds

  • An ionic bond forms when electrons transfer from metal to nonmetal atoms, creating charged ions (cations and anions) held by attractive forces.

  1. Cation: A positively charged ion forming when a metal loses electrons.

  2. Anion: A negatively charged ion forming when a nonmetal gains electrons.

F. Covalent Bonds

  • The strongest bond, formed when nonmetals share electrons.

  1. Atoms can share one (single bond), two (double bond), or three (triple bond) electron pairs.

  2. Electronegativity: An element's ability to attract electrons; increases from bottom left to upper right of the periodic table, making fluorine (F) the most electronegative.

  3. Nonpolar Covalent Bonds: Formed when nonmetals with similar electronegativities share electrons equally.

  4. Polar Covalent Bonds: Formed between nonmetals with differing electronegativities, resulting in unequal electron sharing, creating dipoles with partial charges.

  5. Hydrogen Bonds: Weak attractions between dipoles influencing properties like water's surface tension.

Module 2.3: Chemical Reactions

A. Chemical Notation

  • Symbols and abbreviations representing reaction occurrences; consists of:

  1. Reactants: Ingredients undergoing reactions, found on the left of the equation.

  2. Products: Results of reactants bonding chemically, found on the right side of the equation.

B. Energy and Chemical Reactions

  1. Energy: Capacity to do work or fuel chemical reactions.

  2. Forms of Energy:

    • Potential Energy: Stored energy.

    • Kinetic Energy: Energy in motion.

    • Chemical, Electrical, Mechanical energy exist in the body, either as potential or kinetic.

  3. Reactions involving energy:

    • Endergonic Reactions: Require energy input, products have more energy than reactants.

    • Exergonic Reactions: Release energy, products have less energy than reactants.

C. Homeostasis and Chemical Reactions

  • Processes to maintain homeostasis: 1) breaking down molecules, 2) energy conversion, 3) building new molecules.

  1. Catabolic Reactions: Large substances are broken down into smaller ones (e.g., AB → A+B), usually exergonic.

  2. Exchange Reactions: Atoms in reactants are exchanged (e.g., AB + CD → AD + BC). Includes Redox Reactions: Electrons and energy are exchanged, with oxidation (losing electrons) and reduction (gaining electrons).

  3. Anabolic Reactions: Combine small units into larger substances (e.g., A + B → AB), and are typically endergonic.

D. Reaction Rates and Enzymes

  • Activation Energy (Ea): Required energy for reactions, affected by multiple factors:

  1. Concentration: Increased concentration leads to a higher likelihood of collisions.

  2. Temperature: Increased temperature raises kinetic energy and collision effectiveness.

  3. Particle Size and Phase: Smaller particles have more energy than larger ones; gaseous phase particles generally have higher kinetic energy.

  4. Catalysts: Substances that increase reaction rates by lowering activation energy, include biological catalysts called enzymes.

  5. Enzyme Properties:

    • Speed up reactions by reducing activation energy.

    • Highly specific to substrates.

    • Do not alter reactants/products permanently.

  6. Induced-Fit Mechanism: Enzyme-substrate interactions induce shape changes that lower activation energy.

Module 2.4: Inorganic Compounds

A. Types of Compounds

  1. Inorganic Compounds: Do not contain carbon-hydrogen bonds; include water, acids, bases, and salts.

  2. Organic Compounds: Contain carbon-hydrogen bonds.

B. Water

  • Composes 60-80% of human body mass, possessing several vital properties:

  1. High heat capacity: Absorbs heat without a significant temperature change.

  2. Carries heat during evaporation (liquid to gas).

  3. Cushions and protects structures due to high density.

  4. Acts as a lubricant.

  5. Universal Solvent: Due to polarity, dissolves many solutes (hydrophilic substances) while hydrophobic substances do not dissolve.

C. Acids and Bases

  1. Acid: Hydrogen ion donor, increases H+ ions in water.

  2. Base: Hydrogen ion acceptor, decreases H+ ions in water.

  3. pH Scale: Ranges from 0-14 indicating hydrogen ion concentration.

    • pH 7: Neutral (H+ = Base ions)

    • pH < 7: Acidic (H+ > Base ions)

    • pH > 7: Basic (Base ions > H+)

    • Most body fluids are slightly basic (blood pH 7.35–7.45).

  4. Buffer: A chemical system that resists pH changes, critical for maintaining blood homeostasis.

D. Salts and Electrolytes

  • A salt is formed from metal cations and nonmetal anions via ionic bonds; salts dissolve in water to create electrolytes, which conduct electrical currents.

Module 2.5: Organic Compounds

A. Monomers and Polymers

  1. Dehydration Synthesis: An anabolic reaction linking monomers to form polymers with water byproduct.

  2. Hydrolysis: A catabolic reaction utilizing water to break polymers back into monomers.

B. Carbohydrates

  • Composed of carbon, hydrogen, and oxygen; primarily serve as fuel.

  1. Monosaccharides: Simple sugars with 3 to 7 carbons. Examples include glucose, fructose, galactose.

  2. Disaccharides: Formed by linking two monosaccharides through dehydration synthesis.

  3. Polysaccharides: Long chains of monosaccharides linked via dehydration synthesis.

    • Glycogen: Storage polymer for glucose in muscle/liver cells.

    • Glycoproteins and glycolipids have varied functions when polysaccharides bond with proteins/lipids.

C. Lipids

  • Nonpolar, hydrophobic molecules made of carbon and hydrogen, including fats and oils.

  1. Fatty Acids: Basic lipid monomers (4-20 carbons).

    • Saturated Fatty Acids: Solid at room temperature, no double bonds.

    • Monounsaturated Fatty Acids: One double bond, generally liquid.

    • Polyunsaturated Fatty Acids: Two or more double bonds, liquid.

  2. Triglycerides: Three fatty acids linked to glycerol (storage polymer of fats).

  3. Phospholipids: Composed of glycerol, two fatty acid tails, and a phosphate group. Amphiphilic nature is crucial for cell membrane structure.

  4. Steroids: Nonpolar, four-ring structures where cholesterol is the base for other steroids.

D. Proteins

  • Macromolecules serving various roles, composed of amino acid monomers linked by peptide bonds.

  1. Amino Acids: Twenty types link to form polypeptides.

  2. Peptides: Formed by linking two or more amino acids via peptide bonds.

    • Dipeptides (2 amino acids), Tripeptides (3 amino acids), Polypeptides (10+ amino acids).

  3. Protein Structure:

    • Fibrous Proteins: Long strands that add strength/durability.

    • Globular Proteins: Spherical proteins functioning as enzymes/hormones.

  4. Protein Structure Levels:

    • Primary Structure: Amino acid sequence.

    • Secondary Structure: Folded segments shaped through hydrogen bonding (alpha helix and beta-pleated sheet).

    • Tertiary Structure: 3D shape from twists and folds stabilized by hydrogen bonds.

    • Quaternary Structure: Arrangement of multiple polypeptide chains.

  5. Protein Denaturation: Disruption of a protein's shape due to heat, pH changes, or chemicals, causing loss of function.

E. Nucleotides and Nucleic Acids

  1. Nucleotide Structure: Composed of a nitrogenous base, pentose sugar, and phosphate group.

  2. Nitrogenous Bases: Purines (adenine, guanine) and pyrimidines (cytosine, uracil, thymine).

  3. ATP (Adenosine Triphosphate): Main energy source in the body, forms from ADP and phosphate using energy from glucose oxidation.

  4. DNA and RNA:

    • DNA: Double helix structure with deoxyribose, consisting of A, G, C, T; base pairing (A=T, G≡C).

    • RNA: Single-stranded with ribose, contains uracil (A=U); involved in protein synthesis (transcription and translation).

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