Chapter 2: The Chemistry of Life

Matter, Energy, and Life

  • Matter: Anything that has mass and occupies space.
  • Energy: The ability to do work.
    • Potential Energy: Stored energy, available to do work.
    • Kinetic Energy: Energy of motion.
    • Potential energy can be converted to kinetic energy to do work.

Forms of Energy

  • Mechanical Energy: Energy of movement.
  • Nuclear Energy: Energy from reactions involving atomic nuclei.
  • Electrical Energy: Flow of charged particles.
  • Radiant Energy: Energy in heat, light, x-rays, and microwaves.
  • Chemical Energy: Energy in chemical bonds.

The Nature of Matter

  • Matter: Anything that takes up space and has mass.
    • Can exist as a solid, liquid, or a gas.
  • Element: Substance that cannot be broken down into another substance by ordinary chemical means.
    • 6 elements make up about 98% of the body weight of most living organisms: Carbon, Hydrogen, Nitrogen, Oxygen, Phosphorus, Sulfur.

Structure of Atoms

  • Atomic Symbol: Name of the atom or element (e.g., H for hydrogen, Na for sodium).
  • Subatomic Particles:
    • Neutrons: No electrical charge, found in the nucleus.
    • Protons: Positive charge, found in the nucleus.
    • Electrons: Negative charge, outside the nucleus.
  • Atomic Mass: Equal to the sum of protons + neutrons.
  • Atomic Number: Equal to the number of electrons and protons.

Periodic Table of the Elements

  • Brief overview of the organization of elements.
  • Key groups: Alkali Metals, Alkaline Earth Metals, Transition Metals, Other Metals, Nonmetals, Noble Gases, Inner Transition Metals.
  • States of matter: Gaseous, Liquid, Solid, Synthetically Prepared.

Isotopes

  • Isotopes: Atoms with the same number of protons but a different number of neutrons.
  • Unstable isotopes may decay, emitting radiation (radioactive isotopes).
  • Radioactive isotopes behave essentially the same as stable isotopes of the same element.
    • Can be used as tracers (e.g., PET scan).
    • Can cause damage to cells, potentially leading to cancer.

Arrangements of Electrons

  • Electrons are constantly moving.
  • Models of atoms show energy levels or electron shells.
  • Chemical properties of atoms are largely determined by the arrangement of their electrons.

Electron Shells

  • For atoms up through number 20:
    • 2 electrons fill the first shell.
    • 8 electrons fill each additional shell.
  • Octet Rule for Valence Shell:
    • Valence shell: Outermost shell.
    • Atoms with more than 2 shells are most stable with 8 electrons in the outer shell.
    • Atoms can give up, accept, or share electrons to achieve 8 electrons in their valence shell.
  • Construction of an “energy shell diagram”.

Energy Shell Diagram Examples

  • Phosphorus (P), Oxygen (O), Nitrogen (N), Carbon (C), Hydrogen (H), Sulfur (S) configurations are shown.
  • Examples:
    • Hydrogen: 1H^1
    • Carbon: 6C^{12}
    • Nitrogen: 7N^{14}
    • Oxygen: 8O^{16}
    • Phosphorus: 15P^{31}
    • Sulfur: 16S^{32}

Types of Chemical Bonds

  • Two main types: ionic and covalent bonds.
  • Molecule: Group of atoms bonded together (e.g., O2, H2O, C6H{12}O6, N2).
  • Compound: Molecule containing two or more different elements (e.g., H2O, C6H{12}O6).

Covalent Bonds

  • Formed by sharing of electrons.
  • Structural Formula: Uses straight lines to represent shared electrons (e.g., H-H).
    • 1 line indicates 1 pair of shared electrons.
  • Molecular Formula: Shows the number of atoms involved (e.g., H_2 for hydrogen gas).

Types of Covalent Bonds

  • Double Covalent Bond: sharing 2 pairs of electrons (e.g., Oxygen gas O_2 or O=O).
  • Triple Covalent Bond: sharing 3 pairs of electrons (e.g., Nitrogen gas N_2 or N≡N).
  • Single atom may form bonds with more than one atom (e.g., Oxygen gas O_2).

Ions and Ionic Compounds

  • Ions: Charged atoms (+/-) formed by gaining or losing electrons from the outer shell.
    • Sodium (Na) has 1 more proton than electrons (Na+).
    • Chlorine (Cl) has 1 more electron than protons (Cl-).
  • Ionic compounds are often called salts.

Ionic Bond Formation

  • Forms when 2 atoms are held together by the attraction between opposite charges (+/-).
    • Sodium (Na) has 1 electron in its outer shell and usually gives up an electron.
    • Chlorine (Cl) has 7 electrons in its outer shell and usually accepts an electron from another atom.

Water: The Molecule of Life

  • Single most important molecule on Earth.
  • Organisms are 70-90% water.
  • Unique properties allow it to support life.

Structure and Polarity of Water

  • Polar Covalent Bond: Unequal sharing of electrons.
    • Electrons spend more time around the oxygen nucleus than hydrogen nuclei.
  • Hydrogen Bond: Slightly positive hydrogen of one water molecule attracted to slightly negative oxygen in another water molecule.

Properties of Water Supporting Life

  • Solvency: Other materials easily dissolved.
  • High Surface Tension.
  • High Heat Capacity.
  • High Heat of Vaporization (water into a gas).
  • Varying Density (solid, liquid, and gas).

Water as a Solvent

  • Due to polarity and H-bonding, water dissolves many substances.
    • Hydrophilic: Molecules attracted to water.
    • Hydrophobic: Molecules not attracted to water.
  • Water causes NaCl to dissociate into Na^+ and Cl^-.

Cohesion and Adhesion of Water

  • Cohesion: Ability of water molecules to cling to each other due to hydrogen bonding.
  • Adhesion: Ability of water molecules to cling to other polar surfaces.
  • Allows water to be an excellent transport system in and outside of living organisms.
    • Water evaporates, pulling the water column from the roots to the leaves.
    • Water molecules cling together and adhere to sides of vessels in stems and tree trunks.
    • Water enters a plant at root cells.

Surface Tension of Water

  • Water molecules at the surface cling more tightly to each other than to the air above.
  • Mainly due to hydrogen bonding.

Heat Capacity of Water

  • High Heat Capacity: The many hydrogen bonds linking water molecules allow water to absorb heat without greatly changing its temperature.
  • Temperature of water rises and falls slowly.

Heat of Vaporization of Water

  • High Heat of Vaporization: Takes a great deal of energy to break H bonds for evaporation.
  • Heat is dispelled as water evaporates.

Density of Ice

  • Water is less dense than ice.
  • Unlike other substances, water expands as it freezes.
  • Ice floats rather than sinks, making life possible in water.
  • Ice acts as an insulator during the winter months.

Chemical Reactions

  • Reactants: Molecules that participate in the reaction (shown to the left of the arrow).
  • Products: Molecules formed by reactions (shown to the right of the arrow).
  • Equation is balanced if the same number of each type of atom occurs on both sides of the arrow.

Five Important Chemical Reactions in Biology

  • Oxidation-reduction
  • Dehydration synthesis
  • Hydrolysis
  • Phosphorylation
  • Acid-base reactions

Oxidation-Reduction Reactions

  • Oxidation-reduction reactions: Reactions in which electrons (and their energy) are transferred from one atom to another.
  • Oxidation: An atom loses an electron.
  • Reduction: An atom gains an electron.
  • For oxidation to occur, reduction must also occur.
  • Example: Respiration
    • C6H{12}O6 + 6O2 → 6H2O + 6CO2 + Energy
    • Sugar + oxygen → water + carbon dioxide + energy

Dehydration Synthesis Reactions

  • When two small molecules are joined to form a larger molecule, a molecule of water is released.
  • Example: Joining amino acids to form proteins.
    • NH2CH2CO-OH + H-NHCH2CO-OH → NH2CH2CO-NHCH2CO-OH + H_2O
    • amino acid 1 + amino acid 2 = protein + water

Hydrolysis Reactions

  • When a larger molecule is broken down into smaller parts, a water molecule is split.
  • Opposite of a dehydration synthesis.
  • Example: Digesting proteins into amino acids.
    • NH2CH2CO-NH-CH2CO-OH + H2O → NH2CH2CO-OH + H-NHCH_2CO-OH
    • Protein + water = amino acid 1 + amino acid 2

Phosphorylation Reactions

  • When phosphate groups are added to other molecules, phosphate groups are clusters of oxygen and phosphate atoms.
  • Bonds between phosphate groups and other molecules contain high potential energy.
  • When these bonds are broken, the energy that is released can be used by the cell to do work.
  • Phosphorylation reactions are commonly used to transfer potential energy.

Acid-Base Reactions

  • Occurs when ions from an acid interact with ions from a base.
  • This type of reaction allows harmful acids and bases to neutralize one another.
    • H^+Cl^- + Na^+OH^- → Na^+Cl^- + H^+OH^-
    • Hydrochloric acid + Sodium hydroxide = Sodium chloride + Water

Acids and Bases

  • Water dissociates into an equal number of hydrogen ions (H^+) and hydroxide ions (OH^-), meaning it’s neutral.
    • H_2O → H^+ + OH^-

Acids

  • Substances that dissociate in water, releasing H^+ ions.
  • Common examples: lemon juice, vinegar, tomatoes, and coffee.
  • Example:
    • HCl → H^+ + Cl^-
    • Hydrochloric acid

Bases

  • Substances that either take up hydrogen ions or release hydroxide ions (OH^-).
  • Common bases: ammonia and milk of magnesia.
  • Example:
    • NaOH → Na^+ + OH^-
    • Sodium hydroxide

pH Scale

  • Mathematical way to indicate the number of hydrogen ions in solution.
  • pH scale ranges from 0 to 14.
    • pH below 7: acidic – more [H^+] than [OH^-]
    • pH above 7: basic – more [OH^-] than [H^+]
    • pH of 7: neutral – [H^+] equal to [OH^-]

Buffers

  • Chemical or combination of chemicals that keeps pH within normal limits.
  • Resist pH change by taking up excess H^+ or OH^-.
  • pH of blood is about 7.4 – maintained by a buffer.