BIOL Chapter 2

The Chemical Basis of Life

I: Atoms, Molecules, and Water

Chapter 2

  • Organisms and Matter

    • Organisms are composed of matter.

    • Examples of elements:

    • Na: Sodium

    • Cl: Chlorine gas

  • Atoms

    • Definition: Atoms are the smallest functional units of matter that form all chemical substances.

    • Characteristics:

    • Cannot be further broken down into other substances by ordinary means.

    • Each specific type of atom is recognized as a chemical element.

Structure of Atoms

  • Subatomic Particles:

    • Protons:

    • Charge: Positive (+)

    • Location: Found in the nucleus

    • Neutrons:

    • Charge: Neutral (no charge)

    • Location: Found in the nucleus

    • Property: Number of neutrons can vary among isotopes.

    • Electrons:

    • Charge: Negative (−)

    • Location: Found in orbitals surrounding the nucleus

  • Atomic Structure Influences

    • The number of protons distinguishes one element from another.

    • An element's atomic number equals the number of protons and, in a neutral atom, also equals the number of electrons.

Electron Configuration and Orbitals

  • Orbitals:

    • Definition: Regions surrounding the nucleus where there is a high probability of finding an electron.

    • Types of orbitals:

    • s orbitals: Spherical shape

    • p orbitals: Dumbbell or propeller shaped

    • Each orbital can hold a maximum of 2 electrons.

Electron Shells

  • Electron Shells

    • Atoms with more electrons have orbitals that are further from the nucleus.

    • Shell Structure:

    • 1st Shell: 1 spherical orbital (1s) - holds 2 electrons.

    • 2nd Shell: 1 spherical orbital (2s) - holds 2 electrons; 3 dumbbell-shaped orbitals (2p) - holds 6 electrons (total of 8 electrons).

  • Energy Levels:

    • Electrons can move to higher or lower shells by absorbing or releasing energy.

    • The energy level of each shell increases with distance from the nucleus.

Example of Electron Configuration: Nitrogen Atom

  • Nitrogen Atom

    • Atomic Number: 7

    • Configuration:

    • 1st Shell: 2 electrons (2 in 1s orbital)

    • 2nd Shell: 5 electrons (2 in 2s orbital; 1 in each of the three 2p orbitals)

  • Valence Electrons

    • The outer shell (2nd shell) of nitrogen is not full, making the electrons available for bonding.

Periodic Table

  • Periodic Table Overview

    • Organized by atomic number.

    • Rows correspond to the number of electron shells (periods).

    • Columns indicate the number of valence electrons (groups).

  • Example Elements and Key Data

    • Atomic Mass: Represents the average mass of all isotopes.

    • Valence Electrons and Properties:

    • Elements in the same column have similar chemical bonding properties due to sharing the same number of valence electrons.

Characteristics of Subatomic Particles

  • Atomic Mass Comparison

    • Protons and neutrons: Nearly equal in mass, both more than 1,800 times the mass of an electron.

Mass Number

  • Mass Number

    • Definition: An element's mass number is the sum of protons and neutrons in the nucleus.

    • Calculation: If the atomic number and mass number are known, the number of neutrons can be determined.

Weight vs Mass

  • Weight

    • Defined by the gravitational pull on a given mass.

    • Example:

    • A man weighs 154 pounds on Earth but only 25 pounds on the Moon.

    • On a neutron star, the same person would weigh about 21 trillion pounds.

    • Mass: Remains constant regardless of location.

Units of Measurement

  • Dalton:

    • Unit of measurement for atomic mass, also called atomic mass unit (amu).

    • One Dalton equals 1/12 the mass of a carbon atom.

  • Mole:

    • Definition: 1 mole of any element contains the same number of atoms, known as Avogadro’s number (6.02imes10236.02 imes 10^{23}).

    • Example: 1 mole of NaCl = 58.442 g

Isotopes

  • Definition of Isotopes

    • All atoms of an element have the same number of protons but may differ in the number of neutrons.

    • An example of isotopes:

    • 12C^{12}C (carbon-12) vs. 14C^{14}C (carbon-14); same atomic number (6) but different atomic mass (12 and 14, respectively).

  • Unstable Isotopes:

    • 99% of carbon in nature is stable, while < 1% is unstable.

Radioactive Decay

  • Radioisotopes

    • Definition: Unstable nuclei that break down and emit energy and particles.

    • Decay example: 14C^{14}C decays into 14N^{14}N.

  • Applications of Radioisotopes

    • Used in dating fossils, tracing atoms in metabolism, diagnosing medical disorders.

    • Note: Radiation from decaying isotopes can damage cellular molecules.

    • Example: PET scans utilize Fluorodeoxyglucose (FDG).

Valence Electrons and Chemical Behavior

  • Valence Electrons

    • Definition: Electrons in the outermost shell that determine an atom's chemical behavior.

  • Octet Rule

    • Atoms with complete outer shells (8 electrons) are stable. Exceptions include hydrogen, which only needs 2 electrons to fill its outer shell.

    • Reactivity in atoms is due to unpaired electrons in the valence shell.

Chemical Bonds and Molecules

  • Definition of Molecule

    • Molecules consist of two or more atoms bonded together.

    • Molecular Formula

    • Contains the chemical symbols of the elements in the molecule (e.g., C<em>6H</em>12O6C<em>6H</em>{12}O_6).

    • Subscript indicates the number of each type of atom present (e.g., H2OH_2O has 2 hydrogens and 1 oxygen).

    • Compounds

    • Any molecule composed of two or more elements (e.g., H2OH_2O).

  • Types of Chemical Bonds

    • Covalent

    • Ionic

    • Hydrogen

    • Weak interactions:

    • Van der Waals forces

Covalent Bonds

  • Definition

    • A covalent bond is formed by the sharing of a pair of valence electrons between two atoms.

    • Covalent bonds are strong because shared electrons behave as if they belong to each atom.

  • Types of Covalent Bonds

    • Polar covalent bonds: Unequal sharing of electrons (e.g., O-H).

    • Nonpolar covalent bonds: Equal sharing of electrons (e.g., C-C).

  • Indications of Bonds

    • Single bond: sharing one pair of electrons (e.g., C–H).

    • Double bond: sharing two pairs of electrons (e.g., O=O).

    • Triple bond: sharing three pairs of electrons (e.g., N≡N).

Electronegativity


  • Definition

    • Electronegativity: An atom's ability to attract electrons.


  • Influence on Bonding

    • Atoms in a molecule will pull electrons toward themselves with differing strengths.


  • Electronegativity of Key Elements in Living Organisms:

    Element

    % of Body Mass

    Electronegativity


    Oxygen

    65%

    3.44


    Carbon

    18%

    2.55


    Hydrogen

    9%

    2.20


    Nitrogen

    3%

    3.04

    Polar vs. Nonpolar Covalent Bonds

    • Nonpolar Covalent Bonds:

      • Formed between atoms with similar electronegativities; electrons are shared equally.

      • Examples:

      • C (2.55) – C (2.55): difference = 0.00

      • C (2.55) – H (2.20): difference = 0.35

    • Polar Covalent Bonds:

      • Formed between atoms with different electronegativities; unequal sharing leads to a difference in charge across the molecule.

      • Example: O (3.44) – H (2.20): difference = 1.24

    Ionic Bonds

    • Definition

      • An ion is an atom or molecule that has gained or lost one or more electrons.

      • Cations have a net positive charge (+), and anions have a net negative charge (−).

      • An ionic bond occurs when an electrostatic attraction binds a cation to an anion.

      • Ionic compounds are referred to as salts (e.g., NaCl).

    • Ionic Bonding in Table Salt (NaCl):

      • Sodium (Na) and Chlorine (Cl) form the ionic compound NaCl.

      • Electronegativity comparison:

      • Na = 0.93

      • Cl = 3.16

      • Difference = 2.23 (indicating an ionic bond).

    Weak Chemical Interactions

    • Overview

      • Weak chemical bonds are significant in biological molecules, maintaining their functional form.

      • Types of weak interactions include:

      • Ionic bonds

      • Hydrogen bonds

      • Van der Waals interactions

    • Hydrogen Bonds

      • Formed when a hydrogen atom from one polar molecule is attracted to an electronegative atom of another molecule (commonly O or N).

      • Represented as dashed or dotted lines.

      • Collectively form strong bonds (e.g., holding DNA strands together).

    Van der Waals Forces

    • Definition

      • Occur due to asymmetrical distribution of electrons in molecules resulting in regions of positive or negative charges, facilitating attraction between atoms.

    • Characteristics

      • Effective only when atoms and molecules are very close together.

      • Collectively can create strong interactions, as seen in gecko toe hairs sticking to surfaces.

    Molecular Shape and Function

    • Molecular Binding

      • Biological molecules may bind temporarily through weak interactions if their shapes are complementary.

      • Similar shapes can lead to similar biological effects.

      • The specific shape arises from the arrangement of bonds and the angles formed between atoms.

    Chemical Reactions

    • Definition

      • Chemical reactions involve the transformation of one or more substances into different substances.

      • Reactants and Products:

      • Reactants: Starting molecules; Products: Final molecules in a reaction.

    • Properties of Chemical Reactions:

      • Require energy sources and often enzymes as catalysts to speed up reaction rates.

      • Tend to proceed in a particular direction until reaching equilibrium.

      • Typically occur in liquid environments, such as water.

    Water: An Important Molecule

    • Properties of Water

      • Essential for life; most organisms consist primarily of water and rely on a water-dominated environment.

    Water as the Solvent of Life

    • Solution: A homogeneous mixture of substances.

    • Solvent: The dissolving agent in a solution; in biological contexts, this is typically water.

    • Solute: The substance that is dissolved.

    • Aqueous Solution: A solution where water is the solvent.

      • When ionic compounds dissolve in water, ions are surrounded by water molecules, forming a hydration shell.

    Solubility of Different Compounds in Water

    • Dissolution of Nonionic Polar Molecules:

      • Water can also dissolve nonionic polar molecules, including large polar molecules like proteins if they have ionic and polar regions.

    Hydrophilic and Hydrophobic Substances

    • Hydrophilic Substances:

      • Have an affinity for water (e.g., sugars, amino acids).

    • Hydrophobic Substances:

      • Lack affinity for water (e.g., oils, fats).

    • Amphipathic Molecules:

      • Possess both polar/ionized regions and nonpolar regions; can form micelles in water.

    Concentration in Aqueous Solutions

    • Concentration: Amount of solute dissolved in a unit volume of solution.

      • Example: 1 gram of NaCl dissolved in 1 liter of water results in a concentration of 1 g/L.

    • Molarity: Number of moles of solute per liter of solution.

      • Definition: M=racextmolesofsoluteextlitersofsolutionM = rac{{ ext{{moles of solute}}}}{{ ext{{liters of solution}}}}

      • Example: 1 mole of glucose = 180 g.

    States of Water

    • Three States of Water:

      • Solid (Ice), Liquid (Water), Gas (Water Vapor).

      • Changes in state (solid, liquid, gas) involve an input or release of energy.

    • Heat of Vaporization: Energy required to convert water to vapor.

    • Heat of Fusion: Energy required to melt ice.

    • Specific Heat: Amount of heat energy needed to raise the temperature of a substance.

      • High specific heat results in water's stability as a liquid.

    Colligative Properties of Water

    • Definition: Properties that depend only on the total number of dissolved solute particles.

    • Effects of Dissolved Solutes:

      • Lower freezing point and raise boiling point of water.

      • Example: Antifreeze (ethylene glycol) lowers the freezing point preventing ice formation.

    Role of Water in Biological Systems

    • Functions:

      • Solvent

      • Support

      • Cooling

      • Removal of toxic waste

      • Surface tension and lubrication

      • Participation in chemical reactions

    Ionization of Water

    • Definition: Pure water can ionize to a small extent into hydrogen ions (H+) and hydroxide ions (OH-).

      • Under homeostatic conditions: [H+][OH−]=[10−7M][10−7M]=10−14M[H^+][OH^-] = [10^{-7} M][10^{-7} M] = 10^{-14} M

    Acids and Bases

    • Definitions:

      • Acids release hydrogen ions in solution; strong acids release more H+ than weak acids.

      • Bases lower H+ concentration; can release OH- or bind H+.

    The pH Scale

    • Definition: The pH of a solution is defined as the negative logarithm of H+ concentration ( ext{pH} = - ext{log}_{10}[H^+] ).

      • Acidic solutions: pH < 7; Basic solutions: pH > 7.

      • Most biological fluids have pH values between 6 and 8.

    Effects of pH

    • Changes in pH can affect:

      • Shapes and functions of molecules.

      • Rates of chemical reactions.

      • Binding capacity of molecules.

      • Ability of ions/molecules to dissolve in water.

    Importance of pH Regulation

    • Organisms usually tolerate only small pH changes.

      • Buffers: Systems that help maintain constant pH.

      • Acid-base buffer systems can shift to generate or release H+ ions to adjust for pH changes.

      • Example: Carbonic acid and bicarbonate ion stabilize pH in human blood.