Molecules of Life Vocabulary

Ch. 2 The Molecules of Life

What Are Elements?

  • Definition: A pure substance consisting of one type of atom.
  • Atoms are composed of:
    • Neutrons.
    • Protons.
    • Electrons.

Atoms

  • Diagram shows a carbon atom with:
    • 6 protons.
    • 6 neutrons in the nucleus.
    • Electrons orbiting the nucleus.

Atomic Number and Atomic Mass

  • Elements differ in the number of subatomic particles.
  • Atomic number = number of protons.
  • Mass number = protons + neutrons.
  • Most atoms are neutral: #protons = #neutrons.

Essential Elements of Life

  • Carbon, hydrogen, oxygen, and nitrogen make up 96% of living matter.
  • Other essential elements include:
    • Oxygen.
    • Carbon.
    • Hydrogen.
    • Nitrogen.
    • Calcium.
    • Phosphorous.
    • Potassium.
    • Sulfur.
    • Sodium.
    • Chlorine.
    • Magnesium.
    • Iron, cobalt, copper, zinc, iodine, silicon (varies).

Electron Configuration and Chemical Properties

  • The outer electron shell determines chemical properties like reactivity.
  • Electron orbitals (clouds) have definite shapes.
  • Shapes of electron orbitals define the shapes of molecules.

Reactivity of Atoms

  • Electrons are added across a row until the outer shell contains its complete complement of eight electrons.

Chemical Bonds

  • Atoms share or transfer electrons with certain other atoms.
  • Atoms sharing electrons stay close, connected by chemical bonds.
  • Atoms combine to form molecules.

Types of Chemical Bonds

  • Covalent bonds
    • Polar.
    • Non-polar.
  • Hydrogen bonds.
  • Ionic bonds.
  • Differ in strength and use.

Covalent Bonds

  • Electron sharing.
  • There is one electron in the colored area representing an orbital.
  • Represented by a line between two element names.
  • Classified as polar or non-polar.
  • Sometimes can share two electron pairs (double bond), designated by a double line.

Hydrogen Gas Example

  • Chemical formula: H_2
  • Structural formula: H-H
  • Two atoms of hydrogen combine to form hydrogen gas by sharing electrons in a molecular orbital.

Molecule Stability

  • Most stable with 8 electrons in the outer shell.
  • Periodic table can be used to predict molecules.

Polar Covalent Bonds

  • Unequal sharing of electrons leads to partial charges.
  • Example: Water (H_2O)
    • The bonds linking hydrogen and oxygen atoms are polar.
    • Partial positive charge (\delta+) near the hydrogen atoms.
    • Partial negative charge (\delta-) near the oxygen atom.

Nonpolar Covalent Bonds

  • Equal sharing of electrons.
  • Examples:
    • Hydrogen gas (H_2).
    • Methane (CH_4).

Ionic Bond

  • Forms between 2 charged atoms.
  • Ion: charged atom or molecule.
  • Ionic bond = electron theft.

Sodium Chloride (NaCl) Example

  • Sodium loses an electron and becomes positively charged (Na^+).
  • Chlorine gains an electron and becomes negatively charged (Cl^-).
  • The two ions are attracted to each other.
  • Sodium chloride (NaCl) dissolves in water because the sodium (Na^+) ions and chloride (Cl^-) ions each become surrounded by water molecules.

Hydrogen Bonds

  • A hydrogen atom already in a covalent bond interacts with the electronegative atom of another molecule.
  • Hydrogen bonds form when the partial positive charge of hydrogen atoms are attracted to the partial negative charge of oxygen atoms.

Importance of Weak Chemical Bonds

  • Strength in numbers!
  • Reinforce shapes of large molecules and help molecules adhere to each other.
  • Collectively, such interactions can be strong, as between molecules of a gecko’s toe hairs and a wall surface.

Chemical Reactions

  • Involve the breaking and forming of chemical bonds.
  • Example: Formation of water
    • 2H2 + O2 \rightarrow 2H_2O
    • Reactants: Hydrogen gas and Oxygen.
    • Products: Water.

Water

  • Universal solvent.
  • Polar.
  • pH 7.0.

Unusual Properties of Water

  • Cohesion.
  • Moderation of temperature.
  • Expansion upon freezing.
  • Versatile solvent.

Cohesion

  • Water molecules ‘hold hands’ and stay together.
  • Helps transport water against gravity in plants.
  • Related to surface tension.

Moderation of Temperature

  • Water can absorb or release a large amount of heat with only a slight change in its own temperature.
  • High specific heat minimizes temperature fluctuations within limits that permit life.
  • Heat is absorbed when hydrogen bonds break; heat is released when hydrogen bonds form.

Water as a Universal Solvent

  • Solution: a liquid that is a homogeneous mixture of substances.
  • Solvent: the dissolving agent of a solution.
  • Solute: the substance that is dissolved.
  • Water is a universal solvent due to its polarity.

Hydrophilic vs. Hydrophobic

  • Hydrophilic: Loves water - substance has an affinity for water.
  • Hydrophobic: Hates or fears water - substance does not dissolve readily in water.

Water Dissociation: Acids and Bases

  • A hydrogen atom between two water molecules can shift from one to the other:
    H2O \rightarrow H^+ + OH^- H2O \rightarrow H_3O^+ + OH^-
  • Hydronium ion (H_3O^+).
  • Hydroxide ion (OH^–).

The pH Scale

  • Concentrations of H^+ and OH^- are equal in pure water.
  • Adding acids or bases changes the concentrations of H^+ and OH^-.
  • The pH scale indicates how acidic or basic a solution is.
  • pH is determined by the concentration of hydrogen ions.
  • Acidic solutions have pH values < 7.
  • Basic solutions have pH values > 7.
  • Most biological fluids range pH 6 - 8.

Acids and Bases

  • Acid: molecule that donates H^+ to a solution.
  • Acidic solution has more H^+ than pure water.
  • Base: molecule that accepts H^+ from a solution.
  • Basic solution has less H^+ than pure water.

Buffers

  • Most living cells are close to pH 7.4.
  • Buffers: substances that minimize changes in concentrations of H^+ and OH^- in a solution.
  • Buffers do NOT make the pH to be 7.0.
  • Living cells have “natural” buffers.

Effects of pH Change in a Cell

  • Molecular shapes can be altered.
  • Proteins might unfold or not fold correctly.

The Threat of Acid Precipitation

  • Rain, snow, or fog with a pH lower than 5.6.
  • Caused by the mixing of different pollutants with water in the air.
  • Can damage life in lakes and streams and alter soil chemistry.

Buffering in the Blood

  • Normal blood pH 7.4.
  • Exercise adds H^+ to blood.
  • Blood binds up CO2 to make H2CO_3.

Carbon

  • Forms the backbone of organic molecules.

Carbon Content of a Dehydrated Human

  • Oxygen (O): 30%.
  • Carbon (C): 47%.
  • Hydrogen (H): 9%.
  • Nitrogen (N): 8%.
  • Phosphorus (P): 3%.
  • Potassium (K): 1%.
  • Sulfur (S): 2%.
  • Calcium (Ca): 2%.
  • Sodium (Na): 1%.
  • Chlorine (Cl): 1%.
  • Magnesium (Mg): 1%.
  • Others.

Carbon Bonding

  • Can form single, double, or triple bonds.
  • Versatile in building complex molecules.

Macromolecules Key to Life

  • Proteins.
  • Nucleic acids.
  • Carbohydrates.
  • Lipids.

Proteins

  • Composed of amino acids.
  • Amino acids have an amino group, a carboxyl group, and an R group.

Polypeptide Chains

  • Amino acids are linked by peptide bonds to form polypeptide chains.
  • A polypeptide chain (protein).

Nucleic Acids

  • DNA and RNA.
  • Composed of nucleotides.
  • Nucleotides have a phosphate group, a sugar (ribose or deoxyribose), and a base (A, G, C, T, or U).

Nucleotide Structure

  • Pyrimidines: Thymine (T), Cytosine (C), Uracil (U).
  • Purines: Adenine (A), Guanine (G).
  • In a nucleic acid, each base is attached to either a ribose or a deoxyribose by the bond.

Nucleic Acid Bonding

  • Phosphodiester bonds link nucleotides.
  • Hydrogen bonds between base pairs (A-T, G-C).

Carbohydrates: Sugars

  • Aldoses (e.g., glucose) have an aldehyde group.
  • Ketoses (e.g., fructose) have a ketone group.
  • Sugars -> Carbohydrates - Glycosidic bond

Lipids

  • Defined by hydrophobic property.
  • Broad class of molecular structures.
  • Fatty Acid, Phospholipid & Steroids.

Phospholipid Structure

  • Polar head group (hydrophilic).
  • Glycerol backbone.
  • Fatty acid chains (hydrophobic).

Van der Waals Forces

  • Weak attractions between molecules or parts of molecules that result from transient local partial charges.

Saturated vs. Unsaturated Fats

  • Saturated fats: have no double bonds, solid at room temperature.
  • Unsaturated fats: have one or more double bonds, liquid at room temperature.

Trans Fat

  • Unhealthy type of fat formed through hydrogenation.
  • Cis-double bond & Trans double bond.

Phospholipids

  • Phospholipids have hydrophilic and hydrophobic components.
  • Do bilayers make sense?

Isotopes

  • Isotope: atoms of an element that differ in # neutrons.
  • Most isotopes are stable, but some are radioactive, giving off particles and energy.