Life Chemistry: Atomic Structure, Bonding, and Water Properties

Overview of Life Chemistry and Energy

  • Fundamental Goals of Study
    • Identify specific elements and determine their individual properties.
    • Predict the formation of chemical bonds between atoms.
    • Understand the unique properties of water and its critical importance to living organisms.
    • Determine molecular specific properties based on atomic arrangement and functional groups.

Atomic Structure and Matter

  • Matter and Atoms

    • All living and nonliving matter is composed of atoms.
    • Matter is defined as anything that has mass and volume.
    • An element is a substance made up of only one type of atom (e.g., Aluminum, Carbon, Hydrogen).
    • Atoms are the smallest units of an element and the building blocks of all matter.
    • A molecule is formed when two or more atoms bind together (e.g., H2H_2, O2O_2).
    • A compound is a molecule containing at least two different elements (e.g., H2OH_2O, CO2CO_2, CH4CH_4).

  • Organizational Hierarchy of Life

    • Atoms →\rightarrow Small molecules (Water, Oxygen, Methane, Carbon dioxide) →\rightarrow Large molecules (Proteins, Nucleic acids) →\rightarrow Cells (including Unicellular organisms and Colonial organisms) →\rightarrow Tissues →\rightarrow Organs →\rightarrow Organ systems →\rightarrow Multicellular organism (exemplified by the leopard frog).

  • Essential Elements for Life

    • Living things are primarily composed of six chemical elements:
      1. Carbon (C)
      2. Hydrogen (H)
      3. Phosphorus (P)
      4. Oxygen (O)
      5. Nitrogen (N)
      6. Sulfur (S)

Subatomic Particles and Atomic Properties

  • Composition of Atoms

    • All atoms share a spherical shape and the same internal structure.
    • Protons: Positively charged particles located in the nucleus.
    • Neutrons: Uncharged (neutral) particles located in the nucleus.
    • Electrons: Negatively charged particles that revolve around the nucleus in shells.
    • Atoms are generally neutral particles because the number of protons equals the number of electrons.

  • Numerical and Identifying Characteristics

    • Atomic Number: Identifies an element; it is equivalent to the number of protons in the nucleus.
    • Electrical Neutrality: Achieved when the number of protons (++) equals the number of electrons (−-).
    • Mass Number: The total number of protons and neutrons in an atom.
    • Mass of a Proton: Defined as 1 Dalton1\,Dalton, which is approximately 1.7×10−24 grams1.7 \times 10^{-24}\,grams.
    • Mass of Electrons: Extremely small, approximately 11836\frac{1}{1836} the mass of a proton.

  • Isotopes

    • Isotopes are variants of the same element that have the same number of protons but a different number of neutrons.
    • Isotopes possess the same chemical properties but different physical properties.
    • Examples of Hydrogen isotopes:
      • Deuterium: Contains one proton and one neutron.
      • Tritium: Contains one proton and two neutrons.

Atomic Models and Chemical Bonding

  • The Bohr Model

    • Electrons travel in fixed paths around the nucleus called shells or energy levels (n=1,2,3...n = 1, 2, 3...).
    • First shell: Maximum of 22 electrons.
    • Second shell: Maximum of 88 electrons.
    • Third shell: Maximum of 1818 electrons.
    • Specific atomic nucleus charges (number of protons):
      • Hydrogen (H): 1+1+
      • Carbon (C): 6+6+
      • Nitrogen (N): 7+7+
      • Oxygen (O): 8+8+
      • Phosphorus (P): 15+15+
      • Sulfur (S): 16+16+

  • The Octet Rule

    • Many atoms are most stable when they have 88 electrons in their outermost electron shells (valence shells).
    • Atoms with unfilled outer shells tend to undergo chemical reactions to fill those shells, leading to the formation of molecules through chemical bonds.

Types of Chemical Bonds and Interactions

  • Ionic Bonds

    • Formed by the complete transfer of one or more electrons from one atom to another.
    • Ions: Charged particles resulting from this transfer.
      • Cations: Positively charged ions (e.g., Na+Na^+), formed by losing electrons.
      • Anions: Negatively charged ions (e.g., Cl−Cl^-), formed by gaining electrons.
    • The attraction results from the electrical pull between positive and negative ions. The resulting molecules are called salts.
    • Calculation example: (0)−(−1)=+1(0) - (-1) = +1 (losing electron); (0)+(−1)=−1(0) + (-1) = -1 (gaining electron).
    • Ionic attractions are relatively weak, causing salts to dissolve easily in water.

  • Covalent Bonds

    • Formed when two or more atoms share pairs of electrons to achieve stability.
    • Single covalent bond: Shares 11 pair of electrons.
    • Double covalent bond: Shares 22 pairs of electrons.
    • Triple covalent bond: Shares 33 pairs of electrons.
    • Orientation: The length, angle, and direction of bonds between specific elements are constant. Example: Methane (CH4CH_4) always forms a tetrahedron shape.

  • Electronegativity and Polarity

    • Electronegativity: The attractive force that an atomic nucleus exerts on electrons. It depends on the number of protons and the distance between the nucleus and electrons.
    • Nonpolar Covalent Bond: Atoms have similar electronegativities and share electrons equally.
    • Polar Covalent Bond: Atoms have different electronegativities; electrons are drawn closer to the more attractive atom, creating partial charges (δ−\delta- and δ+\delta+).
    • Example: Oxygen (3.43.4) and Hydrogen (2.22.2). The difference is 1.21.2, causing a polar bond in H2OH_2O.
    • Electronegativity Values (Table 2.2):
      • Oxygen (O): 3.43.4
      • Chlorine (Cl): 3.23.2
      • Nitrogen (N): 3.03.0
      • Carbon (C): 2.62.6
      • Phosphorus (P): 2.22.2
      • Hydrogen (H): 2.22.2
      • Sodium (Na): 0.90.9
      • Potassium (K): 0.80.8

  • Bond Energies (Table 2.1)

    • Measured in kcal/molkcal/mol (energy needed to separate bonded atoms):
      • Covalent bond: 50−110 kcal/mol50-110\,kcal/mol
      • Ionic attraction: 3−7 kcal/mol3-7\,kcal/mol
      • Hydrogen bond: 3−7 kcal/mol3-7\,kcal/mol
      • Hydrophobic interaction: 3−7 kcal/mol3-7\,kcal/mol
      • van der Waals interaction: 1−2 kcal/mol1-2\,kcal/mol

Water Properties and Biological Impact

  • Hydrogen Bonds

    • The attraction between the δ−\delta- end (typically Oxygen) of one molecule and the δ+\delta+ Hydrogen end of another molecule.
    • These bonds can form between molecules (like two water molecules) or within single complex molecules (like DNA base pairs: Adenine and Thymine).

  • Thermal Properties

    • High Heat Capacity: Water requires a large amount of heat energy to change its temperature because that energy must first break the hydrogen bonds. This shields organisms from environmental temperature fluctuations.
    • High Heat of Vaporization: Significant heat is required to transform water from a liquid to a gas. This allows for evaporative cooling, such as sweating, which removes heat from the body.

  • Strength and Interaction

    • Cohesion: The resistance of water molecules to coming apart under tension, allowing columns of water to move through plants from roots to leaves.
    • Adhesion: Interaction between water and other surfaces.
    • Hydrophilic ("water-loving"): Polar molecules that interact with water and become surrounded by water molecules in aqueous solutions.
    • Hydrophobic ("water-hating"): Nonpolar molecules that interact with each other rather than water.

Functional Groups Important to Living Systems

  • General Features

    • Small groups of atoms with specific chemical properties.
    • They confer properties like polarity to larger biological molecules.
    • They determine molecular shape and chemical reactivity.

  • Specific Functional Groups

    • Hydroxyl (R−OHR-OH):
      • Class: Alcohols (e.g., Ethanol).
      • Properties: Polar; forms hydrogen bonds with water; facilitates linkages via condensation.
    • Aldehyde (R−CHOR-CHO):
      • Class: Aldehydes (e.g., Acetaldehyde).
      • Properties: Very reactive C=OC=O group; important in energy-releasing reactions.
    • Keto (R−CO−RR-CO-R):
      • Class: Ketones (e.g., Acetone).
      • Properties: Important in carbohydrates and energy reactions.
    • Carboxyl (R−COOHR-COOH):
      • Class: Carboxylic acids (e.g., Acetate).
      • Properties: Acidic; ionizes to −COO−-COO^- and H+H^+; enters condensation by giving up −OH-OH.
    • Amino (R−NH2R-NH_2):
      • Class: Amines (e.g., Methylamine).
      • Properties: Basic; accepts H+H^+ to form −NH3+-NH_3^+; gives up HH in condensation.
    • Phosphate (R−PO42−R-PO_4^{2-}):
      • Class: Organic phosphates (e.g., 3-Phosphoglycerate).
      • Properties: Negatively charged; releases significant energy upon hydrolysis when bonded to another phosphate.
    • Sulfhydryl (R−SHR-SH):
      • Class: Thiols (e.g., Mercaptoethanol).
      • Properties: Two −SH-SH groups can form a disulfide bridge (S−SS-S) to stabilize protein structures.

Questions and Discussion

  • Discussion on Isotope Reactivity

    • Question: Deuterium is an isotope of hydrogen with one neutron (normal hydrogen has none). Does the neutron change chemical reactivity?
    • Conceptual Answer: Typically, isotopes have the same chemical properties because chemical reactivity is determined by the electron configuration/behavior, not the number of neutrons.

  • Bonding Scenarios for Phosphorus

    • Phosphorus has five electrons in its valence shell. To reach stability/octet configurations, it can form:
      • Three single bonds.
      • One double bond and one single bond.
      • One triple bond.

  • Attendance Credit Activity

    • Requirement: List 3 things learned and 1 remaining question from the session.