Chemistry: Hydrogen Bonds and Chemical Reactions

Foundations of Chemistry in Biology and Microbiology

• Microbial cells, like all living organisms in the universe, are composed of chemical molecules and rely heavily on appropriate chemical reactions to sustain life.

• Approximately $98\%$ of the mass of all living organisms is composed of just six chemical elements:

  • Carbon ($C$)

  • Hydrogen ($H$)

  • Oxygen ($O$)

  • Nitrogen ($N$)

  • Phosphorus ($P$)

  • Sulfur ($S$)

• These six fundamental elements form the key biological molecules essential to life:

  • Water ($H_2O$): The most abundant molecule in living systems and cells. It is essential to all life, and no single organism on Earth can survive without its presence.

  • Four Major Organic Macromolecules ("macro" meaning large), which constitute the remainder of the cell:

  • Lipids

  • Proteins

  • Carbohydrates

  • Nucleic acids

Atomic Structure and Fundamental Properties

• An atom is the smallest unit of matter and makes up everything in the universe.

• Atoms combine with one another to form molecules, which build living cells.

• Subatomic Particles and Anatomy of an Atom:

  • Nucleus: The centrally located core of an atom, possessing an overall positive charge. It consists of two subatomic particles:

  • Protons: Positively charged particles.

  • Neutrons: Electrically neutral particles (carrying no charge).

  • Electron Shells (Orbitals): Concentric outer regions surrounding the positive nucleus where electrons travel.

  • Electrons: Negatively charged particles attracted to the positively charged nucleus.

• Distribution of Mass versus Volume in an Atom:

  • Atomic Mass: Provided almost entirely by the nucleus, where protons and neutrons reside.

  • Atomic Volume: Provided by electrons moving rapidly around the nucleus and occupying space.

  • Scale Comparison Metaphor: If the nucleus were the size of a pencil eraser, the electron shells defining the atom's volume would be the size of a football field.

• Electrical Neutrality:

  • An atom in its baseline state contains an equal number of protons and electrons.

  • Because protons ($+\text{charge}$) and electrons ($-\text{charge}$) carry opposite charges in equal quantities, an atom is electrically neutral.

• Metrics of Atomic Identity:

  • Chemical Element: Classified by atoms sharing the exact same number of protons.

  • Atomic Number: Represents the total number of protons present in the nucleus of an element.

  • Atomic Weight (or Atomic Mass): Represents the combined total number of protons and neutrons in an element.

Isotopes and Radiation

• Definition: Isotopes are alternative structural forms of a chemical element that differ in their number of neutrons.

  • Isotopes of a given element always maintain the exact same atomic number (proton count); changing the proton count would transform the atom into a different element entirely.

  • Because neutron numbers differ, isotopes vary in their atomic weight.

• Case Study of Carbon Isotopes:

  • Carbon naturally exists on Earth in three primary isotopic forms:

  • Carbon-12 ($^{12}C$): A stable isotope form.

  • Carbon-13 ($^{13}C$): A stable isotope form.

  • Carbon-14 ($^{14}C$): An unstable, radioactive isotope form that spontaneously emits radiation.

• Practical Application of Radioisotopes:

  • The predictable radiation emission of Carbon-14 is regularly utilized in scientific fossil dating to determine age by measuring remaining radiation levels.

Electron Shells and Configurations

• Electron Shells (Orbitals): Layers of regions outside the nucleus corresponding to specific energy levels.

• Electron Configuration: The structural arrangement of electrons within these energy shells.

• Maximum Shell Capacities:

  • $1\text{st}$ shell (innermost): Maximum capacity of $2$ electrons.

  • $2\text{nd}$ shell: Maximum capacity of $8$ electrons.

  • $3\text{rd}$ shell: Maximum capacity of $8$ electrons.

  • $4\text{th}$ shell: Maximum capacity of $18$ electrons.

  • $5\text{th}$ shell: Maximum capacity of $18$ electrons.

  • $6\text{th}$ shell: Maximum capacity of $18$ electrons.

• Valence Shells and Chemical Stability:

  • The outermost electron shell is the primary focus of chemical reactivity.

  • Maximum chemical stability is achieved when an atom's outermost shell is completely filled to capacity.

  • Atoms with incomplete outer shells (such as Hydrogen with $1$ outer electron, or Oxygen with $6$ out of $8$ outer electrons) are chemically unstable and naturally react to transfer or share electrons.

Chemical Bonds

• A chemical bond is the attractive force generated between atoms when electrons are shared or transferred to fill their outermost shells and achieve chemical stability.

• Ionic Bonds and Ion Formation:

  • Formation Mechanism: Occurs when electrons are completely transferred from one atom to another.

  • Ion Mechanics:

  • Protons are fixed within the nucleus and cannot move; only electrons move between atoms.

  • Loss or gain of electrons alters an atom's net charge, creating an ion.

  • Cations: Positively charged ions formed when an atom loses electrons.

  • Anions: Negatively charged ions formed when an atom gains electrons.

  • Electrolytes: Ions dissolved in aqueous solution that possess the ability to conduct electricity.

  • Bond Definition: An ionic bond is the electrostatic attraction between oppositely charged ions ($\text{positive} \leftrightarrow \text{negative}$).

  • Example - Sodium ($Na$) and Chlorine ($Cl$):

  • Sodium ($Na$) has $1$ electron in its 3rd (outermost) shell. By donating this $1$ electron, its filled 2nd shell ($8$ electrons) becomes the outer shell, forming a stable $Na^+$ cation.

  • Chlorine ($Cl$) has $7$ electrons in its 3rd shell. By accepting $1$ electron, it completes its 3rd shell ($8$ electrons), forming a stable $Cl^-$ anion.

  • The opposite charges attract $Na^+$ and $Cl^-$ to form a stable ionic compound.

• Covalent Bonds:

  • Formation Mechanism: Occurs when two atoms share one or more pairs of valence electrons.

  • Relative Strength: Covalent bonds are stronger than ionic bonds and are more prevalent in biological organisms.

  • Nonpolar Covalent Bonds:

  • Occur when sharing of electrons between nuclei is equal.

  • Equal attraction ensures the electron cloud is evenly distributed, neutralizing nuclear positive charges and generating no net electrical charge.

  • Polar Covalent Bonds:

  • Occur when one nucleus exerts a stronger attraction on the shared electrons than the other nucleus.

  • Electrons spend more time near the stronger nucleus, creating an unequal distribution of charge.

  • Quantitative Model of Polar Bond Generation:

    • Suppose two nuclei each contain a $+1$ charge.

    • If two shared electrons spend unequal time such that the stronger nucleus averages $1.5$ electrons around it and the weaker nucleus averages $0.5$ electrons:

    • Net charge around weaker nucleus: $+1 - 0.5 = +0.5$ (partially positive, $\delta+$).

    • Net charge around stronger nucleus: $+1 - 1.5 = -0.5$ (partially negative, $\delta-$).

Hydrogen Bonds and Properties of Water

• Nature of Hydrogen Bonds:

  • Hydrogen bonds involve no movement (transfer or sharing) of electrons.

  • They are caused by partial electrostatic attractions between distinct molecules, rather than individual atoms.

  • While an individual hydrogen bond is a relatively weak attraction, thousands or millions of hydrogen bonds acting collectively create an exceptionally strong force.

• Molecular Structure of Water ($H_2O$):

  • Intramolecular Bonding: Within a single water molecule, one oxygen atom and two hydrogen atoms are joined by polar covalent bonds.

  • Oxygen holds electrons more tightly, creating a partial negative charge ($\delta-$) on the oxygen portion.

  • Hydrogen holds electrons more weakly, creating a partial positive charge ($\delta+$) on the hydrogen portions.

  • Intermolecular Bonding: Between separate water molecules, the partial positive hydrogen ($\delta+$) of one molecule is attracted to the partial negative oxygen ($\delta-$) of another molecule, forming a hydrogen bond.

• Life-Sustaining Properties of Water:

  • High Heat Capacity:

  • Significant heat energy must be input to break hydrogen bonds before water's temperature rises or it transitions from liquid to gas.

  • Makes water an ideal temperature buffer in biological systems.

  • Super Solvent ("Solvent of Life"):

  • The unequal partial charge distribution allows water to readily dissolve polar and ionic compounds.

• Hydrophilic versus Hydrophobic Interactions:

  • Hydrophilic ("water-loving"): Ionic or polar compounds that dissolve easily in water.

  • Hydrophobic ("water-fearing"): Non-polar compounds that do not dissolve well in water.

  • Hydrophobic Force: Water molecules preferentially hydrogen-bond with each other, excluding non-polar molecules and forcing them to cluster together in aqueous environments.

Chemical Reactions and Energetics

• Definition of Chemical Reactions:

  • Processes involving the breaking and reforming of chemical bonds between atoms or molecules.

  • Atoms within initial molecules (reactants) are rearranged to form distinct resulting molecules (products).

• Detailed Example - Formation of Water Molecules:

  • Reactants: Oxygen gas ($O_2$) and Hydrogen gas ($H_2$).

  • Process: Chemical bonds holding oxygen atoms together and hydrogen atoms together are broken. Two new covalent bonds are then formed between each oxygen atom and two hydrogen atoms to build $H_2O$ molecules.

• Energetics and Free Energy:

  • All matter possesses chemical energy dictated by its atomic and structural arrangements.

  • Reactions always produce a change in energy levels between reactants and products.

  • Free Energy ($\Delta G$): The difference in energy levels between reactant molecules and product molecules.

  • Exergonic Reactions:

  • Occur when initial reactant molecules possess a higher energy level than the final product molecules.

  • Energy is released from the system into the environment.

  • Endergonic Reactions:

  • Occur when initial reactant molecules possess a lower energy level than the final product molecules.

  • Require a net input of energy into the system to proceed.