Chemistry Basics for Biology

  • Introduction to Chemistry in Biology Chemistry provides the foundational understanding relevant to biological processes, enabling us to comprehend how life functions at a molecular level.
    • Importance of familiarity with the Periodic Table of Elements: It organizes all known elements based on their atomic number, chemical properties, and helps predict their behaviors in chemical reactions.
  • Elements and Atoms Elements and atoms essentially refer to the same substance, with atoms being the building blocks of elements.
    • Atoms may have variations (isotopes) due to differing numbers of subatomic particles (protons, neutrons, electrons).
    • Common elements in the human body:
    • Main Components:
      • Carbon (C), Hydrogen (H), Oxygen (O), Nitrogen (N)
      • These elements together constitute approximately 96% of the body's mass, providing a basis for organic compounds.
    • Micronutrients: Necessary elements in smaller amounts are crucial for various biological functions:
      • Include Calcium (Ca) for bone structure and muscle contraction, Phosphorus (P) for energy transfer via ATP, Potassium (K) for nerve function, Sulfur (S) for protein structure, Chlorine (Cl) for maintaining fluid balance, Sodium (Na) for electrical signaling, Magnesium (Mg) for enzyme activity, Iodine (I) for thyroid function, Iron (Fe) for oxygen transport in blood.
  • Atomic Structure
    • Three subatomic particles: Protons (positive), Neutrons (neutral), Electrons (negative), each playing critical roles in atomic properties.
    • The identity of an element is determined by the number of protons. For example, every hydrogen atom has 1 proton, while carbon has 6, defining their characteristics.
    • Electron configuration: Electrons orbit around a central nucleus (containing protons and neutrons) in specific energy levels.
    • Electrons are responsible for chemical reactions and bonding, influencing how atoms interact to form molecules.
  • Isotopes and Radioactivity
    • Isotopes are variations of an element with differing numbers of neutrons, which can affect their stability.
    • For example, Carbon-12 (6 protons, 6 neutrons) is stable, whereas Carbon-14 (6 protons, 8 neutrons) is unstable and radioactive, used in dating fossils and archaeological samples due to its predictable rate of decay.
  • Ions and Electrolytes
    • Changing the number of electrons produces ions:
    • Cations: Positively charged ions formed by the loss of electrons.
    • Anions: Negatively charged ions formed by the gain of electrons.
    • Examples: Electrolytes in the body include Sodium (Na+), Chloride (Cl-), and Calcium (Ca2+), critical for physiological functions including fluid balance, nerve signaling, and muscle contraction.
  • Chemical Bonds
    • Atoms link through chemical bonds to form molecules, which are vital for building structures and carrying out reactions within organisms.
    • Ionic Bonds: Formed between positively and negatively charged ions due to electrostatic attraction.
    • Example: Table salt (NaCl) is formed from the attraction of Na+ and Cl- ions, resulting in a solid lattice structure.
    • Covalent Bonds: Atoms share electrons, creating stronger bonds compared to ionic bonds.
    • Example: Methane (CH₄) is where carbon shares electrons with four hydrogens, allowing for a stable molecule.
    • Hydrogen Bonds: Weak bonds that occur between polar molecules, arising from partial positive and negative charge regions.
    • These bonds are crucial for properties of water, affecting its fluidity and temperature regulation.
  • Properties of Water
    • Water facilitates chemical reactions and is a universal solvent for ionic and polar molecules, allowing for numerous biological processes.
    • Cohesion (water molecules sticking together) results in surface tension, while adhesion (water molecules sticking to other substances) aids in nutrient transport in plants.
    • Unique properties due to hydrogen bonding include high heat capacity and the ability to stabilize temperatures in living organisms.
  • pH and Homeostasis
    • The pH scale measures hydrogen ion concentration, with solutions being inversely logarithmic
    • Acidic solutions have lower pH; basic solutions have higher pH
    • Acids (substances that donate H+) vs. Bases (substances that accept H+).
    • Human blood normally exhibits a pH of around 7.4 (slightly basic), essential for proper physiological functions such as enzyme activity and metabolic processes.
    • Acidosis: A condition characterized by excess H+, with blood pH dropping below 7.35.
    • Alkalosis: A condition where H+ levels are lower, with blood pH rising above 7.45.
    • Buffering systems in the body, like bicarbonate and phosphate buffers, maintain pH within narrow limits, essential for life.
  • Biological Importance of Molecules
    • Molecules can be hydrophilic (water-loving) or hydrophobic (water-fearing), which influences how they interact in biological systems.
    • Amphipathic molecules (having both hydrophilic and hydrophobic regions) like phospholipids are crucial for forming cell membranes:
    • Phospholipids arrange into bilayers, creating barriers that regulate the movement of substances in and out of cells.
  • Conclusion and Review
    • Understanding these concepts is vital for further studies in biology, including cellular processes, metabolism, and the mechanics of life.
    • Review the relationship between the structure of molecules and their functions in biological systems, as this foundation is critical for advanced topics such as biochemistry and