Chemistry in Water: Bonding, Polarity, and Water’s Role in Biology

Purpose and big idea

• Revisit why chemistry is essential in biology: life is made of atoms that form bonds, build molecules, macromolecules, organelles, and cells. Understanding how atoms interact explains the structure and function of biomolecules.
• Core theme: structure equals function. The type of chemical bonds and the geometry they impose give molecules their shapes and roles in biology.
• Water as a central topic: all life is water-based, and water’s properties arise from chemistry. We’ll use water to illustrate how chemistry underpins metabolism, biomolecule behavior, and the environment.

Elements in biology: CHON, others, and trace elements

• All living things are composed of elements; the relevant scope for biology focuses on key elements and their interactions.
• Major elements by weight: CHON (Carbon, Hydrogen, Oxygen, Nitrogen) together make up about $96\%$ of the weight of living organisms.
• Rest of the weight (about $4\%$) includes seven elements that show up frequently in biology:
  • Phosphorus (P) and Sulfur (S) are common in phosphates and amino acids.
  • Calcium (Ca), Potassium (K), Sodium (Na), Chlorine (Cl), Magnesium (Mg).
• Trace elements: everything else essential but present in <$0.01\%$ abundance; highly potent in small amounts (e.g., iodine).
• Why the big emphasis on CHON and a handful of others: they drive the chemistry of life (biomolecules, energy, signaling).
• Differences across organisms: i.e., humans vs pumpkins differ in elemental composition (oxygen and nitrogen distribution relate to cell types and cell wall composition).
  • Plant cell walls are made of cellulose (a sugar polymer), contributing to higher oxygen content relative to humans.
  • Human cells contain more nitrogen due to proteins and nucleic acids; plants have more oxygen due to carbohydrate-rich cell walls.
• Takeaway: while CHON accounts for most mass, diversity in molecular composition and bonding patterns yields the variety of life.

Atomic structure and chemical reactivity basics

• Elements vs atoms: an element is a collection of atoms with the same number of protons; atoms arrange into elements with unique chemical properties.
• Nucleus and electrons: protons (positive), neutrons (neutral) in the nucleus; electrons surround the nucleus in electron shells. The outermost shell is the valence shell.
• Noble gases: atoms with a complete valence shell are chemically inert (nonreactive).
  • Example: neon (Ne) has a complete valence shell with $8$ outer electrons.
• Stability and reactivity:
  • If an atom’s valence shell is complete, it is typically nonreactive.
  • If it is not complete, atoms tend to become more stable by completing the valence shell:
  • Some elements gain electrons, some lose electrons, and others share electrons to fill their shells.
• Practical takeaway: reactivity and bonding patterns are driven by how far an atom is from a full valence shell.
• Simple but essential example: sodium (Na) with one valence electron loses it to become Na$^+$ (a cation), while chlorine (Cl) with seven valence electrons gains one to become Cl$^-$ (an anion). The electrostatic attraction between Na$^+$ and Cl$^-$ forms an ionic bond and can create crystalline NaCl.
• In biology, ionic bonds occur mainly between metals and nonmetals, but the biological context (especially in water) can alter perceived bond strength.

Ionic and covalent bonds; polar vs nonpolar covalent bonds; electronegativity

• Ionic bonds:
  • Formed by electrostatic attraction between ions created by electron transfer (ionization).
  • Example: Na and Cl form Na$^+$ and Cl$^-$; they attract to make NaCl crystals.
  • In biology, ionic bonds are strong in a vacuum but not necessarily the strongest in aqueous environments because water disrupts electrostatic interactions.
• Covalent bonds:
  • Formed by sharing electrons between two atoms; this bonding creates discrete bond angles and molecular shapes.
  • Carbon is tetra-valent (valence = 4): it can form up to four covalent bonds (e.g., CH$_4$).
  • Covalent bonds can be represented in several ways: electron orbital diagrams, ball-and-stick models, or space-filling models.
  • Covalent bonds determine the three-dimensional shape of biomolecules, which in turn determines function.
• Molecules vs compounds:
  • A molecule is two or more atoms bonded together.
  • A compound is a molecule that contains at least two different elements.
  • All compounds are molecules, but not all molecules are compounds (e.g., O$_2$ is a molecule but not a compound).
• Polar vs nonpolar covalent bonds (bond polarity is key for biology):
  • Nonpolar covalent bonds: electrons shared roughly equally. Example: H$_2$ (two H atoms) with a difference in electronegativity $\Delta\chi \approx 0$1, so a nonpolar bond.
  • Polar covalent bonds: electrons shared unequally, creating partial charges (dipoles). Example: water bonds in H$_2$O show partial negative on oxygen and partial positive on hydrogens.
  • Bond polarity depends on electronegativity differences between the bonded atoms.
• Electronegativity and bond type thresholds (numbers to memorize for biology contexts):
  • If $\Delta\chi \le 0.4$: nonpolar covalent bond.
  • If $0.5 \le \Delta\chi \le 1.7$: polar covalent bond.
  • If \Delta\chi > 1.7: ionic bond (in practice, water can disrupt many ionic interactions, so in biology covalent bonds are often the strongest).
• Quick reference electronegativities (biologically relevant elements):
  • Hydrogen (H): approx. $2.1$
  • Oxygen (O): approx. $3.5$
  • Fluorine (F) is the highest overall, but it’s less central to basic biology courses here.
• Why electronegativity matters for biology:
  • The difference in electronegativity determines bond polarity, which in turn influences molecular interactions with water and other molecules.
  • Polar regions tend to be hydrophilic (water-loving); nonpolar regions tend to be hydrophobic (water-fearing).
  • Dipoles: polar covalent bonds create partial charges ($\delta+$ and $\delta-$) that influence molecular interactions and three-dimensional structure.

Covalent bonds and molecular structure in biomolecules

• Local geometry matters: covalent bonds create specific bond angles, which gives molecules their shapes.
• Methane as a simple example: CH$_4$ forms a tetrahedral geometry with four C–H bonds.
• Macromolecules and shape/function: linking small units (e.g., many CH units) can produce larger structures with different shapes and functions (e.g., cholesterol macromolecule vs. a single methane unit).
• Types of covalent bonds:
  • Single covalent bonds: one pair of electrons shared.
  • Double covalent bonds: two pairs shared.
  • Triple covalent bonds: three pairs shared.
• Common biologically relevant molecules: H$<em>2$O, CO$</em>2$, CH$<em>4$, O$</em>2$, etc.
• Special notes:
  • A molecule can be a molecule without being a compound (same atoms, e.g., O$_2$).
  • A molecule that contains different elements is also a compound (e.g., CO$_2$).
• Hydrogen bonding as a non-covalent interaction:
  • Hydrogen bonds are electrostatic attractions where a partially positive hydrogen is attracted to a partially negative atom (e.g., O or N) in another molecule or within the same molecule.
  • Hydrogen bonds are weaker individually but can be collectively strong and biologically crucial (e.g., in water and biomolecules).

Hydrogen bonds and van der Waals forces

• Hydrogen bonds:
  • Occur between a hydrogen atom covalently bonded to a highly electronegative atom (like O or N) and another electronegative atom with a lone pair.
  • Shown as dotted lines or sometimes as bent lines in diagrams.
  • Individually weak, but many such bonds together produce significant effects in water structure and biology.
• Van der Waals forces (London dispersion forces, in this context):
  • Weak attractions that occur when nonpolar or slightly polar molecules come very close to each other.
  • Distance-dependent: too close leads to repulsion; too far reduces attraction.
  • Important for close-range interactions; underpin phenomena like gecko adhesion through cumulative small attractions.
• Nature of noncovalent interactions in biology:
  • They are essential for the dynamic assembly and disassembly of biomolecular complexes, folding, and membrane behavior.

Water: structure, polarity, and unique properties

• Water’s role in life:
  • Water is essential for metabolism: many cellular reactions require water to make or break chemical bonds (condensation/dehydration synthesis and hydrolysis).
  • Cells are predominantly water, and whole organisms are water-rich; metabolism and ATP production rely on aqueous environments.
• Water’s molecular structure:
  • Chemical formula: H$_2$O; two hydrogens covalently bonded to one oxygen.
  • Bonds: covalent and polar covalent within the molecule due to electronegativity difference, with oxygen more electronegative than hydrogen.
  • Resulting dipole: oxygen bears partial negative charge; hydrogens bear partial positive charges.
• Water-water interactions (cohesion) and water-other interactions (adhesion):
  • Cohesion: water molecules bonding to each other via hydrogen bonds.
  • Adhesion: water bonding to other substances; important for capillary action and plant hydraulics (to be discussed next week).
• Water’s hydrogen-bond network drives several key properties:
  • High density of liquid water relative to ice (water is denser as a liquid than as a solid).
  • High specific heat capacity: large amount of heat required to raise the temperature of water by 1 degree Celsius per unit mass due to breaking/forming many hydrogen bonds.
  • High heat of vaporization: a large amount of heat is required to break hydrogen bonds so water can vaporize.
  • High surface tension: cohesive forces at the air-water interface create a surface that can support small objects and enable certain organisms to move across surfaces.
• Density anomaly and life on Earth: \rho(\text{ice}) < \rho(\text{liquid water})

Metabolism, structure-function connections, and real-world implications

• Metabolic processes rely on aqueous chemistry: hydrolysis, condensation, and other water-involved reactions drive the build-up and break-down of biomolecules.
• Structure-function principle in biology:
  • Molecular shape dictated by covalent bonds and bond angles influences function (binding, catalysis, signaling).
  • Water’s polarity and hydrogen-bonding patterns influence molecular orientation, protein folding, and biomolecular interactions.
• Practical implications and examples mentioned in class:
  • Ionic bonds vs covalent bonds: in chemistry biology, covalent bonds are often the strongest within biological contexts, but their function is modulated by the surrounding water and solutes.
  • Van der Waals forces support fine-tuned interactions and structural stability in biomolecules and in physical phenomena like gecko adhesion.
  • Water’s buffering and thermal properties stabilize environmental temperatures and biological systems.
• Quick quiz-style takeaways:
  • Eye-clicker concepts illustrated the distinctions between bond types and how electronegativity differences determine bond polarity.
  • Memorization of certain electronegativities (e.g., H $\approx 2.1$, O $\approx 3.5$) helps predict polarity and interactions in common biological molecules.

Key numbers, definitions, and formulas (reference)

• Valence and bonding basics:
  • Maximum valence electrons in a shell: $8$ in the outer shell for main-group elements.
  • Carbon valence: $4$ (can form up to four bonds).

Elemental Composition of Living Organisms

Key Electronegativities

Bond Polarity Thresholds (Electronegativity Difference $\Delta\chi$)

• Common molecules discussed:
  • Water: H$_2$O
  • Sodium chloride (table salt): NaCl; ions: Na$^+$, Cl$^-$
  • Methane: CH$_4$
  • Carbon dioxide: CO$_2$
  • Oxygen gas: O$_2$
• Water’s biological importance and properties (conceptual):
  • Specific heat capacity and heat of vaporization reflect energy required to break hydrogen bonds and raise temperature.
  • Cohesion and surface tension arise from intermolecular hydrogen bonds in water.
• Conceptual relationships:
  • Structure (bond types and geometry) determines function in biomolecules.
  • Water’s polarity and hydrogen-bond network underpin metabolism, transport, and environmental stability.

Takeaway connections to prior learning and real-world relevance

• The course emphasizes that chemistry is the foundation for understanding biology at all scales, from atoms to cells to ecosystems.
• Water’s unique properties, driven by hydrogen bonding and polarity, are central to life on Earth, affecting metabolism, climate, and the behavior of biomolecules.
• The distinctions among ionic, covalent (polar and nonpolar), hydrogen bonds, and van der Waals forces explain why molecules fold, interact, and function as they do in living systems.
• Ethical/practical relevance: understanding water chemistry informs fields from medicine to environmental science, and underscores the importance of protecting water quality and aquatic environments for life support.

Quick recap (visible in lecture cues)

• Life’s unity relies on basic chemical principles: atoms, electrons, and bonds shape everything from proteins to membranes.
• CHON plus a handful of key elements are the major players in biology; trace elements, though small in amount, can be essential.
• Bonds determine structure and function; polar vs nonpolar nature guides interactions with water and other molecules.