Comprehensive Study Notes: Bonding, Water, pH, Carbon, and Functional Groups
Bonding and Polarity: electronegativity, polar vs nonpolar
Electronegativity differences determine bond polarity
Polar bonds: electrons pulled toward the more electronegative atom, creating partial charges
Nonpolar bonds: electrons shared more equally
Practical threshold discussed: if the difference is around riangle EN \approx 0.4 or larger, bonds tend to be polar; smaller differences tend toward nonpolar
Examples from lecture:
Oxygen–Hydrogen (O–H): polar bond; oxygen pulls electrons, giving the oxygen a partial negative charge and hydrogen a partial positive charge
Carbon–Hydrogen (C–H): nonpolar bond; electronegativity difference is small, electrons shared evenly
Carbon–Oxygen (C–O): polar bond due to electronegativity difference
Nitrogen–Hydrogen (N–H): polar as well
Visual shorthand for partial charges:
δ+ on the less electronegative end (partial positive)
δ− on the more electronegative end (partial negative)
Practical takeaway for biology:
Carbon–Hydrogen bonds are a hallmark of nonpolarity in many biological hydrocarbons
Oxygen–Hydrogen bonds are highly polar and common in water and organic molecules interacting with water
Hydrogen bonds and water: formation, strength, and consequences
Hydrogen bond definition:
An attraction between a partially positive hydrogen (δ+) in one molecule and a partial negative site (δ−) on another atom (often O or N)
Not a full ionic or covalent bond; individually weak, but collectively strong when many are present
Why hydrogen bonds form in water and biomolecules:
Water is polar; the O atom is δ− and the H atoms are δ+. This drives attraction between water molecules and with other polar/charged species
Comparison of strength:
Hydrogen bonds are weaker than covalent and ionic bonds, but many hydrogen bonds can create strong overall interactions
Dynamic nature:
Hydrogen bonds are constantly forming and breaking; they are transient yet collectively stabilize structures
Consequences of hydrogen bonding in water:
Water as a universal solvent for polar and charged substances
Water provides cohesion (water–water attraction) and adhesion (water–surface attraction)
Surface tension: cohesive water molecules at the surface resist breaking away
Capillary action: water moves through narrow tubes (e.g., plant xylem) via cohesive and adhesive forces
Water–water networks and properties:
When many water molecules hydrogen-bond, they produce unique properties essential for life, such as high heat capacity and temperature moderation
Properties of water arising from hydrogen bonding
Temperature moderation:
High energy is required to break hydrogen bonds to raise temperature or to vaporize
This is why sweating cools the body: energy is consumed to break bonds during evaporation
Coastal climates moderate temperature due to large, stable bodies of water
Cohesion and adhesion:
Cohesion: water molecules attract to each other via hydrogen bonds
Adhesion: water molecules attract to other polar/charged surfaces
These lead to surface tension and capillary action in plants
Ice vs liquid water density:
Water expands upon freezing; ice is less dense than liquid water due to the hydrogen-bond network forming a more open, hexagonal-like structure with greater empty space
This open structure allows ice to float and insulate bodies of water beneath
Dissolving substances in water:
Water readily dissolves polar or charged substances (ions, hydrophilic molecules)
Hydrophilic vs hydrophobic:
Hydrophilic: substances that form/have interactions with water (polar or charged)
Hydrophobic: nonpolar substances that water largely avoids (e.g., oils)
Mechanism example: dissolution of table salt (NaCl) where water molecules surround Na+ (with δ− oxygen) and Cl− (with δ+ hydrogens)
Hydrophilic vs hydrophobic examples:
Hydrophilic: ions (Na+, Cl−), polar molecules, ammonia (NH3)
Hydrophobic: nonpolar hydrocarbons (oil), many carbon–carbon and carbon–hydrogen rich substances
Water as solvent in biology:
Most cellular environments are aqueous; solutes can be dissolved or interact through hydrogen bonding
The concept of aqueous solutions (solvent is water)
Dissociation, pH, and the acid–base framework
Water autoionization:
H_2O \rightleftharpoons H^+ + OH^- (often written as hydronium H3O+ in more complete form)
In pure water, the concentration of H+ (or H3O+) and OH− is very small: [H^+] = [OH^-] = 10^{-7} \text{ M} at neutral pH
pH scale:
Defined as pH = -\log_{10}[H^+]
Neutral water corresponds to pH ≈ 7 because [H^+] = 10^{-7} \text{ M}, so pH = -\log_{10}(10^{-7}) = 7
Relationship between pH and hydrogen ion concentration:
[H^+] = 10^{-pH}
Magnitude of pH changes:
Each unit change in pH corresponds to a tenfold change in hydrogen ion concentration
Example: going from pH 7 to pH 4 increases [H+] by a factor of 10^{3} = 1000 (three pH units downwards)
Conversely, going from pH 7 to pH 10 decreases [H+] by a factor of 1000 (three pH units upwards)
Buffers:
Substances that resist changes in pH by binding or releasing hydrogen ions as needed
Example concept: if a solution becomes too acidic, a buffer accepts H+; if it becomes too basic, a buffer donates H+ to raise acidity
Ammonia in water as a buffer example:
Reaction: NH3 + H2O \rightleftharpoons NH_4^+ + OH^-
The added ammonia shifts equilibrium to form ammonium and hydroxide, increasing basicity by effectively removing H+ from solution
Practical problem framing from case study:
If ammonia addition reduces free hydrogen ion concentration by a factor of 100, the pH increases by 2 units (from pH 7 to pH 9) due to the log relationship
For a 100× decrease in [H+], \Delta pH = +2 (solution becomes more basic)
Acidic vs basic within biological contexts:
pH can critically affect protein structure, binding, and reaction rates; small pH changes can have large biological consequences
Ocean acidification is an example of broad ecological impact due to changes in pH from increased atmospheric CO2
Carbon, organic backbones, and functional groups
Why carbon is the backbone of organic molecules:
Carbon has four valence electrons, allowing up to four covalent bonds
This enables flexible, diverse backbones: chains, branches, rings, multiple bonds, and varied geometries
Backbone chemistry with carbon and hydrogen:
Carbon–hydrogen interactions form hydrocarbon backbones
Hydrogen tends to fill remaining valence spots on carbon, giving a stable skeleton
Polar vs nonpolar bonds in carbon chemistry:
Carbon–hydrogen: typically nonpolar
Carbon–oxygen and carbon–nitrogen bonds: can be polar depending on electronegativity differences
Carbon–carbon bonds can be nonpolar or polar depending on substituents
Functional groups:
Seven common functional groups (per Fig. 4.9 in the course materials) are repeatedly encountered in biology
Purpose: endow molecules with characteristic reactivity and properties
Flashcards are recommended to memorize the name, structure, and behavior of these groups
Visual representations vary: line-angle drawings, ball-and-stick models, or full letter notation
Isomers:
Isomers have the same chemical formula but different arrangements of atoms
Differences can be in bond connectivity or spatial orientation
Isomers can dramatically affect biological function and properties despite identical formulas
Practical examples emphasizing functional groups and isomerism:
Estrogen vs testosterone differ by functional group positioning; small changes yield large physiological differences
THC vs CBD: similar skeleton but different functional group arrangement or ring closure alters biological activity
Practical examples: polarity, solubility, and molecular interactions
How polarity guides solubility:
Hydrophilic molecules dissolve or disperse in water due to polar bonds and/or ionic charges
Hydrophobic molecules (nonpolar) tend to separate from water (e.g., oil in water)
Molecular orientation in water around solutes:
Water molecules orient their dipoles to interact with charged or polar regions of solutes
The surrounding water shell around ions or polar groups stabilizes solutes in solution
Case study ideas:
For a given molecule, inspect individual bonds to assess overall polarity rather than trying to judge the whole molecule at once
If in water, a molecule shows water interaction via hydrogen bonds; it is hydrophilic and polar
If water largely ignores a molecule, that molecule is hydrophobic and nonpolar
Review of study and exam-related notes (from the lecturer)
Chapter and reading workflow:
Chapter 3 (water) discussed in class; Pearson pre-reading due before class; Chapter 4 contains a table (Figure 4.9) with the seven functional groups
Students should prepare flashcards for these functional groups to recognize their behavior across multiple molecules
Case study expectations:
First case study posted; some information provided earlier; remaining questions to be completed in class
A short follow-up video will be posted to aid students who need help starting questions
Exam logistics and study strategy:
First exam scheduled one week from Thursday; format: 50-minute exam with ~40 multiple-choice questions (scantrons provided)
Academic accommodations (DRC) should be scheduled in advance
Practice exam will be posted to illustrate general exam format; optional additional practice quizzes may be provided for focused topics
Study guidance: use learning objectives and study guides; study progressively (e.g., 10–20 minutes daily) rather than cramming
Emphasis on understanding core concepts (bond types, water properties, pH concepts) rather than memorizing broad content
Final notes on the carbon-centric approach:
Carbon’s versatility underpins biology; the focus is on recognizing how polarity and functional groups influence molecular behavior
The instructor highlighted that recognizing functional groups helps explain large-scale molecular behavior and biological function
Bonding and Polarity: electronegativity, polar vs nonpolar
Electronegativity differences determine bond polarity: Electronegativity is the measure of an atom's attraction for electrons in a covalent bond. The greater an atom's electronegativity, the more strongly it pulls shared electrons toward itself.
Polar bonds: occur when there is an uneven sharing of electrons between two atoms due to a significant difference in their electronegativity. Electrons are pulled toward the more electronegative atom, creating an unequal distribution of charge, resulting in partial positive (\delta+) and partial negative (\delta-) charges on the atoms.
Nonpolar bonds: occur when electrons are shared more equally between two atoms, typically when their electronegativity difference is very small. This results in little to no partial charge development.
Practical threshold discussed: if the difference in electronegativity (\triangle EN) is approximately 0.4 or larger, bonds tend to be polar; smaller differences tend toward nonpolar. For example, a 0.2 difference would still technically be polar, but often considered functionally nonpolar in many biological contexts, while a 1.7 difference or more often indicates an ionic bond.
Examples from lecture:
Oxygen–Hydrogen (O–H): a highly polar bond. Oxygen is significantly more electronegative than hydrogen, pulling the shared electrons closer to itself. This gives the oxygen a partial negative charge (\delta-) and the hydrogen a partial positive charge (\delta+).
Carbon–Hydrogen (C–H): a nonpolar bond. The electronegativity difference between carbon and hydrogen is small (\approx 0.35), leading to a nearly even sharing of electrons. This characteristic makes hydrocarbons largely nonpolar.
Carbon–Oxygen (C–O): a polar bond due to the greater electronegativity of oxygen compared to carbon.
Nitrogen–Hydrogen (N–H): also a polar bond, as nitrogen is more electronegative than hydrogen (\triangle EN \approx 0.49).
Visual shorthand for partial charges:
\delta+ on the less electronegative end (partial positive), indicating a slight electron deficiency.
\delta- on the more electronegative end (partial negative), indicating a slight electron excess.
Practical takeaway for biology:
Carbon–Hydrogen bonds are a hallmark of nonpolarity in many biological hydrocarbons, like lipids and fatty acid tails.
Oxygen–Hydrogen bonds and Nitrogen–Hydrogen bonds are highly polar and are common in water, alcohols, amines, and other organic molecules that readily interact with water (hydrophilic).
Hydrogen bonds and water: formation, strength, and consequences
Hydrogen bond definition:
An attraction between a partially positive hydrogen (\delta+) atom already bonded to a highly electronegative atom (like O, N, or F) in one molecule, and a partial negative site (\delta-) on another electronegative atom (often O or N) in a different molecule (or sometimes within the same large molecule).
This interaction specifically involves the lone electron pairs on the \delta- atom.
It is not a full ionic or covalent bond; individually, hydrogen bonds are relatively weak (about 5-10% the strength of a covalent bond), but they are collectively strong when many are present, contributing significantly to the stability of macromolecules and the unique properties of water.
Why hydrogen bonds form in water and biomolecules:
Water is a highly polar molecule due to the bent structure and the two O-H polar bonds. The oxygen atom bears a \delta- charge, and the hydrogen atoms bear \delta+ charges. This strong polarity drives strong attractions between individual water molecules and with other polar or charged species, enabling extensive hydrogen bonding.
Each water molecule can form up to four hydrogen bonds with other water molecules (two through its hydrogens and two through its oxygen's lone pairs).
Comparison of strength:
Hydrogen bonds are significantly weaker than both covalent (which involves sharing electrons) and ionic bonds (which involves a complete transfer of electrons and strong electrostatic attraction). However, the vast number of hydrogen bonds that can form in systems like liquid water or between strands of DNA creates powerful overall interactions and structural stability.
Dynamic nature:
Hydrogen bonds are constantly forming, breaking, and reforming within liquid water, with each bond lasting only a few picoseconds (10^{-12} seconds). This dynamic nature allows for fluidity and rapid molecular rearrangements, yet the collective presence of many such bonds maintains overall structural integrity and stability over time.
Consequences of hydrogen bonding in water:
Water as a universal solvent for polar and charged substances (hydrophilic molecules). Its polarity allows it to surround and dissolve various solutes.
Water provides cohesion (water–water attraction) and adhesion (water–surface attraction).
Surface tension: a phenomenon where the cohesive forces between water molecules at the surface are stronger than those below the surface, creating a