BIO Exam One LO's
Fundamentals: Elements, Atoms, Bonding, and Basic Biochemical Principles
Essential elements in living organisms
Most organisms require a small set of essential elements often summarized as carbon (C), hydrogen (H), nitrogen (N), oxygen (O), phosphorus (P), sulfur (S) (the CHNOPS elements), plus trace elements needed in smaller amounts (e.g., iron, iodine, calcium, potassium, magnesium, zinc, copper, manganese).
These elements form the chemical basis of biology and drive all macromolecular structures and metabolic processes.
The cell vs. the atom
The cell is the smallest unit of life capable of carrying out all life processes.
The atom is the smallest unit of an element that retains its properties and is the basic unit of chemistry.
Atomic structure: given protons, electrons, and neutrons
Mass number: A = Z + N where Z = ext{protons (atomic number)} and N = ext{neutrons}.
If the atom is neutral, the number of electrons equals the number of protons: E = Z.
If the atom is an ion, E
eq Z and the charge reflects the difference.Common reference: atomic number increases left-to-right, protons define the element identity; mass number is a sum of protons and neutrons.
Valence electrons and bonding capacity (outer shell electrons)
The number of valence electrons determines how many covalent bonds an atom can form (to complete its outer shell).
General rules (for biologically common elements):
Hydrogen: can form up to 1 bond (1 valence electron).
Carbon: up to 4 bonds (4 valence electrons).
Nitrogen: up to 3 bonds (5 valence electrons, typically forms 3 bonds with one lone pair).
Oxygen: up to 2 bonds (6 valence electrons, typically forms 2 bonds).
Fluorine: up to 1 bond (7 valence electrons).
Note: actual bonding patterns can include double bonds (shared pairs) and resonance, but the valence concept helps predict bonding capacity.
What the numbers around an element's symbol mean
The standard isotopic notation uses mass number A and atomic number Z (often written as
^{A}_{Z}X, where X is the element symbol).Example: 2H (deuterium) indicates mass number 2 for isotope of hydrogen; both mass number and atomic number help identify isotopes.
The same element (same Z) can have different mass numbers (A) if neutrons differ.
Isotopes: definition and differences
Isotopes are variants of the same element with the same number of protons (Z) but different numbers of neutrons.
Chemical properties are largely similar (because Z and electron count are similar), but physical properties (mass, stability, decay) differ.
Examples: carbon-12 vs carbon-14; they differ in N but not in Z; mass number A differs.
Practical distinctions: medical imaging (e.g., PET tracers like ${}^{18}\text{F}$), radiometric dating (e.g., ${}^{14}\text{C}$), etc.
Real-life uses of isotopes
Carbon-14 dating for archaeological and geological samples.
Stable isotopes as tracers in metabolic studies.
Medical imaging and diagnostic tracers (e.g., ${}^{18}\text{F}$ in PET scans, ${}^{131}\text{I}$ for thyroid imaging/treatment).
Predicting the number of bonds from valence electrons
Atoms tend to form as many bonds as needed to reach a stable, often octet (or duet for H) in the valence shell.
Examples: C forms up to 4 bonds; O forms up to 2 bonds; N forms up to 3 bonds; H forms 1 bond.
Multiple bonds (double/triple) count as more than one bond per atom involved in the sharing.
Types of atomic bonds: overview
Nonpolar covalent bonds: equal sharing of electrons; typically electronegativity difference ~0.
Polar covalent bonds: unequal sharing of electrons; moderate electronegativity difference.
Ionic bonds: transfer of electrons creating ions that attract; large electronegativity difference.
Formation rules in brief:
Nonpolar covalent: e.g.,
H–H, C–H, C–C in many hydrocarbons.
Polar covalent: e.g.,
C–O, C–N, O–H, N–H (differences in electronegativity create partial charges).
Ionic: e.g., NaCl (sodium donates an electron to chlorine).
Examples: NaCl (ionic), H2O (polar covalent with strong hydrogen bonding potential).
Bond types among biologically common elements (C, H, N, O, F)
Relative electronegativities: F > O > N > C > H.
Predicted bond types with carbon:
C–H: largely nonpolar covalent.
C–N: polar covalent.
C–O: polar covalent.
C–F: highly polar covalent.
Hydrogen and hydroxyl bonds: O–H and N–H bonds are typically polar covalent and participate in hydrogen bonding.
Molecular properties based on bonding
Nonpolar covalent molecules: typically hydrophobic, little to no charge, little interaction with water.
Polar covalent molecules: have partial charges, can dissolve in water (hydrophilic) and form dipole interactions.
Ionic compounds: exist as ions in solution, highly soluble in water, conduct electricity when dissolved or melted.
Overall polarity and charge influence solubility, interaction with membranes, and reactivity.
Hydrogen bonding: definition and origin
A hydrogen bond is a strong dipole-dipole interaction formed when a hydrogen atom covalently bonded to a highly electronegative atom (N, O, or F) experiences attraction to a lone pair on another electronegative atom.
Partial charges drive the interaction: δ+ on H, δ− on the electronegative donor/acceptor atom.
Hydrophilic vs. hydrophobic properties
Hydrophilic: polar or charged groups that interact favorably with water (e.g., hydroxyl, carboxyl, amino groups).
Hydrophobic: nonpolar groups (e.g., long hydrocarbon chains) that minimize contact with water.
Water: unique properties and life science implications
Water as a solvent: dissolves many ionic and polar substances, enabling chemistry of life.
Adhesion and cohesion: due to hydrogen bonding, supports capillary action and surface tension.
Temperature moderation: high specific heat and high heat of vaporization help stabilize temperatures in organisms and environments.
Expansion upon freezing: ice is less dense than liquid water due to hydrogen-bonded lattice, allowing ice to float.
Hydrogen bonding underlies these properties and the solvent behavior of water in biology.
Hydrogen bonding and water properties in detail
Hydrogen bonds create an extensive network in liquid water, providing cohesion and high surface tension.
The network requires energy to break, giving water a high heat capacity and heat of vaporization.
The polarity of water enables solvent capability for many dissolved solutes.
Acids, bases, and pH without chemistry-specific expertise
Acids donate protons (H⁺) to the solution; bases accept protons or donate OH⁻ in aqueous solutions.
In water, acids increase H⁺ concentration, lowering pH; bases increase OH⁻ concentration, raising pH.
Brønsted–Lowry definitions emphasize proton transfer; Arrhenius definitions emphasize H⁺ or OH⁻ production.
Buffer systems in biology and their importance
Buffers resist pH changes by neutralizing added acids or bases; typically consist of a weak acid and its conjugate base (or a weak base and its conjugate acid).
Henderson–Hasselbalch equation (common in biology):
\mathrm{pH} = \mathrm{p}Ka + \log{10}\left(\dfrac{[\mathrm{A^-}]}{[\mathrm{HA}]}\right)Biological buffers help maintain stable intracellular and extracellular pH, critical for enzyme activity and metabolic processes.
Calculating pH, pOH, and ion concentrations
Definitions and relationships:
\mathrm{pH} = -\log_{10}([\mathrm{H^+}])
\mathrm{pOH} = -\log_{10}([\mathrm{OH^-}])
[\mathrm{H^+}][\mathrm{OH^-}] = K_w = 1.0\times 10^{-14}\;\text{at }25^\circ\mathrm{C}
\mathrm{pH} + \mathrm{pOH} = 14 (at 25°C)
Inverse relationships:
[\mathrm{H^+}] = 10^{-\mathrm{pH}}
[\mathrm{OH^-}] = 10^{-\mathrm{pOH}}
Carbon’s bonding versatility
Carbon forms up to four covalent bonds, enabling a vast diversity of molecular frameworks (tetravalence).
Carbon-based bonding supports single, double, and triple bonds, enabling a variety of shapes and stability.
Hydrocarbon structure variations (summary)
Chain length (number of carbon atoms).
Branching (isomers with different branching patterns).
Rings (cyclic hydrocarbons).
Unsaturation (double or triple bonds).
These variations lead to vast diversity in lipid and carbohydrate structures as well as in degradation and metabolism.
Macromolecules: the four major classes
Proteins, nucleic acids, carbohydrates, lipids.
Monomers vs polymers
Monomer: the basic building block (e.g., amino acids, monosaccharides, nucleotides).
Polymer: a long chain of monomers linked by covalent bonds (e.g., proteins, polysaccharides, nucleic acids).
Given a molecule, determine whether it is a monomer or a polymer based on whether it consists of repeating monomer units connected by covalent bonds.
Dehydration (condensation) vs. hydrolysis reactions
Dehydration synthesis: links monomers by removing a water molecule, forming a covalent bond.
Hydrolysis: breaks covalent bonds by adding a water molecule, releasing monomers.
Monomer–polymer relationships for macromolecules
Proteins: monomer = amino acids; polymer = polypeptides/proteins.
Carbohydrates: monomer = monosaccharides (e.g., glucose); polymer = disaccharides and polysaccharides (e.g., sucrose, starch, glycogen, cellulose).
Nucleic acids: monomer = nucleotides; polymer = DNA or RNA.
Lipids: not true polymers; composed of components (e.g., glycerol + fatty acids in triglycerides; phospholipids = glycerol + two fatty acids + phosphate group + head group; steroids = cholesterols with fused rings).
How a small set of monomers yields many polymers
Different sequences, arrangements, and linkages of monomers produce diverse polymers (e.g., 26 letters forming many words; similarly, a finite set of monomers yields diverse macromolecules through sequence and structural diversity).
Identifying macromolecules from a molecule’s features
If the molecule shows repeating amino-acid units linked by peptide bonds, it is a protein.
If it consists of nucleotides linked by phosphodiester bonds, it is a nucleic acid.
If it consists of monosaccharide units linked by glycosidic bonds, it is a carbohydrate.
If it is composed of glycerol backbones with fatty acid chains (and possibly phosphate groups), it is a lipid.
Page 2: Lipids, Carbohydrates, and Proteins – Key Concepts
Why lipids are not true polymers (yet belong to lipid family)
Lipids lack repeating monomer units that form polymers in the same sense as proteins, nucleic acids, or carbohydrates.
They are grouped together by hydrophobicity and diverse functions (energy storage, membranes, signaling).
Are lipids built/broken down by dehydration/hydrolysis?
Yes, lipids can be assembled and broken down via condensation/dehydration reactions and hydrolysis, but they do not form true polymers with repeating units.
Three most important lipids and their cellular roles
Fats (triglycerides): energy storage, insulation, and cushioning.
Key components: glycerol backbone + three fatty acids.
Steroids (e.g., cholesterol): membrane fluidity regulation; precursors to steroid hormones.
Key component: four fused carbon rings; functional groups include cholesterol structure.
Phospholipids: main components of cell membranes; amphipathic (hydrophobic tails, hydrophilic heads).
Key components: glycerol, two fatty acids, phosphate group, head-group such as choline.
Carbohydrate classes and roles
Monosaccharides: simple sugars (e.g., glucose, fructose, galactose).
Disaccharides: two monosaccharides linked (e.g., lactose, sucrose, maltose).
Polysaccharides: many monosaccharides linked (e.g., starch, glycogen, cellulose).
Roles: energy storage (starch in plants, glycogen in animals); structural (cellulose in plants).
Comparing forms of glucose polymers
Cellulose: β-1,4-glycosidic bonds; structural in plants; largely indigestible by most animals due to lack of cellulase.
Starch: α-1,4 (and some α-1,6) glycosidic bonds; energy storage in plants; digestible by humans via amylases.
Glycogen: highly branched α-1,4 and α-1,6 linkages; energy storage in animals; highly soluble.
Linkage types dictate digestibility and structural roles.
Important monosaccharides and disaccharides
Monosaccharides: glucose, fructose, galactose.
Disaccharides: lactose (glucose + galactose), sucrose (glucose + fructose), maltose (glucose + glucose).
Protein diversity and functions: overview
Proteins perform a wide array of functions: enzymes, structural components, transport, signaling, immune defense, storage, and more.
Protein structure determines function: shape, charge distribution, and surface properties influence binding and catalysis.
Amino acid structure and sketch (text description)
A typical amino acid contains:
An amino group: (-NH_2)
A central carbon (α-carbon)
A carboxyl group: (-COOH)
A hydrogen atom: (-H)
A side chain (R group): variable, determines identity and properties.
In physiological conditions, the amino group is often protonated ((-NH_3^+)) and the carboxyl group deprotonated ((-COO^-)).
Diversity of R groups and protein function
R groups vary in size, charge, polarity, and shape, giving rise to diverse interactions: hydrophobic packing, hydrogen bonding, ionic interactions, van der Waals forces.
Amino acid substitutions and protein function
Substituting one amino acid for another can affect folding, structure, and function; substitutions in critical regions are more likely to disrupt activity (nonconservative substitutions tend to have larger effects).
Interactions with R groups: charge/polarity considerations
Positively charged side chains (e.g., Lys, Arg) interact with negatively charged groups; negatively charged side chains (Asp, Glu) interact with positively charged groups; polar uncharged side chains participate in hydrogen bonding.
Directionality of polypeptides: N-terminus and C-terminus
A polypeptide has a defined direction: N-terminus (amino end) and C-terminus (carboxyl end).
Protein structure levels and what stabilizes them
Primary structure: sequence of amino acids held together by peptide bonds.
Secondary structure: backbone—stabilized by hydrogen bonds between the carbonyl and amide groups; forms alpha helices and beta pleated sheets.
Tertiary structure: overall 3D shape stabilized by interactions among R groups (hydrophobic interactions, ionic bonds, hydrogen bonds, disulfide bridges).
Quaternary structure: arrangement and interactions between multiple polypeptide subunits.
Protein folding in different environments
In aqueous (polar) environments, hydrophobic residues tend to be buried inside the protein; hydrophilic residues tend to be on the surface.
In nonpolar environments, hydrophobic residues may be exposed on the surface, altering folding patterns.
Importance of proper protein folding
Correct folding is essential for function; misfolding can lead to loss of function or gain of toxic function (e.g., prions, certain neurodegenerative diseases).
Cellular helpers (chaperone proteins) assist in proper folding.
Protein denaturation and structure disruption
Denaturation disrupts non-covalent interactions (hydrogen bonds, ionic interactions, hydrophobic packing) and can unfold proteins.
Heat, extreme pH, and high salt concentrations commonly denature proteins, typically affecting secondary, tertiary, and quaternary structures while the primary sequence remains.
DNA and RNA: roles and functions in cells
DNA stores genetic information; serves as the template for replication and transcription.
RNA acts in transcription (mRNA), translation (tRNA and rRNA), and other regulatory and catalytic roles in cells.
Structure influences function: DNA is double-stranded, highly stable; RNA is usually single-stranded and more reactive.
Nucleotide structure: sketch described
Each nucleotide consists of:
A five-carbon sugar (deoxyribose in DNA; ribose in RNA).
A phosphate group.
A nitrogenous base (pyrimidines: C, T/U; purines: A, G).
In text form: sugar–phosphate backbone with attached bases.
DNA and RNA directionality: 5' and 3' ends
Strands are directional: 5' end has a phosphate group attached to the 5' carbon of the sugar; 3' end has a hydroxyl group on the 3' carbon.
Strands are read and synthesized in the 5'→3' direction in biology.
Chargaff's rules and DNA complementarity
In DNA, the amount of adenine equals thymine, and the amount of cytosine equals guanine: \%A = \%T, \; \%C = \%G.
This proportionality underpins the Watson–Crick base pairing (A with T via two hydrogen bonds; C with G via three hydrogen bonds).
Antiparallel nature of DNA
The two strands run in opposite directions: one strand is 5'→3' while its complement runs 3'→5'.
The ends of a double-stranded DNA molecule can be identified as 5' ends on one strand and 3' ends on the complementary strand.
Predicting a DNA strand given its complementary strand
Use base pairing: A ↔ T, C ↔ G (in DNA) or A ↔ U (in RNA) to determine the partner strand.
Directionality is opposite between the two strands (anti-parallel): if one strand runs 5'→3' left to right, the complementary runs 3'→5' left to right.
DNA nucleotides vs RNA nucleotides: key differences
Sugar: DNA uses deoxyribose (lacks 2' OH); RNA uses ribose (has 2' OH).
Bases: DNA uses adenine (A), thymine (T), cytosine (C), guanine (G); RNA uses adenine (A), uracil (U) in place of thymine, cytosine (C), guanine (G).
Structure: DNA is typically double-stranded and forms a stable double helix; RNA is typically single-stranded and can fold into various shapes.
Formula/notation recap (quick reference)
Mass number: A = Z + N
Neutrons: N = A - Z
pH/pOH: \mathrm{pH} = -\log{10}([\mathrm{H^+}]),\quad \mathrm{pOH} = -\log{10}([\mathrm{OH^-}]),\quad [\mathrm{H^+}][\mathrm{OH^-}] = K_w = 1.0 \times 10^{-14}\
\text{(at }25^{\circ}\text{C),}\quad \mathrm{pH} + \mathrm{pOH} = 14.
Page 3: DNA vs RNA Nucleotides (Summary)
Key distinction: DNA nucleotides vs RNA nucleotides
DNA nucleotides: deoxyribose sugar, bases A, T, C, G, phosphate group, typically form double-stranded molecules; designed for long-term storage of genetic information.
RNA nucleotides: ribose sugar, bases A, U, C, G, phosphate group, typically single-stranded; roles include transcription, translation, regulation, and catalysis in some contexts.
Quick Reference: Core Formulas and Concepts (Inline)
Mass number: A = Z + N
Atomic number: Z = ext{protons}
Neutrons: N = A - Z
Valence and bonding capacity: carbon (4), nitrogen (3), oxygen (2), hydrogen (1)
Bond types: nonpolar covalent, polar covalent, ionic
Hydrogen bonding criterion: H attached to N, O, or F forms H-bonds with lone pairs on other electronegative atoms
pH/pOH relationships: \mathrm{pH} = -\log{10}([\mathrm{H^+}]), \quad \mathrm{pOH} = -\log{10}([\mathrm{OH^-}]), \quad [\mathrm{H^+}][\mathrm{OH^-}] = K_w = 1.0 \times 10^{-14}, \quad \mathrm{pH} + \mathrm{pOH} = 14.
Henderson–Hasselbalch (buffer): \mathrm{pH} = \mathrm{p}Ka + \log{10}\left(\dfrac{[\mathrm{A^-}]}{[\mathrm{HA}]}\right)$$
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