Chapter 02 Notes: Atoms, Ions, and Molecules
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- Chapter 02: Atoms, Ions, and Molecules
- Focus: foundational chemistry for physiology.
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- Body chemistry underlies all physiological reactions (movement, digestion, heart pumping, nervous system).
- Chemistry split into: 1) Basic chemistry 2) Biochemistry.
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- Matter has mass and occupies space; 3 forms: solid, liquid, gas.
- Weight = matter × gravity.
- An atom is the smallest particle exhibiting chemical properties of an element.
- Elements are organized in the periodic table.
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- All matter is composed of elements.
- Four elements make up ~96% of the body: C, O, H, N.
- About 20 others are present; 118 elements recognized; 92 occur in nature.
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- Periodic table features: Chemical symbol, Atomic number, Average atomic mass.
- Trends: Electronegativity increases from left to right and from bottom to top within a group.
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- Most Common Elements of the Human Body
- Major elements (≈99% body weight): O 65 ext{%}, C 18.5 ext{%}, H 9.5 ext{%}, N 3.0 ext{%}, Ca 1.5 ext{%}, P 1.0 ext{%}, S 0.25 ext{%}, K 0.20 ext{%}, Na 0.15 ext{%}, Cl 0.15 ext{%}, Mg 0.05 ext{%}, Fe 0.006 ext{%}.
- Minor elements: remaining trace amounts.
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- Atoms have three subatomic particles:
- Protons (p): mass 1 ext{ amu}, charge +1, in nucleus.
- Neutrons (n): mass 1 ext{ amu}, charge 0, in nucleus.
- Electrons (e⁻): mass rac{1}{1800} ext{ amu}, charge -1, in electron orbitals.
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- Chemical symbol: unique to element (often first letter, sometimes with another letter).
- Atomic number Z: number of protons (located above symbol).
- Atomic mass (mass number) A: protons + neutrons; shown below symbol.
- Elements arranged by atomic number.
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- Subatomic particle counts:
- Proton number = atomic number Z.
- Neutron number =
A - Z. - Electron number (neutral atom) = Z.
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- Atoms have electron shells surrounding the nucleus.
- Each shell has a specific energy level and capacity:
- 1st shell: up to 2 electrons.
- 2nd shell: up to 8 electrons.
- Inner shells fill first.
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- Nuclear and cloud views of atomic structure:
- Nucleus: protons (p⁺) and neutrons (n⁰).
- Electron cloud/shells: electrons in orbitals around nucleus.
- Protons and neutrons ≈ 1 ext{ amu} each; electrons much lighter.
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- Isotopes: same element (same Z) but different neutron number; same chemical properties, different atomic masses.
- Example: Carbon has ^{12} ext{C}, ^{13} ext{C}, ^{14} ext{C}.
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- Half-lives:
- Physical half-life: time for 50 ext{%} of a radioisotope to decay.
- Biological half-life: time for half of a substance to be eliminated from the body; applies to hormones/drugs as well.
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- Molecules vs. compounds:
- Molecule: 2+ atoms bonded together.
- Compound: molecule composed of 2+ different kinds of atoms.
- Diatomic molecules are made of two atoms of the same element (e.g., ext{H}2, ext{O}2).
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- Chemical stability and the Octet Rule:
- Periodic table columns reflect number of valence electrons in outer shell.
- Column IA (e.g., H, Li, Na, K) has one electron in valence shell.
- Column VIIA has a full valence shell; noble gases are inert.
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- Octet rule:
- Elements tend to lose, gain, or share electrons to achieve a full outer shell of eight electrons (an octet).
- Some elements already have full octets and are stable/unreactive (noble gases in VIIIA).
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- Periodic table organization by valence electrons (visual concept).
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- 2.2 Ions and Ionic Compounds:
- Ionic compounds are stable associations of ions in a lattice bound by ionic bonds.
- Distinct from covalently bonded molecules.
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- Ions:
- Cations (positive) and anions (negative) formed by loss/gain of electrons.
- Important physiological roles (electrolyte balance, nerve/m muscle function).
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- Formation of cations: loss of electrons creates positive charge (e.g., Na → Na⁺).
- Example: Na has 11 protons and 10 electrons after losing one electron; charge +1.
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- Formation of anions: gain of electrons creates negative charge (e.g., Cl → Cl⁻).
- Example: Cl has 17 protons and 18 electrons after gaining one electron; charge -1.
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- Ionic bonds:
- Attraction between cations and anions.
- Salts (e.g., ext{NaCl}, ext{MgCl}_2) form lattices.
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- Formation of an ionic bond (NaCl): Na donates an electron to Cl; resulting ions Na⁺ and Cl⁻ attract and form NaCl.
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- 2.3 Covalent Bonding, Molecules, and Molecular Compounds:
- Covalent bonds: atoms share electrons.
- Molecular compounds: molecules with two or more different kinds of atoms.
- Example: ext{CO}2; ext{O}2 is a molecule but not a molecular compound (same element).
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- Chemical formulas:
- Molecular formula indicates number and type of atoms (e.g., ext{H}2 ext{CO}3).
- Structural formula shows arrangement (e.g., ext{O=C=O} for carbon dioxide).
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- Isomers:
- Glucose, galactose, and fructose share ext{C}6 ext{H}{12} ext{O}_6 but differ in arrangement.
- Isomers can have different chemical properties.
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- Covalent bonds form when atoms need electrons; common in biology: H, O, N, C.
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- Covalent bond capacities:
- H forms 1 bond; O forms 2; N forms 3; C forms 4.
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- Types of covalent bonds:
- Single bond: one pair of electrons (H–H).
- Double bond: two pairs (O=O).
- Triple bond: three pairs (N≡N).
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- Visuals of single, double, and triple covalent bonds.
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- Carbon needs four electrons to satisfy the octet; can form multiple covalent bonds to achieve this.
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- Carbon skeleton formation:
- Carbon bonds form straight chains, branched chains, or rings.
- Hydrogens saturate remaining valence.
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- Carbon skeleton arrangements (visual): consistent theme across forms.
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- Nonpolar vs. polar covalent bonds:
- Electronegativity determines sharing.
- Equal sharing → nonpolar covalent bond.
- Unequal sharing → polar covalent bond.
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- Electronegativity trends (common in biology):
- Hydrogen < Carbon < Nitrogen < Oxygen (increasing across a row and up a column).
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- Partial charges: more electronegative atom → partial negative charge
eg; less electronegative → partial positive + - Polar bonds may have a dipole moment; exception: C–H bond is often treated as nonpolar.
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- 2.3c Nonpolar, Polar, and Amphipathic Molecules:
- Nonpolar: bonds are nonpolar covalent (e.g., ext{O}- ext{O}, ext{C}- ext{H}).
- Polar: polar covalent bonds (e.g., ext{O}- ext{H} in water).
- CO₂: nonpolar overall due to canceling polar bonds.
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- Amphipathic molecules: large molecules with both polar and nonpolar regions (e.g., phospholipids).
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- Illustrations contrasting nonpolar, polar, and amphipathic regions in molecules (phospholipid / glycerol examples).
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- 2.3d Intermolecular Attractions:
- Intermolecular forces are weak but influence shape and behavior of large molecules (DNA, proteins).
- Hydrogen bonds form between polar molecules and contribute to water properties.
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- Hydrogen bonding example: glucose–water interactions produce multiple H-bonds that stabilize structure.
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- Other intermolecular attractions:
- Induced dipoles in nonpolar molecules.
- Hydrophobic interactions drive aggregation in polar environments.
- Intramolecular attractions occur within a large molecule.
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- 2.4 Molecular Structure and Properties of Water:
- Distinction: organic vs inorganic molecules; water is inorganic.
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- Water structure:
- Polar molecule: ext{H}_2 ext{O} with two partial positive hydrogens and a partial negative oxygen.
- Can form four hydrogen bonds per molecule.
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- Water illustration: shows polarity and hydrogen bonding network.
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- 2.4b Properties of Water:
- Phases of water: gas, liquid, solid depend on temperature.
- Water is liquid at room temperature due to hydrogen bonding.
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- Functions of liquid water:
- Transports dissolved substances; lubricates; cushions; excretes wastes.
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- Cohesion, surface tension, adhesion:
- Cohesion due to hydrogen bonding.
- Surface tension at air-water interface.
- Adhesion is water adhering to other substances.
- Surfactants prevent alveolar collapse in lungs.
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- High specific heat and high heat of vaporization:
- Water resists temperature change; helps stabilize body temperature.
- Heat of vaporization is high due to H-bonds; sweating cools the body.
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- Continued: heat of vaporization; sweating as cooling mechanism.
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- 2.4c Water as the Universal Solvent:
- Water dissolves many substances; termed universal solvent.
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- Substances that dissolve in water:
- Hydrophilic: polar molecules and ions; hydration shells form around solutes.
- Nonelectrolytes dissolve but do not dissociate (e.g., glucose, alcohol).
- Electrolytes dissolve and dissociate (e.g., NaCl); conduct electricity.
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- Substances that do not dissolve in water:
- Hydrophobic (nonpolar) molecules; require carrier proteins for transport in blood.
- Fats and cholesterol are poorly soluble in water.
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- Hydration shells and amphipathic interactions illustrated with phospholipid bilayers and micelles.
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- 2.5a Water: A Neutral Solvent:
- Water autoionizes: ext{H}_2 ext{O}
ightleftharpoons ext{H}^+ + ext{OH}^- - Typical ion concentration: ~10^{-7} ext{ M}, yielding neutral pH.
- Hydronium ion: ext{H}_3 ext{O}^+.
- Water autoionizes: ext{H}_2 ext{O}
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- 2.5b Acids and Bases 1:
- Acid: dissociates to produce ext{H}^+ and an anion; proton donor.
- Strong vs weak acids (e.g., ext{HCl} vs carbonic acid in blood).
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- 2.5b Acids and Bases 2:
- Base: accepts ext{H}^+; proton acceptor.
- Strong vs weak bases (e.g., ammonia, bicarbonate).
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2.5c pH, Neutralization, and Buffers 1:
- pH measures relative amount of ext{H}^+ in solution; scale 0–14.
- Plain water pH = 7; relation: pH = -
\log_{10}[ ext{H}^+ ].
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- 2.5c pH scale 2:
- Neutral: pH =7; Acidic: pH < 7; Basic: pH > 7.
- Each one-unit pH change represents a 10-fold change in [ ext{H}^+].
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- pH scale examples and common values (e.g., lemon juice ~3, blood ~7.4, NaOH ~14).
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- 2.5c Buffers and Neutralization:
- Buffers resist pH changes by buffering excess acid/base; carbonic acid/bicarbonate buffer in blood maintains pH 7.35–7.45.
- Neutralization occurs when acids/bases are countered by opposite species.
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- Mixtures (Colloids/Emulsions): heterogeneous mixtures with non-uniform distribution.
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- 2.6 Water Mixtures:
- Mixtures are formed without chemical changes; components can be separated by physical means.
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- 2.6a Categories: Suspension – large solutes, settles out; Blood cells in plasma; appears cloudy.
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- Colloid: smaller particles; remains mixed; scatters light.
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- Solution: true solution – homogeneous, dissolved solutes < 1 nm; does not scatter light; may not settle.
- Emulsion: water and nonpolar liquid (e.g., oil) that do not mix unless shaken.
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- Definitions: Solvent is the greatest amount; solvent is usually water in body; solutes are dissolved substances.
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- Visual: mixtures, emulsions, and micelles illustrating light scattering and phase behavior.
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- 2.6b Expressions of Solution Concentration 1:
- Concentration via Mass/Volume (solutes per volume).
- Mass/volume percent (g solute per 100 mL solution).
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- 2.6b Expressions of Solution Concentration 2:
- Molarity M = rac{ ext{moles solute}}{V( ext{L})}$$; temperature-dependent; common in lab use and physiology.
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- Organic Compounds:
- All contain carbon.
- Major classes: carbohydrates, lipids, proteins, nucleic acids.
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- 2.7a Biological Macromolecules – General:
- Large organic molecules synthesized by the body.
- Contain C, H, O; may contain N, P, S.
- Carbon skeletons include hydrocarbons and functional groups; many polar and capable of hydrogen bonding.
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- Polymers:
- Monomers linked covalently form polymers.
- Carbohydrates (sugar monomers), nucleic acids (nucleotide monomers), proteins (amino acid monomers).
- Dimers form when two monomers bond.
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- 2.7a Dehydration synthesis and hydrolysis:
- Dehydration synthesis (condensation): subunits join with loss of water.
- Hydrolysis: water is used to split polymers.
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- 2.7b Lipids – Overview:
- Nonpolar, water-insoluble; energy storage, membranes, hormones.
- Four classes: triglycerides, phospholipids, steroids, eicosanoids.
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- Triglycerides: long-term energy storage; formed from glycerol + 3 fatty acids; variability in chain length and saturation.
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- Triglyceride structure and synthesis/hydrolysis (lipogenesis/lipolysis):
- Dehydration synthesis builds triglycerides; hydrolysis breaks them down.
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- 2.7b Lipids 3: Phospholipids – amphipathic; polar (hydrophilic) head and nonpolar (hydrophobic) tails; form membranes.
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- 2.7b Lipids 4: Steroids – four fused rings; cholesterol as membrane component and hormone precursor.
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- 2.7b Lipids 5: Eicosanoids – 20-carbon fatty acids; local signaling in inflammation and nervous system.
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- Clinical View: Fatty Acids — Saturated, Unsaturated, and Trans fats
- Most animal fats are saturated (solid at room temperature).
- Most vegetable fats are unsaturated (usually liquid).
- Partial hydrogenation can create trans fats; associated with increased cardiovascular risk.
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- 2.7c Carbohydrates 1:
- Carbohydrates contain carbon, hydrogen, and oxygen in a typical ratio; general formula is often a multiple of CH₂O.
- Monosaccharides, disaccharides, polysaccharides.
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- 2.7c Carbohydrates 2: Glucose and glycogen
- Glucose: primary energy source.
- Glycogen: stored glucose in liver and muscle; glycogenesis and glycogenolysis; gluconeogenesis in liver.
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- Figure: Glucose and glycogen including glycogenesis and glycogenolysis processes.
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- 2.7c Carbohydrates 3: Other carbohydrates
- Hexoses (glucose, galactose, fructose) – isomers.
- Pentose sugars (ribose, deoxyribose).
- Disaccharides (sucrose, lactose, maltose).
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- 2.7c Carbohydrates 4: Polysaccharides and plant storage
- Glycogen in animals; starch in plants; cellulose in plants (fiber).
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- Visuals of simple carbohydrates and disaccharides (glucose, galactose, fructose; ribose, deoxyribose; sucrose, lactose, maltose).
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- 2.7d Nucleic Acids 1: DNA and RNA
- Polymers of nucleotides; phosphodiester bonds.
- Functions: store and transfer genetic information.
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- 2.7d Nucleic Acids 2: Nucleotide structure
- Sugar (pentose), phosphate group, nitrogenous base.
- Bases bind to sugar at the 1′ position; phosphate attaches at 5′.
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- 2.7d Nucleic Acids 3: Nitrogenous bases
- Pyrimidines: C, U (RNA), T (DNA).
- Purines: A, G.
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- 2.7d Nucleic Acids 4: DNA
- Double-stranded, in nucleus and mitochondria.
- Bases pair: A with T; G with C; deoxyribose; phosphate backbone; hydrogen bonds between bases.
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- 2.7d Nucleic Acids 5: RNA
- Single-stranded; ribose; bases A, G, C, U; no thymine.
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- 2.7d Nucleic Acids 6: ATP
- Adenosine triphosphate; energy carrier; three phosphate groups; last two phosphates with high-energy bonds; hydrolysis releases energy.
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- 2.7d Nucleic Acids 7: Other nucleotide-containing molecules
- NAD⁺ and FAD participate in ATP production.
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- 2.7e Proteins 1: Protein functions
- Catalysis (enzymes), structure, movement, transport, membranes, protection (antibodies).
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- 2.7e Proteins 2: General protein structure
- Polymers of amino acids; 20 different amino acids; amino group and carboxyl group on each amino acid; side chain (R) distinguishes amino acids.
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- 2.7e Proteins 3: Peptide bonds and protein formation
- Amino acids join via peptide bonds during dehydration synthesis.
- N-terminal end has free amine; C-terminal end has free carboxyl.
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- 2.7e Proteins 4: Oligopeptides and polysaccharides
- Oligopeptides (3–20 amino acids); longer chains form polypeptides/proteins.
- Glycoproteins: proteins with carbohydrate attached; ABO blood groups are glycoproteins on erythrocytes.
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- 2.8b Amino Acid Sequence and Protein Conformation 1
- Primary structure: linear sequence of amino acids.
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- 2.8b Amino Acid Sequence and Protein Conformation 2
- Conformation: 3D shape essential for function; folding guided by chaperones.
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- 2.8b Amino Acid Sequence and Protein Conformation 3
- Intramolecular interactions:
- Hydrophobic exclusion
- Hydrogen bonds (polar R groups)
- Ionic bonds (opposite charges)
- Disulfide bonds (cysteine)
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- 2.8b Amino Acid Sequence and Protein Conformation 4
- Secondary structures: α-helix and β-sheet.
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- 2.8b Amino Acid Sequence and Protein Conformation 5
- Tertiary structure: 3D shape of single poly-peptide; globular vs fibrous proteins.
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- 2.8b Amino Acid Sequence and Protein Conformation 6
- Quaternary structure: multiple polypeptide chains in a protein (e.g., hemoglobin).
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- 2.8b Denaturation
- Loss of conformation; often irreversible; temperature rise common cause.
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- 2.8b Denaturation (continued)
- pH changes can denature proteins; disrupts electrostatic interactions and bonds; severe pH changes can be lethal.
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- Summary note: Protein structure levels (primary to quaternary) determine function; denaturation disrupts function.
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- Additional emphasis: denaturation can result from heat or pH shifts; stability is essential for biological activity.
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- Concluding idea: understanding molecular structure helps explain function and pathology in physiology.