Basic Chemistry and Biochemistry: Key Concepts for Anatomy & Physiology (Ch. 2)
2.1 Matter and Energy
Matter definition: anything that has mass and occupies space; can be seen, smelled, or felt.
Weight vs mass: weight = mass × gravity; mass is constant, weight changes with gravity (e.g., on the Moon vs Earth).
States of matter: solid (definite shape and volume), liquid (definite volume, variable shape), gas (variable shape and volume).
Energy definition: the capacity to do work or move matter; energy has no mass or volume.
Relationship between work and energy: the greater the work done, the more energy used.
Energy in biological systems comes in multiple forms, and energy can be transformed from one form to another.
Energy forms (4 major types mentioned):
Chemical energy: stored in bonds of chemical substances.
Electrical energy: movement of charged particles.
Mechanical energy: energy directly involved in moving matter.
Radiant (electromagnetic) energy: travels in waves (e.g., heat, visible light, UV, X-rays).
Energy concept in cells: chemical bonds store potential energy; when bonds are broken or rearranged, energy can be released to power cellular processes.
Energy conversions are not 100% efficient; some energy is lost as heat during transformations.
Practical relevance: understanding chemistry and biochemistry is essential for medical decisions (e.g., treating dehydration and fluid loss) because the body is a huge mixture of chemicals that underlie physiology.
Connection to physiology: physiology depends on chemistry; body functions (movement, digestion, heart pumping, nervous system) are driven by chemical reactions.
Quick example snippet: a stored chemical energy (e.g., in a battery or chemical bonds) can be converted to kinetic energy (movement) and electrical energy in a circuit or biological systems (nerve signaling).
2.2 Atoms and Elements (Part 1 of 3)
All matter is composed of elements; elements cannot be broken down into simpler substances by ordinary chemical methods.
Four elements make up about 96% of the body by mass: Oxygen (O), Carbon (C), Hydrogen (H), Nitrogen (N).
O = ~65.0%
C = ~18.5%
H = ~9.5%
N = ~3.2%
Lesser elements (~3.9%) include: Calcium (Ca) ~1.5%, Phosphorus (P) ~1.0%, Potassium (K) ~0.4%, Sulfur (S) ~0.3%, Sodium (Na) ~0.2%, Chlorine (Cl) ~0.2%, Magnesium (Mg) ~0.1%, Iodine (I) ~0.1%, Iron (Fe) ~0.1%.
Trace elements exist at <0.01% body mass (examples listed as Cr, Co, Cu, F, Mn, Mo, Se, Si, Sn, V, Zn); required in very small amounts.
Atomic symbols: one- or two-letter shorthand (e.g., O for oxygen, C for carbon); some come from Latin names (Na = natrium, K = kalium).
A table in the transcript summarizes major, lesser, and trace elements with approximate body mass percentages and functional notes.
Practical note: memorize the major elements and their approximate percent body mass for a foundational understanding of biochemistry in the human body.
2.2 Atoms and Elements (Part 2 of 3)
Atoms are the building blocks of elements; they are the smallest units that retain the properties of the element.
Protons (positive charge, +1), neutrons (neutral, 0 charge), and electrons (negative charge, −1).
Atomic nucleus contains protons and neutrons; electrons orbit the nucleus.
Protons and neutrons together make up about 1 atomic mass unit (amu) each; electrons have negligible mass.
Atoms are electrically neutral overall due to the balance of protons and electrons.
Historical/educational models of atomic structure:
Planetary model (older, simplified): electrons in fixed circular orbits around the nucleus.
Orbital model (current): electrons occupy regions called orbitals where they are most likely to be found; electron density forms an electron cloud around the nucleus.
Figure references in the slides illustrate the differences between these models and the concept of electron density.
2.2 Atoms and Elements (Part 3 of 3)
Atomic number (Z): number of protons in the nucleus; defines the identity of the element.
Denoted as a subscript to the left of the atomic symbol (e.g., _3Li).
Mass number (A): total number of protons and neutrons in the nucleus; written as a superscript to the left of the atomic symbol (e.g., ^7Li).
Isotopes: same atomic number (same number of protons) but different mass numbers (different number of neutrons).
Atomic weight: average of the masses of all isotopes of an element, weighted by their natural abundance.
Example concepts: hydrogen, helium, and lithium isotopes with varying neutron counts; atomic weights shift with isotopes.
Practical note: isotopes share chemical behavior with their stable forms, but their physical properties (mass, radioactivity) differ; some isotopes are radioactive and have medical/industrial applications.
2.3 Combining Matter (Molecules, Compounds, and Mixtures)
Molecule: two or more atoms bonded together.
Compound: a molecule composed of two or more different kinds of atoms bonded together; e.g., glucose, a molecule with formula
where the same element may appear in different ratios.Some molecules contain only one type of atom (e.g., , ) and are still called molecules but not compounds.
Most matter exists as mixtures: a physical combination of two or more components that are not chemically bound together.
Three basic types of mixtures:
Solutions: homogeneous mixtures; particles evenly distributed; solvent is usually the substance present in greatest amount (often water); solute(s) are dissolved substances (e.g., glucose in blood plasma).
Colloids (emulsions): heterogeneous mixtures with large solute particles that do not settle out; can look cloudy or milky; examples include cytosol sol-gel transformations and some gels like Jell-O.
Suspensions: heterogeneous mixtures with large solutes that do settle out (e.g., water and sand; blood cells in plasma if left undisturbed).
Concentration expressions for true solutions:
Percent concentration:
mg/dL: milligrams per deciliter (common in clinical measurements; e.g., normal fasting glucose around ).
Molarity (M): number of moles of solute per liter of solvent, , where 1 mole equals the molecular weight in grams, e.g., glucose with molecular weight ≈ 180.12 g/mol has a 1 M solution when 180.12 g is dissolved in enough water to make 1 liter.
Avogadro's number:
particles per mole; 1 mole contains this many molecules.Practical note: in the body, molar concentrations are typically small (often in the mM range).
Example molecular weights and symbols: glucose (C$6$H${12}$O$_6$) and common ions are used to illustrate solubility and buffers in physiology.
2.4 Chemical Bonds
Chemical bonds are not physical structures but energy relationships between electrons of reacting atoms; electrons are the key players in bonding and reactions.
Electron energy and shells: electrons occupy shells around the nucleus; shells have discrete energy levels and can hold a limited number of electrons.
Shell capacities (approximate): 1st shell up to 2 electrons, 2nd shell up to 8, 3rd shell up to 18; atoms can have up to 7 shells in theory.
Valence shell: outermost electron shell; electrons here have the highest potential energy and participate most in chemical bonding.
Octet rule (8-8 rule): atoms strive to achieve 8 electrons in their valence shell for stability. Exceptions: hydrogen (H) and helium (He) want 2 electrons in the first shell.
Bonding outcomes: atoms share, transfer, or pool electrons to reach a stable valence configuration.
Types of chemical bonds:
Ionic bonds:
Formed by transfer of valence electrons from one atom to another, creating ions (cations and anions).
Atoms with opposite charges attract to form ionic bonds; compounds are often salts.
Example: NaCl (sodium chloride) forming crystals when dry.
Covalent bonds:
Formed by sharing of valence electrons between atoms.
Single bonds (share 2 electrons), double bonds (4 electrons), triple bonds (6 electrons).
Purpose: each atom can fill its valence shell at least part of the time.
Types: polar covalent bonds and nonpolar covalent bonds.
Polar covalent bonds:
Unequal sharing of electrons due to atoms having different electronegativities.
Result in dipole moments; molecules become electrically polar (e.g., H–O in water).
Nonpolar covalent bonds:
Equal sharing of electrons; electrically balanced molecules (e.g., CO₂ in a linear arrangement).
Hydrogen bonds:
Not true bonds; attractive force between an electropositive hydrogen of one molecule and an electronegative atom of another molecule.
Common in dipoles like water; contribute to water’s liquid state; can also stabilize the 3D shapes of large molecules.
Visual models:
Covalent bonding often illustrated with structural formulas showing shared electron pairs and bond types (single, double, triple).
Polar vs nonpolar molecules have different shapes and properties due to unequal electron sharing and electronegativity.
Practical implications: bonding dictates molecule polarity, structure, reactivity, and function in physiology (e.g., water’s properties, macromolecule folding).
2.5 Chemical Reactions
Chemical reactions involve making, breaking, or rearranging chemical bonds to form products from reactants.
Chemical equations: symbolic representations of reactions with reactants on the left and products on the right; coefficients indicate relative amounts and must be balanced.
Molecular formulas: compounds represented as formulas like or ; subscripts show the number of atoms in a molecule; prefixes can denote non-bonded atoms or units when needed.
Common reaction types (3 main categories):
Synthesis (combination) reactions: atoms or molecules combine to form larger, more complex molecules; anabolic processes in biology.
General form: A + B → AB
Decomposition (bond-breaking) reactions: a molecule breaks down into smaller molecules or its constituent atoms; often catabolic (energy-releasing or requiring hydrolysis).
General form: AB → A + B
Exchange (displacement) reactions: bonds are both made and broken; components trade partners.
General forms: AB + C → AC + B or AB + CD → AD + CB
Redox reactions (oxidation-reduction): a subset of exchange reactions where electrons are transferred.
Oxidized = gains or loses electrons? (Oxidation is loss of electrons; reduction is gain of electrons).
Example: C₆H₁₂O₆ + 6 O₂ → 6 CO₂ + 6 H₂O + ATP; glucose is oxidized, oxygen is reduced; energy stored as ATP.
Energy flow in chemical reactions:
Exergonic (energy-releasing): products have less potential energy than reactants; often catabolic or oxidative processes.
Endergonic (energy-absorbing): products have more potential energy than reactants; often anabolic processes.
Reversibility and equilibrium:
Reactions are theoretically reversible: A + B ⇄ AB.
Chemical equilibrium occurs when forward and reverse reactions are balanced, and no net change occurs.
In biology, many reactions are not easily reversible due to energy requirements or product removal/denaturation of proteins.
Rate of chemical reactions (factors that influence speed):
Temperature: higher temperatures generally increase reaction rate.
Concentration of reactants: higher concentrations typically increase rate.
Particle size: smaller particles react faster due to larger surface area.
Catalysts: substances that increase reaction rate without being consumed; enzymes are biological catalysts that lower the activation energy of reactions.
Enzymes:
Biological catalysts; essential for most physiological reactions; they work by lowering the activation energy needed for a reaction to proceed.
Practical context: understanding reaction types and rates helps explain metabolism, energy production, and physiological regulation.
Key equations and concepts to remember:
Energy and work:
Idealized molarity definition: where n is moles and V is liters.
Avogadro’s number: molecules per mole.
1 mole of a substance has a mass equal to its molar mass in grams (e.g., glucose has molar mass ).
For chemical reactions: A + B → AB; AB → A + B; AB + C → AC + B; AB + CD → AD + CB.
Redox example:
Energy flow: exergonic vs endergonic definitions as described above.
Ethical, philosophical, and practical implications (brief):
Ethical: understanding biochemical reactions is essential for safe medical practices, including drug design and interpretation of treatments (e.g., dehydration therapies).
Philosophical: life is organized chemistry; emergent properties arise from molecular interactions and energy transformations.
Practical: knowledge of buffers, solutions, and reaction rates informs laboratory techniques, clinical chemistry, and diagnostic methods.
Quick connections to foundational principles:
Matter and energy underpin all physiology; the behavior of atoms and molecules governs cellular processes.
Bond formation and energy storage drive metabolism, muscle contraction, neural signaling, and homeostasis.
The concept of balance and equilibrium mirrors homeostatic regulation in the body.
Notes on terminology and examples from the transcript (glossary-style quick reference):
Solvent vs solute: solvent is the dissolving medium (often water); solute is dissolved substance (e.g., glucose in blood).
Solutions are typically homogeneous and transparent; colloids are heterogeneous and may appear cloudy; suspensions separate over time.
Colloids can undergo sol–gel transformations (e.g., cytosol as a sol–gel system).
Hydration and electrolyte balance tie directly to ionic bonds and the behavior of ions (Na⁺, K⁺, Cl⁻) in body fluids.
The octet rule guides most bonding behavior, with noble gases already possessing a full valence shell and thus low reactivity.