The Chemistry of Life

2.1 Atoms and Elements

• Atom: the smallest unit of matter that retains its original properties. It is the fundamental building block of all chemical substances.

• Atoms are composed of three basic subatomic particles with distinct charge, location, and mass:

  • Protons: Positively charged ($+1$), located in the nucleus, with a mass of approximately $1$ atomic mass unit (amu).

  • Neutrons: No charge ($0$), located in the nucleus, with a mass of approximately $1$ amu. They contribute to the atomic mass but not the charge.

  • Electrons: Negatively charged ($-1$), orbit the nucleus in specific electron shells, with a very small mass (approximately $1/1836$ amu, often considered negligible in terms of atomic mass).

• Electron shells:

  • Electrons occupy distinct energy levels or shells around the nucleus.

  • Each shell can hold a specific maximum number of electrons:

  • First shell (closest to the nucleus) = $2$ electrons.

  • Second shell = $8$ electrons.

  • Third shell can theoretically hold $18$ electrons, but in many biologically relevant atoms, it is considered "satisfied with only $8$" when it is the outermost shell, following the octet rule.

  • The outermost shell is called the valence shell, and the electrons in it (valence electrons) determine the atom's chemical reactivity.

• The atom is electrically neutral when the number of positively charged protons equals the number of negatively charged electrons. If these numbers differ, the atom becomes an ion.

• Element: a substance composed of one or more identical atoms, all having the same number of protons. The number of protons uniquely defines an element.

• Chemical notation:

  • Atomic Number ($Z$) = number of protons in the nucleus. This number is unique to each element and determines its identity.

  • Atomic Symbol = one- or two-letter shorthand for the element (e.g., H for Hydrogen, O for Oxygen, Na for Sodium).

  • Mass Number ($A$) = total number of protons + neutrons in the nucleus. It represents the approximate mass of the atom in amu.

  • Atomic Weight = the average mass of all naturally occurring isotopes of an element, weighted by their natural abundance. This is typically what's seen on the periodic table and is often expressed in amu or g/mol.

• Isotopes:

  • Variants of the same element that have different numbers of neutrons, but the same number of protons (and thus the same atomic number). This means they have the same chemical properties but different atomic masses.

  • Examples for carbon:

  • Carbon-12: $^{12}\text{C}$ (6 protons, 6 neutrons). Most common isotope.

  • Carbon-13: $^{13}\text{C}$ (6 protons, 7 neutrons). Stable isotope, often used in labeling studies.

  • Carbon-14: $^{14}\text{C}$ (6 protons, 8 neutrons). A radioisotope used in carbon dating due to its radioactive decay.

  • Isotopes share the same atomic number but differ in mass number due to the varying neutron count.

• Atomic weight concept:

  • The atomic weight listed on the periodic table is the average mass of all naturally occurring isotopes of an element, reflecting their relative abundance in nature. For example, chlorine has two major isotopes, $^{35}\text{Cl}$ and $^{37}\text{Cl}$, leading to an atomic weight of approximately $35.45$ amu.

• Radioisotopes:

  • Some isotopes are unstable because their nuclei contain an unfavorable proton-to-neutron ratio. These unstable isotopes undergo radioactive decay, releasing energy and/or subatomic particles (like alpha particles, beta particles, or gamma rays) to attain a more stable configuration.

  • Widely used in nuclear medicine as tracers (e.g., PET scans use fluorine-18) to visualize metabolic processes or as agents for radiation therapy (e.g., cobalt-60 for cancer treatment).

• Periodic table (overview):

  • Organizes elements systematically based on increasing atomic number and recurring chemical properties.

  • Elements are characterized by atomic number, atomic symbol, and atomic weight, and are grouped into periods (rows) and groups/families (columns).

  • It distinguishes metals (generally shiny, conductive, malleable, good electron donors) from nonmetals (tend to be poor conductors, brittle, good electron acceptors), and metalloids (intermediate properties).

  • The periodic table includes familiar groupings such as alkali metals (Group 1), alkaline earth metals (Group 2), halogens (Group 17), noble gases (Group 18, inert due to full valence shells), and transition metals.

• Summary concepts:

  • You should be able to describe atomic structure, including the location, charge, and relative mass of protons, neutrons, and electrons. You should also be able to identify the four major elements in the human body (Oxygen, Carbon, Hydrogen, Nitrogen).

  • Recognize metals vs nonmetals on the periodic table based on their general location and properties.

  • Define valence shell, atomic symbol, atomic number, mass number, and atomic weight accurately.

  • Understand what an isotope is and provide examples, describing how they differ in mass number but not atomic number.

  • You should also be able to determine the numbers of protons, neutrons, and electrons for a given atom or isotope, using its atomic and mass numbers.

2.3 Combining Matter: Mixtures & Chemical Bonds

• Mixtures: physical intermixture of two or more components (elements or compounds) that retain their original properties and can be physically separated. The chemical nature of the atoms involved does not change, only their physical association.

• Types of mixtures:

  • Suspensions: Heterogeneous mixtures with large, visible solutes that tend to settle out over time (e.g., blood cells in plasma, sand in water). They are often cloudy or opaque.

  • Colloids: Heterogeneous mixtures with smaller, generally invisible solutes that do not settle out. They often appear milky or translucent (e.g., milk, gelatin, cytoplasm in cells).

  • Solutions: Homogeneous mixtures consisting of very small, evenly distributed solutes dissolved in a solvent. They are clear and do not separate (e.g., saline solution, glucose in blood plasma).

• Solutions, solvents & solutes:

  • Solute: The substance being dissolved in a solution (e.g., salt in saltwater).

  • Solvent: The substance in which the solute dissolves, typically present in the largest amount (e.g., water in saltwater).

  • Solubility: The maximum amount of solute that can dissolve in a given amount of solvent at a specific temperature and pressure. It reflects the degree to which the solute dissolves in a solution.

  • A major factor in body transport and biological processes is a chemical’s solubility in water, as water is the universal biological solvent.

  • Concentration: The amount of solute present in a given volume or mass of a solution. It is often expressed as a percentage, molarity, or parts per million (ppm).

  • Example: A $40\%$ salt solution means there are $40$ grams of salt dissolved in $60$ grams of water (totaling $100$ grams of solution).

• Chemical bonds, molecules, and compounds:

  • Chemical bonds: Forces of attraction that hold atoms together to form molecules or compounds.

  • Molecules: Two or more atoms bonded together. These atoms can be of the same element (e.g., $O_2$) or different elements.

  • Compounds: Two or more atoms from different elements bonded together (e.g., $H_2O$). All compounds are molecules, but not all molecules are compounds.

  • Molecular formulas show the kinds and numbers of atoms in a molecule (e.g., $O<em>2$ for oxygen gas, $H</em>2O$ for water, $C<em>6H</em>{12}O_6$ for glucose).

• Valence electrons & chemical bonds:

  • Bonds form when valence electrons (electrons in the outermost shell) interact between atoms.

  • Atoms bond to achieve greater stability and minimize their overall energy. They do this by gaining, losing, or sharing electrons to fill their valence shells.

  • Octet rule: Atoms are most stable when they have $8$ electrons in their valence shell. (An exception is hydrogen and helium, which follow the duet rule, being stable with $2$ electrons in their first and only shell).

  • Atoms with filled valence shells are considered inert (non-reactive), like the noble gases (e.g., Neon has $8$ valence electrons, Helium has $2$).

  • Atoms not having a full valence shell will tend to react with other atoms to satisfy the octet rule (or duet rule for atoms with five or fewer electrons in their outer shell).

• Types of chemical bonds:

  • Ionic bonds: Occur between metals (which tend to lose electrons) and non-metals (which tend to gain electrons). Electrons are completely transferred from the metal atom to the non-metal atom, resulting in the formation of oppositely charged ions.

  • Cation: A positively charged ion formed when an atom loses one or more electrons (e.g., $Na^+$).

  • Anion: A negatively charged ion formed when an atom gains one or more electrons (e.g., $Cl^-$).

  • The electrostatic attraction between cations and anions forms the ionic bond (e.g., NaCl).

  • Covalent bonds: Form predominantly between non-metals. Electrons are shared between atoms to fulfill their duet or octet rules.

  • Can be single (one shared electron pair), double (two shared electron pairs), or triple (three shared electron pairs), depending on the number of electrons shared.

  • Single bond: C-C (sharing 2 electrons)

  • Double bond: C=C (sharing 4 electrons)

  • Triple bond: C$\equiv$C (sharing 6 electrons)

• Polar vs. non-polar bonds:

  • Nonpolar covalent bond: Electrons are shared equally between two atoms because their electronegativities are very similar or identical (e.g., $O<em>2$, $CH</em>4$ due to the similar electronegativity of C and H).

  • Polar covalent bond: Electrons are shared unequally between two atoms because one atom has a significantly higher electronegativity than the other, creating a partial positive charge ($\delta+$) on one end and a partial negative charge ($\delta-$) on the other (e.g., $H_2O$).

  • Electronegativity: The measure of an atom's ability to attract shared electrons in a chemical bond. Generally, electronegativity increases from the bottom-left to the top-right of the periodic table (Fluorine is the most electronegative element), with noble gases being an exception as they typically do not form bonds.

  • Metals typically have low electronegativity and tend to lose electrons, while non-metals generally have high electronegativity and tend to gain or strongly attract electrons.

• Non-polar vs polar molecules:

  • Non-polar molecules: Molecules where either the bonded atoms are of the same element (resulting in nonpolar covalent bonds, e.g., $O<em>2$), or the molecule is symmetrical, causing any bond dipoles to cancel each other out (e.g., $CO</em>2$). Carbon-hydrogen bonds also tend to be non-polar, making hydrocarbons non-polar molecules.

  • Polar molecules: Molecules that contain two or more nonmetals with substantially different electronegativities, and the molecule has an asymmetrical shape, leading to a net molecular dipole moment. Groups with highly electronegative atoms like Oxygen (O), Nitrogen (N), and Chlorine (Cl) often contribute to molecular polarity.

  • Examples: Ethanol ($CH<em>3CH</em>2OH$) is polar due to the presence of the -OH group. Octanol ($CH<em>3(CH</em>2)_7OH$) is also polar but less polar than ethanol due to its longer nonpolar hydrocarbon chain, which influences its solubility.

• Hydrogen bonds:

  • Special, weak intermolecular attractions that occur between a partially positively charged hydrogen atom (covalently bonded to a highly electronegative atom like O, N, or F) and a partially negatively charged electronegative atom (O, N, or F) in another polar covalent compound. These are not true chemical bonds where electrons are shared or transferred but rather electrostatic attractions.

  • Important for water properties (responsible for its high surface tension, cohesion, and adhesion), protein folding (stabilizing secondary and tertiary structures), and DNA structure (holding the two strands of the double helix together).

• Why polarity matters:

  • Polarity significantly affects how molecules interact with each other and their environment, influencing their movement, transport across biological membranes, and solubility within the body.

  • The principle of “like dissolves like” is critical: polar solvents (like water) dissolve polar and ionic solutes well, while non-polar solvents dissolve non-polar solutes.

• Summary concepts:

  • Distinguish mixtures (physical intermixture, components retain properties, separable) from chemical bonds (chemical combination, new substance formed, inseparable without chemical reaction).

  • Distinguish suspensions (large particles, settle), colloids (medium particles, do not settle, appear cloudy/milky), and solutions (small particles, homogeneous, clear), and give body examples for each.

  • Contrast ionic bonds (electron transfer, forming ions, metal-nonmetal) and covalent bonds (electron sharing, nonmetal-nonmetal), and describe the electron behavior in each.

  • Explain the octet/duet rule and how it applies to bond formation, affecting atomic stability.

  • Define solvent, solute, solubility, molecule, compound, ion, cation, anion, polarity, and electronegativity, providing clear distinctions.

  • Distinguish non-polar and polar covalent bonds based on electron sharing equality and determine the polarity of simple molecules based on bond polarity and molecular geometry.

  • Explain why polarity matters for biological systems, particularly regarding solubility and transport, and define hydrogen bonds, clarifying how they differ from ionic and covalent bonds.

2.3 Energy & Chemical Reactions

• Energy: The capacity to do work, which is defined as putting matter into motion.

• Two interconvertible classes of energy in the body:

  • Potential energy (energy of position): Stored energy that has the potential to do work. Examples include the chemical energy stored in molecular bonds (like ATP or glucose) or the energy stored in a stretched spring or a concentration gradient across a membrane.

  • Kinetic energy (energy of motion): Energy of an object in motion. Examples include muscle contraction, heat (random molecular motion), and the flow of ions across a membrane.

• Forms of energy in the body that are interconvertible:

  • Chemical energy: Stored in the bonds of chemical substances (e.g., ATP, glucose, fats). When these bonds are broken, energy is released, often converted into other forms.

  • Electrical energy: Results from the movement of charged particles (ions). Critical for nerve impulse transmission and muscle contraction in the body.

  • Mechanical energy: Directly involved in moving matter. For example, when muscles contract, they apply mechanical energy to move bones or pump blood.

  • Radiant (electromagnetic) energy: Travels in waves, like light, X-rays. Essential for vision and involved in vitamin D synthesis in the skin.

• Energy conservation:

  • The Law of Conservation of Energy (First Law of Thermodynamics) states that energy cannot be created or destroyed, only converted from one form to another.

  • Conversions are not $100\%$ efficient in biological systems; some usable energy is always lost as heat with each conversion, following the Second Law of Thermodynamics, which states that systems tend towards increased entropy (disorder).

• Chemical reactions:

  • Processes that involve the rearrangement of the atomic structure of molecules. Chemical bonds are broken in reactants, and new bonds are formed to create products.

  • Reactants are the starting substances, and products are the substances formed. They are shown in chemical equations, with an arrow indicating the direction of the reaction.

  • Example (simplified, an addition reaction for ethylene hydration): $CH<em>2=CH</em>2 + H<em>2O \rightarrow CH</em>3CH_2OH$ (Ethylene + Water $\rightarrow$ Ethanol)

• Three important types of chemical reactions in physiology:

  • Anabolic (synthesis) reactions: Build larger, more complex molecules from smaller ones. They typically require energy input (endergonic). Represented as: $A + B \rightarrow AB$. Example: protein synthesis from amino acids.

  • Catabolic (decomposition) reactions: Break down larger molecules into smaller ones. They typically release energy (exergonic). Represented as: $AB \rightarrow A + B$. Example: digestion of food, cellular respiration.

  • Exchange reactions: Involve both synthesis and decomposition. Parts of reacting molecules trade places, forming new products. Represented as: $AB + CD \rightarrow AD + BC$. Example: Neutralization reactions (acid + base $\rightarrow$ salt + water).

• Oxidation-reduction (redox) reactions:

  • A type of chemical reaction that involves the transfer of electrons between reactants. These reactions are fundamental to energy metabolism in the body.

  • Oxidation: The process where a reactant loses one or more electrons (or gains oxygen, or loses hydrogen). Oxidation often results in an increase in the oxidation state of an atom.

  • Reduction: The process where a reactant gains one or more electrons (or loses oxygen, or gains hydrogen). Reduction often results in a decrease in the oxidation state of an atom.

  • LEO says GER (Lose Electrons Oxidation, Gain Electrons Reduction). Oxidation and reduction always occur simultaneously.

• Energetics of reactions:

  • Endergonic reactions: Energy input is greater than the energy of the reactants; the products have more potential energy than the reactants. These reactions absorb energy from their surroundings and typically occur in anabolic processes (e.g., building complex molecules like proteins or ATP synthesis).

  • Exergonic reactions: Products have less potential energy than reactants; energy is released during the reaction. These reactions release energy into their surroundings and typically occur in catabolic processes (e.g., breakdown of glucose, ATP hydrolysis).

• Reversible reactions:

  • Reactions that can proceed in both the forward and reverse directions simultaneously. The net direction of the reaction depends on the concentrations of reactants and products, and the presence of catalysts.

  • They eventually reach a state of equilibrium where the rate of the forward reaction equals the rate of the reverse reaction.

  • Example: The carbonic acid–bicarbonate buffer system in blood is a crucial reversible reaction for pH regulation: $CO<em>2 + H</em>2O \rightleftharpoons H<em>2CO</em>3 \rightleftharpoons H^+ + HCO_3^-$. (Carbon dioxide + Water $\rightleftharpoons$ Carbonic acid $\rightleftharpoons$ Hydrogen ion + Bicarbonate ion).

• Reaction rates and activation energy:

  • All chemical reactions require an initial input of energy to start, known as activation energy ($E_a$). This energy is needed to break existing bonds in the reactants before new bonds can form in the products.

  • Factors that increase reaction rate by lowering activation energy or increasing the likelihood of collisions between reactant molecules:

  • Concentration: Higher concentration of reactants increases the frequency of collisions, thus increasing the reaction rate.

  • Temperature: Increased temperature typically raises the kinetic energy of molecules, leading to more forceful and frequent collisions, accelerating the reaction.

  • Properties of reactants: The physical state (phase), particle size, and chemical nature of reactants influence how easily they can interact. Smaller particles and gaseous/liquid reactants generally react faster.

  • Presence or absence of a catalyst: Catalysts are substances that increase the rate of a chemical reaction without being consumed in the process, primarily by lowering the activation energy.

• Enzymes (biological catalysts):

  • Proteins that function as highly specialized biological catalysts, enormously accelerating the rate of biochemical reactions within cells.

  • They reduce the activation energy required for a reaction to proceed and thus increase the reaction rate, often by millions of times compared to uncatalyzed reactions.

  • Enzymes are highly specific for particular substrates (the molecules they act upon) and reactions. Their specificity arises from the unique three-dimensional shape of their active site.

  • They are not consumed or altered in the reaction; after converting substrate to product, the enzyme is freed to catalyze another reaction.

  • Enzymes do not change the overall nature of the reaction (i.e., they do not change the energy of reactants or products, or the equilibrium constant); they only affect how quickly equilibrium is reached.

• Enzyme examples:

  • Ethanol metabolic pathway: When ethanol is consumed, it is metabolized in the body through a two-step process involving specific enzymes:

  • Enzymes: Alcohol dehydrogenase (ADH) efficiently converts ethanol to acetaldehyde.

  • Aldehyde dehydrogenase (ALDH2) then converts acetaldehyde to acetate.

  • Pathway: Ethanol $\rightarrow$ Acetaldehyde $\rightarrow$ Acetate. Acetaldehyde is toxic and causes many symptoms of hangover; a fast ADH and slow ALDH2 can lead to accumulation of acetaldehyde.

  • Acetylcholinesterase (AChE): An enzyme crucial for nervous system function, found in the synaptic cleft. It degrades the neurotransmitter acetylcholine (ACh) into acetate and choline to terminate neuronal signaling, allowing the postsynaptic membrane to repolarize and be ready for the next signal.

• Summary concepts:

  • Define activation energy, anabolic reaction, catabolic reaction, exchange reaction, and redox reaction, explaining their core characteristics and energy changes.

  • Explain the reversibility of chemical reactions and provide a relevant example like the carbonic acid-bicarbonate buffer system.

  • Explain how enzymes work to increase the rate of a chemical reaction (by lowering activation energy, increasing collision likelihood) and give body enzyme examples such as ADH, ALDH2, and AChE, describing their functions.

2.4 Inorganic Compounds: Water, Acids, Bases, and Salts

• Inorganic compounds: broad category of chemical compounds that typically do not contain carbon-hydrogen bonds. This category includes water, acids, bases, and salts. (Organic compounds, in contrast, are generally defined by the presence of carbon-hydrogen bonds).

• Water functions in the body:

  • Constitutes approximately $50-65\%$ of total body mass, making it the most abundant inorganic compound.

  • Roles:

  • Transport: Acts as a universal solvent, facilitating the transport of nutrients, gases, hormones, and waste products throughout the body (e.g., blood plasma is mostly water).

  • Lubrication: Reduces friction between body parts (e.g., synovial fluid in joints, serous fluid in body cavities).

  • Cushioning: Protects organs from trauma (e.g., cerebrospinal fluid cushions the brain and spinal cord, amniotic fluid cushions a fetus).

  • Excretion of wastes: Helps dissolve and flush out metabolic waste products via urine.

  • Regulation of body temperature: Due to its high heat capacity and high heat of vaporization, water helps maintain a stable internal body temperature.

• Properties of water:

  • Hydrogen bonding strongly influences water’s unique properties. The polarity of water molecules (oxygen is partially negative, hydrogens are partially positive) allows them to form hydrogen bonds with each other.

  • Surface tension: The inward pulling of cohesive forces at the surface of a liquid, creating a film-like effect. This is due to strong hydrogen bonding between water molecules. It is important for processes like breathing in the lungs (surfactant reduces surface tension).

  • Cohesion: The attraction between water molecules themselves, due to hydrogen bonding, allowing them to stick together.

  • Adhesion: The attraction between water molecules and other polar or charged substances, allowing water to cling to surfaces (e.g., water in blood vessels).

• Thermal properties of water:

  • High heat capacity: The amount of energy (in calories or joules) needed to raise the temperature of $1$ gram of a substance by $1\text{\degree C}$. Water has a high heat capacity, meaning it can absorb or release large amounts of heat with only a small change in its own temperature, which is crucial for thermoregulation in the body.

  • High heat of vaporization: The large amount of energy required for molecules to escape from a liquid to the gaseous phase. Again, due to strong hydrogen bonding, a lot of energy is needed to break these bonds and allow water to evaporate, which provides an effective cooling mechanism for the body through sweating.

  • Comparative specific heats (examples, in cal/g/°C):

  • Silver: $0.06$

  • Granite/Sand: $0.2$

  • Aluminum: $0.22$

  • Alcohol: $0.3$

  • Acetone: $0.51$

  • Pure liquid water: $1.00$ (benchmark, one of the highest among common substances)

  • Ammonia: $1.13$

• Water as solvent:

  • Hydrophilic solutes (literally “water-loving”): Substances that are polar or ionic and readily dissolve in water. Water molecules surround and separate these solutes (e.g., salt, sugar, ions).

  • Hydrophobic solutes (literally “water-fearing”): Substances that are non-polar and do not dissolve in water. Instead, they tend to aggregate together in water (e.g., oils, fats).

  • Amphipathic/amphiphilic molecules: Molecules that possess both polar (hydrophilic) and non-polar (hydrophobic) regions. These molecules do not dissolve completely in water but rather arrange themselves in specific ways, often forming micelles or bilayers (like phospholipids in cell membranes).

• Acids, bases, and pH:

  • Water auto-dissociation: Even pure water undergoes a slight dissociation into hydrogen ions (protons, $H^+$) and hydroxide ions ($OH^-$). In pure water, the concentration of both $[H^+]$ and $[OH^-]$ is approximately $1 \times 10^{-7}\text{ M}$ (moles per liter) at $25\text{\degree C}$.

  • pH scale:

  • pH measures the relative concentration of $H^+$ in a solution. It is inversely related to $[H^+]$: as $[H^+]$ increases, pH decreases.

  • Defined as $\text{pH} = -\log_{10}[H^+].$

  • The scale runs from $0$ to $14$ at $25\text{\degree C}$.

  • Neutral solutions have a pH of $7$ ($[H^+] = 1 \times 10^{-7}\text{ M}$), like pure water.

  • Acidic solutions have a pH less than $7$ ([H^+] > 1 \times 10^{-7}\text{ M}).

  • Basic (alkaline) solutions have a pH greater than $7$ ([H^+] < 1 \times 10^{-7}\text{ M}).

  • Each unit change on the pH scale corresponds to a tenfold change in $[H^+]$ concentration (e.g., pH $6$ has $10$ times more $H^+$ than pH $7$).

  • Acids: Proton donors; they dissociate in water to release $H^+$ ions, thereby increasing the $[H^+]$ concentration of the solution.

  • Stronger acids dissociate more completely in water (e.g., $HCl \rightarrow H^+ + Cl^-$), releasing a large amount of $H^+$.

  • Weaker acids dissociate less completely (e.g., carbonic acid), releasing fewer $H^+$.

  • Bases: Proton acceptors; they decrease $H^+$ concentration in a solution either by directly accepting $H^+$ ions or by releasing $OH^-$ ions that then combine with $H^+$ to form water.

  • Stronger bases dissociate more and bind more $H^+$ (e.g., $NaOH \rightarrow Na^+ + OH^-$).

  • Important weak base: bicarbonate ion ($HCO_3^-$), which can accept $H^+$ to form carbonic acid.

• Buffers:

  • Buffers are chemical systems (typically a weak acid and its conjugate base, or a weak base and its conjugate acid) that resist sudden or large changes in pH by either releasing $H^+$ when pH rises or binding $H^+$ when pH falls.

  • The urinary and respiratory systems play critical roles in helping to maintain the body's pH balance through physiological buffering mechanisms.

  • The carbonic acid–bicarbonate buffer system ($H<em>2CO</em>3/HCO_3^-$) in blood is the most important buffer, working to maintain blood pH within a narrow, life-sustaining range of $7.35-7.45$.

• Salts & electrolytes:

  • Salts: Ionic compounds formed from a cation and an anion, generally produced from the reaction of an acid and a base. When dissolved in water, salts dissociate into their constituent ions.

  • Electrolytes: Ions (cations and anions) formed when salts, acids, or bases dissociate in water. These ions are capable of conducting an electrical current in solution, making them vital for many physiological processes.

  • Roles of electrolytes:

  • Calcium salts (e.g., calcium phosphate) are major structural components of teeth and bones, providing hardness and rigidity.

  • Dissolved ions like sodium ($Na^+$), potassium ($K^+$), calcium ($Ca^{2+}$), and chloride ($Cl^-$) are critical for muscle contractions, nerve impulse transmission, fluid balance, and various enzymatic reactions within the body.

• Summary concepts:

  • Define acid, base, salt, and electrolyte, explaining their chemical nature and behavior in water.

  • Explain the pH scale, relating it to $H^+$ concentration, and identifying the values for neutral, acidic, and basic solutions.

  • Explain how buffers work chemically to maintain pH within a narrow range and mention the physiological systems (urinary, respiratory, bicarbonate buffer) that contribute.

2.5 Organic Compounds: Macromolecules

• Macromolecules overview:
  The primary classes of organic macromolecules essential for life are:

  • Carbohydrates

  • Lipids

  • Proteins

  • Nucleic Acids

• Representing hydrocarbons:

  • Hydrocarbons are organic compounds composed only of carbon and hydrogen atoms, forming the basic skeleton of many organic molecules. They can exist as straight chains, branched chains, or ring forms.

  • Examples of structural representations include linear structural formulas (e.g., $CH<em>3CH</em>2CH_3$ for propane), condensed formulas, and ring structures (e.g., benzene ring).

• Monomers and polymers:

  • Monomers: Single, repeating subunits that serve as the building blocks for larger molecules.

  • Polymers: Large, complex structures (macromolecules) composed of many identical or similar monomers covalently linked together.

  • Polymers are built by anabolic dehydration synthesis (also known as condensation reaction): In this process, a covalent bond is formed between two monomers with the removal of a water molecule ($H_2O$).

  • Polymers are broken down by hydrolysis: This catabolic process involves the addition of a water molecule ($H_2O$) across a covalent bond, breaking the polymer into its constituent monomers.

  • Each type of organic macromolecule (except for some lipids) has its own specific monomer and polymer forms.

Carbohydrates

• Structure of monosaccharides:

  • Monosaccharides are the simplest carbohydrates, often called simple sugars. They are characterized by a sweet taste and typically have a formula of ($CH_2O$)n.

  • Pentoses: Five-carbon sugars. Crucial components of nucleic acids (e.g., deoxyribose in DNA, ribose in RNA).

  • Hexoses: Six-carbon sugars. The most important energy source for cells (e.g., glucose, fructose, and galactose, which are isomers with the same molecular formula $C<em>6H</em>{12}O_6$ but different structural arrangements).

• Disaccharide formation:

  • Disaccharides are formed when two monosaccharides are covalently linked together via dehydration synthesis. A glycosidic bond is formed.

  • Example: Sucrose (table sugar) is formed from one glucose and one fructose molecule, with the removal of water.

  • Hydrolysis breaks disaccharides back into their constituent monosaccharides by adding a water molecule.

• Glycogen (a polysaccharide):

  • Glycogen is a complex polymer of many glucose monomers linked together, serving as the primary long-term energy storage molecule in animals, especially in the liver and muscles.

Lipids

• Structure of fatty acids:

  • Fatty acids are long hydrocarbon chains with a carboxyl group ($-COOH$) at one end. They are a primary component of many lipids.

  • Saturated fatty acids: Contain no carbon-carbon double bonds in their hydrocarbon chain. They are 'saturated' with hydrogen atoms, resulting in a straight chain structure, allowing them to pack tightly (e.g., palmitic acid, stearic acid).

  • Unsaturated fatty acids: Contain one or more carbon-carbon double bonds in their hydrocarbon chain. These double bonds create 'kinks' or bends in the chain, preventing tight packing.

  • Monounsaturated: One double bond (e.g., oleic acid).

  • Polyunsaturated: Two or more double bonds (e.g., linoleic acid, linolenic acid).

• Triglycerides:

  • Formed from a glycerol backbone chemically linked to three fatty acid chains via dehydration synthesis (ester bonds).

  • They are the major form of energy storage in adipose tissue (fat cells) and provide insulation and cushioning for organs.

• Phospholipids:

  • Consist of a glycerol backbone with two fatty acid tails and a phosphate-containing head group. The fatty acid tails are hydrophobic, while the phosphate head is hydrophilic.

  • This amphipathic nature is crucial for forming cellular membranes: phospholipids spontaneously arrange into a phospholipid bilayer in aqueous environments, with hydrophilic heads facing out and hydrophobic tails forming the core.

  • Example: Phosphatidylcholine, where the head group contains choline, is a common phospholipid in biological membranes.

• Steroids:

  • Characterized by a distinctive four-ring hydrocarbon structure called the steroid nucleus (or steroid core).

  • Cholesterol is a key steroid: it serves as a vital component of animal cell membranes, influencing membrane fluidity, and acts as a precursor molecule from which other important steroids (e.g., steroid hormones like testosterone, estrogen, cortisol, and vitamin D) are synthesized.

Proteins

• Structure of amino acids:

  • Amino acids are the monomers of proteins. Each amino acid possesses a central carbon atom (alpha-carbon) bonded to four groups:

  • An amino group ($-NH_2$)

  • A carboxyl group ($-COOH$)

  • A hydrogen atom ($-H$)

  • A variable side chain (R group): The unique chemical structure of the R group (of which there are 20 common types) determines the specific properties of each amino acid (e.g., nonpolar, polar, charged, acidic, basic).

  • Examples: Tryptophan (nonpolar due to its bulky aromatic R group), Cysteine (polar due to -SH group, important for disulfide bonds), Glycine (the simplest amino acid with an -H as its R group).

• Dipeptide formation:

  • Peptides are formed when two or more amino acids are linked together by peptide bonds.

  • Dehydration synthesis forms a peptide bond between the carboxyl group of one amino acid and the amino group of another, with the release of a water molecule.

  • Hydrolysis breaks peptide bonds to release individual amino acids by adding a water molecule.

• Fibrous and globular proteins:

  • Fibrous proteins: Typically long, insoluble, and stable, serving structural roles in the body (e.g., collagen in connective tissue, keratin in hair and nails, actin and myosin in muscle).

  • Globular proteins: Generally compact, soluble, and sensitive to environmental changes (temperature, pH). They often have functional roles, acting as enzymes, hormones, antibodies, and many regulatory proteins (e.g., hemoglobin, insulin).

• Levels of protein structure:

  • Primary structure: The unique linear sequence of amino acids in the polypeptide chain, determined by the genetic code. This sequence dictates all higher levels of structure.

  • Secondary structure: Localized folding patterns within sections of the polypeptide chain, primarily stabilized by hydrogen bonds between the backbone atoms (not R groups). Common forms include alpha helices (a coiled structure) and beta-pleated sheets (a folded, zig-zag structure).

  • Tertiary structure: The overall three-dimensional shape of a single polypeptide chain, resulting from interactions between the R groups of amino acids (e.g., hydrogen bonds, ionic bonds, disulfide bridges, hydrophobic interactions) and interactions between R groups and the surrounding solvent. This level is crucial for protein function.

  • Quaternary structure: The assembly of two or more separate polypeptide chains (subunits) into a larger, functional protein complex. Not all proteins possess a quaternary structure; many functional proteins consist of a single polypeptide chain.

• Enzyme-substrate interaction (active site):

  • Enzymes function by binding to specific substrate molecules at a region called the active site. The active site is a uniquely shaped pocket or groove that complements the shape of the substrate.

  • Upon substrate binding, the enzyme often undergoes a slight conformational change (induced fit) to optimize the interaction and promote the reaction.

  • The enzyme then facilitates the formation of a transition state, lowering the activation energy for the reaction.

  • Products are formed from the modified substrate.

  • The product(s) and enzyme then dissociate; the enzyme returns to its original shape, ready to bind another substrate molecule and catalyze the reaction again.

Nucleotides & Nucleic Acids

• Structure of nucleotides:

  • Nucleotides are the monomers of nucleic acids (DNA and RNA) and are also involved in energy transfer (ATP). Each nucleotide consists of three main components:

  • A nitrogenous base: Carbon-nitrogen ring structures. Classified into:

    • Purines: Double-ring structures (Adenine (A), Guanine (G)).

    • Pyrimidines: Single-ring structures (Cytosine (C), Thymine (T), Uracil (U)).

  • A pentose sugar (five-carbon sugar):

    • Deoxyribose in DNA (lacks an oxygen atom on the 2' carbon).

    • Ribose in RNA (has a hydroxyl group on the 2' carbon).

  • One, two, or three phosphate groups: Attached to the sugar (e.g., adenosine monophosphate (AMP), adenosine diphosphate (ADP), adenosine triphosphate (ATP)). These phosphate bonds store significant energy.

• ATP structure:

  • ATP (Adenosine Triphosphate) is a crucial energy currency of the cell. It consists of an adenosine molecule (adenine base + ribose sugar) attached to three phosphate groups.

  • ATP formation from ADP (adenosine diphosphate) + an inorganic phosphate ($P_i$) requires energy input (endergonic reaction), typically from cellular respiration.

  • Hydrolysis of ATP (breaking the outermost phosphate bond) releases a large amount of energy to power various cellular processes (exergonic reaction), forming ADP and $P_i$.

• Nucleic acids connectivity:

  • Nucleotides connect to one another through phosphodiester bonds, which form between the phosphate group of one nucleotide and the sugar of another. This creates a sugar-phosphate backbone, with the nitrogenous bases projecting outwards.

• DNA and RNA structure:

  • DNA (Deoxyribonucleic Acid): Typically exists as a double helix, resembling a twisted ladder. It stores genetic information.

  • The two polynucleotide strands are held together by hydrogen bonds between complementary nitrogenous bases.

  • RNA (Ribonucleic Acid): Typically exists as a single strand and plays various roles in gene expression (e.g., mRNA, tRNA, rRNA), carrying genetic information from DNA and synthesizing proteins.

  • In RNA, the nitrogenous base thymine (T) found in DNA is replaced by uracil (U).

• Base pairing:

  • In DNA: Adenine (A) always pairs with Thymine (T) via two hydrogen bonds. Guanine (G) always pairs with Cytosine (C) via three hydrogen bonds.

  • In RNA: Adenine (A) pairs with Uracil (U) via two hydrogen bonds. Guanine (G) pairs with Cytosine (C) via three hydrogen bonds (in regions where RNA folds back on itself or interacts with other nucleic acids).

• Summary concepts:

  • Briefly describe the major functions of carbohydrates (energy, structure), lipids (energy storage, membranes, hormones), proteins (structure, enzymes, transport, defense), nucleotides (energy transfer, signaling), and nucleic acids (genetic information storage and expression) in the body.

  • Identify monomer and polymer structures for each macromolecule (e.g., monosaccharide/polysaccharide, amino acid/protein).

  • Describe the structure and formation of ATP from ADP and relate the cleavage of its terminal phosphate bond to cellular energy release.

  • Understand dehydration synthesis as a bond-forming process (removal of water) and hydrolysis as a bond-breaking process (addition of water), and recognize their importance in macromolecule metabolism.

"Like dissolves like" concept is reinforced throughout polarity discussions; water polarity and hydrogen bonding underpin many of the body’s transport and reaction processes.