Biochemistry and Chemical Foundations Notes

Energy and States of Matter in Biology

• Energy in the body is present as kinetic and potential energy. Potential energy is stored energy; kinetic energy is energy in motion. Example: holding a clicker up stores potential energy; dropping it converts to kinetic energy.
• Energy can be transferred from potential to kinetic within biological systems; ATP is the cell’s energy currency and will be discussed later.
• Elements are composed of atoms with unique physical and chemical properties and unique chemical symbols (often based on Latin or historical usage).
• On the test, you should know the four elements that make up about 96% of the mass of living things: $ext{C, H, O, N}$ (carbon, hydrogen, oxygen, nitrogen).
• Mnemonics and quick recall: a mnemonic story may be used in class (HONC1234) to remember the bonding behavior of the major elements, with the explicit note that the four are $ext{H, O, N, C}$; others like calcium, phosphorus, potassium, sulfur, sodium, chlorine, magnesium, iodine, iron are present in smaller amounts.
• Trace elements exist in very small amounts but are essential; excessive amounts of some trace elements (e.g., zinc or others) can be toxic.
• Calcium is found in bones and is needed for contraction and nerve function; phosphorus is important for energy phosphates (like ATP); potassium is essential for muscle contraction and is a critical electrolyte; iron is needed for transporting oxygen in blood.
• Some elements are needed in very small quantities and are called trace elements; many are toxic at high levels.
• An atom is composed of a nucleus (protons and neutrons) and electrons in electron shells around the nucleus. Protons are positively charged; neutrons are neutral; electrons are negatively charged.
• The nucleus contains protons and neutrons; electrons orbit the nucleus in shells (electron clouds).
• Electron shell capacities (as presented in the transcript): the first shell holds $2$ electrons; the next shell holds $8$ electrons; the following shell is described as holding $8$ and $16$ in a way that is simplified for this course.
• An atom is stable when its outer (valence) shell is full; if not full, the atom is unstable and tends to bond to become stable.
• Quick visualization: in the classroom analogy, the nucleus is dense and the nearest electron is conceptually far away (e.g., if the nucleus were the size of a basketball, the closest electron would be the size of a golf ball and about a mile away).
• Protons are positively charged; electrons are negatively charged; neutrons are neutral.
• The “planetary” atomic model is not strictly used here; chemistry is suggested as the better place to learn the details if you want to go deeper.
• A molecule is two or more atoms bonded together; a compound is two or more atoms of different kinds bonded together.
• In solutions, a solute dissolved in a solvent forms a solution; a salt is an ionic compound often formed when ionic bonds dissolve and separate into ions.
• Colloids and suspensions are two terms used to describe particle size and dispersion in a medium; colloids remain dispersed and are not fully dissolved; suspensions are larger particles suspended in a medium.
• Avogadro’s number is often referenced in chemistry and biochemistry; for reference, $N_A = 6.022 imes 10^{23}$ particles per mole.

Ionic, Covalent, and Hydrogen Bonds

• Atoms that are not stable will bond with other atoms to become stable.
• Ionic bonds involve electron transfer: one atom donates an electron, the other accepts it, forming ions (cations and anions).
• Covalent bonds involve sharing electrons between atoms.
• Hydrogen bonds are weak interactions between a partial positive hydrogen (often attached to a highly electronegative atom) and a partial negative atom elsewhere; hydrogen bonds are not true covalent bonds but are important for structures like water.
• Key terms:
  • Cations: positively charged ions formed when atoms lose electrons.
  • Anions: negatively charged ions formed when atoms gain electrons.
• Illustrative example: sodium (Na) donates an electron (Na → Na^+ + e^-); chlorine (Cl) gains an electron (Cl + e^- → Cl^-). The Na^+ and Cl^- form an ionic bond, yielding NaCl (table salt).
• Mnemonic to remember charge-direction: cat ion (cation) is positive; an ion is the negative ion (anion).
• The steps for forming an ionic bond using Na and Cl:
  • Sodium loses one electron to become Na^+ (a cation).
  • Chlorine gains one electron to become Cl^- (an anion).
  • Opposite charges attract, forming the ionic bond in NaCl.
• Ionic compounds form salts (e.g., sodium chloride).
• Covalent bonds—remember the four major elements that account for most bonds in biology: $H, O, N, C$. The pattern of bonding: H forms one covalent bond; O forms two; N forms three; C forms four. The mnemonic HONC1234 helps with the basic bonding pattern.
• Carbon commonly forms four covalent bonds, contributing to complex macromolecules.
• How chemists depict covalent bonds: one shared bond is shown as a single bond; a double bond is two shared pairs; a triple bond is three shared pairs.
• Nonpolar covalent bonds share electrons equally; polar covalent bonds share electrons unequally, creating partial positive and partial negative ends (dipoles).
• Water as an example of polar covalent bonding within a molecule:
  • Oxygen is more electronegative than hydrogen, so electrons spend more time around oxygen, giving it a partial negative charge and hydrogen a partial positive charge.
  • Water can form up to four hydrogen bonds (two hydrogens bonding with other electronegative atoms in neighboring molecules).

Water, Solvents, and Hydrophilic vs Hydrophobic Effects

• Water is the most important inorganic molecule for life and a versatile solvent; the cells are about two thirds water; fat cells are less hydrated; muscle cells are more hydrated.
• Properties of water:
  • High heat capacity: can absorb a lot of energy before changing temperature.
  • High heat of vaporization: can dissipate heat through evaporation.
  • Water acts as a solvent for polar and ionic substances; nonpolar molecules are not easily dissolved by water.
  • Hydrophilic substances dissolve in water; hydrophobic substances do not.
  • Polar molecules and ionic substances dissolve well in water; oil and water do not mix.
• Important terms:
  • Hydrophilic: water-loving and water-mixing.
  • Hydrophobic: water-fearing and water-repelling.
• Water enables many biological processes, including hydrolysis and dehydration synthesis, and acts as a transport medium in the body.
• Water also provides cushioning in the body (e.g., cerebrospinal fluid around the brain).

Salts, Electrolytes, and pH Basics

• Salts are ionic compounds formed when ionic bonds dissociate in water, producing ions; electrolytes are ions in solution that conduct electricity.
• Replenishing electrolytes is important in cases of dehydration or illness (examples: Gatorade, Pedialyte; in severe cases, IV fluids).
• Acids and bases are electrolytes that ionize in water.
• Acids are proton (H^+) donors; bases are proton (H^+) acceptors.
• pH is a measure of hydrogen ion concentration; the pH scale runs from 0 to 14; 7 is neutral; below 7 is acidic; above 7 is basic (alkaline).
• Water self-ionizes slightly: H2O ⇌ H^+ + OH^-; the relative concentrations determine the pH.
• Common pH examples:
  • pH 2: strong acid.
  • pH 8: weak base.
  • pH 13: strong base.
  • Household acids (like stomach acid) vs bases (like certain cleaners) have different pH values.
• Buffer systems in the blood and body help minimize pH changes by neutralizing added acids or bases; buffers include bicarbonate systems and other natural buffers; lungs and kidneys help regulate pH by excreting hydrogen ions or adjusting bicarbonate, and buffers operate like a chemical “Alka-Seltzer” in the bloodstream.
• The pH scale is logarithmic: each unit change represents a 10-fold change in hydrogen ion concentration. For example, a pH difference of 4 units corresponds to a factor of $10^4$ in [H^+], and a difference of 5 units corresponds to a factor of $10^5$, etc. Example: changing from $pH=4$ to $pH=8$ is 10,000 times more acidic at pH 4 than at pH 8.
• Blood pH: $7.35 ext{ to } 7.45$ is the normal range; deviations can lead to acidosis or alkalosis and can be fatal if not compensated.
• Urine pH: approximately $4 ext{ to } 8$; the kidneys help regulate acid-base balance by excreting hydrogen ions or conserving bicarbonate.
• Neutralization concept: adding a base to an acid neutralizes excess hydrogen ions and forms water and a salt (e.g., NaOH + HCl → NaCl + H2O).
• pH regulation in the body relies on kidneys, lungs, and buffers to maintain homeostasis, defined as maintaining a steady state within parameters.

Acids, Bases, and pH in Detail

• Definitions:
  • Acid: proton donor.
  • Base: proton acceptor.
• Examples:
  • Hydrochloric acid (HCl) is an acid.
  • Sodium hydroxide (NaOH) is a base because it produces hydroxide ions (OH^-).
• Acid-base reactions can be represented in chemical equations; neutralization results in water and a salt.
• pH is a measure of hydrogen ion concentration and is influenced by buffers and physiological processes; pH is crucial for protein structure and enzyme activity; small pH changes can be life-threatening due to the logarithmic scale.

Organic vs Inorganic Molecules and Biochemistry Context

• Inorganic molecules are small and often do not contain carbon; carbon dioxide is a notable exception.
• Organic molecules contain carbon and are central to biochemistry; examples include carbohydrates, lipids, proteins, and nucleic acids.
• Water is an important inorganic molecule that participates in many biochemical reactions.
• Biochemistry is the study of chemical processes in living organisms; it bridges chemistry and biology and includes how molecules interact in living systems.

Quick Reference: Common Concepts and Test Tips

• The four major elements: $ext{H, O, N, C}$; these account for about 96 ext{ ext%} of the mass of living things.
• The remaining ~4 ext{ ext%} includes elements like calcium, phosphorus, potassium, sulfur, sodium, chlorine, magnesium, iodine, iron, and trace elements; some are essential in trace amounts, others can be toxic if in excess.
• Bond types recap:
  • Ionic bonds: transfer of electrons; formation of ions; ions attract to form ionic compounds (salts).
  • Covalent bonds: sharing of electrons; can be nonpolar (equal sharing) or polar (unequal sharing, creating partial charges).
  • Hydrogen bonds: weak, non-covalent interactions between a partial positive hydrogen and a partial negative atom in another molecule.
• HONC1234 mnemonic helps you remember how many covalent bonds the four major atoms typically form:
  • H: 1 bond
  • O: 2 bonds
  • N: 3 bonds
  • C: 4 bonds
• Water-specific notes:
  • Polar covalent bonds within water create a bent molecule with a partial negative charge on oxygen and partial positive charges on hydrogens.
  • Water forms up to four hydrogen bonds per molecule.
  • Water’s properties (high heat capacity, high heat of vaporization, solvent versatility) are essential for temperature regulation, chemical reactions, and transport systems in organisms.
• pH and buffers:
  • pH measures hydrogen ion concentration; scale is logarithmic and bounded 0–14.
  • pH < 7: acidic; pH > 7: basic; pH = 7: neutral.
  • Buffers minimize pH changes; kidneys and lungs help regulate pH in the body; buffers are present in blood and other fluids.
• Test-style ideas to practice:
  • Given two pH values, determine how many times more acidic one is than the other by inserting 10 between each step and counting powers of 10.
  • Predict whether a given formula corresponds to an acid or a base (e.g., HCl is an acid; NaOH is a base).
  • Understand how an ionic bond forms from electron transfer (Na → Na^+; Cl + e^- → Cl^-; NaCl formation).
• Real-world relevance:
  • Understanding ATP as the energy currency connects chemistry to cellular respiration and metabolism.
  • Electrolytes and hydration are critical in diagnosing and treating dehydration and electrolyte disturbances.
  • Buffers protect proteins and enzymes from denaturation due to pH changes; blood and urine pH are clinically relevant indicators.

Summary: Foundational Principles to Remember

• Energy flow in cells centers on potential-to-kinetic transitions; ATP is the chief energy currency.
• The vast majority of mass in living organisms comes from four elements: $ext{C, H, O, N}$, with others present in smaller amounts.
• Atoms bond to achieve stability; ionic bonds transfer electrons (forming ions); covalent bonds share electrons; hydrogen bonds provide temporary, weak attractions that help stabilize structures like water networks.
• Water’s properties as a solvent and regulator of temperature are central to biochemistry and physiology; pH control and buffers are essential for maintaining homeostasis.
• Distinguishing organic vs inorganic chemistry helps categorize biological molecules and reactions; organic molecules (carbons-based) form the backbone of life’s macromolecules.
• Core biochemistry topics for study include carbohydrates, lipids, proteins, and nucleic acids, which will be explored in more detail in later chapters.

Important Formulas and Concepts (in LaTeX)

• Ionization and water equilibrium: ext{H}_2 ext{O}
  ightleftharpoons ext{H}^+ + ext{OH}^-
• pH concept (logarithmic relation): $ext{pH} = -\log [ ext{H}^+]$
• Avogadro’s number: $N_A = 6.022 imes 10^{23} ext{ particles per mole}$
• Salt formation (example): ext{NaOH} + ext{HCl}
  ightarrow ext{NaCl} + ext{H}_2 ext{O}
• Ionic bond electron-transfer visualization (summary): Na → Na^+ + e^-; Cl + e^- → Cl^-; resulting in NaCl
• Hydrogen bond (definition): a weak bond between a partial positive hydrogen and a partial negative atom in another molecule
• Electron shell capacities (as presented): first shell $2$; second shell $8$; subsequent shells described in class as $8/16$ in simplified terms
• pH ranges and categories: acids (pH < 7), bases (pH > 7), neutral (pH = 7); common ranges: blood $7.35 o 7.45$; urine $4 o 8$
• Calorie concept (educational note): a calorie is the amount of heat needed to raise the temperature of $1 ext{ g}$ of water by $1^ ext{o}C$; $1 ext{ cal} = 4.184 ext{ J}$ (conversion used in broader contexts)