Chemistry Notes: Matter, Ions, Bonding, Water, Biomolecules, and Energy

HLTH115 Notes: Matter, Ions, Bonding, Water, Biomolecules, and Energy

Matter and Atoms

  • Matter has mass and occupies space.

  • 3 forms of matter:

    • Solid (e.g., bone)

    • Liquid (e.g., blood)

    • Gas (e.g., oxygen)

  • An atom is the smallest particle exhibiting the chemical properties of an element.

  • There are 92 naturally occurring elements that make up matter.

  • Elements are organized in the periodic table of elements.

Ions and Ionic Compounds

  • Ionic compounds are stable associations between two or more elements combined in a fixed ratio.

  • Ionic compounds are classified as ionic or molecular.

  • Ionic bonds are electrostatic forces that hold ions together.

  • Commonly form salts (e.g., table salt, NaCl).

  • Ionic lattice crystal structures arise when oppositely charged ions are arranged in a repeating pattern.

  • Ions are atoms with positive or negative charges formed by the loss or gain of one or more electrons.

  • Physiological examples:

    • Potassium ions (K+) are used to replace K+ lost in sweat (e.g., sports drinks).

    • In very large doses, K+ is used in lethal injections.

Ions and Ion Formation (Examples)

  • Cations (positive charge) form when atoms lose electrons.

    • Example: Sodium (Na) loses one outer-shell electron to chlorine to satisfy the octet rule.

    • Result: Na+ with 11 protons and 10 electrons; net charge +1.

  • Anions (negative charge) form when atoms gain electrons.

    • Example: Chlorine (Cl) gains an electron to satisfy the octet rule.

    • Result: Cl− with 17 protons and 18 electrons; net charge −1.

  • Polyatomic ions are ions composed of more than one atom (e.g., bicarbonate HCO3− and phosphate PO4^3−).

Ionic Bonds

  • Cations and anions are bound by ionic bonds to form salts.

  • Ionic bonds form lattice crystal structures in ionic compounds (e.g., NaCl and MgCl2).

  • Formation example: In NaCl, sodium loses one electron and chlorine gains one electron; the ions are held together by electrostatic forces.

  • Figure reference: Formation of an Ionic Bond Involving Sodium and Chloride (Figure 2.5).

Covalent Bonding, Molecules, and Molecular Compounds

  • Covalently bonded molecules result from electrons being shared between atoms of two or more elements.

  • These are termed molecular compounds (e.g., CO2, but not O2 which is considered a diatomic molecule).

Covalent Bonds

  • Covalent bonds form when both atoms require electrons and share them.

  • Commonly form in the human body using:

    • Hydrogen (H)

    • Oxygen (O)

    • Nitrogen (N)

    • Carbon (C)

  • Number of covalent bonds an atom can form:

    • Simplest case: two hydrogens form a single bond.

    • Oxygen needs two electrons to complete its outer shell and forms two bonds.

    • Nitrogen forms three bonds.

    • Carbon forms four bonds.

Types and Strength of Covalent Bonds

  • Single covalent bond: one pair of electrons shared (e.g., H–H).

  • Double covalent bond: two pairs of electrons shared (e.g., O=O).

  • Triple covalent bond: three pairs of electrons shared (e.g., N≡N).

  • Nonpolar covalent bonds vs polar covalent bonds (electronegativity):

    • Equal sharing of electrons yields a nonpolar covalent bond when atoms have the same electronegativity.

    • Unequal sharing yields a polar covalent bond.

Intermolecular Attractions

  • Intermolecular attractions are weak chemical attractions between molecules but are important for the shape and behavior of complex molecules (e.g., DNA and proteins).

  • Hydrogen bonds form between polar molecules when a partially positive hydrogen atom is attracted to a partially negative atom (usually O, N, or F).

  • Individually weak, collectively strong; influence properties of water and biological macromolecules.

Water and Hydrogen Bonding

  • Water is a polar molecule: one oxygen atom bonded to two hydrogen atoms.

    • Oxygen has partial negative charges; hydrogens have partial positive charges.

  • Water can form four hydrogen bonds with adjacent molecules.

  • Water is central to many properties and processes in biology.

Molecular Structure of Water

  • Water composition: two-thirds of the human body by weight is water ( frac{2}{3}).

  • Water’s role is fundamental to transport, lubrication, cushioning, and waste excretion.

Properties of Water

  • Phases of water: gas (water vapor), liquid (water), solid (ice).

  • Functions of liquid water:

    • Transports dissolved substances through the body.

    • Lubricates joints and tissues to decrease friction.

    • Cushions tissues by absorbing shock.

    • Excretes wastes by dissolving them for elimination.

  • Cohesion: attraction between water molecules due to hydrogen bonding.

  • Surface tension: inward pulling of cohesive forces at the water surface; surfactant prevents alveolar collapse in the lungs.

  • Adhesion: attraction between water and other substances.

  • Water has high specific heat and high heat of vaporization due to hydrogen bonding.

  • Water as a universal solvent: many substances dissolve in water; solubility depends on chemical properties of the solute.

Acids, Bases, pH, Neutralization, and Buffers

  • Acids dissociate in water to produce H+ and anions; proton donors; increase free H+ concentration.

    • Stronger acids dissociate more; e.g., HCl in the stomach.

    • Weaker acids dissociate less; e.g., carbonic acid in the blood.

    • General representation: Substance A (an acid) → H+ + A−.

  • Bases accept H+ when added to solution; proton acceptors; decrease free H+ concentration.

    • Stronger bases accept more H+; e.g., ammonia and bleach.

    • Weaker bases absorb less H+; e.g., bicarbonate in blood and in secretions released into the small intestine.

    • General representation: Substance B (a base) + H+ → HB+ or similar.

  • pH, Neutralization, and Buffers:

    • pH is a measure of H+ concentration; range 0–14.

    • Pure water has pH = 7; [H+] = [OH−] = 1 imes 10^{-7} ext{ M}.

    • pH and H+ concentration are inversely related: pH = −log10([H+]).

    • Moving from one pH unit to the next represents a 10-fold change in H+ concentration: e.g., pH 6 has 10× more H+ than pH 7.

    • Solutions with equal H+ and OH− are neutral (pH = 7).

    • Acidic solutions have pH < 7; basic solutions have pH > 7.

    • Neutralization: acids neutralized by adding base; bases neutralized by adding acid.

    • Buffers help prevent pH changes when excess acid or base is added; they can accept H+ from excess acid or donate H+ to neutralize base.

    • In the blood, a critical buffer system is the carbonic acid/bicarbonate buffer, maintaining pH in the range 7.35 ext{ to } 7.45.

Biological Macromolecules: General Characteristics

  • Organic molecules contain carbon; most are components of living organisms.

  • Biological macromolecules (biomolecules) are a subset of organic molecules.

  • Four classes of organic biomolecules in living systems:

    • Lipids

    • Carbohydrates

    • Nucleic acids

    • Proteins

  • Inorganic molecules are all other molecules.

Lipids

  • Lipids are not polymers.

  • They are a diverse group of fatty, water-insoluble molecules.

  • Functions:

    • Stored energy

    • Components of cellular membranes

    • Hormones

  • Four primary classes:

    • Triglycerides

    • Phospholipids

    • Steroids

    • Eicosanoids

Triglycerides

  • Most common lipid in living things.

  • Long-term energy storage in adipose tissue; structural support, cushioning, insulation.

  • Formed from a glycerol molecule and three fatty acids.

    • Formation occurs via dehydration synthesis.

Phospholipids

  • Amphipathic molecules forming chemical barriers of cell membranes.

  • Amphipathic: containing both water-soluble (polar) and nonpolar portions.

  • Structure is similar to a triglyceride except one end of the glycerol has a polar phosphate group with attached organic groups instead of a fatty acid.

  • Components:

    • Glycerol, phosphate group, organic groups (polar head)

    • Two fatty acid tails (nonpolar)

  • Form hydrophilic heads and hydrophobic tails; arrange into bilayers in membranes.

Steroids

  • Composed of four fused hydrocarbon rings; three six-membered rings and one five-membered ring.

  • Side chains extend from the rings.

  • Examples:

    • Cholesterol: component of animal plasma membranes; precursor to other steroids.

    • Steroid hormones (e.g., testosterone and estrogen).

    • Bile salts.

Eicosanoids

  • Modified 20-carbon fatty acids synthesized from arachidonic acid in membranes.

  • Local signaling molecules with key roles in inflammatory response and nervous system communication.

  • Four classes:

    • Prostaglandins

    • Prostacyclins

    • Thromboxanes

    • Leukotrienes

Fatty Acids: Saturated, Unsaturated, and Trans Fats

  • Most animal fats are saturated (solid at room temperature).

  • Most vegetable fats are unsaturated (liquid at room temperature; generally healthier).

  • Hydrogenation can convert unsaturated fats to saturated fats.

  • Partial hydrogenation may produce trans fats, which increase risk of heart attack and stroke.

Carbohydrates

  • General formula: (CH2O)n where n = number of carbon atoms.

  • Monosaccharides: simple sugar monomers.

  • Disaccharides: formed from two monosaccharides.

  • Polysaccharides: formed from many monosaccharides.

  • Glucose:

    • A six-carbon carbohydrate; the most common monosaccharide.

    • Primary energy source for cells; its concentration must be carefully maintained.

  • Glycogen:

    • Stored in liver and skeletal muscle; excess glucose is stored as glycogen (glycogenesis).

    • Liver can hydrolyze glycogen back into glucose as needed (glycogenolysis).

    • Liver can also synthesize glucose from noncarbohydrate sources (gluconeogenesis).

  • Other carbohydrates:

    • Polysaccharides such as glycogen (animals), starch and cellulose (plants).

    • Plant starch is a major nutritional source of glucose for humans.

    • Cellulose is a source of dietary fiber (nondigestible).

Nucleic Acids

  • Nucleic acids store and transfer genetic information.

  • Two classes:

    • Deoxyribonucleic acid (DNA)

    • Ribonucleic acid (RNA)

  • Both are polymers of nucleotide monomers joined by phosphodiester bonds.

Deoxyribonucleic Acid (DNA)

  • DNA is double-stranded.

  • Location: chromosomes in the nucleus and in mitochondria.

  • Backbone components: deoxyribose sugar, phosphate, and one of four bases: Adenine (A), Guanine (G), Cytosine (C), Thymine (T).

  • No uracil in DNA.

  • Base pairing via hydrogen bonds:

    • Thymine pairs with Adenine (A–T)

    • Guanine pairs with Cytosine (G–C)

Ribonucleic Acid (RNA)

  • RNA is single-stranded.

  • Location: nucleus and cytoplasm of the cell.

  • Backbone components: ribose sugar, phosphate, and four bases: A, G, C, Uracil (U).

  • No thymine in RNA; uracil pairs with adenine.

Nucleic Acids: ATP (Adenosine Triphosphate)

  • ATP is a nucleotide composed of adenine, a ribose sugar, and three phosphate groups.

  • Central molecule for transfer of chemical energy within the cell.

  • The covalent bonds between the last two phosphate groups are energy-rich (phosphoanhydride bonds).

  • Energy is released when these bonds are broken.

  • Figure reference: Figure 2.22.

Proteins

  • Functions of proteins:

    • Act as catalysts (enzymes) in metabolic reactions.

    • Defend the body (immunity).

    • Aid in transport.

    • Contribute to structural support.

    • Cause movement.

    • Perform regulation.

    • Provide storage.

Energy, Chemical Reactions, and Cellular Respiration

  • All living organisms require energy to power muscle, pump blood, absorb nutrients, exchange respiratory gases, synthesize new molecules, and establish cellular ion concentrations.

  • Glucose breakdown through metabolic pathways forms ATP, the “energy currency” of cells.

Classes of Energy

  • Energy: the capacity to do work.

  • Two main classes:

    • Potential energy (stored energy; energy of position)

    • Kinetic energy (energy of motion)

  • Both can be converted from one class to the other.

Forms of Energy

  • Chemical energy: stored in chemical bonds; the most important form of energy in the human body; used for movement, molecule synthesis, and establishing concentration gradients.

  • Present in: triglycerides (long-term energy storage in adipose tissue), glucose (glycogen stores in liver and muscle), ATP (stored in all cells; produced continuously and used immediately).

  • Proteins also store chemical energy but have more important structural and functional roles.

  • Kinetic energy forms:

    • Electrical energy: movement of charged particles (e.g., nerves: movement of ions across the plasma membrane).

    • Mechanical energy: motion of objects due to applied force (e.g., muscle contraction for walking).

    • Sound energy: molecule compression caused by a vibrating object (e.g., sound waves vibrating the eardrum).

  • Radiant energy: energy of electromagnetic waves (vary in wavelength and frequency; higher frequency means higher energy; frequencies higher than visible light can penetrate the body and mutate DNA; cells are protected by melanin; visible light is detected by retinal cells and relayed to the brain).

  • Heat: kinetic energy of random motion; usually not available to do work; measured as the temperature of a substance.

Key Formulas and Numbers

  • pH and hydrogen ion concentration:

    • \text{pH} = -\log_{10}[H^+]

    • Pure water at neutral pH: [H^+] = [OH^-] = 1.0\times 10^{-7} \text{ M}

    • pH 7 corresponds to equal H+ and OH− concentrations; acidic solutions have pH < 7; basic solutions have pH > 7.

    • Moving from one pH unit to the next represents a 10-fold change in [H+].

  • Water composition: \tfrac{2}{3} of the body by weight is water.

  • Lipid types: e.g., four classes of lipids; triglycerides, phospholipids, steroids, eicosanoids.

  • Steroids: four fused rings (three 6-member, one 5-member).

  • Nucleic acids: nucleotides linked by phosphodiester bonds; A–T, G–C base pairing in DNA; A–U and G–C pairing in RNA.

Connections to Foundational Principles

  • Chemical bonds determine molecular structure and function (ionic vs covalent vs hydrogen bonds).

  • Water properties underpin solvent behavior, temperature regulation, and biochemical reactions.

  • Biomolecule classes provide energy storage, genetic information, catalysis, signaling, and structural roles essential for cellular life.

  • Energy transformations (chemical energy → ATP; ATP hydrolysis) drive all cellular processes.

  • Buffers and pH homeostasis are critical for maintaining enzyme activity and metabolic stability (e.g., blood pH in a narrow range).

Practical and Real-World Implications

  • Lethal injections and electrolyte balance (K+ in sports drinks) illustrate physiological importance of ions.

  • Hydrogenation and trans fats connect chemistry to nutrition and cardiovascular risk.

  • Buffers in blood (carbonic acid/bicarbonate) demonstrate human body’s regulation of pH for homeostasis.

  • Surfactants prevent alveolar collapse and are essential for respiration.

Summary

  • Matter consists of atoms arranged in various bonds and structures.

  • Ionic and covalent bonds create a diverse array of compounds with distinct properties.

  • Water’s unique properties support life and physiological processes.

  • Biologically important macromolecules (lipids, carbohydrates, nucleic acids, proteins) provide energy storage, information, structure, and catalysis.

  • Energy flow in biology is governed by conversions between potential and kinetic energy, primarily through chemical bonds and ATP.