Lecture #3

Chemistry and Physiological Reactions

Introduction to Chemistry in Anatomy & Physiology

Chemistry forms the foundation of all physiological processes in the human body, including movement, digestion, heart function, and nervous system activity. Understanding basic chemistry and biochemistry is essential for studying anatomy and physiology.

• Basic Chemistry: Focuses on the structure and properties of matter, including atoms, elements, and chemical bonds.

• Biochemistry: Examines the chemical processes and substances that occur within living organisms.

Basic Chemistry Concepts

Atoms and Elements

All matter is composed of elements, which are listed in the periodic table. The atom is the smallest unit of an element and consists of subatomic particles:

• Neutrons: No charge, located in the nucleus.

• Protons: Positive charge, located in the nucleus.

• Electrons: Negative charge, found in a cloud around the nucleus.

• Atomic Number: Number of protons in an atom; unique to each element.

• Atomic Weight: Sum of protons and neutrons in an atom.

Common Elements in the Human Body

The human body is primarily composed of a few key elements, each with specific functions:

Element	Symbol	Body Mass (%)	Functions
Oxygen	O	65	Component of organic and inorganic molecules; essential for ATP production.
Carbon	C	18.5	Component of all organic molecules (carbohydrates, lipids, proteins, nucleic acids).
Hydrogen	H	9.5	Component of organic molecules; as an ion (proton), influences pH of body fluids.
Nitrogen	N	3.2	Component of proteins and nucleic acids.
Calcium	Ca	1.5	Found in bones/teeth; required for muscle contraction, nerve conduction, blood clotting.
Phosphorus	P	1.0	Part of bones/teeth, nucleic acids, ATP, phospholipids.
Potassium	K	0.4	Major cation in cells; necessary for nerve impulses and muscle contraction.
Sulfur	S	0.3	Component of proteins, especially muscle proteins.
Sodium	Na	0.2	Major cation in extracellular fluids; important for water balance, nerve impulses, muscle contraction.
Chlorine	Cl	0.2	Major anion in extracellular fluids.
Magnesium	Mg	0.1	Present in bone; cofactor in metabolic reactions.
Iodine	I	<0.01	Needed for thyroid hormone production.
Iron	Fe	<0.01	Component of hemoglobin and some enzymes.

Trace Elements: Elements like chromium, cobalt, copper, fluorine, manganese, molybdenum, selenium, silicon, tin, vanadium, and zinc are required in minute amounts, often as enzyme cofactors.

Combining Matter: Molecules, Compounds, and Mixtures

Molecules vs. Compounds

• Molecule: Two or more atoms of the same element bonded together (e.g., O2).

• Compound: Molecule with two or more different kinds of atoms bonded together (e.g., C6H12O6).

Mixtures

Most matter exists as mixtures, which are physical combinations of two or more components. Types include:

• Solutions: Homogeneous mixtures with evenly distributed particles. Example: blood plasma.

• Colloids (Emulsions): Heterogeneous mixtures with large particles that do not settle out. Example: cytosol.

• Suspensions: Heterogeneous mixtures with large, visible particles that settle out. Example: blood (cells settle if left standing).

Differences Between Mixtures and Compounds

• Mixtures do not involve chemical bonding; compounds do.

• Mixtures can be separated by physical means; compounds require chemical methods.

• Mixtures can be heterogeneous or homogeneous; compounds are always homogeneous.

Measurement of Solution Concentration

• Percent (%): Parts of solute per 100 parts solution (e.g., 10g salt in 100ml water = 10% solution).

• Milligrams per deciliter (mg/dl): Mass of solute per 100ml solution (e.g., normal fasting blood glucose ≈ 80 mg/dl).

• Molarity (M): Moles of solute per liter of solution. A mole contains $6.02 \times 10^{23}$ molecules (Avogadro’s number).

Example: 180.12g of glucose (C6H12O6) in 1L water = 1M solution.

Chemical Bonds

Types of Chemical Bonds

• Ionic Bonds: Formed by transfer of electrons from one atom to another, creating ions (cations and anions) that attract each other.

• Covalent Bonds: Formed by sharing of valence electrons between atoms. Can be single, double, or triple bonds.

• Hydrogen Bonds: Weak attractions between a hydrogen atom (partially positive) and an electronegative atom (e.g., oxygen in water).

Types of Covalent Bonds

• Nonpolar Covalent Bonds: Equal sharing of electrons (e.g., O2, CO2).

• Polar Covalent Bonds: Unequal sharing of electrons, resulting in partial charges (e.g., H2O).

Comparison of Bond Types

Bond Type	Description	Relative Strength
Covalent	Sharing of electron pairs; can be polar or nonpolar	Strongest
Ionic	Attraction between oppositely charged ions	Intermediate
Hydrogen	Attraction between H (partial +) and electronegative atom	Weakest

Chemical Reactions

Chemical Equations

Chemical reactions involve the formation, rearrangement, or breaking of chemical bonds. They are represented by chemical equations:

• Reactants: Substances entering the reaction.

• Products: Substances produced by the reaction.

Example Equations:

• $\mathrm{H}_2 + \mathrm{O}_2 \rightarrow 2\mathrm{H}_2\mathrm{O}$

• $\mathrm{NaOH} + \mathrm{HCl} \rightarrow \mathrm{NaCl} + \mathrm{H}_2\mathrm{O}$

• $6\mathrm{CO}_2 + 6\mathrm{H}_2\mathrm{O} \leftrightarrow \mathrm{C}_6\mathrm{H}_{12}\mathrm{O}_6 + 6\mathrm{O}_2$

Types of Chemical Reactions

• Synthesis (Combination): Smaller particles combine to form larger, more complex molecules. Example: Amino acids forming proteins.

• Decomposition: Breakdown of a molecule into smaller molecules or atoms. Example: Glycogen breaking down to glucose.

• Exchange (Displacement): Bonds are both made and broken; atoms are exchanged. Example: ATP transferring a phosphate group.

• Redox (Oxidation-Reduction): Involves electron transfer; atoms are oxidized (lose electrons) or reduced (gain electrons).

Factors Affecting Reaction Rates

• Temperature: Higher temperature increases reaction rate.

• Concentration: Higher concentration increases reaction rate.

• Particle Size: Smaller particles increase reaction rate.

• pH: Can affect enzyme activity and reaction rates.

• Catalysts: Increase reaction rate without being consumed. Enzymes are biological catalysts.

Biochemistry: Organic and Inorganic Compounds

Inorganic Compounds

• Water: Most abundant inorganic compound (60-80% of cell volume). Key properties:

  • High heat capacity: Absorbs/releases heat with little temperature change.

  • High heat of vaporization: Evaporation cools the body.

  • Polar solvent: Dissolves ionic substances, forms hydration layers.

  • Reactivity: Involved in hydrolysis and dehydration synthesis.

  • Cushioning: Protects organs (e.g., cerebrospinal fluid).

• Salts: Ionic compounds that dissociate in water to form electrolytes (conduct electricity). Examples: NaCl, CaCO3, KCl.

• Acids and Bases: Both are electrolytes that ionize in water.

  • Acids: Proton donors; release H+ ions (e.g., HCl, acetic acid, carbonic acid).

  • Bases: Proton acceptors; release OH- ions (e.g., NaOH, bicarbonate, ammonia).

pH and Buffers

• pH Scale: Measures hydrogen ion concentration.

  • Acidic: pH 0-6.99 (high [H+])

  • Neutral: pH 7 (equal [H+] and [OH-])

  • Alkaline (Basic): pH 7.01-14 (low [H+])

• Neutralization: Acid + base → salt + water (e.g., NaOH + HCl → NaCl + H2O)

• Buffers: Resist changes in pH by releasing or binding H+ ions. Example: Carbonic acid-bicarbonate buffer system in blood.

Summary Table: Major Chemical Bond Types

Bond Type	Description	Strength
Covalent	Sharing of electron pairs (polar or nonpolar)	Strongest
Ionic	Attraction between oppositely charged ions	Intermediate
Hydrogen	Attraction between H (partial +) and electronegative atom	Weakest

Additional info:

• Organic compounds (carbohydrates, lipids, proteins, nucleic acids) are covered in later sections.

• Enzyme mechanisms and detailed buffer systems are essential for understanding metabolic regulation.