Basic Chemistry Concepts for Anatomy & Physiology

Basic Chemistry Concepts

Elements and Atoms

Elements are the simplest forms of matter that cannot be broken down by chemical reactions. They are fundamental to understanding the chemical basis of life and are organized in the Periodic Table of Elements. In biological systems, elements are classified based on their abundance and importance.

• Major elements: Carbon (C), Oxygen (O), Hydrogen (H), Nitrogen (N)

• Minor elements: Calcium (Ca), Sodium (Na), Potassium (K), Iron (Fe), Chloride (Cl), Phosphorus (P), Sulfur (S), Iodine (I), Magnesium (Mg)

• Trace elements: Present in very small amounts, e.g., Fluorine (F)

Each element is made up of atoms, which consist of three types of subatomic particles:

• Protons: Positively charged particles located in the nucleus

• Neutrons: Neutral particles also found in the nucleus

• Electrons: Negatively charged particles that move around the nucleus in orbitals or an electron cloud

The atomic number (top number in the Periodic Table) equals the number of protons. The atomic weight (bottom number) approximates the sum of protons and neutrons (mass number). The number of electrons in a neutral atom equals the number of protons.

• Number of neutrons = Mass number - Atomic number

Example: For carbon (atomic number 6, atomic weight ≈ 12), there are 6 protons, 6 neutrons, and 6 electrons.

Ions and Isotopes

Atoms can gain or lose electrons, forming ions:

• Cation: Positively charged ion (e.g., Na+), formed by losing electrons

• Anion: Negatively charged ion (e.g., Cl-), formed by gaining electrons

• Some ions have charges greater than one (e.g., Ca2+ has lost two electrons)

Isotopes are atoms of the same element with different numbers of neutrons (same atomic number, different mass number). Some isotopes are unstable and radioactive.

• Example: Carbon-12 (6 protons, 6 neutrons), Carbon-14 (6 protons, 8 neutrons) – Carbon-14 is used in carbon dating.

• Medical isotopes: Iodine-131 (thyroid imaging/treatment), Sodium-24 (electrolyte studies)

Molecules and Compounds

Definitions and Examples

A molecule forms when two or more atoms join chemically (e.g., H2O, O2, NaCl). A compound contains at least two different elements joined together (e.g., NaCl, H2O). All compounds are molecules, but not all molecules are compounds.

Chemical Bonds

Covalent Bonds

Covalent bonds involve the sharing of electrons between atoms and are the strongest type of chemical bond. The sharing can be equal (nonpolar) or unequal (polar).

• Single bond: One pair of electrons shared (e.g., CH4)

• Double bond: Two pairs of electrons shared (e.g., O2)

• Triple bond: Three pairs of electrons shared (e.g., N2)

Polar covalent bonds: Electrons are shared unequally, resulting in partial charges (e.g., H2O – oxygen is partially negative, hydrogens are partially positive).

Nonpolar covalent bonds: Electrons are shared equally (e.g., CO2).

Ionic Bonds

Ionic bonds are formed by the attraction between cations and anions (opposite charges). They are intermediate in strength.

• Example: Sodium chloride (NaCl) – Na+ (cation) and Cl- (anion)

• Dietary note: Iodized salt contains NaI to prevent iodine deficiency

Hydrogen Bonds

Hydrogen bonds are weak attractions between the positive end of a polar molecule and the negative end of another. They are crucial for the structure of DNA and proteins, and for the unique properties of water (e.g., high surface tension).

• Example: Hydrogen bonding between water molecules allows insects like water striders to walk on water.

Solutions

A solution is a homogeneous mixture of a solute dissolved in a solvent. In biological systems, water is the main solvent.

• Solute: Substance being dissolved (e.g., salt in saline solution)

• Solvent: Substance doing the dissolving (e.g., water)

Example: In saline, salt is the solute and water is the solvent.

Chemical Reactions

All chemical reactions in the body are collectively called metabolism. Reactions that build molecules are anabolic, while those that break down molecules are catabolic. When these processes are balanced, the system is at equilibrium.

• Reactants (substrates): Starting materials

• Products: Substances formed

Chemical reactions are represented as:

• $\text{Substrates} \rightarrow \text{Products}$

Types of Chemical Reactions

• Synthesis (Anabolic): Building larger molecules from smaller ones $\mathrm{A} + \mathrm{B} \rightarrow \mathrm{AB}$

• Decomposition (Catabolic): Breaking down molecules $\mathrm{AB} \rightarrow \mathrm{A} + \mathrm{B}$

• Exchange: Rearrangement of components between molecules $\mathrm{AB} + \mathrm{CD} \rightarrow \mathrm{AD} + \mathrm{CB}$

Reactions can be reversible (can proceed in both directions):

• $\mathrm{A} + \mathrm{B} \rightleftharpoons \mathrm{AB}$

Catabolic reactions are usually exergonic (release energy), while anabolic reactions are endergonic (absorb energy).

Summary Table: Types of Chemical Bonds

Bond Type	Description	Relative Strength	Example
Covalent	Electrons are shared between atoms	Strongest	H2O, O2, CH4
Ionic	Attraction between oppositely charged ions	Intermediate	NaCl
Hydrogen	Attraction between polar molecules	Weakest	Between water molecules, DNA strands

Key Terms

• Element: Pure substance made of only one kind of atom

• Atom: Smallest unit of an element

• Ion: Atom with a net electric charge due to loss or gain of electrons

• Isotope: Atoms of the same element with different numbers of neutrons

• Molecule: Two or more atoms bonded together

• Compound: Molecule containing at least two different elements

• Solute: Substance dissolved in a solvent

• Solvent: Substance that dissolves the solute

• Metabolism: All chemical reactions in the body

• Anabolic: Building up molecules

• Catabolic: Breaking down molecules

• Exergonic: Reactions that release energy

• Endergonic: Reactions that absorb energy

Additional info: This guide expands on the original notes by providing definitions, examples, and a summary table for clarity and exam preparation.