General Biology Study Notes: Chemistry Foundations for Biology

Terminology

Matter

Matter is anything that has mass and occupies space. It is composed of atoms and molecules and exists in three states: solid, liquid, and gas.

• Mass: The quantity of matter in an object.

• Weight: The force with which an object is attracted by gravity. $\text{Wt} = \text{mass} \times \text{acceleration due to gravity}$

Atoms, Elements, Molecules, and Compounds

• Atom: The smallest unit of an element (e.g., O, H).

• Element: A pure substance that cannot be broken down into simpler substances by chemical reactions (e.g., oxygen, hydrogen).

• Molecule: The smallest unit of a compound (e.g., H2O, C12H22O11).

• Compound: Material formed from two or more elements in fixed proportions (e.g., water, sucrose).

Atomic Structure

Atomic Theory

John Dalton (Brit., 1766–1844) proposed that matter is composed of atoms, which are indivisible and combine in fixed ratios to form compounds.

Major Subatomic Particles

• Proton: Located in the nucleus, positive charge.

• Neutron: Located in the nucleus, no charge.

• Electron: Orbits the nucleus, negative charge.

Periodic Table of Elements

Organization and Use

• Lists all known elements.

• Chemical symbol: Abbreviation for element (e.g., H for hydrogen, C for carbon).

• Atomic number: Number of protons in an atom.

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

• Elements are listed in order of increasing atomic number and grouped by similar properties.

• Row (period): By energy level (shell).

• Column (group): By number of valence electrons.

Common Elements in Organisms

• C – carbon

• H – hydrogen

• N – nitrogen

• O – oxygen

• P – phosphorus

• S – sulfur

• Other: K, Na, Mg, Ca (often as electrolytes)

Isotopes

Atoms of the same element with different numbers of neutrons, resulting in different atomic masses.

Radioactive Isotopes

• Unstable isotopes that disintegrate over time, releasing subatomic particles.

• Half-life: Time required for half of isotope to disintegrate.

  • $\text{Half-life of } ^{3}\text{H} = 12 \text{ years}$

  • $\text{Half-life of } ^{125}\text{I} = 60 \text{ days}$

  • $\text{Half-life of } ^{238}\text{U} = 4.51 \text{ billion years}$

Electron Arrangement

• Energy shells (levels): K, L, M, N

• Orbitals: s, p, d, f (max 2 electrons per orbital)

• Subshells: s (1 orbital), p (3 orbitals), d (5 orbitals), f (7 orbitals)

• Shells: K shell (1st) holds 2 electrons max; L shell (2nd) holds 8 electrons max

• Octet rule: Atoms tend to fill outer shell with 8 electrons

• Valence electrons: Electrons in the outermost shell, possess highest energy

Electron Configurations

• C: 1s2 2s2 2p2

• O: 1s2 2s2 2p4

• Ne: 1s2 2s2 2p6

Bohr Models

Bohr models depict electrons in discrete energy levels (shells) around the nucleus.

Ions

• Atoms that gain or lose electrons, resulting in a net charge.

• Cation: Positively charged ion (e.g., Na+, Ca2+).

• Anion: Negatively charged ion (e.g., Cl-, F-).

Ionization

Formation of ions by loss or gain of one or more electrons.

Energy, Bonding & Electrons

Chemical Bonds

The number of electrons in the outer shell (energy level) determines how an atom reacts with other atoms.

• Atoms "want" complete sets of electrons.

• Atoms will donate, accept, or share electrons to fill valence shells.

• Noble gases have full valence shells and are generally inert.

Types of Chemical Bonds

• Ionic bond: Formed between ions with opposite charges (e.g., Na+ + Cl- → NaCl).

• Covalent bond: Formed when atoms share electrons to complete valence shells (e.g., H2, O2).

• Nonpolar covalent bond: Electrons are shared equally.

• Polar covalent bond: Electrons are shared unequally (e.g., H2O).

Electronegativity

Electronegativity is the measure of an atom's ability to attract electrons in a chemical bond.

• Higher value = greater ability to attract electrons.

• Difference in electronegativity determines bond polarity.

• General range: 0.7 to 4.0

• Difference ≥ 2.0: Ionic bond

• Difference between 0.5 and 2.0: Polar covalent bond

• Difference < 0.5: Nonpolar covalent bond

Hydrogen Bonds

Hydrogen bonds are weak attractions between molecules, not atoms, often involving hydrogen and electronegative atoms (e.g., O, N).

Water

Properties of Water

• Covers ~75% of Earth's surface

• Exists in all three states of matter

• Organisms are about 60–90% water

Biological Solvent

• Facilitates chemical reactions in the body

• "Like dissolves like" – polar molecules dissolve polar molecules

• Example: Glucose dissolves in water

Adhesion/Cohesion

• Transport medium due to hydrogen bonds

• Capillary action: Water moves up plant stems due to adhesion and cohesion

High Surface Tension

• Due to hydrogen bonds

High Specific Heat (High Heat Capacity)

• Water resists temperature changes due to hydrogen bonds

High Heat of Vaporization

• Energy required to change liquid water to gas is high

• Evaporative cooling

Less Dense as Solid

• Ice floats on water due to hydrogen bonds

Chemical Reactions

Balancing Chemical Equations

• Law of Conservation of Mass: Total mass of reactants equals total mass of products.

Chemical Calculations

• Atomic weight: Mass of atom (amu)

• Molecular mass: Sum of atomic weights of atoms in molecule

• Mole concept: 1 mole = $6.02 \times 10^{23}$ atoms or molecules (Avogadro's number)

• Molar mass: Mass of one mole of atoms/molecules in grams

Solutions & Concentrations

• Solution: Homogeneous mixture of two or more substances

• Solvent: Substance that dissolves the solute

• Solute: Substance dissolved in the solvent

Molar Concentrations

• Molarity (M): $M = \frac{\text{moles solute}}{\text{liters solution}}$

• % Concentrations (w/v): $\% = \frac{\text{grams solute}}{\text{ml solution}} \times 100$

Osmolarity

• Sum of molarity of all particles in solution

• Important for osmotic pressure calculations

• Osmolarity of body fluids regulated by kidneys

pH and Acids/Bases

pH

• Measure of H+ ion concentration in solution

• Scale ranges from 0 to 14 (logarithmic scale)

• $\text{pH} = -\log[\text{H}^+]$

Importance of pH

• Proper body functioning and medications

• Industry: manufacturing, agriculture, etc.

Acids and Bases

• Acid: Compound that dissociates into ions in water, "proton donor"

• Strong acids: 100% dissociation in water (e.g., HCl, H2SO4, HNO3)

• Weak acids: Partial dissociation (e.g., acetic acid)

• Base: Compound that dissociates into ions in water, "proton acceptor"

• Strong bases: 100% dissociation (e.g., NaOH, KOH)

• Weak bases: Partial dissociation (e.g., NH3)

Salts and Electrolytes

• Salt: Ionic compound formed from acid-base reaction (e.g., NaOH + HCl → NaCl + H2O)

• Electrolyte: Compound that dissociates in water and conducts electricity

Buffers

• Mixture of compounds that maintain constant pH

• Buffers resist drastic changes in pH

• Example: Carbonic acid buffer in blood

Buffer Equations

• Carbonic acid buffer: $\text{H}_2\text{CO}_3 \rightleftharpoons \text{HCO}_3^- + \text{H}^+$

• LeChatelier's Principle: System shifts equilibrium to relieve stress

• If acid is added: $\text{HCO}_3^- + \text{H}^+ \rightarrow \text{H}_2\text{CO}_3$

• If base is added: $\text{H}_2\text{CO}_3 + \text{OH}^- \rightarrow \text{HCO}_3^- + \text{H}_2\text{O}$

Buffer Capacity

• Measure of ability of solution to resist pH changes

• Most buffers work within a specific pH range

Summary Table: Types of Chemical Bonds

Bond Type	Definition	Example
Ionic	Transfer of electrons between atoms	NaCl
Covalent	Sharing of electrons between atoms	H2, O2
Polar Covalent	Unequal sharing of electrons	H2O
Hydrogen Bond	Weak attraction between molecules	Between H2O molecules

Summary Table: Properties of Water

Property	Description	Biological Importance
Solvent	Dissolves polar molecules	Facilitates chemical reactions
Adhesion/Cohesion	Water molecules stick to each other and surfaces	Transport in plants
High Surface Tension	Resists external force	Supports small organisms
High Specific Heat	Resists temperature change	Stabilizes climate
High Heat of Vaporization	Requires energy to evaporate	Evaporative cooling
Less Dense as Solid	Ice floats on water	Aquatic environments

Additional info: Some explanations and examples have been expanded for clarity and completeness.