Chapter 2: Chemistry Fundamentals for Cell & Molecular Biology

Introduction to Cell & Molecular Biology

Course Overview

This section introduces the foundational chemical principles that underlie biological processes, focusing on atomic structure, chemical bonding, and the properties of water essential for life.

Chemistry Fundamentals: Atoms

Basic Atomic Structure

Atoms are the fundamental units of matter, composed of subatomic particles that determine their chemical properties.

• Nucleus: Contains protons (positively charged) and neutrons (neutral).

• Electrons: Negatively charged particles that orbit the nucleus in defined energy levels (shells).

• Atomic Number: Number of protons in the nucleus; defines the element.

• Mass Number: Total number of protons and neutrons.

Example: Carbon has 6 protons, 6 neutrons, and 6 electrons.

Periodic Table: Elements are organized by increasing atomic number.

Atomic Stability

Atoms are most stable when their outermost electron shell is filled. This can be achieved by forming chemical bonds.

• Covalent Bonds: Atoms share electrons to fill their valence shells.

• Number of unpaired electrons: Determines bonding capacity.

Example: Helium and Argon are stable due to filled outer shells; Sodium and Chlorine are reactive due to unfilled shells.

Chemistry Fundamentals: Bonds

Covalent Bonding

Covalent bonds form when two atoms share one or more pairs of electrons, resulting in a filled valence shell and increased stability.

• Single Covalent Bond: One pair of shared electrons (e.g., H2 molecule).

• Double/Triple Bonds: Two or three pairs of shared electrons (e.g., O2, N2).

Example: Two hydrogen atoms each share one electron to form H2.

Nonpolar and Polar Covalent Bonds

The type of covalent bond depends on the electronegativity difference between the atoms involved.

• Nonpolar Covalent Bond: Electrons are shared equally (e.g., H2).

• Polar Covalent Bond: Electrons are shared unequally, creating partial charges (e.g., H2O).

Electronegativity: The tendency of an atom to attract electrons in a bond. Increases up and to the right on the periodic table.

Example: In water, oxygen is more electronegative than hydrogen, resulting in a polar molecule.

Electron-Sharing Continuum

The degree to which electrons are shared in chemical bonds forms a continuum from equal sharing (nonpolar) to complete transfer (ionic).

• Nonpolar Covalent: Equal sharing (e.g., H2).

• Polar Covalent: Unequal sharing (e.g., H2O).

• Ionic Bond: Complete transfer of electrons (e.g., NaCl).

Ionic Bonds

Some atoms achieve stability by transferring electrons, forming ions that are held together by electrostatic attraction.

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

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

• Ionic Bond: Attraction between oppositely charged ions.

Example: Sodium donates an electron to chlorine, forming NaCl.

Properties of Water

Structure and Polarity

Water is a key molecule in biology, with unique properties due to its polar covalent bonds and ability to form hydrogen bonds.

• Polarity: Oxygen is more electronegative, creating partial negative (δ-) and positive (δ+) charges.

• Hydrogen Bonding: Weak attractions between the δ+ hydrogen of one water molecule and the δ- oxygen of another.

Example: Extensive hydrogen bonding in water leads to its unique properties.

Water as a Solvent

Water's polarity makes it an excellent solvent for ionic and polar substances.

• Hydrophilic: Substances that dissolve in water (e.g., salts, sugars).

• Hydrophobic: Substances that do not dissolve in water (e.g., oils).

Example: NaCl dissolves in water due to interaction with water's partial charges.

High Specific Heat

Water has a high specific heat, meaning it can absorb or release large amounts of energy with little temperature change.

• Specific Heat: Amount of energy required to raise the temperature of 1 gram of a substance by 1°C.

• Hydrogen Bonds: Energy is used to break many weak hydrogen bonds before temperature increases.

Substance	Specific Heat (J/g°C)
Water (H2O)	4.18
Ethanol (C2H5OH)	2.44
Glycerol (C3H8O3)	2.38
Benzene (C6H6)	1.74
Xylene (C8H10)	1.72

Application: Water's high specific heat stabilizes temperatures in cells and environments.

High Heat of Vaporization

Water requires significant energy to change from liquid to gas, due to hydrogen bonding.

• Heat of Vaporization: Energy required to convert 1 gram of liquid into vapor.

• Biological Importance: Evaporation of water (e.g., sweating) cools organisms.

Example: Sweating in mammals helps regulate body temperature.

Key Equations and Definitions

• Atomic Number: $\text{Atomic Number} = \text{Number of Protons}$

• Mass Number: $\text{Mass Number} = \text{Number of Protons} + \text{Number of Neutrons}$

• Specific Heat: $q = m \cdot c \cdot \Delta T$ where $q$ = heat energy, $m$ = mass, $c$ = specific heat, $\Delta T$ = temperature change

Summary Table: Types of Chemical Bonds

Bond Type	Electron Sharing	Example	Properties
Nonpolar Covalent	Equal	H2	No charge separation
Polar Covalent	Unequal	H2O	Partial charges, dipole
Ionic	Transferred	NaCl	Full charges, strong attraction

Additional info: These notes expand on the slide content with definitions, examples, and equations for clarity and completeness.