Basic Definitions in Chemistry

Matter and Energy

Understanding chemistry is essential for biology, as all living things are composed of matter and governed by chemical principles. This section introduces the foundational concepts of matter and energy.

• Matter: Anything that occupies space and has mass.

• Energy: The capacity to put matter into motion. Energy exists in various forms, such as light, heat, and sound, and can be categorized as kinetic energy (energy of motion) or potential energy (stored energy due to position or structure).

• Atoms: The smallest particles of matter that retain the properties of an element and cannot be separated by chemical means.

All matter is made up of atoms, and over 100 different types of atoms exist, called elements. Of these, six elements make up about 99% of all living tissues by dry weight:

Example: Water (H2O) is composed of hydrogen and oxygen atoms bonded together.

Structure of Atoms

Subatomic Particles

Atoms consist of three main types of subatomic particles:

• Protons (p+): Positive charge, mass ≈ 1 atomic mass unit (amu).

• Neutrons (n0): No charge, mass ≈ 1 amu.

• Electrons (e-): Negative charge, mass ≈ 1/1800 amu (much lighter than protons or neutrons).

Protons and neutrons are found in the atomic nucleus; electrons orbit around the nucleus.

Atomic Number, Mass Number, and Isotopes

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

• Mass Number: Total number of protons and neutrons in an atom.

• Isotopes: Atoms of the same element with different numbers of neutrons.

Example: Hydrogen has three isotopes: Protium (1H), Deuterium (2H), and Tritium (3H).

Atomic Weight and Gram Atomic Weight

• Atomic Weight: The average mass of all the atoms of an element, taking into account isotopes.

• Gram Atomic Weight: The atomic weight expressed in grams; 1 mole of an element contains Avogadro's number () of atoms.

Electron Configuration and the Periodic Table

Electron Shells and Orbitals

Electrons occupy energy levels (shells) around the nucleus. Each shell can hold a specific number of electrons:

The valence shell is the outermost electron shell and determines chemical reactivity. Atoms are most stable when their valence shell is full (the Octet Rule).

Periodic Table Organization

• Period (row): Indicates the number of electron shells.

• Group (column): Indicates the number of electrons in the valence shell.

Example: Helium (He) is in period 1, group 18 (noble gas), with a full valence shell (2 electrons).

Chemical Bonds

Ionic Bonds

Ionic bonds form when electrons are transferred from one atom to another, resulting in oppositely charged ions that attract each other.

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

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

Example: Sodium (Na) donates an electron to chlorine (Cl), forming Na+ and Cl-, which attract to form NaCl (table salt).

Covalent Bonds

Covalent bonds form when two atoms share one or more pairs of electrons. These are the strongest and most stable bonds in biological molecules.

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

• Double Covalent Bond: Two pairs of shared electrons (e.g., O=O).

• Nonpolar Covalent Bond: Electrons are shared equally between atoms (e.g., C–H).

• Polar Covalent Bond: Electrons are shared unequally, resulting in partial charges (e.g., O–H in water).

Example: In a water molecule (H2O), the oxygen atom is more electronegative than hydrogen, creating a polar covalent bond.

Electronegativity and Bond Polarity

Electronegativity (EN): The ability of an atom to attract electrons in a bond. EN increases across a period (left to right) and decreases down a group (top to bottom) in the periodic table.

• Within any group, smaller elements (fewer energy shells) have greater EN.

• Within any period, EN increases as the number of protons increases.

Example: Oxygen (EN = 3.5) is more electronegative than hydrogen (EN = 2.1), resulting in polar bonds in water.

Hydrogen Bonds

Hydrogen bonds are weak interactions between an electronegative atom (such as oxygen or nitrogen) and a hydrogen atom covalently bonded to another electronegative atom. These bonds are responsible for many unique properties of water and play a crucial role in the structure of biological molecules.

• Example: Hydrogen bonds between water molecules contribute to water's high boiling point and surface tension.

Summary Table: Types of Chemical Bonds

Key Points for Biology Students

• Understanding atomic structure and chemical bonding is essential for studying biological molecules and processes.

• The properties of water and organic molecules depend on the types of bonds and interactions between atoms.

• Electronegativity and bond polarity are crucial for predicting molecular behavior in biological systems.

Additional info: The notes emphasize the importance of mastering these chemistry concepts before advancing in biology, as they underpin many biological phenomena, including enzyme function, DNA structure, and cellular processes.