Elements, Atoms, and Chemical Bonds: Foundations of General Biology

Elements and Compounds

Introduction to Matter

All living organisms are composed of matter, which is anything that occupies space and has mass. Matter is made up of elements, the fundamental substances that cannot be broken down by chemical reactions.

• Element: A pure substance consisting of only one type of atom. Examples include Hydrogen (H), Oxygen (O), and Carbon (C).

• Compound: A substance formed when two or more elements combine in a fixed ratio. Compounds have properties different from their constituent elements. Example: Water (H2O).

Subatomic Particles

Structure of Atoms

Atoms are the smallest units of matter that retain the properties of an element. They are composed of three types of subatomic particles:

• Neutrons: Particles with no electrical charge, found in the nucleus.

• Protons: Positively charged particles, also located in the nucleus.

• Electrons: Negatively charged particles, orbiting the nucleus in energy levels.

• Number of protons determines the identity of the element.

Atomic Number and Atomic Mass

Defining Elements

Each element is defined by its atomic number and atomic mass, which relate to the number of subatomic particles in its atoms.

• Atomic Number (Z): The number of protons in the nucleus of an atom.

• Mass Number (A): The sum of protons and neutrons in the nucleus.

• Atomic Mass: The total mass of an atom, approximately equal to the mass number.

Formula:

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

Isotopes

Variation in Atomic Structure

Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons.

• Isotopes: Atoms with identical atomic numbers but different mass numbers.

• Radioactive Isotopes: Unstable isotopes that decay spontaneously, emitting particles and energy.

The Energy Levels of Electrons

Electron Shells and Potential Energy

Electrons in an atom occupy specific energy levels, or shells, which determine their potential energy and chemical behavior.

• Energy: The capacity to cause change.

• Potential Energy: Energy due to position or structure; electrons farther from the nucleus have higher potential energy.

• Electron Shell: A region around the nucleus where electrons are likely to be found.

Electron Distribution and Chemical Properties

Valence Electrons and Reactivity

The distribution of electrons in an atom's electron shells determines its chemical properties and reactivity.

• Valence Electrons: Electrons in the outermost shell; they participate in chemical bonding.

• Chemical Behavior: Mostly determined by the number of valence electrons.

• Periodic Table: Shows the electron distribution for each element.

• Elements with a full valence shell are chemically inert (e.g., noble gases).

Covalent Bonds

Sharing of Electrons

Covalent bonds form when two atoms share pairs of valence electrons. The sharing can be equal or unequal, leading to nonpolar or polar covalent bonds.

• Covalent Bond: Sharing of a pair of valence electrons between atoms.

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

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

• Electronegativity: An atom's attraction for electrons in a covalent bond.

Electronegativity Difference and Bond Type:

• 0 to 0.4 = Covalent (Nonpolar)

• 0.5 to 1.7 = Polar Covalent

• 1.8 and up = Ionic

Examples of Covalent Bonds

Example	Molecular Formula	Electron Distribution Diagram	Lewis Dot Structure	Ball-and-Stick Model
Hydrogen	H2	H:H	H–H	Two spheres joined
Oxygen	O2	O::O	O=O	Two spheres joined
Water	H2O	H–O–H	O with two H attached	Three spheres, bent shape
Methane	CH4	H–C–H (tetrahedral)	C with four H attached	Five spheres, tetrahedral

Ionic Bonds

Transfer of Electrons

Ionic bonds form when one atom transfers electrons to another, resulting in the formation of ions with opposite charges that attract each other.

• Ion: A charged atom or molecule.

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

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

• Ionic Bond: The attraction between a cation and an anion.

• Ionic Compounds (Salts): Compounds formed by ionic bonds, often found as crystals in nature (e.g., sodium chloride, NaCl).

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

Additional info: The periodic table is a useful tool for predicting the types of bonds elements will form based on their position and electronegativity values.