Atoms and Matter

Definition and Structure of Matter

Matter is anything that takes up space and has mass. All matter is composed of chemical elements, which are pure substances made of only one type of atom. The atom is the smallest unit of an element and, therefore, the smallest unit of matter.

• Matter: Includes organisms, rocks, oceans, etc.

• Chemical Element: Pure substance consisting of one type of atom.

• Atom: Smallest unit of an element; makes up both living and non-living matter.

• Example: Atoms are the fundamental building blocks of all substances.

Atomic Structure and Subatomic Particles

Atoms are composed of three main subatomic particles: protons, neutrons, and electrons. Each has a characteristic charge, mass, and location within the atom.

• Proton: Positively charged, mass of 1 atomic mass unit (AMU), located in the nucleus.

• Neutron: No charge, mass of 1 AMU, located in the nucleus.

• Electron: Negatively charged, almost no mass, located in electron shells orbiting the nucleus.

Elements and Atomic Properties

Elements of Life and the Periodic Table

Only a small subset of elements is found in living organisms. The periodic table arranges all known elements based on their chemical properties. About 97% of the mass of most life is composed of Carbon, Hydrogen, Nitrogen, Oxygen, Phosphorus, and Sulfur (CHNOPS).

• Major Elements: Required for life in large amounts.

• Trace Elements: Required in small amounts.

Atomic Number, Mass Number, and Atomic Mass

Each atom of an element has unique properties:

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

• Mass Number: Sum of protons and neutrons in the nucleus.

• Atomic Mass: Weighted average mass of all isotopes of an element.

Electron Configuration and Energy Shells

Electron Orbitals and Energy Shells

Electrons occupy three-dimensional regions called orbitals, envisioned as energy shells. Shells closer to the nucleus are lower in energy, while distant shells are higher in energy. Valence electrons are found in the outermost energy shell.

• 1st shell: Holds up to 2 electrons.

• 2nd shell: Holds up to 8 electrons.

• Valence Electrons: Electrons in the outermost shell; important for chemical reactivity.

Octet Rule

The octet rule states that atoms are more stable (less reactive) when their valence shells are fully occupied, typically with eight electrons. Atoms are most reactive when their outer valence shells are not full.

• Stability: Achieved when valence shell is full.

• Example: Neon is unreactive because its valence shell is full.

Isotopes and Radioactivity

Isotopes

Isotopes are atoms of the same element that differ in the number of neutrons. They have the same atomic number but different mass numbers. Atomic mass is the average mass of all isotopes.

• Example: Carbon-12, Carbon-13, and Carbon-14 are isotopes of carbon.

Radioactive Isotopes and Half-Life

Radioactive isotopes are unstable and break down, emitting energy in the form of rays or particles. The half-life is the time it takes for half of the radioactive atoms in a sample to decay. Radioactive isotopes are used in medicine and radiometric dating.

• Half-life: Key concept in radiometric dating.

• Example: Carbon-14 has a half-life of 5,730 years.

Chemical Bonding

Introduction to Chemical Bonding

Chemical bonds are attractive forces between atoms, holding them together to form molecules and compounds. A molecule contains two or more chemically bound atoms, while a compound is a molecule composed of two or more different elements.

• Chemical Formula: Reveals the types and numbers of atoms in a molecule (e.g., C6H12O6).

Intramolecular vs. Intermolecular Bonds

Bonds between atoms can be intramolecular (within a molecule) or intermolecular (between molecules).

• Intramolecular Bonds: Hold atoms together within a molecule.

• Intermolecular Bonds: Occur between atoms of different molecules.

Covalent Bonds

Types of Covalent Bonds

Covalent bonds are interactions between two atoms resulting from the sharing of electrons. There are two types: nonpolar covalent and polar covalent, which arise due to differences in electronegativity.

• Electronegativity: Measure of an atom’s attraction to electrons (scale 0-4).

Nonpolar Covalent Bonds

Nonpolar covalent bonds involve equal sharing of electrons between atoms with similar electronegativities.

• Example: Bond between two hydrogen atoms.

Polar Covalent Bonds

Polar covalent bonds involve unequal sharing of electrons between atoms with different electronegativities, resulting in partial (δ) charges.

• Example: Bond between hydrogen and oxygen in water.

Noncovalent Bonds

Types of Noncovalent Bonds

Noncovalent bonds are interactions resulting from full or partial charges, without sharing of electrons. Types include ionic bonds, hydrogen bonds, van der Waals forces, and electrostatic bonds.

• Hydrogen Bonds: Occur between a highly electronegative atom (F, O, N) and a hydrogen atom.

• Van der Waals Bonds: Very weak interactions.

Ionic Bonding

Ions: Anions vs. Cations

Ions are atoms or molecules with a net electrical charge, resulting from the gain or loss of electrons.

• Anion: Negatively charged ion (gains electron).

• Cation: Positively charged ion (loses electron).

Ionic Bonds

Ionic bonds are electrical attractions between oppositely charged ions (cations and anions). The transfer of electrons fills the valence shells of both atoms and creates charges.

• Example: Formation of NaCl (sodium chloride).

Hydrogen Bonding

Hydrogen Bonds

Hydrogen bonds are interactions between a highly electronegative atom (F, O, N) and a hydrogen atom. Individually, hydrogen bonds are weak, but collectively, they can be strong and are important in biology, including the properties of water and the structure of macromolecules.

• Example: Hydrogen bonding between water molecules.

Key Equations

• Mass Number:

• Atomic Mass (Weighted Average):

• Half-Life Decay: , where is the number of half-lives