Atoms & Elements

Definition and Structure of Matter

All matter is anything that takes up space and has mass. 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 that retains the properties of that element.

• Matter: Anything with mass and volume (e.g., organisms, rocks, water).

• Chemical Element: A pure substance consisting of only one kind of atom.

• Atom: The smallest unit of an element, composed of subatomic particles.

Subatomic Particles

Atoms are made up 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 (neutral), mass of 1 AMU, located in the nucleus.

• Electron: Negatively charged, negligible mass, orbits the nucleus in electron shells.

Atomic Structure and Examples

Atoms consist of a dense nucleus containing protons and neutrons, surrounded by electrons in defined energy shells. The arrangement of these particles determines the atom's identity and properties.

• Nucleus: Contains protons and neutrons.

• Electron Shells: Regions where electrons are likely to be found.

Atomic Properties

Atomic Number, Mass Number, and Atomic Mass

Each atom of an element has unique properties defined by its atomic number, mass number, and atomic mass.

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

• Mass Number (A): Total number of protons and neutrons in the nucleus.

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

Formula:

Elements Essential for Life

Of all known elements, only a small subset is found in living organisms. The majority of biological mass is composed of six elements: Carbon, Hydrogen, Nitrogen, Oxygen, Phosphorus, and Sulfur (CHNOPS).

• Major Elements: Required in large amounts (e.g., CHNOPS).

• Trace Elements: Required in minute amounts for life.

Electron Configuration and Energy Shells

Electron Orbitals and Energy Shells

Electrons occupy three-dimensional regions called orbitals, grouped into energy shells around the nucleus. The arrangement of electrons in these shells determines the atom's chemical behavior.

• Shells closer to the nucleus are lower in energy; those farther away are higher in energy.

• Valence Electrons: Electrons in the outermost shell, crucial for chemical bonding.

• 1st shell holds up to 2 electrons; 2nd shell up to 8 electrons; higher shells hold more.

Octet Rule

The octet rule states that atoms are most stable when their valence shell is fully occupied, typically with 8 electrons (except for the first shell, which is full with 2 electrons). Atoms will gain, lose, or share electrons to achieve a full valence shell, driving chemical reactivity.

• Atoms with incomplete valence shells are more reactive.

• Noble gases are unreactive due to full valence shells.

Isotopes

Definition and Properties

Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons, resulting in different mass numbers. The chemical properties of isotopes are nearly identical, but their physical properties (such as mass and stability) can differ.

• Stable Isotopes: Do not undergo radioactive decay.

• Radioactive Isotopes: Unstable, decay over time, emitting radiation.

Radioactive Isotopes and Half-Life

Radioactive isotopes decay at a predictable rate, characterized by their half-life—the time required for half of the atoms in a sample to decay. These isotopes are used in medicine (e.g., cancer treatment, imaging) and in dating ancient materials (e.g., radiocarbon dating).

• Half-life: A constant for each radioactive isotope.

• Applications: Medical imaging, cancer therapy, radiometric dating.

Chemical Bonding

Types of Chemical Bonds

Chemical bonds are attractive forces that hold atoms together in molecules and compounds. They can be classified as intramolecular (within a molecule) or intermolecular (between molecules).

• Molecule: Two or more atoms chemically bonded (e.g., O2).

• Compound: Molecule composed of two or more different elements (e.g., H2O).

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

Intramolecular vs. Intermolecular Bonds

• Intramolecular Bonds: Hold atoms together within a molecule (e.g., covalent bonds in H2O).

• Intermolecular Bonds: Hold molecules together (e.g., hydrogen bonds between water molecules).

Covalent Bonds

Definition and Types

Covalent bonds involve the sharing of electron pairs between atoms. The type of covalent bond depends on the difference in electronegativity between the atoms involved.

• Nonpolar Covalent Bond: Equal sharing of electrons (similar electronegativities).

• Polar Covalent Bond: Unequal sharing of electrons (different electronegativities), resulting in partial charges (δ+ and δ−).

Electronegativity: A measure of an atom's ability to attract shared electrons (scale: 0–4).

Examples

• Nonpolar: H2, O2 (equal sharing).

• Polar: H2O (oxygen is more electronegative than hydrogen, creating partial charges).

Noncovalent Bonds

Types and Importance

Noncovalent bonds are interactions that do not involve the sharing of electrons. They include ionic bonds, hydrogen bonds, and van der Waals forces. These interactions are generally weaker than covalent bonds but are essential for the structure and function of biological molecules.

• Ionic Bonds: Attraction between oppositely charged ions.

• Hydrogen Bonds: Attraction between a hydrogen atom covalently bonded to a highly electronegative atom (F, O, N) and another electronegative atom.

• Van der Waals Forces: Weak attractions due to transient dipoles.

Ionic Bonding

Ions: Anions and Cations

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

• Anion: Negatively charged ion (gains electrons).

• Cation: Positively charged ion (loses electrons).

Ionic Bonds

Ionic bonds are formed by the electrostatic attraction between cations and anions. This transfer of electrons allows both atoms to achieve full valence shells, resulting in stable ionic compounds (e.g., NaCl).

Hydrogen Bonding

Definition and Biological Importance

Hydrogen bonds are weak interactions between a hydrogen atom covalently bonded to a highly electronegative atom (such as O, N, or F) and another electronegative atom. While individually weak, hydrogen bonds are collectively strong and play a crucial role in the properties of water and the structure of biological macromolecules (e.g., DNA, proteins).

• Responsible for water's high boiling point, surface tension, and solvent properties.

• Stabilize the three-dimensional structures of proteins and nucleic acids.