Atoms: The Smallest Unit of Matter

Definition and Structure

Atoms are the fundamental building blocks of matter, forming the basis of all living and nonliving things. Understanding atomic structure is essential for grasping the chemical properties of elements and molecules.

• Atom: The smallest unit of an element, retaining the properties of that element.

• Composed of three primary subatomic particles: protons (positive charge), neutrons (no charge), and electrons (negative charge).

• Protons and neutrons are located in the nucleus; electrons orbit the nucleus in electron shells.

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

• Atomic mass (atomic weight): Total number of protons and neutrons in the nucleus.

• Electrons have negligible mass compared to protons and neutrons.

Example: A carbon atom has 6 protons, 6 neutrons, and 6 electrons.

Elements of Life

Major and Trace Elements

Living organisms are primarily composed of a few key elements, with trace elements required in smaller amounts.

• About 97% of living matter is made up of CHONPS: Carbon, Hydrogen, Oxygen, Nitrogen, Phosphorus, and Sulfur.

• Trace elements (e.g., iron, iodine) are required in minute quantities but are essential for life.

Atomic Properties

Key Atomic Numbers and Masses

Understanding atomic number, mass number, and atomic mass is fundamental in chemistry.

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

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

• Atomic mass (atomic weight): Weighted average mass of all isotopes of an element.

Formula:

Example: Carbon-12 has 6 protons and 6 neutrons; mass number = 12.

Electron Orbitals and Energy Shells

Organization of Electrons

Electrons occupy specific regions around the nucleus called shells or orbitals, which determine the atom's chemical behavior.

• Shells closer to the nucleus have lower energy; outer shells have higher energy.

• Valence electrons: Electrons in the outermost shell; determine chemical reactivity.

• The first shell holds up to 2 electrons, the second up to 8, and so on.

Octet Rule

Atoms are most stable when their valence shell is full, typically with 8 electrons (except for the first shell, which is full with 2).

Electronegativity

Definition and Trends

Electronegativity is the tendency of an atom to attract electrons in a chemical bond. It influences bond type and molecular polarity.

• Increases across a period (left to right) and decreases down a group (top to bottom) in the periodic table.

• Oxygen and fluorine are among the most electronegative elements.

• Differences in electronegativity between atoms determine bond polarity.

Example: In water (H2O), oxygen is more electronegative than hydrogen, resulting in a polar covalent bond.

Isotopes and Radioactive Isotopes

Definition and Applications

Isotopes are atoms of the same element with different numbers of neutrons, resulting in different mass numbers.

• Stable isotopes: Do not decay over time.

• Radioactive isotopes: Unstable; decay spontaneously, emitting radiation.

• Half-life: Time required for half the atoms in a sample to decay.

• Used in medicine (e.g., cancer treatment) and dating fossils (radiometric dating).

Molecules and Compounds

Definitions

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

• Compound: Two or more different elements chemically bonded (e.g., H2O).

• INTRA-molecular bonds: Bonds within a molecule (e.g., covalent bonds in H2O).

• INTER-molecular bonds: Bonds between different molecules (e.g., hydrogen bonds between water molecules).

Chemical Bonding Types

Covalent Bonding

Covalent bonds form when atoms share electrons. They can be non-polar (equal sharing) or polar (unequal sharing).

• Non-polar covalent bond: Electrons shared equally (e.g., O2).

• Polar covalent bond: Electrons shared unequally, leading to partial charges (e.g., H2O).

Non-covalent Bonding

Non-covalent bonds include ionic bonds and hydrogen bonds, which are generally weaker than covalent bonds but crucial for biological processes.

• Ionic bond: Formed when one atom donates an electron to another, resulting in oppositely charged ions (e.g., NaCl).

• Hydrogen bond: Weak interaction between a hydrogen atom (bonded to F, O, or N) and another electronegative atom.

Ionic Bonding

Ionic bonds occur between atoms with large differences in electronegativity, resulting in the transfer of electrons.

• Cation: Positively charged ion (loss of electron).

• Anion: Negatively charged ion (gain of electron).

• Strength of ionic bond depends on the difference in electronegativity (usually >1.7).

Hydrogen Bonding

Hydrogen bonds are weak individually but collectively strong, playing a key role in the properties of water and the structure of macromolecules.

• Occurs between a hydrogen atom covalently bonded to F, O, or N and another electronegative atom.

• Important in DNA structure and protein folding.

Summary Table: Types of Chemical Bonds

Additional info:

• Electronegativity values are highest for fluorine and oxygen, which is why these elements often participate in hydrogen bonding.

• Radioactive isotopes are also used as tracers in biological research.