Atoms and Electron Energy

Introduction

This study guide covers the fundamental structure of atoms, the nature of subatomic particles, atomic mass, electron configuration, and periodic trends such as atomic size, ionization energy, and metallic character. These concepts are essential for understanding the behavior of elements in general, organic, and biological (GOB) chemistry.

The Atom

Definition and Atomic Theory

• Atom: The smallest particle of an element that retains the chemical properties of that element.

• Dalton's Atomic Theory (1808):

  1. All matter is made up of tiny particles called atoms.

  2. All atoms of a given element are the same and different from atoms of other elements.

  3. Atoms of two or more elements combine to form compounds. A particular compound is always made up of the same kinds and numbers of atoms.

  4. Chemical reactions involve the rearrangement, separation, or combination of atoms. Atoms are neither created nor destroyed during chemical reactions.

• Modifications to Dalton's Theory:

  • Atoms of the same element are not completely identical (isotopes exist).

  • Atoms consist of even smaller particles (subatomic particles), but the atom is the smallest unit that retains element properties.

Subatomic Particles and Electrical Charges

Types and Properties of Subatomic Particles

• Proton (p or p+): Positive charge (+1), mass ≈ 1.007 amu, located in the nucleus.

• Neutron (n or n0): No charge (0), mass ≈ 1.008 amu, located in the nucleus.

• Electron (e-): Negative charge (–1), mass ≈ 0.00055 amu, located outside the nucleus.

Key Points:

• Like charges repel; opposite charges attract.

• Atoms are electrically neutral when the number of protons equals the number of electrons.

Table: Subatomic Particles

Atomic Mass Units (amu)

• Defined as 1/12th the mass of a carbon-12 atom.

• Very small unit, used to express the mass of subatomic particles.

• Also called the Dalton (Da) in biology.

Atomic Number, Mass Number, and Isotopes

Definitions and Calculations

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

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

• Isotopes: Atoms of the same element (same Z) with different numbers of neutrons (different A).

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

Example: Magnesium has isotopes Mg-24, Mg-25, and Mg-26, each with 12 protons but different numbers of neutrons.

Formula:

• Number of neutrons = Mass number – Atomic number

• Atomic mass (average) =

Electron Arrangement

Shells, Subshells, and Orbitals

• Electrons occupy regions of space called shells (energy levels), subshells, and orbitals.

• Shells: Numbered 1, 2, 3, etc., outward from the nucleus.

• Maximum electrons per shell: 1st shell = 2, 2nd shell = 8, 3rd shell = 18, etc.

• Subshells: Regions within shells, labeled s, p, d, f.

• Orbitals: Specific regions within subshells where electrons are likely to be found; each orbital holds up to 2 electrons.

Order of Filling: Electrons fill lower energy levels first, following the Aufbau principle.

Example Electron Configurations:

• Carbon (Z=6):

• Chlorine (Z=17):

• Chromium (Z=24):

Valence Electrons and the Periodic Table

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

• Group number (for main group elements) indicates the number of valence electrons.

• Lewis Dot Symbols: Dots placed around the element symbol to represent valence electrons.

Examples:

• Group 1A: 1 valence electron (e.g., Na)

• Group 2A: 2 valence electrons (e.g., Mg)

• Group 7A: 7 valence electrons (e.g., Cl)

Periodic Trends

Atomic Size (Atomic Radius)

• Atomic size increases down a group (top to bottom) due to additional electron shells.

• Atomic size decreases across a period (left to right) due to increased nuclear charge pulling electrons closer.

Ionization Energy

• Ionization energy: The energy required to remove an electron from a neutral atom.

• Decreases down a group (electrons are farther from the nucleus, easier to remove).

• Increases across a period (greater nuclear charge, harder to remove electrons).

• Equation:

Metallic Character

• Metallic character: Tendency of an element to lose valence electrons easily.

• Increases down a group; decreases across a period from left to right.

• Metals (left side of periodic table) have high metallic character; nonmetals (right side) have low metallic character; metalloids are intermediate.

Table: Periodic Trends Summary

Applications and Examples

• Identifying Elements: Use group and period to determine properties (e.g., Group 7A, Period 4 is Bromine).

• Classifying Elements: Sodium (metal), Bromine (nonmetal), Boron (metalloid), Manganese (metal), Carbon (nonmetal).

• Isotope Notation: For carbon isotopes:

  • C-12: 6 protons, 6 neutrons, 6 electrons

  • C-13: 6 protons, 7 neutrons, 6 electrons

  • C-14: 6 protons, 8 neutrons, 6 electrons

• Comparing Atomic Size and Ionization Energy:

  • Atomic size: Larger down a group, smaller across a period.

  • Ionization energy: Lower down a group, higher across a period.

Additional info: These notes are based on Timberlake, K. (2018). Chemistry: Introduction to General, Organic and Biological Chemistry (13th ed.). Pearson Education.