Study Guide: Electrons and the Periodic Table

Atoms & Elements

Atomic Structure and Electron Behavior

The structure of atoms and the behavior of electrons are foundational concepts in general chemistry. Understanding how electrons are gained, lost, and arranged helps explain chemical properties and reactivity.

• Atoms gaining or using electrons: Atoms can gain or lose electrons to form ions. - Negative ion (anion): Gains electrons. - Positive ion (cation): Loses electrons.

• Atomic number: The number of protons in the nucleus of an atom determines its identity.

• Valence electrons: Electrons in the outermost shell of an atom. These are responsible for chemical reactivity.

• Electron movement: Electrons move between energy levels. - Closer to the nucleus: Lower energy. - Farther from the nucleus: Higher energy.

• Excited state: When an electron absorbs energy and moves to a higher energy level.

• Relaxation: Electrons return to lower energy levels, often emitting light of specific colors (different atoms react to different colors/wavelengths).

Example: Sodium (Na) has 11 protons and 11 electrons. If it loses one electron, it becomes Na+ (cation).

Periodic Properties of the Elements

The Periodic Table Organization

The periodic table organizes elements by atomic number and groups elements with similar properties together. Columns are called groups or families, and rows are called periods.

• Groups: Elements in the same column have the same number of valence electrons and similar chemical properties.

• Periods: Elements in the same row have the same number of electron shells.

• Metals, nonmetals, and metalloids: The periodic table is divided into these categories based on properties.

Example: Lithium (Li) and Sodium (Na) are both in Group 1 and have similar reactivity due to having one valence electron.

Bonding & Molecular Structure

Valence Electrons and Chemical Reactivity

Valence electrons determine how atoms bond and react with each other. Atoms with the same number of valence electrons are found in the same group and exhibit similar chemical behavior.

• Lewis Dot Diagrams: Visual representations of valence electrons for elements. - Example for H, Li, O, and Ne: H: 1 dot, Li: 1 dot, O: 6 dots, Ne: 8 dots.

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

• Similar properties: Elements with the same number of valence electrons have similar chemical properties.

Example: Oxygen (O) has 6 valence electrons and is in Group 16; Neon (Ne) has 8 valence electrons and is in Group 18.

Periodic Trends

Atomic Radius, Ionization Energy, and Electronegativity

Periodic trends describe how certain properties of elements change across the periodic table.

• Atomic radius: The size of an atom; increases down a group and decreases across a period (left to right).

• Ionization energy (IE): The energy required to remove an electron from an atom. - Increases across a period (left to right). - Decreases down a group.

• Electronegativity (EN): The ability of an atom to attract electrons in a chemical bond. - Increases across a period. - Decreases down a group.

• Trends: - Atomic radius: Smallest at top right, largest at bottom left. - Ionization energy and electronegativity: Highest at top right (Fluorine), lowest at bottom left.

Example: Fluorine (F) has the highest electronegativity; Cesium (Cs) has one of the largest atomic radii.

Trend	Direction of Increase	Example Element
Atomic Radius	Down a group, left across a period	Cs (large), F (small)
Ionization Energy	Up a group, right across a period	F (high), Cs (low)
Electronegativity	Up a group, right across a period	F (high), Cs (low)

Key Equations

• Atomic number: $Z = \text{number of protons}$

• Maximum electrons per shell: $2n^2$ (where n = shell number)

• Ionization energy: $\text{IE} = \text{energy required to remove an electron}$

Additional info: These notes expand on the handwritten content by providing definitions, examples, and a summary table of periodic trends for clarity and completeness.