Atoms and Elements: Structure, Periodic Table, and Periodic Trends

Chapter 4: Atoms and Elements

Introduction

This chapter introduces the fundamental concepts of atoms and elements, the structure and organization of the periodic table, and the periodic trends that govern the properties of elements. Mastery of these topics is essential for understanding chemical behavior and predicting the properties of substances.

Elements and Symbols

Definition of Elements

• Element: A pure substance from which all matter is composed. Elements cannot be broken down into simpler substances by chemical means.

• Element names are often derived from planets, mythological figures, minerals, colors, geographic locations, or famous people.

Examples: Sulfur (S), Carbon (C), Bromine (Br)

Element Symbols

• Element symbols are one- or two-letter abbreviations for an element’s name. The first letter is always capitalized; the second letter, if present, is lowercase (e.g., Co for cobalt, CO for carbon and oxygen).

The Periodic Table

Development and Structure

• The periodic table was first organized by Dmitri Mendeleev in 1869, who arranged elements by increasing atomic mass and grouped elements with similar properties into columns.

• The modern periodic table is organized by increasing atomic number (number of protons).

Periods and Groups

• Periods: Horizontal rows numbered 1–7. Each period corresponds to the filling of a principal energy level with electrons.

• Groups: Vertical columns numbered 1–18. Elements in the same group have similar chemical properties due to similar valence electron configurations.

Group Names

• Group 1A (1): Alkali metals

• Group 2A (2): Alkaline earth metals

• Group 7A (17): Halogens

• Group 8A (18): Noble gases

Classification of Elements

• Metals: Located to the left of the zigzag line; shiny, ductile, malleable, good conductors, usually solid at room temperature (except mercury).

• Nonmetals: Located to the right of the zigzag line; not shiny, not ductile or malleable, poor conductors, often gases or brittle solids.

• Metalloids: Border the zigzag line; have properties intermediate between metals and nonmetals.

Examples of Metals, Nonmetals, and Metalloids

• Metals: Copper (Cu), Gold (Au), Silver (Ag), Mercury (Hg)

• Nonmetals: Hydrogen (H), Carbon (C), Nitrogen (N), Oxygen (O), Chlorine (Cl), Sulfur (S)

• Metalloids: Boron (B), Silicon (Si), Germanium (Ge), Arsenic (As), Antimony (Sb), Tellurium (Te)

The Atom

Dalton’s Atomic Theory

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

• All atoms of a given element are identical and different from atoms of other elements.

• Atoms of two or more elements combine to form compounds in fixed ratios.

• Chemical reactions involve rearrangement, separation, or combination of atoms; atoms are not created or destroyed in chemical reactions.

Subatomic Particles

• Protons (p+): Positively charged, located in the nucleus.

• Neutrons (n0): Neutral, located in the nucleus.

• Electrons (e-): Negatively charged, move in regions around the nucleus.

Atomic Number and Mass Number

• Atomic Number (Z): Number of protons in the nucleus; unique for each element.

• Mass Number (A): Sum of protons and neutrons in the nucleus; applies to a specific isotope.

Formulas:

• Number of neutrons = Mass number – Atomic number

• Number of electrons (neutral atom) = Atomic number

Isotopes and Atomic Mass

Isotopes

• Atoms of the same element with different numbers of neutrons (and thus different mass numbers).

• Isotopes are represented as: $^A_ZX$ where A = mass number, Z = atomic number, X = element symbol.

Atomic Mass

• The atomic mass of an element is the weighted average of the masses of all naturally occurring isotopes, measured in atomic mass units (amu).

• Atomic mass is listed below the symbol of each element on the periodic table.

Formula for average atomic mass:

$\text{Average atomic mass} = (\text{fractional abundance}_1 \times \text{mass}_1) + (\text{fractional abundance}_2 \times \text{mass}_2) + \ldots$

Electron Energy Levels and Arrangements

Energy Levels (Quantum Levels)

• Electrons occupy specific energy levels (n = 1, 2, 3, ...), with lower levels closer to the nucleus.

• Each energy level can hold a maximum number of electrons:

Energy Level (n)	Maximum Electrons
1	2
2	8
3	18
4	32

Electron Arrangements

• Electrons fill the lowest available energy levels first (Aufbau principle).

• Electron arrangements for the first 20 elements are important for understanding chemical properties.

Valence Electrons and Lewis Symbols

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

• The group number for representative elements (1A–8A) equals the number of valence electrons.

• Lewis symbols: Dots around the element symbol represent valence electrons.

Periodic Trends

Atomic Size

• Atomic size increases down a group and decreases across a period (left to right).

• Measured in picometers (pm) or angstroms (Å).

Ionization Energy

• Energy required to remove an outermost electron from a neutral atom.

• Ionization energy decreases down a group and increases across a period.

Metallic Character

• Metallic character increases down a group and decreases across a period.

• Elements with high metallic character lose valence electrons easily.

Summary Table: Periodic Trends

Property	Down a Group	Across a Period (Left to Right)
Atomic Size	Increases	Decreases
Ionization Energy	Decreases	Increases
Metallic Character	Increases	Decreases

Additional info: Mastery of these periodic trends is essential for predicting and explaining the chemical behavior of elements and their compounds.