Atoms, Elements, and Chemical Composition: Pre-Test 2 Study Notes

Atoms, Elements, and Chemical Composition

Introduction

This study guide covers foundational topics in introductory chemistry, including atomic theory, the periodic table, isotopes, chemical formulas, and calculations involving chemical composition. These concepts are essential for understanding the structure of matter and the quantitative relationships in chemical reactions.

Atomic Theory and Structure

Development of Atomic Theory

• Dalton's Atomic Theory: Proposed that all matter is composed of indivisible atoms, each element consists of identical atoms, and chemical reactions involve rearrangement of these atoms.

• Discovery of the Electron: J.J. Thomson discovered the electron, a negatively charged subatomic particle, using cathode ray experiments. This led to the "plum pudding" model of the atom.

• Modern Atomic Model: Atoms consist of a dense nucleus containing protons and neutrons, surrounded by electrons in defined energy levels.

Example: Dalton's theory explained the law of constant composition, while Thomson's work led to the understanding of subatomic structure.

Subatomic Particles

• Proton (p+): Positively charged particle found in the nucleus.

• Neutron (n0): Neutral particle found in the nucleus.

• Electron (e-): Negatively charged particle found outside the nucleus.

Ions

• Cation: An atom that has lost one or more electrons, resulting in a positive charge.

• Anion: An atom that has gained one or more electrons, resulting in a negative charge.

Example: When a sodium atom loses an electron, it becomes a Na+ cation.

The Periodic Table

Organization and Classification

• Groups (Columns): Elements with similar chemical properties.

• Periods (Rows): Elements with the same number of electron shells.

• Metals, Nonmetals, Metalloids: Nonmetals are typically found on the right side of the periodic table, metals on the left, and metalloids along the zig-zag line.

Example: Silver (Ag) has atomic number 47 and is a transition metal.

Isotopes and Atomic Mass

• Isotopes: Atoms of the same element with different numbers of neutrons.

• Abundance: The average atomic mass is calculated based on the relative abundance of each isotope.

ISOTOPE	ABUNDANCE
16	50.0%
18	50.0%

Example: An element with two isotopes of equal abundance will have an average atomic mass midway between the two isotope masses.

Chemical Nomenclature and Formulas

Naming Compounds

• Ionic Compounds: Named using the cation (metal) first, followed by the anion (nonmetal) with an -ide ending.

• Polyatomic Ions: Some compounds contain polyatomic ions, such as BrO3- (bromate).

• Molecular Compounds: Use prefixes to indicate the number of each atom (e.g., dinitrogen pentasulfide: N2S5).

Example: Ca(OH)2 is named calcium hydroxide.

Empirical and Molecular Formulas

• Empirical Formula: The simplest whole-number ratio of atoms in a compound.

• Molecular Formula: The actual number of atoms of each element in a molecule.

Example: A compound with 26.68% C, 2.24% H, and 71.08% O has the empirical formula C2H2O4.

Chemical Composition Calculations

Formula Mass and Moles

• Formula Mass: The sum of the atomic masses of all atoms in a chemical formula.

• Mole: A quantity representing $6.022 \times 10^{23}$ particles (Avogadro's number).

• Calculating Moles: $\text{Moles} = \frac{\text{Mass (g)}}{\text{Molar Mass (g/mol)}}$

Example: The formula mass of copper(II) fluoride (CuF2) is calculated by adding the atomic masses of one Cu and two F atoms.

Percent Composition

• Percent Composition: The percentage by mass of each element in a compound.

• Calculation: $\%\,\text{Element} = \frac{\text{Mass of Element in 1 mol}}{\text{Molar Mass of Compound}} \times 100$

Example: To find the percent of iron in an ore sample, divide the mass of iron by the total mass of the sample and multiply by 100.

Key Laws and Principles

Law of Constant Composition

• All samples of a given compound have the same proportions of their constituent elements by mass.

Example: Water (H2O) always contains 2 hydrogen atoms for every 1 oxygen atom, regardless of the source.

Practice Problems and Applications

• Identify the correct chemical symbol for mercury: Hg.

• Calculate the formula mass for given compounds using atomic masses from the periodic table.

• Determine the empirical formula from percent composition data.

• Apply the law of constant composition to explain the uniformity of compounds.

Summary Table: Common Polyatomic Ions (Selected)

Ion	Formula	Name
Bromate	BrO3-	Bromate ion
Hydroxide	OH-	Hydroxide ion
Nitrate	NO3-	Nitrate ion

Additional info: Table inferred from context to support nomenclature questions.

Conclusion

Understanding atomic structure, the periodic table, chemical nomenclature, and composition calculations is fundamental to success in introductory chemistry. Mastery of these topics enables students to analyze chemical formulas, interpret data, and solve quantitative problems in chemistry.