Atoms and Elements: Fundamental Concepts in General Chemistry

Atoms and Elements

Introduction

This chapter introduces the foundational concepts of atoms and elements, including the laws governing chemical combinations, atomic theory, subatomic particles, isotopes, and the organization of the periodic table. Understanding these principles is essential for further study in chemistry.

Law of Conservation of Mass

Definition and Application

• Law of Conservation of Mass: The total mass of reactants in a chemical reaction is equal to the total mass of the products.

• Mass is neither created nor destroyed in ordinary chemical reactions.

• This law is fundamental to all chemical processes and is demonstrated by careful measurement of mass before and after reactions.

• Example: Burning of magnesium in air: the mass of magnesium plus oxygen before reaction equals the mass of magnesium oxide produced.

Law of Definite Proportions

Definition and Example

• Law of Definite Proportions: All samples of a given compound have the same proportions of their constituent elements by mass, regardless of the source.

• Example: Calcium oxide (CaO) always contains calcium and oxygen in a fixed mass ratio.

• Calculation for CaO:

  • 56.08 g of CaO decomposes to form 40.08 g of Ca and 16.00 g of O.

  • Mass ratio: $\frac{40.08}{16.00} = 2.505$

Law of Multiple Proportions

Definition and Example

• Law of Multiple Proportions: When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in ratios of small whole numbers.

• Example: Carbon monoxide (CO) and carbon dioxide (CO2):

  • CO: 1.33 g O reacts with 1 g C

  • CO2: 2.67 g O reacts with 1 g C

  • Ratio: $\frac{2.67}{1.33} = 2$ (a small whole number)

Applying the Law of Multiple Proportions

• To verify the law, calculate the ratio of masses of one element that combine with a fixed mass of another in different compounds.

• Example: If compound 1 has 1.50 g O per 1 g N, and compound 2 has 3.00 g O per 1 g N, the ratio is $\frac{3.00}{1.50} = 2$.

Dalton's Atomic Theory

Postulates

• 1. All matter is composed of tiny, indestructible particles called atoms.

• 2. All atoms of a given element are identical in mass and properties.

• 3. Atoms combine in simple, whole-number ratios to form compounds (e.g., AB, A2B).

• 4. Atoms are neither created nor destroyed in chemical reactions; they are rearranged.

Properties of Electric Charge

Basic Properties

• Opposite charges (positive and negative) attract each other.

• Like charges repel each other.

• An object with no net charge is termed neutral.

Behavior in Magnetic Fields

• Charged particles are deflected by magnetic fields; the direction depends on the sign of the charge.

• Positively charged particles are deflected in one direction, negatively charged in the opposite.

Discovery of the Electron

J. J. Thomson's Experiments

• Used a cathode ray tube to discover the electron, a subatomic particle with negative charge.

• Measured the charge-to-mass ratio ($\frac{e}{m}$) of the electron by balancing electric and magnetic fields.

Millikan's Oil Drop Experiment

Determining the Electron's Charge

• Robert Millikan measured the charge of a single electron by observing the motion of oil droplets in an electric field.

• Combined with Thomson's results, this allowed calculation of the electron's mass.

The Plum-Pudding Model of the Atom

J. J. Thomson's Model

• Proposed that the atom is a sphere of positive charge with negatively charged electrons embedded throughout, like "plums" in a "pudding".

Radioactivity and the Nuclear Atom

Discovery and Types of Radiation

• Henri Becquerel discovered radioactivity: spontaneous emission of radiation from certain elements.

• Types of radiation:

  • Alpha (α) particles: positively charged, high mass

  • Beta (β) particles: negatively charged, low mass

  • Gamma (γ) rays: high energy, no charge

Rutherford's Gold Foil Experiment

Experimental Design and Results

• Alpha particles were directed at a thin gold foil.

• Most particles passed through, but some were deflected at large angles.

• This led to the nuclear model of the atom: a small, dense, positively charged nucleus surrounded by electrons.

The Nuclear Theory of the Atom

Key Points

• Most of the atom's mass and all positive charge are concentrated in the nucleus.

• Electrons occupy the vast space outside the nucleus.

• Atoms are electrically neutral: number of protons equals number of electrons.

The Third Subatomic Particle: The Neutron

Discovery and Properties

• Mass measurements indicated the presence of a neutral particle, the neutron, in the nucleus.

• Neutrons have a mass similar to protons but no charge.

Summary of Subatomic Particles

Properties Table

Particle	Symbol	Charge	Mass (amu)
Proton	p+	+1	1.0073
Neutron	n0	0	1.0087
Electron	e-	-1	0.00055

Additional info: The mass of electrons is negligible compared to protons and neutrons.

Mass Spectrometer

Purpose and Use

• Measures precise masses and relative amounts of isotopes in a sample.

• Can be used to detect isotopic composition and atomic mass.

Terminology of Isotopes

Definitions

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

Names & Symbols for Isotopes

Notation

• Isotopes are represented as: Mass number AZ Chemical symbol

• Example: Carbon-12 is written as 12C

Atomic Mass

Definition and Calculation

• Atomic mass is the weighted average of the masses of all naturally occurring isotopes of an element.

• Measured in atomic mass units (amu).

• Formula: $\text{Atomic mass} = \sum (\text{fractional abundance} \times \text{isotope mass})$

Example of Weighted Averages

• Weighted averages are used to calculate atomic mass, similar to calculating a final grade from exam scores with different weights.

• Example: If an element has two isotopes with masses 10 amu (90% abundance) and 11 amu (10% abundance): $\text{Atomic mass} = (0.90 \times 10) + (0.10 \times 11) = 9.0 + 1.1 = 10.1$ amu

Example Atomic Mass Calculation

Isotope	Mass (amu)	Natural abundance (%)
Ne-20	19.992	90.48
Ne-21	20.994	0.27
Ne-22	21.991	9.25

• Calculation:

  • Convert % to decimal: 90.48% = 0.9048, etc.

  • Atomic mass = (19.992 × 0.9048) + (20.994 × 0.0027) + (21.991 × 0.0925) = 20.179 amu

Mendeleev's Periodic Table

Development and Organization

• Dmitri Mendeleev arranged elements in order of increasing atomic mass, grouping elements with similar properties.

• Predicted the existence and properties of undiscovered elements.

The Modern Periodic Table

Structure and Classification

• Elements are arranged by increasing atomic number (number of protons).

• Horizontal rows are called periods.

• Vertical columns are called groups or families; elements in a group have similar properties.

• Metals: Shiny, good conductors, malleable, ductile; found on the left and center.

• Nonmetals: Dull, poor conductors; found on the right.

• Metalloids: Properties intermediate between metals and nonmetals; found along the "stair-step" line.

Major Divisions of the Periodic Table

• Alkali metals (Group 1), Alkaline earth metals (Group 2), Transition metals (Groups 3-12), Halogens (Group 17), Noble gases (Group 18).

• Lanthanides and actinides are placed below the main table.