Atoms, Molecules, and Ions: Structure, Experiments, and the Periodic Table

Atoms, Molecules, and Ions

Introduction

This chapter introduces the fundamental concepts of atomic structure, the discovery of subatomic particles, the development of atomic models, isotopes, and the organization of elements in the periodic table. It also covers the calculation of atomic and molecular masses, and the use of the mole concept in chemical calculations.

Discovery of Subatomic Particles

J. J. Thomson's Cathode Ray Tube Experiment

J. J. Thomson's experiments with cathode ray tubes in 1897 provided the first evidence for the existence of subatomic particles, specifically the electron.

• Cathode rays were observed to be deflected by electric and magnetic fields, indicating they were composed of negatively charged particles (electrons).

• Thomson measured the charge-to-mass ratio of the electron, showing it was much lighter than atoms.

• Conclusion: Atoms contain negatively charged electrons.

Example: The path of cathode rays bends toward a positively charged plate, confirming the negative charge of the particles.

Robert Millikan's Oil Drop Experiment

Robert Millikan (1909) determined the charge and mass of the electron using the oil drop experiment.

• Measured the charge on tiny oil droplets suspended in an electric field.

• Calculated the fundamental charge of the electron: $e = -1.602 \times 10^{-19}\ \text{C}$

• Determined the mass of the electron: $m_e = 9.109 \times 10^{-31}\ \text{kg}$

Example: By balancing gravitational and electrical forces on oil droplets, Millikan deduced the quantized nature of electric charge.

Atomic Models

Thomson's Plum-Pudding Model

Thomson proposed that the atom was a sphere of diffuse positive charge with negatively charged electrons embedded throughout, like "plums" in a "pudding."

• Electrons are distributed within a positively charged sphere.

• This model could not explain certain experimental results, such as the scattering of alpha particles.

Example: The model predicted that alpha particles would pass through atoms with minimal deflection.

Rutherford's Gold Foil Experiment

Ernest Rutherford and his colleagues tested the plum-pudding model by bombarding thin gold foil with alpha particles.

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

• Concluded that the atom has a small, dense, positively charged nucleus where most of the mass is concentrated.

• Electrons surround the nucleus in mostly empty space.

Example: The unexpected deflection of some alpha particles led to the nuclear model of the atom.

Comparison of Atomic Models

Model	Main Features	Experimental Evidence
Plum-Pudding Model	Diffuse positive charge with embedded electrons	Could not explain large-angle deflections
Nuclear Model	Dense nucleus with electrons in surrounding space	Explained alpha particle scattering

Structure of the Atom

Subatomic Particles

Atoms are composed of three main subatomic particles:

• Protons (p+): Positively charged, located in the nucleus, mass ≈ 1 u

• Neutrons (n0): Electrically neutral, located in the nucleus, mass ≈ 1 u

• Electrons (e-): Negatively charged, located outside the nucleus, mass ≈ 0.0005 u

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

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

Number of neutrons: $N = A - Z$

Isotopes and Nuclide Symbols

Isotopes

Isotopes are atoms of the same element (same Z) with different numbers of neutrons (different A).

• Isotopes have nearly identical chemical properties but different physical properties (e.g., mass).

• Nuclide symbol: $^{A}_{Z}\text{X}$, where X is the element symbol, A is the mass number, and Z is the atomic number.

Example: Neon has three isotopes: $^{20}_{10}\text{Ne}$, $^{21}_{10}\text{Ne}$, $^{22}_{10}\text{Ne}$

Isotopic Abundance and Average Atomic Mass

The average atomic mass of an element is the weighted average of the masses of its naturally occurring isotopes.

• Calculated using the formula: $\text{Average atomic mass} = \sum (\text{isotopic mass} \times \text{fractional abundance})$

Example (Neon):

Isotope	Mass (u)	Natural Abundance (%)
Neon-20	19.9924	90.4838
Neon-21	20.9940	0.2696
Neon-22	21.9914	9.2465

Calculation:

$\text{Average atomic mass} = (19.9924 \times 0.904838) + (20.9940 \times 0.002696) + (21.9914 \times 0.092465) = 20.1799\ \text{u}$

The Periodic Table

Organization and Classification

The periodic table arranges elements in order of increasing atomic number (Z). Elements in the same column (group) have similar chemical and physical properties.

• Groups (columns): Numbered 1–18 (IUPAC) or with A/B notation (e.g., 1A, 2A, 3B, etc.).

• Periods (rows): Numbered 1–7.

• Metals: Left and center; good conductors, malleable, ductile.

• Nonmetals: Right side; poor conductors, brittle (if solid), some are gases or liquids.

• Metalloids (semimetals): Border between metals and nonmetals; properties intermediate between metals and nonmetals.

Special Groups:

• Group 1: Alkali metals

• Group 2: Alkaline earth metals

• Group 16: Chalcogens

• Group 17: Halogens

• Group 18: Noble gases

Formulas and Mass Calculations

Chemical Formulas

Chemical formulas represent the composition of molecules and compounds using element symbols and subscripts.

• Molecular formula: Shows the exact number and type of atoms in a molecule (e.g., $\text{H}_2\text{O}$).

• Formula unit: The simplest ratio of ions in an ionic compound (e.g., $\text{NaCl}$).

Molecular and Formula Mass

The molecular mass (for molecules) or formula mass (for ionic compounds) is the sum of the atomic masses of all atoms in the chemical formula.

• Molecular mass: $\sum$ (atomic masses of all atoms in a molecule)

• Formula mass: $\sum$ (atomic masses of all atoms in a formula unit)

Example: The molecular mass of $\text{SO}_2$ is $32.07 + 2 \times 16.00 = 64.07\ \text{u}$.

The Mole Concept and Avogadro's Number

The Mole and Avogadro's Number

The mole (mol) is the SI unit for amount of substance. One mole contains Avogadro's number ($6.022 \times 10^{23}$) of particles (atoms, molecules, or formula units).

• Avogadro's number: $6.022 \times 10^{23}$ particles/mol

• Molar mass: The mass (in grams) of one mole of a substance; numerically equal to the molecular or formula mass in u.

Example: 1 mol of $\text{H}_2\text{O}$ has a mass of 18.02 g and contains $6.022 \times 10^{23}$ molecules.

Conversions Using the Mole

Conversions between mass, moles, and number of particles use the following relationships:

• $\text{Moles} = \dfrac{\text{Mass (g)}}{\text{Molar mass (g/mol)}}$

• $\text{Number of particles} = \text{Moles} \times 6.022 \times 10^{23}$

Example: To find the number of molecules in 10 g of $\text{H}_2\text{O}_2$:

• Calculate moles: $\dfrac{10\ \text{g}}{34.02\ \text{g/mol}} = 0.294\ \text{mol}$

• Calculate molecules: $0.294\ \text{mol} \times 6.022 \times 10^{23} = 1.77 \times 10^{23}$ molecules

Summary Table: Subatomic Particles

Particle	Symbol	Charge	Mass (u)
Proton	p+	+1	1.007276
Neutron	n0	0	1.008665
Electron	e-	-1	0.00054858

Key Concepts and Applications

• Atoms are composed of protons, neutrons, and electrons.

• Isotopes differ in neutron number but have the same number of protons.

• The periodic table organizes elements by atomic number and groups elements with similar properties.

• The mole allows chemists to count atoms and molecules by weighing them.

• Average atomic mass reflects the weighted average of all isotopes of an element.

Additional info: Some details, such as the exact process of calculating isotopic abundance or the full classification of periodic table groups, have been expanded for clarity and completeness.