Structure of Atoms: Subatomic Particles, Atomic Mass, and the Mole

Structure of Atoms

Introduction to Atomic Structure

The atom is the fundamental unit of matter, composed of a dense central nucleus surrounded by a cloud of electrons. Understanding the structure of atoms is essential for explaining chemical properties and reactions.

• Atom: The smallest unit of an element that retains its chemical properties.

• Nucleus: The tiny, dense center of the atom containing protons and neutrons.

• Scale Analogy: If an atom were the size of a stadium, the nucleus would be the size of a pea at the center, illustrating the vast empty space within atoms.

Discovery of Subatomic Particles

J.J. Thomson's Cathode-Ray Tube Experiment (1890s)

J.J. Thomson's experiments led to the discovery of the electron, the first identified subatomic particle.

• Experiment: Cathode rays were observed to be deflected by electric and magnetic fields, indicating they were composed of charged particles.

• Key Finding: The direction of deflection showed the particles were negatively charged, later named electrons.

• Conclusion: Atoms contain negatively charged electrons.

Charge Balance in Atoms

• Since atoms are electrically neutral, they must also contain positively charged particles to balance the negative charge of electrons.

Rutherford’s Scattering Experiment (1911)

Ernest Rutherford's gold foil experiment revealed the existence of the atomic nucleus and the proton.

• Experiment: Alpha particles (positively charged) were directed at a thin gold foil.

• Observation: Most alpha particles passed straight through, but a few were deflected at large angles.

• Conclusion: The atom is mostly empty space, with a small, dense, positively charged nucleus at its center.

• Discovery: The nucleus contains protons (positively charged particles).

Size of the Atom and Nucleus

• Atomic diameter: Approximately $10^{-10}$ m (1 angstrom, Å).

• Nuclear diameter: Approximately $10^{-15}$ m.

• Scale comparison: The nucleus is about 1/100,000 the size of the atom.

• Analogy: If the atom is a stadium (about 500 m), the nucleus is a pea (about 5 mm).

Subatomic Particles

Types, Location, Charge, and Mass

Atoms are composed of three main subatomic particles: electrons, protons, and neutrons.

Particle	Location	Charge	Mass (u)
Electron (e-)	Outside nucleus	-1	~0.0005
Proton (p+)	Nucleus	+1	~1
Neutron (n0)	Nucleus	0	~1

• Electrons: Negatively charged, very small mass, occupy the space around the nucleus.

• Protons: Positively charged, located in the nucleus, determine the atomic number.

• Neutrons: No charge, located in the nucleus, contribute to atomic mass.

Atomic Number, Mass Number, and Isotopes

Definitions and Notation

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

• Mass Number (A): Total number of protons and neutrons in the nucleus ($A = Z + N$).

• Isotopes: Atoms of the same element (same Z) with different numbers of neutrons (different A).

Example: Isotopes of Carbon

Isotope	Protons (Z)	Neutrons (N)	Electrons	Mass Number (A)
Carbon-12	6	6	6	12
Carbon-13	6	7	6	13

Atomic Mass and Atomic Weight

Definitions

• Atomic Mass: The mass of a specific atom, measured in unified atomic mass units (u).

• Atomic Weight: The weighted average of the atomic masses of all naturally occurring isotopes of an element.

• 1 u (atomic mass unit): Defined as 1/12 the mass of a carbon-12 atom ($1\,\text{u} = 1.6605 \times 10^{-24}\,\text{g}$).

Example: Why is the atomic weight of carbon 12.011 u?

• Carbon-12: atomic mass = 12 u; natural abundance = 98.89%

• Carbon-13: atomic mass = 13.0034 u; natural abundance = 1.11%

Calculation:

$\text{Atomic weight} = (12\,\text{u})(0.9889) + (13.0034\,\text{u})(0.0111) = 11.867\,\text{u} + 0.144\,\text{u} = 12.011\,\text{u}$

• No single carbon atom has a mass of 12.011 u; this is an average value.

The Mole and Avogadro's Number

Definition and Use

• Mole (mol): The SI unit for the amount of substance; 1 mole contains $6.022 \times 10^{23}$ entities (Avogadro's number).

• Avogadro's Number ($N_A$): $6.022 \times 10^{23}$ particles/mol.

• Molar Mass: The mass of one mole of a substance, numerically equal to its atomic or molecular weight in grams.

Example Calculation: How many Mn atoms are in 30.5 g of Mn? (Molar mass = 54.938 g/mol)

• Number of moles: $\frac{30.5\,\text{g}}{54.938\,\text{g/mol}} = 0.55517\,\text{mol}$

• Number of atoms: $0.55517\,\text{mol} \times 6.022 \times 10^{23}\,\text{atoms/mol} = 3.34 \times 10^{23}\,\text{atoms}$

Summary Table: Subatomic Particles

Particle	Symbol	Charge	Mass (u)	Location
Electron	e-	-1	0.0005	Outside nucleus
Proton	p+	+1	1.0073	Nucleus
Neutron	n0	0	1.0087	Nucleus

Key Equations

• Mass number: $A = Z + N$

• Number of particles: $\text{Number of particles} = \text{moles} \times N_A$

• Molar mass: $\text{Molar mass (g/mol)} = \text{Atomic or molecular weight (u)}$

Additional info:

• Electrons determine the chemical properties of an element, while protons determine its identity.

• Neutrons contribute to the mass and stability of the nucleus but do not affect chemical properties.