Unit 3: Atomic Structure, Nuclear Chemistry, and Electronic Structure Study Notes

Atomic Structure

Subatomic Particles

The atom is composed of three primary subatomic particles: protons, neutrons, and electrons. Each particle has distinct properties and plays a specific role in atomic structure.

• Proton: Positively charged particle found in the nucleus. Mass ≈ 1 amu.

• Neutron: Neutral particle found in the nucleus. Mass ≈ 1 amu.

• Electron: Negatively charged particle found in orbitals around the nucleus. Mass ≈ 0.0005 amu.

Atomic number (Z) is the number of protons in the nucleus and defines the element. Mass number (A) is the sum of protons and neutrons.

Example: Carbon-12 has 6 protons, 6 neutrons, and 6 electrons.

Isotopes and Ions

Isotopes are atoms of the same element with different numbers of neutrons, resulting in different mass numbers. Ions are atoms or molecules that have gained or lost electrons, resulting in a net charge.

• Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

Example: $^{12}C$ and $^{13}C$ are isotopes of carbon. Na+ is a sodium cation.

Isotope Abundance and Average Atomic Mass

The average atomic mass of an element is calculated using the relative abundances and masses of its isotopes.

• Formula:

$\text{Average Atomic Mass} = \sum (\text{Fractional Abundance} \times \text{Isotope Mass})$

Example: If $^{35}Cl$ (75.77%) and $^{37}Cl$ (24.23%), then:

$\text{Average Atomic Mass} = (0.7577 \times 34.969) + (0.2423 \times 36.966)$

Mass Number vs. Atomic Mass

Mass number is a whole number representing the sum of protons and neutrons in a specific isotope. Atomic mass is the weighted average mass of all isotopes of an element, usually not a whole number.

• Mass number: Specific to each isotope.

• Atomic mass: Reflects natural abundance of isotopes.

Atomic Theory and Experimental Evidence

Atomic Theory

The modern atomic theory describes atoms as the fundamental building blocks of matter, composed of a nucleus (protons and neutrons) and electrons in defined energy levels.

• Dalton's Atomic Theory: Matter is made of indivisible atoms; atoms of the same element are identical.

• Modern updates: Atoms can be divided (subatomic particles); isotopes exist.

Experimental Evidence

Key experiments led to the development of atomic theory:

• Cathode Ray Tube: Discovery of the electron (J.J. Thomson).

• Gold Foil Experiment: Discovery of the nucleus (Ernest Rutherford).

• Oil Drop Experiment: Measurement of electron charge (Robert Millikan).

Nuclear Chemistry

Nuclear Equations

Nuclear equations represent changes in atomic nuclei during radioactive decay or nuclear reactions. They must be balanced for both mass number and atomic number.

• Alpha decay: Emission of an alpha particle ($^4_2He$).

• Beta decay: Emission of a beta particle ($^0_{-1}e$).

• Gamma decay: Emission of gamma radiation ($\gamma$).

Example: $^{238}_{92}U \rightarrow ^{234}_{90}Th + ^{4}_{2}He$

Fission and Fusion

Fission is the splitting of a heavy nucleus into smaller nuclei, releasing energy. Fusion is the combining of light nuclei to form a heavier nucleus, also releasing energy.

• Fission: Used in nuclear reactors and atomic bombs.

• Fusion: Powers stars and hydrogen bombs.

Example: $^{235}U + n \rightarrow ^{141}Ba + ^{92}Kr + 3n$ (fission)

$^2H + ^3H \rightarrow ^4He + n$ (fusion)

Half-Life Problems

Half-life is the time required for half of a radioactive sample to decay. Calculations often involve exponential decay equations.

• Formula:

$N = N_0 \left(\frac{1}{2}\right)^{\frac{t}{t_{1/2}}}$

Where $N$ is the remaining amount, $N_0$ is the initial amount, $t$ is elapsed time, and $t_{1/2}$ is the half-life.

Electronic Structure

Light Equations

Light exhibits both wave and particle properties. Key equations relate energy, frequency, and wavelength.

• Speed of light: $c = \lambda \nu$

• Energy of a photon: $E = h \nu$

Where $c$ is the speed of light, $\lambda$ is wavelength, $\nu$ is frequency, $E$ is energy, and $h$ is Planck's constant.

Electron Configuration

Electron configuration describes the arrangement of electrons in an atom's orbitals, following the Aufbau principle, Pauli exclusion principle, and Hund's rule.

• Aufbau principle: Electrons fill lowest energy orbitals first.

• Pauli exclusion principle: No two electrons in an atom can have the same set of quantum numbers.

• Hund's rule: Electrons occupy degenerate orbitals singly before pairing.

Example: Oxygen: 1s2 2s2 2p4

Orbital Diagrams

Orbital diagrams visually represent electron configurations using arrows for electrons and boxes for orbitals.

• Each box represents an orbital; arrows indicate electron spin.

• Helium: ↑↓ in 1s box.

• Nitrogen: ↑↓ in 1s, ↑↓ in 2s, ↑ ↑ ↑ in 2p.

Summary Table: Key Concepts

Concept	Definition	Example
Isotope	Atoms of same element, different neutrons	$^{12}C$, $^{13}C$
Ion	Atom with net charge (lost/gained electrons)	Na+, Cl-
Half-life	Time for half of sample to decay	Carbon-14: 5730 years
Electron Configuration	Arrangement of electrons in orbitals	1s2 2s2 2p4
Fission	Splitting of heavy nucleus	$^{235}U$ fission
Fusion	Combining of light nuclei	$^2H + ^3H$ fusion

Additional info: Some explanations and examples have been expanded for clarity and completeness beyond the original test outline.