Atomic Structure: Radioactivity & Electron Orbitals – Study Notes

Atomic Structure

Subatomic Particles

Atoms are composed of three main subatomic particles: protons, neutrons, and electrons. Each has distinct properties and locations within the atom.

• Proton (p+): Mass = 1 u, Charge = +1, Location = Nucleus

• Neutron (n0): Mass = 1 u, Charge = 0, Location = Nucleus

• Electron (e-): Mass ≈ 1/1837 u, Charge = -1, Location = Outside nucleus

Particle	Symbol	Mass (u)	Charge	Location in Atom
Proton	p+	1	1+	Nucleus
Neutron	n	1	0	Nucleus
Electron	e-	1/1837	1-	Outside nucleus

Atomic Number and Mass Number

• Atomic number (Z): The number of protons in the nucleus of an atom. Determines the element's identity.

• Mass number (A): The sum of protons and neutrons in the nucleus.

Example: Silicon (Si) has atomic number 14 and mass number 28.085.

Isotopes

Definition and Examples

Isotopes are atoms of the same element (same atomic number) that have different numbers of neutrons (different mass numbers).

• Example: Hydrogen has three isotopes: Protium (1 proton), Deuterium (1 proton, 1 neutron), Tritium (1 proton, 2 neutrons).

• Carbon has two stable isotopes: 12C and 13C.

Key Point: Isotopes have the same number of protons but different numbers of neutrons.

Nuclear Symbol Notation

The nuclear symbol is used to represent isotopes:

• Format: AZX, where X = element symbol, A = mass number, Z = atomic number.

• Example: 16O ("light" oxygen), 18O ("heavy" oxygen, 2 extra neutrons)

Radioactivity

Definition and Types

Radioactivity is the spontaneous emission of radiation or particles from unstable atomic nuclei.

• Radioactive isotopes: Undergo radioactive decay; they are not stable.

• Half-life: The time required for half of the radioactive atoms in a sample to decay into a different isotope.

Stable vs. Radioactive Isotopes

• Stable isotopes: Do not undergo radioactive decay (e.g., 12C, 13C).

• Radioactive isotopes: Decay over time (e.g., 14C, used in radiocarbon dating).

Types of Radioactive Emissions

Name	Greek Letter	Mass (u)	Charge
Alpha particles	α	4	2+
Beta particles	β	1/1837	1-
Gamma rays	γ	0	0

• Alpha decay: Emission of an alpha particle (2 protons, 2 neutrons).

• Beta decay: A neutron turns into a proton, emitting an electron (beta particle).

• Gamma decay: Emission of high-energy photons (no mass or charge).

Example: Carbon-14 Decay

• 14C undergoes beta decay: a neutron becomes a proton, and an electron is emitted.

• This transforms 14C into 14N (nitrogen-14).

Electron Arrangement

Bohr Model

The Bohr model describes electrons as occupying specific energy levels (shells) around the nucleus.

• Ground state: Electrons are in the lowest energy state.

• Excited state: Electrons absorb energy and move to higher energy levels.

• When returning to a lower energy state, electrons emit a photon of energy (light).

Quantum Model of the Atom

The quantum model is a probability-based model that describes the likely location of electrons in orbitals.

• Principal energy levels (shells): Indicate the distance from the nucleus.

• Sublevels (subshells): Each shell is divided into sublevels (s, p, d, f).

• Orbitals: Regions of high probability for finding an electron; each orbital can hold up to 2 electrons.

Electron Orbitals and Subshells

• s orbital: Spherical, 1 per shell, holds 2 electrons.

• p orbitals: Dumbbell-shaped, 3 per shell (from n=2), hold 6 electrons total.

• d orbitals: 5 per shell (from n=3), hold 10 electrons total.

• f orbitals: 7 per shell (from n=4), hold 14 electrons total.

Orbital Type	Number of Orbitals	Maximum Electrons
s	1	2
p	3	6
d	5	10
f	7	14

Electron Configurations

Electron configurations describe the arrangement of electrons in an atom using subshell notation.

• Notation: The format is nℓx, where n = shell number, ℓ = subshell letter, x = number of electrons.

• Example: Nitrogen (N): 1s2 2s2 2p3

Electrons fill orbitals in order of increasing energy, following the Aufbau principle.

Order of Filling Orbitals

Electrons fill lower-energy orbitals first, following a specific sequence (1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc.).

• Example: Chlorine (atomic number 17): 1s2 2s2 2p6 3s2 3p5

Applications and Examples

Molar Mass and Molar Ratios

Molar mass is the mass of one mole of a substance, expressed in grams per mole (g/mol). Molar ratios are used to convert between amounts of substances in chemical reactions.

• Example: Methane (CH4): 1 mole contains 4 moles of hydrogen atoms.

• Practice: If a wetland releases 37 g CH4 per m2 per year, calculate the grams of carbon released:

$\text{Molar mass of CH}_4 = 12.01 + 4 \times 1.008 = 16.042\,\text{g/mol}$ $\text{Moles of CH}_4 = \frac{37\,\text{g}}{16.042\,\text{g/mol}} = 2.31\,\text{mol}$ $\text{Moles of C} = 2.31\,\text{mol}$ $\text{Mass of C} = 2.31\,\text{mol} \times 12.01\,\text{g/mol} = 27.7\,\text{g}$

Flame Tests and Line Spectra

• Flame tests: Different elements emit characteristic colors when heated in a flame due to electron transitions.

• Line spectra: Each element produces a unique set of spectral lines, serving as a "fingerprint" for identification.

Quantum and Photons

• Quantum: The smallest unit of energy, often associated with the energy absorbed or emitted when electrons change energy levels.

• Photon: A particle of light emitted during electron transitions.

Summary Table: Key Concepts

Concept	Definition	Example
Isotope	Atoms with same protons, different neutrons	12C, 13C
Radioactivity	Spontaneous emission from unstable nuclei	14C decay
Electron Configuration	Arrangement of electrons in shells/subshells	1s2 2s2 2p6
Orbital	Region of high probability for electrons	s, p, d, f orbitals

Takeaways

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

• Isotopes differ in neutron number; some are radioactive.

• Radioactive decay changes one element into another.

• Electrons occupy orbitals defined by quantum mechanics.

• Understanding electron configuration is essential for predicting chemical bonding and reactivity.