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	n0	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) but with different mass numbers due to varying numbers of neutrons.

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

• Carbon has two stable isotopes: 12C (99%) and 13C (1%).

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) vs. 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 and are not stable.

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

Stable vs. Radioactive Isotopes

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

• Radioactive isotopes: Undergo decay (e.g., 14C, used in radiocarbon dating).

Example: Carbon-14 decays via beta decay, turning a neutron into a proton and emitting an electron.

Types of Radioactivity

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 a beta particle (electron).

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

Electron Arrangement and Orbitals

Flame Tests and Line Spectra

Different elements emit characteristic colors when heated in a flame due to electron transitions between energy levels. These emissions produce line spectra, which serve as "fingerprints" for elements.

• Line spectrum: Produced when light from excited atoms passes through a prism.

• Quantum: The smallest unit of energy, often released as a photon when an electron transitions between energy levels.

The Bohr Model

The Bohr model describes electrons as orbiting the nucleus in fixed energy levels (shells). When electrons absorb energy, they move to higher energy levels (excited state) and emit photons when returning to lower levels (ground state).

• Ground state: Lowest energy state of an electron.

• Excited state: Higher energy state after absorbing energy.

Electron Arrangement (Shells and Subshells)

Electrons are arranged in shells (energy levels) around the nucleus. Each shell can hold a specific number of electrons:

Element	1st shell	2nd shell	3rd shell
H	1
He	2
Li	2	1
Na	2	8	1

Each shell contains subshells (s, p, d, f), which are further divided into orbitals. Each orbital can hold a maximum of 2 electrons.

The Quantum Model

The quantum model describes electrons as existing in orbitals, which are regions of high probability for finding an electron. This model is probability-based and more accurate than the Bohr model.

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

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

• Orbitals: Regions in space where electrons are likely to be found.

Example: The second main shell contains one s orbital and three p orbitals (total 4 orbitals).

Electron Configurations

Electron configurations describe the arrangement of electrons in an atom using subshell notation. The order of filling is determined by increasing energy levels.

• Notation: For example, 1s2 2s2 2p3 for nitrogen.

• Maximum electrons per orbital: 2

• Maximum electrons per subshell: s (2), p (6), d (10), f (14)

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

Ions and Charge Balance

Cation Retention in Soils

Cations (positively charged ions) are retained in soils due to attraction to negatively charged soil particles. Charge balance is maintained as the total positive and negative charges add up to zero.

• Positive-negative charges attract

• Positive-positive and negative-negative charges repel

• Charge balance: Total charges in a system sum to zero

Practice Problems and Applications

Molar Mass and Molar Ratios

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

• Example: How many grams of carbon are in 37 g CH4? - Molar mass of CH4 = 12.01 (C) + 4 × 1.008 (H) = 16.042 g/mol - 1 mole CH4 contains 1 mole C - Grams of C = (37 g CH4) × (12.01 g C / 16.042 g CH4) ≈ 27.7 g C

Summary Table: Electron Capacity of Orbitals

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

Key Takeaways

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

• Isotopes have the same number of protons but different numbers of neutrons.

• Radioactive isotopes decay, emitting alpha, beta, or gamma radiation.

• Electron arrangement determines chemical properties and bonding.

• Electron configurations follow specific rules based on energy levels and orbitals.

Additional info: Understanding electron configurations is essential for predicting chemical bonding and reactivity.