Atoms, Isotopes, and Radioactivity: Foundations for GOB Chemistry

Atoms and Their Components

Subatomic Particles

Atoms are the fundamental building blocks of matter, composed of three main types of subatomic particles: protons, neutrons, and electrons. Each type of particle has distinct properties and locations within the atom.

• Protons: Positively charged particles found in the nucleus.

• Neutrons: Neutral particles (no charge) also located in the nucleus.

• Electrons: Negatively charged particles that move rapidly in the space surrounding the nucleus, forming the electron cloud.

• Atoms are electrically neutral because the number of protons equals the number of electrons.

Atomic Structure and Mass

The nucleus, containing protons and neutrons, is extremely small and dense compared to the overall size of the atom. The electrons occupy a much larger region called the electron cloud, which accounts for most of the atom's volume.

• The atomic mass unit (amu) is used to express the mass of subatomic particles.

• 1 amu is defined as one-twelfth the mass of a carbon-12 atom.

• Proton mass ≈ 1 amu; Neutron mass ≈ 1 amu; Electron mass ≈ 0 amu (negligible).

Atomic Number and Mass Number

Definitions and Calculations

Each element is defined by its atomic number (Z), which is the number of protons in its nucleus. The mass number (A) is the sum of protons and neutrons in the nucleus.

• Atomic Number (Z): Number of protons; determines the element's identity.

• Mass Number (A): Number of protons + number of neutrons.

• For a neutral atom, number of electrons = number of protons.

Symbolic Notation: Atoms are often represented as AZX, where X is the element symbol, A is the mass number, and Z is the atomic number.

Example: For an atom with atomic number 8 (oxygen) and 8 neutrons: Mass number = 8 (protons) + 8 (neutrons) = 16. Symbolic notation: 168O.

Steps to Determine Subatomic Particles:

1. Find the element on the periodic table.

2. Use atomic number for protons and electrons.

3. Subtract atomic number from mass number to get neutrons.

4. Write the symbolic notation.

Isotopes and Atomic Mass

Isotopes

Isotopes are atoms of the same element (same number of protons) but with different numbers of neutrons, resulting in different mass numbers. Isotopes are represented either by symbolic notation or by stating the mass number after the element name (e.g., carbon-12, carbon-13, carbon-14).

• Isotopes have nearly identical chemical properties but different physical properties (such as mass and stability).

Atomic Mass

The atomic mass listed on the periodic table is a weighted average of all naturally occurring isotopes of an element. The most common isotope is usually the one whose mass number is closest to the atomic mass.

Example: Magnesium has three naturally occurring isotopes: magnesium-24, magnesium-25, and magnesium-26. The most abundant isotope is the one closest to the average atomic mass of magnesium.

Radioactivity and Radioisotopes

Radioactivity

Radioactivity is the spontaneous emission of energy from the nucleus of an unstable atom. Isotopes that emit radiation are called radioisotopes. The process by which unstable nuclei lose energy is called radioactive decay.

• Most naturally occurring isotopes are stable; unstable isotopes emit radiation to become stable.

Forms of Ionizing Radiation

• Alpha (α) particles: Positively charged, low penetration (stopped by paper or skin).

• Beta (β) particles: Negatively charged, moderate penetration (stopped by heavy clothing or plastic).

• Gamma (γ) rays: Neutral, high penetration (requires lead or concrete for shielding).

• X-rays: Penetrate tissues but not bone; shielded by lead aprons.

• Neutrons: High penetration; shielded by concrete or water.

Type	Travel Distance Through Air	Tissue Penetration	Protective Shielding
Alpha (α)	A few centimeters	Stops at skin surface	Paper, clothing
Beta (β)	A few meters	Will not penetrate past skin layer	Heavy clothing, plastic, aluminum foil, gloves
Gamma (γ)	Several hundred meters	Fully penetrates body	Thick lead, concrete, water layer
X-ray	Several meters	Penetrates tissues but not bone	Lead apron, concrete barrier
Neutron (n)	Hundreds to thousands of meters	Fully penetrates body	Concrete barrier, water layer

Biological Effects of Radiation

• Ionizing radiation can remove electrons from atoms, making them more reactive and less stable.

• Exposure is measured in sieverts (Sv) or millisieverts (mSv); 1 mrem = 1/100,000 Sv.

• Biological effects depend on dose; high doses can cause radiation sickness or death.

Exposure (mrem)	Clinical Effect
5,000–20,000	Cannot be detected
20,000–50,000	Temporary decrease in white blood cells
50,000–100,000	Mild radiation sickness, nausea, vomiting
100,000–200,000	Light radiation poisoning, hair loss, recovery in 3 months
200,000–300,000	Moderate poisoning, permanent sterility possible
300,000–400,000	Lethal dose for half the population after 30 days (LD50)

Nuclear Equations and Radioactive Decay

Writing Nuclear Equations

Nuclear equations represent the transformation of one nucleus into another during radioactive decay. The sum of atomic numbers and mass numbers must be equal on both sides of the equation.

• Alpha decay: Emission of an alpha particle (He nucleus, 42He).

• Beta decay: Emission of a beta particle (electron, 0-1e); a neutron becomes a proton.

• Positron emission: Emission of a positron (0+1e); a proton becomes a neutron.

• Gamma decay: Emission of gamma rays (energy only, no change in atomic or mass number).

Example (Alpha Decay):

$238\ \rightarrow\ 234\ +\ 4$ $92\ \rightarrow\ 90\ +\ 2$ $\mathrm{U}\ \rightarrow\ \mathrm{Th}\ +\ \mathrm{He}$

Producing Radioisotopes: Stable isotopes can be converted to radioisotopes by bombarding them with particles (e.g., neutrons, protons).

Example:

$98\ +\ 1\ \rightarrow\ 99$ $42\ +\ 0\ \rightarrow\ 42$ $\mathrm{Mo}\ +\ n\ \rightarrow\ \mathrm{Mo}$

Radiation Units and Half-Lives

Units of Radioactivity

• Becquerel (Bq): 1 disintegration per second.

• Curie (Ci): $1\ \mathrm{Ci} = 3.7 \times 10^{10}$ disintegrations per second.

• Millicurie (mCi): $1\ \mathrm{Ci} = 1000\ \mathrm{mCi}$

• Microcurie (μCi): $1\ \mathrm{Ci} = 1,000,000\ \mu\mathrm{Ci}$

Half-Life Calculations

The half-life of a radioisotope is the time required for half of the radioactive atoms in a sample to decay. The amount remaining after n half-lives is given by:

$\text{isotope remaining} = \left(\frac{1}{2}\right)^n \times \text{starting amount}$

Where $n = \frac{\text{time elapsed}}{t_{1/2}}$

Example: If a 25 g sample has a half-life of 2 days, after 4 half-lives (8 days), $\left(\frac{1}{2}\right)^4 \times 25 = 1.56$ g remains.

Medical Applications for Radioisotopes

Diagnostic Uses

Radioisotopes are widely used in medicine for both diagnosis and treatment. In diagnostics, a small amount of a radioisotope (tracer) with a short half-life is administered to the patient. The emitted radiation is detected to create images of organs or tissues.

• Technetium-99m (99mTc): Used to diagnose blood flow in the lungs (half-life 4.8 h).

• Fluorine-18 (18F): Used in PET scans to detect functional abnormalities, especially in the brain (half-life 110 min).

Therapeutic Uses

• External beam radiation therapy: Gamma rays from cobalt-60 are directed at tumors to destroy cancerous tissue.

• Brachytherapy: Radioactive seeds are implanted in tumors for localized treatment.

• Iodine-131 (131I): Used to treat thyroid cancer; rapidly dividing cells are more susceptible to ionizing radiation.

Radioisotopes are chosen for medical use based on their type of radiation, half-life, and ability to target specific tissues.