Atoms & Elements

2.1 Atoms and Their Components

Atoms are the fundamental building blocks of matter, composed of smaller subatomic particles. Understanding the structure and location of these particles is essential for studying chemical properties and reactions.

• Subatomic Particles: Atoms are made up of protons, neutrons, and electrons.

• Protons (p+): Positively charged particles located in the nucleus.

• Neutrons (n0): Neutral particles (no charge) also found in the nucleus.

• Electrons (e-): Negatively charged particles that move in regions called orbitals around the nucleus.

• Location: The nucleus (center of the atom) contains protons and neutrons, while electrons occupy the surrounding space.

• Relative Mass: Protons and neutrons have similar masses (~1 atomic mass unit, amu), while electrons are much lighter (~1/1836 amu).

• Example: A carbon atom has 6 protons, 6 neutrons, and 6 electrons.

Atoms & Elements

2.2 Atomic Number and Mass Number

The identity and properties of an atom are determined by its atomic number and mass number, which relate to the numbers of subatomic particles present.

• Atomic Number (Z): The number of protons in the nucleus of an atom. It defines the element.

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

• Distinction: Atomic number is always an integer and unique to each element; mass number can vary for atoms of the same element (isotopes).

• Calculating Subatomic Particles:

  • Number of protons = atomic number (Z)

  • Number of neutrons = mass number (A) - atomic number (Z)

  • Number of electrons = number of protons (for a neutral atom)

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

• Example: represents a sodium atom with 11 protons and 12 neutrons.

Atoms & Elements

2.3 Isotopes and Atomic Mass

Isotopes are atoms of the same element with different numbers of neutrons, leading to different mass numbers. The atomic mass of an element is a weighted average of its naturally occurring isotopes.

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

• Mass Number (A): The sum of protons and neutrons in a specific isotope.

• Atomic Mass: The weighted average mass of all naturally occurring isotopes of an element, expressed in atomic mass units (amu).

• Distinction: Mass number refers to a single isotope; atomic mass is an average for the element.

• Predicting Most Common Isotope: The isotope with a mass number closest to the atomic mass is usually the most abundant.

• Example: Chlorine has two main isotopes: and . The atomic mass of chlorine is about 35.45 amu, so $ ^{35}_{17} \mathrm{Cl} $ is more abundant.

Nuclear Chemistry

2.4 Radioactivity and Radioisotopes

Some isotopes are unstable and spontaneously emit radiation, a process known as radioactivity. The types and properties of ionizing radiation are important for understanding nuclear reactions and their effects.

• Radioactivity: The spontaneous emission of particles or energy from an unstable atomic nucleus.

• Types of Ionizing Radiation:

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

  • Beta (β) particles: High-speed electrons (), moderate penetration, stopped by aluminum.

  • Gamma (γ) rays: High-energy electromagnetic radiation, high penetration, requires thick lead or concrete to stop.

  • Positron (β+): Positively charged electron (), similar penetration to beta particles.

• Penetrating Power: Gamma > Beta > Alpha.

• Biological Exposure: Radiation exposure is measured in millirem (mrem) units. High doses can be harmful; safe exposure limits depend on the type and duration of exposure.

• Example: Alpha emitters are dangerous if ingested but not hazardous outside the body due to low penetration.

Nuclear Chemistry

2.5 Nuclear Equations and Radioactive Decay

Nuclear equations represent the changes that occur during radioactive decay. Each type of emission has a characteristic nuclear equation that must be balanced for both mass and atomic numbers.

• Alpha Decay: Emission of an alpha particle ().

  • Example:

• Beta Decay: Emission of a beta particle ().

  • Example:

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

  • Example:

• Positron Emission: Emission of a positron ().

  • Example:

• Production of Radioisotopes: Nuclear reactions can be used to produce radioisotopes for medical or industrial use.

Nuclear Chemistry

2.6 Radiation Units and Half-Lives

Radiation is measured using specific units, and the concept of half-life describes how quickly a radioactive substance decays. These concepts are essential for calculating exposure and safe handling of radioactive materials.

• Units of Radioactivity:

  • Curie (Ci): Measures the rate of radioactive decay.

  • Becquerel (Bq): SI unit for decay rate (1 Bq = 1 disintegration per second).

  • Rem (Roentgen equivalent man): Measures biological damage.

  • Sievert (Sv): SI unit for biological effect (1 Sv = 100 rem).

• Dosing Calculations: Used to determine the amount of radiation exposure or activity.

  • Example: If a sample has an activity of 2.0 mCi, and the half-life is 3 days, the activity after 6 days is mCi.

• Half-Life (t1/2): The time required for half of a radioactive sample to decay.

  • Formula:

  • Where N = remaining amount, N0 = initial amount, t = elapsed time, t1/2 = half-life.

Nuclear Chemistry

2.7 Medical Applications for Radioisotopes

Radioisotopes have important applications in medicine, both for diagnosis and treatment. Their use relies on the controlled application of radioactive decay and detection of emitted radiation.

• Diagnostic Uses: Radioisotopes such as technetium-99m are used in imaging to trace biological processes (e.g., PET scans, thyroid scans).

• Treatment: Radioisotopes can target and destroy diseased tissue, such as using iodine-131 for thyroid cancer or cobalt-60 for radiotherapy.

• Safety: The choice of isotope and dose is carefully controlled to maximize benefit and minimize harm.

• Example: A PET scan uses positron-emitting isotopes to image metabolic activity in tissues.