Atoms, Elements, and Nuclear Chemistry: Study Guide for GOB Chemistry

Atoms, Elements, and the Periodic Table

Element Names and Symbols

Each element is represented by a unique chemical symbol, usually one or two letters. The first letter is always capitalized, and the second (if present) is lowercase. For example, Na stands for sodium, and Cl stands for chlorine.

• Identify elements from symbols and vice versa using the periodic table.

• Examples: Na = Sodium, Ca = Calcium, K = Potassium.

The Periodic Table: Structure and Classification

The periodic table organizes elements by increasing atomic number and similar chemical properties.

• Groups (columns): Vertical columns numbered 1–18. Elements in the same group have similar chemical properties.

• Periods (rows): Horizontal rows numbered 1–7. Elements in the same period have the same number of occupied energy levels.

• Alkali metals: Group 1 (e.g., Li, Na, K)

• Alkaline earth metals: Group 2 (e.g., Mg, Ca)

• Halogens: Group 17 (e.g., F, Cl)

• Noble gases: Group 18 (e.g., He, Ne, Ar)

• Metals: Left and center of the table (e.g., Na, Fe)

• Nonmetals: Upper right (e.g., O, Cl)

• Metalloids: Along the staircase (e.g., Si, B)

Atoms and Subatomic Particles

An atom is the smallest unit of an element that retains its chemical properties. Atoms are composed of three main subatomic particles:

• Proton: Positive charge (+1), located in the nucleus, mass ≈ 1 amu

• Neutron: No charge (0), located in the nucleus, mass ≈ 1 amu

• Electron: Negative charge (–1), located outside the nucleus in energy levels, mass ≈ 0.0005 amu

Dalton's Atomic Theory

• All matter is composed of tiny particles called atoms.

• Atoms of the same element are identical; atoms of different elements are different.

• Atoms cannot be created or destroyed in chemical reactions (later modified due to discovery of subatomic particles and nuclear reactions).

Discovery of Subatomic Particles

• J.J. Thomson: Discovered the electron using the cathode ray experiment, showing atoms contain negatively charged particles.

• Rutherford: Gold foil experiment revealed the existence of a small, dense, positively charged nucleus at the center of the atom.

Atomic Number, Mass Number, and Isotopes

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

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

• Number of neutrons: $\text{Number of neutrons} = A - Z$

• Isotopes: Atoms of the same element with different numbers of neutrons (different mass numbers).

• Nuclear symbol: $^{A}_{Z}\text{X}$, where X is the element symbol.

Example: $^{14}_{6}\text{C}$ is carbon-14 (6 protons, 8 neutrons).

Atomic Mass Unit (amu)

The atomic mass unit is defined as 1/12 the mass of a carbon-12 atom. Protons and neutrons each have a mass of approximately 1 amu; electrons are much lighter.

Valence Electrons

Valence electrons are the electrons in the outermost energy level of an atom. They determine chemical reactivity and bonding.

• Group number (for main-group elements) often equals the number of valence electrons.

Electron Arrangement and Energy Levels

• Electrons occupy energy levels (shells) around the nucleus.

• Maximum electrons per level (simplified): 1st = 2, 2nd = 8, 3rd = 8

• Electron configuration shows the distribution of electrons among energy levels and sublevels.

Example: Oxygen (Z=8): 2 electrons in 1st level, 6 in 2nd level.

Electron Sublevels and Orbitals

• s sublevel: 1 orbital, spherical shape, holds 2 electrons

• p sublevel: 3 orbitals, dumbbell shape, holds 6 electrons

• Each orbital can hold 2 electrons with opposite spins.

Electron Configurations (s and p orbitals)

• Electrons fill the lowest energy levels first (Aufbau principle).

• Example for Oxygen (Z=8): $1s^2\ 2s^2\ 2p^4$

• Example for Chlorine (Z=17): $1s^2\ 2s^2\ 2p^6\ 3s^2\ 3p^5$

Periodic Trends

• Atomic size: Increases down a group, decreases across a period (left to right).

• Ionization energy: Decreases down a group, increases across a period.

• Metallic character: Increases down a group, decreases across a period.

Electron Dot Symbols (Lewis Symbols)

Show valence electrons as dots around the element symbol.

Nuclear Chemistry and Radioactivity

Natural Radioactivity

• Radiation: Energy or particles emitted from unstable atomic nuclei.

• Radioisotopes: Isotopes that are radioactive (e.g., $^{14}\text{C}$, $^{238}\text{U}$).

Types of Radiation

Type	Symbol	Charge	Penetrating Power
Alpha	$^4_2\text{He}$ or $\alpha$	+2	Low
Beta	$^0_{-1}\text{e}$ or $\beta$	-1	Moderate
Positron	$^0_{+1}\text{e}$	+1	Moderate
Gamma	$\gamma$	0	High

Biological Effects and Protection

• Radiation can damage living cells, cause mutations, or lead to radiation sickness.

• Protection: Increase distance, use shielding (lead, concrete), minimize exposure time.

Nuclear Reactions and Decay

• Radioactive decay: Spontaneous change of an unstable nucleus into a more stable one, emitting radiation.

• Alpha decay: Emission of an alpha particle ($^4_2\text{He}$).

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

• Positron emission: Emission of a positron ($^0_{+1}\text{e}$).

• Gamma emission: Emission of high-energy photons ($\gamma$), usually with no change in atomic or mass number.

Measuring Radiation

• Instruments: Geiger-Müller counter, scintillation counter.

• Units: Becquerel (Bq), Curie (Ci), Sievert (Sv), Gray (Gy).

• Exposure depends on intensity, time, distance, and shielding.

Half-Life of Radioisotopes

• Half-life (t1/2): Time required for half the atoms in a sample to decay.

• After one half-life: $N = N_0/2$

• After two half-lives: $N = N_0/4$

• Decay curve: Exponential decrease of radioisotope amount over time.

Medical Applications of Radioactivity

• PET scan: Uses positron-emitting isotopes to image metabolic activity.

• CT scan: Uses X-rays to create cross-sectional images.

• MRI: Uses magnetic fields and radio waves (not ionizing radiation) to image soft tissues.

Nuclear Fission and Fusion

• Nuclear fission: Splitting a heavy nucleus into smaller nuclei, releasing energy (e.g., in nuclear reactors).

• Nuclear fusion: Combining small nuclei to form a larger nucleus, releasing even more energy (e.g., in stars).

• Chain reaction: A self-sustaining series of fission reactions.

Sample Calculations and Practice Problems

Isotopic Abundance and Average Atomic Mass

• Average atomic mass is calculated using the masses and abundances of isotopes:

$\text{Average atomic mass} = (\text{mass}_1 \times \text{abundance}_1) + (\text{mass}_2 \times \text{abundance}_2) + \ldots$

Example: If isotope X (10 amu, 80%) and isotope Y (11 amu, 20%): $10 \times 0.80 + 11 \times 0.20 = 8.0 + 2.2 = 10.2$ amu

Radioactive Decay Calculations

• After n half-lives, remaining amount: $N = N_0 \times (1/2)^n$

Example: 80 g sample, half-life 10 years. After 20 years (2 half-lives): $N = 80 \times (1/2)^2 = 80 \times 1/4 = 20$ g

Summary Table: Subatomic Particles

Particle	Symbol	Charge	Location	Relative Mass (amu)
Proton	p+	+1	Nucleus	1
Neutron	n0	0	Nucleus	1
Electron	e-	-1	Outside nucleus	~0

Key Concepts and Relationships

• Elements in the same group have similar chemical properties due to similar valence electron configurations.

• Atomic number determines the identity of an element.

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

• Radioactive decay changes the nucleus, emitting radiation and forming new elements or isotopes.

• Periodic trends (atomic size, ionization energy, metallic character) help predict element behavior.

Additional info: This guide expands on the study guide outline by providing definitions, explanations, and representative examples for each topic, as would be found in a modern GOB Chemistry textbook.