Foundations of Atomic Theory and Structure

Atomic Theory and Structure

John Dalton's Atomic Theory of Matter

John Dalton's atomic theory, proposed in the early 19th century, laid the groundwork for our modern understanding of matter. Dalton's postulates explained the composition and behavior of atoms and molecules, forming the basis for chemical reactions and the law of definite proportions.

• All matter is composed of extremely small particles called atoms. Atoms are indivisible and indestructible in chemical processes.

• All atoms of a given element are alike and differ from atoms of any other element. Each element is characterized by its unique type of atom.

• Compounds are formed when atoms of different elements combine in fixed proportions. The ratio of atoms in a compound is always constant.

• A chemical reaction involves the rearrangement of atoms. Atoms are not created or destroyed, but are simply rearranged to form new substances.

Example: Water (H2O) always consists of two hydrogen atoms and one oxygen atom.

Ernest Rutherford's Atomic Model

Ernest Rutherford's experiments in the early 20th century led to the nuclear model of the atom. He demonstrated that atoms consist of a dense, positively charged nucleus surrounded by electrons.

• The nucleus is extremely small (~0.01–0.1 Å) compared to the overall size of the atom (1–5 Å).

• Most of the atom's mass is concentrated in the nucleus, which contains protons and neutrons.

• Electrons occupy the space around the nucleus, defining the atom's size.

Example: The gold foil experiment showed that most alpha particles passed through the atom, but some were deflected, indicating a small, dense nucleus.

Subatomic Particles

Properties of Subatomic Particles

Atoms are composed of three fundamental subatomic particles: electrons, protons, and neutrons. Each has distinct properties that determine the behavior and identity of atoms.

Particle	Charge	Charge Unit	Mass
Electron (e-)	-1.602 × 10-19 C	-1	9.109 × 10-31 kg
Proton (p+)	+1.602 × 10-19 C	+1	1.672 × 10-27 kg
Neutron (n)	0	0	1.674 × 10-27 kg

• Electrons are negatively charged and have very small mass.

• Protons are positively charged and reside in the nucleus.

• Neutrons are neutral and also found in the nucleus.

Element Annotation and Atomic Notation

Standard Atomic Notation

Chemists use a standard notation to represent elements and their isotopes. This notation provides information about the atomic number, mass number, and element symbol.

• X: Element symbol (e.g., Fe for iron)

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

• Z: Atomic number (number of protons)

Notation: AZX

Example: 126C represents carbon with 6 protons and 6 neutrons.

Isotopes

Definition and Examples

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

• Protium (1H): 1 proton, 0 neutrons (ordinary hydrogen)

• Deuterium (2H): 1 proton, 1 neutron (heavy hydrogen)

• Tritium (3H): 1 proton, 2 neutrons (radioactive hydrogen)

Example: Hydrogen has three isotopes, each with different numbers of neutrons but the same number of protons.

Isotopes of Iron

Iron (Fe) has several naturally occurring isotopes, each with a different number of neutrons and varying abundance.

# protons	# electrons	# neutrons	abundance
26	26	28	0.05845
26	26	30	0.91754
26	26	31	0.02119
26	26	32	0.00282

• Average atomic mass is calculated based on the relative abundance of each isotope.

Example: The atomic mass of iron (55.845) reflects the weighted average of its isotopes.

Typical Test Question (TTQ): Atomic Structure

Practice Table: Determining Subatomic Particles

Students are often asked to determine the number of protons, neutrons, and electrons in an atom or ion using atomic notation.

Symbol	Protons	Neutrons	Electrons
120Sn	50	70	50
74	74	30	25

• Protons = atomic number

• Neutrons = mass number - atomic number

• Electrons = protons (for neutral atoms); adjust for charge in ions

Example: 120Sn has 50 protons, 70 neutrons, and 50 electrons.

Additional info: The second row in the table appears to be incomplete or a placeholder; in practice, students should use atomic and mass numbers to determine subatomic particle counts.