Models of the Atom: Structure, Subatomic Particles, and Electron Arrangement

Models of the Atom

Dalton Atomic Theory

The Dalton Atomic Theory, proposed in 1803, was the first comprehensive model describing the nature of atoms. Dalton's theory laid the foundation for modern chemistry by introducing the concept of atoms as fundamental building blocks of matter.

• Postulate 1: An element is composed of tiny, indivisible, indestructible particles called atoms.

• Postulate 2: All atoms of an element are identical and have the same properties.

• Postulate 3: Atoms of different elements combine to form compounds.

• Postulate 4: Compounds contain atoms in small whole-number ratios.

• Postulate 5: Atoms can combine in more than one ratio to form different compounds.

Thomson Model of the Atom

J. J. Thomson's experiments with cathode rays led to the discovery of the electron in 1897. He proposed the "plum pudding" model in 1903, which described the atom as a sphere of positive charge with electrons embedded throughout.

• Discovery: Electrons are negatively charged particles found within atoms.

• Plum Pudding Model: Electrons are distributed evenly in a positively charged sphere.

Robert Millikan later determined the charge of the electron in 1911, further refining the understanding of atomic structure.

Rutherford Model of the Atom

Ernest Rutherford's gold foil experiment in 1906 revealed that atoms are mostly empty space, with a dense, positively charged nucleus at the center. This experiment disproved the plum pudding model and led to the nuclear model of the atom.

• Gold Foil Experiment: Alpha particles were fired at thin gold foil; most passed through, but some were deflected.

• Conclusion: Atoms have a small, dense nucleus containing protons; electrons occupy the surrounding space.

Rutherford predicted the existence of neutral particles (neutrons), which were discovered by James Chadwick in 1932.

Subatomic Particles

Types and Properties

Atoms are composed of three main subatomic particles: electrons, protons, and neutrons. Each has distinct properties and locations within the atom.

Subatomic Particle	Symbol	Location	Relative Charge	Relative Mass
electron	e-	outside nucleus	-1	1/1836
proton	p+	inside nucleus	+1	1
neutron	n0	inside nucleus	0	1

Atomic Notation and Isotopes

Atomic Notation

Atomic notation is used to represent the composition of an atom, including its atomic number (Z), mass number (A), and element symbol.

• Atomic Number (Z): Number of protons in the nucleus.

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

• Element Symbol: One or two letters representing the element.

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons, resulting in different mass numbers. Isotopes have identical chemical properties but differ in mass.

• Example: Carbon-12 and Carbon-14 are isotopes of carbon.

• Notation: Isotopes are often written as the element name followed by the mass number (e.g., Carbon-14).

Atomic Number	Mass Number	Atomic Notation	Atomic Number	Mass Number	Atomic Notation
1	1	11H	11	23	1123Na
2	3	23He	12	24	1224Mg
3	6	36Li	13	27	1327Al
4	9	49Be	14	28	1428Si
5	10	510B	15	31	1531P
6	12	612C	16	32	1632S
7	14	714N	17	35	1735Cl
8	16	816O	18	40	1840Ar
9	19	919F	19	39	1939K
10	20	1020Ne	20	40	2040Ca

Atomic Mass

The atomic mass of an element is the weighted average mass of all its naturally occurring isotopes, taking into account their percent abundance.

• Formula: $\text{Average atomic mass} = \sum (\text{isotope mass} \times \text{fractional abundance})$

• Example: Gallium has two isotopes: $69\text{Ga}$ (68.926 amu, 60.11%) and $71\text{Ga}$ (70.925 amu, 39.89%). $\text{Average atomic mass} = (68.926 \times 0.6011) + (70.925 \times 0.3989) = 69.72$ amu

The Wave Nature of Light

Wavelength and Frequency

Light exhibits wave-like properties, characterized by wavelength and frequency. The speed of light is constant at $3.00 \times 10^8$ m/s.

• Wavelength (λ): The distance a wave travels in one cycle.

• Frequency (ν): The number of cycles per second.

• Relationship: $\text{Speed of light} = \lambda \nu$

• Inverse Relationship: Longer wavelength means lower frequency; shorter wavelength means higher frequency.

Continuous Spectrum

The radiant energy spectrum is a continuous band, with visible light ranging from 400 to 700 nm. Light can be separated into its component wavelengths using a prism.

The Quantum Concept

Quantization of Energy

The quantum concept states that energy is not continuous, but exists in discrete bundles called quanta. This is analogous to a tennis ball losing energy in steps when rolling down stairs, rather than continuously down a ramp.

• Energy Levels: Electrons in atoms occupy specific energy levels; energy changes are quantized.

Bohr Model of the Atom

Energy Levels

Niels Bohr proposed that electrons orbit the nucleus in fixed energy levels. Electrons can only exist in these quantized levels, not between them.

• Quantized Levels: Electrons move between levels by absorbing or emitting energy.

• Electron Configuration: Electrons fill the lowest available energy levels first.

Energy Levels and Sublevels

Sublevel Designations

Each main energy level contains sublevels designated as s, p, d, and f. The number of sublevels increases with the energy level.

• First energy level: 1s

• Second energy level: 2s, 2p

• Third energy level: 3s, 3p, 3d

Electron Occupancy

The maximum number of electrons in each sublevel is fixed:

• s sublevel: 2 electrons

• p sublevel: 6 electrons

• d sublevel: 10 electrons

• f sublevel: 14 electrons

Energy Level	Energy Sublevel	Electrons/Sublevel	Electrons/Energy Level
1	1s	2	2
2	2s, 2p	2, 6	8
3	3s, 3p, 3d	2, 6, 10	18
4	4s, 4p, 4d, 4f	2, 6, 10, 14	32

Electron Configurations

Filling Order and Notation

Electrons fill sublevels in order of increasing energy, starting with the lowest. The electron configuration is written as the sublevel followed by a superscript indicating the number of electrons.

• Example: Bromine (Br) has 35 electrons: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5

• Order of filling: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d ...

Quantum Mechanical Model of the Atom

Atomic Orbitals

Orbitals are regions of space with a high probability of finding electrons. The shape and size of orbitals depend on their energy and type.

• s orbitals: Spherical shape

• p orbitals: Dumbbell shape; three orientations

• d orbitals: Clover-leaf and complex shapes; five orientations

Chapter Summary

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

• Protons and neutrons are located in the nucleus; electrons are outside the nucleus.

• Atoms are mostly empty space.

• The atomic number is the number of protons in the atom.

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

• The mass number is the sum of protons and neutrons.

• The atomic mass is the weighted average of all naturally occurring isotopes.

• Light has properties of both waves and particles; photons are quantized energy packets.

• Electrons exist in discrete, quantized energy levels and fill sublevels starting with the lowest energy.