001 Atomic Models
The Atomic Model
Dalton – 1809 (Billiard Ball Model)
Fundamental Concepts:
Matter is composed of tiny atoms.
Atoms are indestructible.
All atoms of an element are identical.
Atoms of different elements differ in mass and properties.
Atoms combine in whole number ratios.
Limitation:
Dalton's model could not explain why atoms of elements combine in the specific ratios they do.
Thomson – 1897 (Plum Pudding Model)
Key Discovery:
Discovered the electron.
Model Description:
Proposed that electrons are situated within a positively charged sphere.
Limitation:
This model was disproved by Rutherford's gold foil experiment.
Rutherford – 1909 (Planetary Model)
Gold Foil Experiment:
Conducted an experiment by shooting alpha particles at a thin gold foil.
Observed that 1 in 8000 particles were significantly deflected, contrary to Thomson's model expectations.
Findings:
The atom is mainly empty space with a concentrated positive charge at its center, known as the nucleus.
Limitations:
Positive charges should repel each other.
The model did not account for the entire weight of the atom.
Electrons in motion would continuously emit radiation, leading to the implosion of atoms.
Atomic Numbers and Mass Numbers
Periodic Table Representation:
Elements are often symbolized with their mass number (A) and atomic number (Z).
Example: Oxygen is represented as O16 (8 protons, 8 neutrons).
Key Relationships:
Number of protons = number of electrons = atomic number (Z).
Number of neutrons = mass number (A) - atomic number (Z).
Example for Calculation:
Calculate the number of electrons (e-), neutrons (n0), and protons (p+) for Calcium (Ca), Argon (Ar), and Bromine (Br).
Bohr – 1913 (Energy Orbits)
Key Insights:
Energy (light) emitted by excited electrons should be continuous, however, it is not.
Electrons can occupy only specific energy levels.
Within an energy level, electrons do not emit energy.
Electrons can change levels by absorbing or emitting energy (light).
Quantum Concept:
Planck described discrete packets of energy as QUANTUM, akin to Einstein's concept of photons.
Limitation:
This model could only accurately explain single electron atoms such as H, He+, and Li2+.
Emission Spectra Homework
Assignment:
Read section 3.1.
Complete problems on page 142, questions #1-4.
The Atomic Model
Dalton – 1809 (Billiard Ball Model)
Fundamental Concepts:
Matter is composed of tiny indivisible particles called atoms.
Atoms are indestructible during chemical reactions; they cannot be created or destroyed.
All atoms of a given element are identical in mass and properties, fundamentally distinguishing them from other elements.
Atoms of different elements differ in mass and other characteristic properties.
Atoms combine in whole number ratios to form compounds, which leads to the formation of distinct molecules.
Limitations:
Dalton's model could not explain why atoms of different elements combine in specific ratios, nor could it interpret the behavior of atoms at a subatomic level including the presence of charged particles.
Thomson – 1897 (Plum Pudding Model)
Key Discovery:
Discovered the electron, establishing the first subatomic particle.
Model Description:
Proposed a model where electrons are embedded in a positively charged sphere, resembling a plum pudding or chocolate chip cookie, where the positive charge holds the negatively charged electrons in place.
Limitations:
This model was disproved by Rutherford's gold foil experiment, which indicated a more structured internal arrangement within the atom, leading to the conclusion that there is a nucleus.
Rutherford – 1909 (Planetary Model)
Gold Foil Experiment:
Conducted an experiment by bombarding a thin foil of gold with alpha particles.
Observed that 1 in 8000 particles were significantly deflected at angles that could not be explained by Thomson's model.
Findings:
Concluded that the atom is primarily empty space with a centralized, positively charged nucleus, around which electrons orbit, similar to planets orbiting the sun.
Limitations:
The model could not explain the stability of the atom; positive charges should repel each other and the model did not consider isotopes or the entire weight of the atom.
Additionally, moving electrons would continuously emit radiation, which should theoretically lead to the implosion of the atom over time.
Atomic Numbers and Mass Numbers
Periodic Table Representation:
Each element is represented by its mass number (A) and atomic number (Z). For example, Oxygen is represented as O16, indicating it has 8 protons and 8 neutrons.
Key Relationships:
The number of protons equals the number of electrons, which defines the atomic number (Z).
The number of neutrons is calculated by subtracting the atomic number from the mass number (A).
Example for Calculation:
To calculate the number of electrons (e-), neutrons (n0), and protons (p+) for elements such as Calcium (Ca, atomic number 20), Argon (Ar, atomic number 18), and Bromine (Br, atomic number 35).
Bohr – 1913 (Energy Orbits)
Key Insights:
Contrary to classical physics, Bohr observed that energy (light) emitted by excited electrons is not a continuous spectrum but consists of specific wavelengths, represented quantizably.
Electrons can only occupy specified energy levels at which they do not emit energy.
Electrons can move between these levels by either absorbing or emitting energy (light), which aligns with Planck's quantum theory.
Quantum Concept:
Introduced the idea of discrete packets of energy called QUANTUM, similar to Einstein's notion of photons in light.
Limitations:
Bohr's model only accurately explained the hydrogen atom and a few simple ions, failing to account for multi-electron atoms and complex behaviors observed in larger systems.
Emission Spectra Homework
Assignment:
Read section 3.1 thoroughly for deeper understanding.
Complete the problems listed on page 142, specifically questions #1-4 to practice calculation of atomic data and comprehension of the atomic model concepts.