Atomic Structure and Models: Dalton's Theory

Atomic Structure of Matter

• Matter consists of atoms

Atomic Models of Matter

• Democritus, in the 5th century BC, proposed that atoms are eternal, indivisible, homogeneous, indestructible, and invisible.
• Atoms differ only in shape and size, but not in internal qualities.
• Material properties vary with atom grouping.
• The concept of Átomos is defined as indestructible or indivisible.
• John Dalton in 1807 proposed the Solid Atom and the Law of Definite Proportions.

Dalton's Atomic Theory

• In 1808, John Dalton started modern chemistry with his atomic theory.
• Elements are composed of tiny particles called atoms.
• Atoms of the same element are identical in size, mass, and chemical properties but differ from atoms of other elements.
• Compounds consist of atoms of more than one element.
• The ratio of the number of atoms between any two elements in any compound is always a whole number.
• A chemical reaction only involves the separation, combination, or rearrangement of atoms.
• Atoms cannot be created or destroyed during a chemical reaction.
• The Law of Multiple Proportions indicates that if two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in a ratio of small whole numbers.

J.J. Thomson's Atomic Theory

• J.J. Thomson in 1904 discovered the electron using a cathode ray tube.
• Cathode rays are attracted to a positive charge and repelled by a negative charge and consist of negatively charged particles known as electrons.
• Thomson proposed the "Plum Pudding" model of the atom.

Rutherford's Gold Foil Experiment

• In 1911, Ernest Rutherford proposed the planetary model of the atom.
• Most of the atom is empty space, which explains why most alpha particles passed through the gold foil with little to no deflection.
• Positive charges in atoms are concentrated in a dense central area called the nucleus.
• Positively charged particles in the nucleus are called protons.
• Limitations of the Rutherford model include classical mechanics and atomic stability issues.

Rutherford's Model Problem

• Rutherford's atomic structure model had an unresolved problem in the hydrogen atom and helium atom.
• Hydrogen atom has one proton, and helium has two, so the mass ratio should be 2:1, but it is 4:1.
• Later experiments found that rays consist of subatomic neutral particles called neutrons with somewhat more mass than protons.

Electromagnetic Spectrum

• The electromagnetic spectrum consists of the categorization of photons based on the properties of wavelength, frequency and energy.
• Shortest to longest wavelengths: gamma rays, X-rays, ultraviolet, visible, infrared, microwaves, radio.
• Energy and frequency increase from radio waves to gamma rays.

Quantum Principles

• Max Planck in 1900 said radiation is emitted or absorbed in discrete amounts called quanta: E = hf.
• In 1905, Albert Einstein discovered the photoelectric effect, where electromagnetic radiation (visible or ultraviolet light) causes a material to emit electrons.

Pre Quantum Models

• Niels Bohr in 1913 suggested that electrons move around the nucleus in circular orbits without emitting energy.
• Only certain orbits are allowed and defined by a mathematical expression with an integer (n = 1, 2, 3,...).
• An electron emits or absorbs energy by jumping from one allowed orbit to another in discrete amounts or photons (emission spectra).
• Limitations of the Bohr model include deficient spectral predictions with large Z.
• Arnold Sommerfeld in 1916 refined the model by introducing elliptical orbits and energy sublevels.

Quantum Principles

• Electrons and matter particles can behave as waves, explained by the De Broglie wave-particle duality principle.
• The Heisenberg uncertainty principle says that the more accurately the position of a particle is determined, the less accurately its momentum is known.

Quantum Model

• Erwin Schrödinger in 1913 developed the quantum wave mechanical model.
• A wave function ψ(x,t) describes a particle's quantum state.

Wave Functions

• A wave function ψ² represents the probability of finding an electron around the nucleus.
• Wave functions describe an atom's electron behavior probabilistically.
• Diagrams show the surface boundary of hydrogen's 1s, 2s, and 3s orbitals, each sphere containing about 90% of the total electron density.
• The size of an orbital is proportional to n², where n is the principal quantum number in hydrogen atoms.

Atomic Orbitals

• Orbitals are described by quantum numbers n, l, ml, and s.

Quantum Numbers

• n: principle quantum number, represents energy.
• l: angular momentum quantum number, orbital shape.
• m: magnetic quantum number, orbital orientation.
• s: spin quantum number, electron spin orientation.

Quantum Numbers and Electron Distribution

• Quantum mechanics specifies three quantum numbers for electron distribution in hydrogen and other atoms.
• These numbers are derived from solving. Schrödinger's equation for the hydrogen atom which are principal, angular momentum and magnetic quantum numbers.
• The spin quantum number describes electron behavior and completes the description of electrons in atoms.
• The principal quantum number (n) can be any integer, n in a hydrogen atom defines orbital energy.
• Principal quantum numbers are related to the average distance between the electron and nucleus; a greater n means the orbital size increases.

Angluar Quantum Number

• The angular momentum quantum number (l) explains the shape of the orbitals with subshells.
• Possible l values are integers beginning from 0 up to n-1.
• If n=1, l=0; if n=2, l=0 and 1; if n=3, l=0, 1, and 2.
• Designation of l value with letters, s, p, d and so on.
• If l=0, then orbital s; if l=1, then orbital p, and so on.

Magnetic Quantum Number

• For each l value, there exist (2l+1) number of ml, and m₁ indicates the orbital orientation.
• If l=1, then number of ml = 3 (2×1+1) is −1, 0, 1
• If l=2, the number of ml = 5 (2×2+1) is −2, −1, 0, 1 and 2.
• The 2p subshell has three 2p orbitals, and the ml have three values, which are −1, 0 and 1.

Spin Quantum Number

• Electron's spin can explain observed phenomena.
• Electrons behaving like tiny magnets spinning on their axis.
• The electron's spin generates a magnetic field, so the electron has a magnetic field that orients up or down, relating to +1/2 and -1/2.