The Atomic Structure

History

Democritus

Democritus described an indivisible object that comprises all of matter, he called it atomos.

atomos; from the greek work tomo, meaning cut; and the prefix a-, meaning not

John Dalton

John Dalton proposed the Billiard Ball model. He imagined atoms as tiny solid balls; in his model, atoms are indivisible and indestructible, meaning they cannot be split into small parts and they don’t break.

He also proposed the law of multiple proportions.

J.J. Thomson

J.J. Thomson proposed the Plum pudding model, which related atoms to the English desert of the same name, mainly as a means to refer to electrons.

He proposed that electrons were embedded in a uniform sphere of positive charge, like blueberries stuck into a muffin.

He also conducted the Cathode Ray Experiment which led to the discovery of the existence of subatomic particles.

Ernest Rutherford

Ernest Rutherford proposed the atomic model which modifies Thomson’s plum pudding model by stating that protons were concentrated within the central nucleus of the atom.

Later on, Chadwick added that neutrons were also in the nucleus as a means to stabilize it and prevent the protons from repelling from one another.

Protons and neutrons are bonded via virtual particles known as gluons.

Rutherford conducted the Gold Film Experiment that proved atoms are mostly empty space with a positive charge concentrated in the nucleus, which further backs up his Nuclear Model.

Neils Bohr

Neils Bohr proposed that electrons travel around the nucleus in a distinct circular orbit or shell, which led to him figuring out the number of electrons in each shell which is the basis of the Electron configuration.

Erwin Schrodinger

Erwin Schrodinger proposed the Quantum Model which treats electrons as matter waves which has specific positions where they might be.

His model is heavily related to Heisenberg’s Uncertainty Principle which states that knowing both the position and momentum of a particle is impossible with perfect accuracy.

His Quantum Model elevated Neils Bohr’s orbit and electron configuration proposals into the concept of orbitals, which gave us the $s$, $p$, $d$, and $f$ orbitals.

Subatomic Particles

Subatomic particles are what comprises and defines an atom’s identity, charge and stability.

	Symbol	Charge	Mass	Effect	Discoverer
Proton	$p^{+}$	$+1$	$1$	Identity	E. Rutherford
Electron	$e^{-}$	-1	$0$ (negligible)	Charge	J.J. Thomson (particle) & R.A. Milikan (charge)
Neutron	$n^o$	0	1	Stability	J. Chadwick

Terms for Elements with Varying Subatomic Particles

Term	Same	Different
Isotopes	# of Protons	Mass Number & # of Neutron
Isobar	Mass Number	# of Proton & # of Neutron
Isotone	# of Neutrons	# of Proton & Mass Number

Atomic Mass Units

Atomic Mass Units is a weighted average mass of naturally occurring isotopes of an atom

Nuclide Writing

${}_Z^A{X}_n^m$

• $X$ - Element
• $A$ - Mass Number
• $Z$ - Atomic Number (# of Protons)
• $m$ - Charge
• $n$ - Number of moles of the element

Rules in Electron Configuration

• Aufbau Principle - Lower energy levels are filled up first
• Hund’s Rule - orbital are filled up singly before pairing up
• Pauli’s Exclusion Principle - No two electron can have the same quantum number
  • This is the reason why orbitals can only occupy 2 electrons because ms should only have 2 values ( $+\frac{1}{2}$ or$-\frac{1}{2}$ )

Molecular Bonding

Intramolecular Forces

Also known as Chemical bonds.

Intramolecular forces come in three distinct types dictated by their difference in electronegativity.

• Non-Polar Covalent (Nonmetal to Nonmetal)
• Polar Covalent (Nonmetal to Nonmetal)
• Ionic (Metal to Nonmetal)

Bonding Theories

Valence or Lewis Bond Theory

This states that unpaired electrons will pair up to complete their octet and the atomic orbitals of the reactants will overlap forming molecular orbitals.

Sigma bonds ( $\sigma$ ) - single bonds. Pi bonds ( $\pi$ ) - double bonds, this refers to the consecutive bonds after the initial sigma bond.

Molecular Orbital Theory

This states that bonding electrons are not thought to be between atoms but are shared across the entire molecule.

This theory involves concepts such as antibonding.

Charge

Partial Charge

due to bonded atoms having a significant difference in electronegativity values ( $\Delta E$ )

• Neutral Covalent = $0.0–0.4$
• Polar Covalent = $0.5–2$
• Ionic (Formal Charges) = $>2$

Formal Charge

A whole integer charge.

Formal Charge Formula

• Use case:
  • For determining formal charge
• Variables:
  • $\text{formal charge}$ - The integer charge of the element.
  • ${\text{Ve}}^{-}$ - The valence electron of the element.
  • $\text{bond}$ - The number of bonds of the element.
  • ${\text{unpaired e}}^{-}$ - The number of unpaired electrons of the element
• Formula: $\text{formal charge}={\text{Ve}}^{-}-\text{bond}-{\text{unpaired e}}^{-}$

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Intermolecular Forces

These are forces between molecules or compounds, these are influenced by charge interaction and polarizability.

These forces classify molecules or compounds as dipoles or induced-dipoles

These forces are generally called Van der Waals.

H-Bonding	Keesom	Debye	London dispersion
Keesom	D to D	D to ID	ID to ID

Chemical Formulas

Organic Compounds

Formula Type	Description
Kekule / Lewis	All atoms, bonds & lone electrons
Structural	All atoms & bonds
Skeletal	Heteroatoms & bonds
Condensed	All atoms & bonds (double, triple)

Inorganic Compounds

Formula Type	Description	Compound Type	Method
Molecular	Summary of atoms present	Covalent	Valence Bonds Theory
Empirical	Subscripts are reduced	Ionic	Criss-cross