Atomic Structure

Democritis

400 B.C. , stated that nature's basic particle was the atom ("indivisible" in Greek)

Dalton

proposed Atomic Theory, accounted for all three laws (mass, definite proportions, and multiple proportions). Consisted that all matter is composed of atoms, atoms of a given element are identical in size, mass, and all other properties, cannot be subdivided, created, or destroyed, atoms of different elements combine in simple whole-number ratios to form chemical compounds, and in chemical reactions, atoms are separated, combined, or rearranged.l

Thomson

discovered the charge to mass ratio for the electron, concluded that the electron has a very large charge-to-mass ratio, conducted Cathode-ray experiments to prove that atoms are divisible, and proposed plum pudding model for the atom

Milikan

oil drop experiment measured mass of droplets, applied charge to droplets through x-rays, applied electric voltage to top and bottom pales, measured the amount of voltage it took to keep the oil droplets suspended, and was able to measure the charge of a single electron

Rutherford

discovered the nucleus of an atom with gold-foil experiment, concluded that the volume of the nucleus was very small compared to the total volume of the atom

Hyphen notation

name of element - mass #

Nuclear symbol

Mass # SYMBOL OF

Atomic # ELEMENT

Average atomic mass formula

(mass x abundance)isotope 1 + (mass x abundance)isotope 2 + … = # amu

Electromagnetic radiation

form of energy that exhibits wavelike behavior as it travels through space

Wavelike behavior

reflection, refraction, interference. repetitive nature

Wavelength

distance from 2 equivalent points in a wave, λ, m or nm

Frequency

number of waves that pass a given point in a specific time, s^-1 or Hertz (Hz)

Crest and trough

highest point/peak of a wave ; lowest point/valley of a wave

Frequency and wavelength formula

c=λv

c=speed of light (2.998×10^8)

λ=wavelength in meters

v=frequency in s^-1 or Hz

Photoelectric effect

emission of electrons from a metal when light shines on the metal. light must be at a minimum frequency to knock an electron loose from the metal

Quantum

minimum quantity of energy that can be lost or gained by an atom

Photon

particle of electromagnetic radiation having zero mass and carrying a quantum of energy

Energy formula

E=hv

E=J

h=Planck’s constant (6.626×10^-34 Js)

v=frequency

Ground state and Excited state

lowest energy state of an atom ; state in which an atom has a higher potential energy than it has in its ground state

How line emission spectra are generated

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Bond’s model of the atom + drawbacks

electrons can only circle nucleus in paths or orbits (like planets around the sun), but could only calculate for one-electron system (hydrogen) and not multi-electron atoms

Louis de Broglie’s dual wave-particle nature of electron

suggested that electrons behaved as waves in that they could only exist at specific frequencies

Heisenburg uncertainty principle

it is impossible to determine simultaneously both the position and velocity of an electron or any other particle

Schrodinger orbitals and quantum model of the atom

electrons exist in regions called orbitals, or a 3D region around the nucleus that indicates the probable location of an electron

Hund’s rule

orbitals of equal energy are each occupied by one electron before any orbitalis occupied by a second electron, and all electrons in singly occupied orbitals must bear the same spin state

Aufbau principle

an electron occupies the lowest energy orbital that can receive it

Pauli exclusion principle

no two electrons in the same atom can have the same set of four quantum numbers

Types of electron configurations

full electron configuration, noble gas notation, orbital notation (know how to do all)

Subshells

s-1 orbital, up to 2 electrons

p-3 orbitals, up to 6 electrons

d-5 orbitals, up to 10 electrons (-1 energy)

f-7 orbitals, up to 14 electrons (-2 energy)