ATOM MODEL HISTORY Democritus Fifth century B C

ATOM MODEL HISTORY

Democritus � Fifth century B. C. Greek philosopher � All matter composed of indivisible and indestructible particles called atoms (Greek for uncuttable). � Though later challenged, many ideas agreed with later theory

Dalton’s Atomic Theory � � � Early 1800’s - Billiard Ball Model Viewed the atom as a small, solid, indivisible sphere. Atoms of each element identical in mass and properties. Atoms of one element differ from another atom. Got the "ball" rolling for modern chemistry!

J. J. Thomson � � � Late 1800’s-Plum Pudding Model Atom was a sphere of positive electricity which was diffuse with negative particles imbedded throughout Discovered the electron Experiments that passed electric currents through gases at low pressure Attracted to + charge so particles of current must be – Nobel Prize in physics in 1906.

What he did…. .

Rutherford � Solar System Model � Discovered the nucleus- Gold Foil Experiment � Atom mostly empty space with dense positively charged nucleus surrounded by negative electrons. � Nobel Prize in chemistry in 1908

What this shows… � � � flying dots = Alpha particles( +) emitted green wall = detecting screen (where particles hit) yellow window in the middle = gold foil as you see, 1/8000 alpha particles are repelled backwards, that's because the nucleus is VERY small, only few alpha particles hit it completely and repel back. Thus Rutherford knew that the nucleus was super duper small!

Bohr � Modified Rutherford model- why electrons don’t crash into the nucleus if + attract – � Energy levels - electrons traveled in circular orbits and that only certain orbits were allowed. � Close to nucleus- lower energy level e- occupies � Far from nucleus- higher energy level e- occupies � Electron in one energy level or another not in between – like a ladder � Electrons gain or lose energy by changing energy level � Difference in energy between levels (to move e-) = quantum of energy � Nobel Prize in physics in 1922

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Electrons and Light � Light- moving waves and particles � Frequency- occurrences per time (1/s or s-1) (v measured in Hz or cycles/second) � Speed- of light ( c ) 2. 998 x 10 8 m/s � Wavelength- distance wave repeats – 2 consecutive peaks (λ measured in m) Stream of particles – energy Is determined by light’s frequency To remove e-, particle of light needs min. energy and a min frequency

Electromagnetic Spectrum – broad range of wavelengths Visible Spectrum: 400 nm (violet) to 700 nm (red) Red: low frequency/ long wavelength Violet: high frequency/ short wavelength � The product of frequency and wavelength always equals a constant, c, the speed of light � C = λv � Frequency and wavelength are inversely proportional

Light provides info… � � Ground State – lowest possible energy Excited State - electron absorbs energy & moves levels � A quantum of energy in the form of light is emitted when the electron drops back to a lower energy level – an abrupt step. � The light emitted by an electron from a higher to lower energy level has a frequency directly proportional to the energy change of the electron. Or…. �

E=hxf Where E = energy measured in Joules (J) h = Planck’s constant in J/Hz (J-s) = 6. 63 x 10 -34 J-s f = frequency of light measured in Hz If, c = λv Then f = c λ And, E = h c λ

Quantum Models � How likely to find an electron in various locations – probability � Uncertainty Principle: -exact position and momentum of e- is unknown; instead look at probability … think of blades on an airplane propeller or ceiling fan � � Quantum Number – specifies the properties of electrons � Electron Configuration- specific dispersal of electrons among subshells (or sublevels)

Follow 3 principles: 1. Aufbrau- electrons occupy lowest energy orbital first 1. Pauli Exclusion- orbital may describe at most two electrons 1. Hund’s Rule – orbital of a sublevel fill up by a single electron before pairing

Electron Capacity � 2 n 2 where n = quantum number designation and indicates energy level Quantum Number (n) Shell Capacity (2 n 2 ) 1 2 2 8 3 18 4 32

The periodic table is structured so that elements with the same type of valence electron configuration are arranged in columns.

� The left-most columns include the alkali metals and the alkaline earth metals. In these elements the valence s orbitals are being filled � On the right hand side, the right-most block of six elements are those in which the valence p orbitals are being filled � These two groups comprise the main-group elements � In the middle is a block of ten columns that contain transition metals. These are elements in which d orbitals are being filled � Below this group are two rows with 14 columns. These are commonly referred to the f-block metals. In these columns the f orbitals are being filled

Quantum Number Quantum # Principle Angular Momentum (Azimuthal) Magnetic Spin Symbol n l ml ms s, p, d, f … -1, 0, +1… +, - 1/2 Orientation in space Magnetic spin Possible Values 1, 2, 3, 4, … Characteristics Size and energy Shape level * *bigger number, higher energy level n=1, ground state

Electron Configuration Details � Address of Atom � Periodic Table is the Map � Ions-Isoelectronic Cations Previous noble gas as core Anions next noble gas as core Paramagnetic-one or more unpaired electrons Diamagnetic – all electrons paired � �