The Atom Chapters 4-5

Atomic Theories Democritus ~ 400 BC Dalton ~ 1800’s believed that atoms were indivisible and indestructible Dalton ~ 1800’s Developed through experiments First Atomic Model

Dalton’s Atomic Model All elements are composed of tiny indivisible particles called atoms Atoms of the same element are identical. Different elements have different atoms.

Discovery of Electron 1897 – JJ Thomson, using cathode ray tube, discovered negatively charged particles called electrons 1909 – Robert Millikan - Oil Drop Experiment Determined charge on an electron.

“Plum Pudding” Model Uniform positive sphere with negatively charged electrons embedded within.

Radiation Late 1800’s – discovery of radiation Three Types Alpha Beta Gamma

Rutherford Gold Foil Experiment - 1909 Shot alpha particles at gold foil Most went through foil with little or no deflection. Some were deflected at large angle and some straight back.

Rutherford Gold Foil Experiment - 1909

Rutherford Model Conclusions from Gold Foil Experiment Atom is Mostly Empty Space Dense positive nucleus Electrons moving randomly around nucleus

The Atom An Atom is made of 3 Subatomic Particles Electrons Protons Neutrons

Subatomic Particles Electron Discovered in 1897 by JJ Thomson Negative charge (-1) Mass = 9.109389*10-28g Approx mass ~ 0 Found outside of nucleus

Subatomic Particles Proton Discovered in 1919 by Rutherford Positive charge (+1) Mass = 1.672623*10-24g Approx mass ~ 1 atomic mass unit (u) Found inside nucleus

Subatomic Particles Neutron Discovered in 1932 by James Chadwick No charge (0) Mass = 1.6749286*10-24g Approx mass ~ 1 atomic mass unit (u) Just slightly larger than a proton Found inside nucleus

Atomic Structure Atoms have no net charge # of electrons = # of protons # of electrons around nucleus = # of protons in nucleus

Atomic Structure Atomic Number Mass Number Number of protons in an element All atoms of the same element have the same number of protons Mass Number Number of protons and neutrons in an atom

Atomic Structure # of Neutrons = Mass Number – Atomic Number Atoms of the same elements can have different numbers of neutrons Isotope – atoms of the same element with different number of neutrons

Ion Atom or group of atoms that have gained or lost one or more electrons Have a charge Example: H+, Ca2+, Cl-, OH-

Chemical Symbols Cl-35 Chlorine-35 Mass Number Charge Atomic Number Number of Atoms

Atomic Theories Rutherford’s model could not explain the chemical properties of elements Niels Bohr believed Rutherford’s model needed to be improved Bohr proposed that electrons are found only in circular paths around the nucleus

Bohr Model Dense positive nucleus Electrons in specified circular paths, called energy levels These energy levels gave results in agreement with experiments for the hydrogen atom.

Bohr Model

Bohr Model Each energy level can only hold up to a certain number of electrons Level 1  2 electrons Level 2  8 electrons Level 3  18 electrons Level 4  32 electrons

Electron Configuration The way in which electrons are arranged in the atom Example: Na 2-8-1 Valence Electrons Electrons in the outermost energy level

Wave Mechanical Model More detailed view of the Bohr Model Schrödinger Wave Equation and Heisenberg Uncertainty provides region of high probability where electron COULD be. Orbital Modern Model AKA Quantum Mechanical Model, Electron Cloud Model

Wave Mechanical Model Orbital Regions of space where there is a high probability of finding an electron

M&M’s Demo What colors are found in a regular M&M’s bag? Green Yellow Orange Blue Red Brown

M&M’s Demo Do you get an equal amount of each color in each bag? If we opened up all the regular M&M bags in the world would we get an equal number of each color? Are you supposed to?

M&M’s Demo Color 1 bag World Blue % 24% Green 16% Yellow 14% Orange 20% Red 13% Brown

M&M’s Demo M&M’s come in certain abundances (percentages) So do isotopes of each element Relative Abundance Percent of each naturally occurring isotope found in nature

Average Atomic Mass Atomic Mass Example Weighted average based on the relative abundance and mass number for all naturally occurring isotopes Example C-12 98.9% 12.011u C-13 1.1%

Atomic Mass C-12 98.9% C-13 1.1% Carbon = 0.989*12 + 0.011*13 = 12.011u