H CHEM - WED, 9/7/16. Do Now Be ready for notes. Sigfig review problem. Agenda Atomic Theory. Homework. Error Analysis
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1 H CHEM - WED, 9/7/16 Do Now Be ready for notes. Sigfig review problem Agenda Atomic Theory Error Analysis Homework Possibly atomic theory paragraph
2
3 THE ATOM
4 DEFINITION TO START Atom smallest particle of an element that retains its identity They are tiny! Electron microscope allows us to observe individual atoms
5
6 DEMOCRITUS THE PHILOSOPHER ~490 BC First credited with proposing the existence of an atom Indivisible and indestructible Shortcomings? Did Aristotle agree?
7 ARISTOTLE ~380BC Did not agree! Earth, water, air and fire
8 THEN CAME JOHN DALTON Experimental methods -> scientific theory Dalton s atomic theory Matter is composed of indivisible atoms Atoms of same element are identical Combine in whole number ratios to form compounds Rxns occur when atoms separate, bond, or rearrange. Atoms of one element never become atoms of another thru chemical rxns
9 JJ THOMSON S EXPERIMENT Cathode Ray Experiment Gas filled glass tube fitted with electrodes Electricity -> cathode ray (travels from cathode to anode) Conclusion: electrons - negatively charged subatomic particles Further tests: Mass to charge ratio Different gases Conclusion: Electrons are part of atoms of all elements.
10 THOMSON S MODEL Thomson plum pudding (1897) Chocolate chips in cookie dough
11 MILIKAN Oil Drop Experiment Determined the charge of an electron
12 OTHER SUBATOMIC PARTICLES? We know things are neutral so where s the positive Actually discovery is disputed. Goldstein: anode ray (1886) Detected rays that contained positively charged particles Rutherford: 1920 Coined proton after his work with H nuclei Conclusion: protons positively charged subatomic particles 1840x mass of an electron
13 THOMSON S MODEL TO RUTHERFORD S MODEL Gold Foil Experiment (1911) Alpha particles thru gold foil Predictions: only slight deflection Results: Most: straight thru or slight deflection Some: large deflection or bounced back toward the source
14 RUTHERFORD MODEL/NUCLEAR MODEL Atom is mostly empty space explains why the alpha particles could pass straight thru All the positive charge and most of mass is located in a small region explains the large deflections Nucleus - protons and neutrons Electrons are around the nucleus and account for most of the volume Still not quite right!
15 IT S ALL ABOUT COLOR In terms of atomic models, so far: Dalton (1803) = Tiny, solid particle Thomson (1897) = Plum Pudding model Electrons stuck on the outside of a big positive charge Rutherford (1911) = Positively-charged nucleus with electrons moving around it Rutherford s model of the atom not quite right Could not explain chemical properties of elements Could not explain color changes when metal is heated
16 BOHR MODEL OF THE ATOM-1913 Niels Bohr s model of the atom Electron found only on specific, circular paths around nucleus Each orbit has fixed energy level Hypothesis: When electrons are excited (added energy), jump into higher energy levels. When they moved back into lower energy levels - gave off light. Electrons do not exist between levels (think of rungs on a ladder) Electrons absorb and emit only certain quanta (amounts) of energy Quantum of energy = fixed amount of energy required to move from one energy level to another energy level
17 BOHR S MODEL Nucleus Electron Orbit Energy Levels Chapter 5
18 BOHR S PLANETARY MODEL OF THE ATOM Electrons must have enough energy to keep moving around the nucleus Electrons orbit nucleus in defined energy levels, just like planets orbit the sun Each energy level assigned a principal quantum number n. Lowest energy level called ground state (n=1) Higher energy levels (n=2, 3, 4...) excited states Model worked OK for hydrogen but not so good for other elements n = 1 Nucleus n = 2
19 Increasing energy BOHR S MODEL Fifth Fourth Third Further away from the nucleus means more energy. There is no in between energy Energy Levels Chapter 5 Second First Nucleus
20 Electron starts on lowest energy level (ground state) Add energy to electron moves to excited state Lowest energy level = ground state Higher energy levels = excited states Energy levels are not evenly spaced Energy Level 3 Energy Level 2 Energy Level 1 Nucleus
21 Electron starts on lowest energy level (ground state) Add energy to electron moves to excited state Energy Lowest energy level = ground state Higher energy levels = excited states Energy levels are not evenly spaced Electron returns to lower state emits/gives off quantum of energy Energy Level 3 Energy Level 2 Energy Level 1 Nucleus
22 Bohr used this theory to explain the lines in the Chapter 5 atomic emission spectra for hydrogen
23 Each of these lines corresponds to different energy changes Chapter nm 656 nm 410 nm 486 nm
24 So is the Bohr model correct? For H, it does a wonderful job explaining things For everything else.not so much
25 CHADWICK THE LAST SUBATOMIC PARTICLE Chadwick (1932) Chased evidence for a particle that Rutherford predicted to exist Neutrons neutral subatomic particles ~same mass as proton
26 History lesson over.for now
27 WHAT CAN THE PERIODIC TABLE TELL US? Atomic number the number of protons in an element Elements are defined by their atomic number Mass number total number of neutrons and protons
28 ISOTOPES! Same number of protons Different number of neutrons Used to calculate the atomic weight of an element Mass of an atom is tiny so knowing the actual mass is a bit impractical
29 ATOMIC MASS UNIT Isotope Carbon 12 was assigned the mass of 12 atomic mass units. So then, 1/12 of the mass of a carbon 12 atom is 1 amu. So then helium 4 has 1/3 the mass of carbon 12. Carbon has 6 protons and 6 neutrons this accounts for the bulk of the mass so one proton or 1 neutron has a mass of ~1 amu.
30 ATOMIC WEIGHT Atomic weight weighted average of the atomic masses of the isotopes of an element in a naturally occurring sample Example problem.grades!
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