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Chem 1045

by: Elizabeth Wallington

Chem 1045 CHM 1045

Elizabeth Wallington
GPA 3.75

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About this Document

These go over atomic numbers and the scientists and their experiments she is requiring us to know
General Chemistry 1
Dr. Vu
Class Notes
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This 4 page Class Notes was uploaded by Elizabeth Wallington on Tuesday September 20, 2016. The Class Notes belongs to CHM 1045 at Florida State University taught by Dr. Vu in Fall 2016. Since its upload, it has received 83 views. For similar materials see General Chemistry 1 in Chemistry at Florida State University.


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Date Created: 09/20/16
Chemistry 9/7/16 2.3 Atomic Number, atomic mass, & isotopes The entire mass of the atom is based on protons and neutrons 1 1H  The 1 superscript is the atomic mass  The 1 subscript is the atomic number (the protons) 12 6 C  6 protons  6 neutrons (12-6; mass-protons; mass-#) 91F  9 protons  10 neutrons 148Si  14 protons  14 neutrons The mass indicates the average of all isotopes of this element; so the majority of the element found is the isotope with that mass. The protons will always stay the same of an element; that is what defines the element, which is why it is the atomic number. 2.2 Evolution of Atomic Theory 1. Dalton (1803) atoms are tiny individual particles (section 2.1) 2. Mendeleev (1897) organized periodic table by atomic mass 3. *JJ Thomson (1897) determined electrons are negatively charged; mass to charge ratio of an electron  Cathode Ray Experiment o Beam of electrons  The beam of electrons was “shot” at a screen that was coated in a substance that, when struck by the electrons, became excited and lit up; this was used as the detector.  When the beam of electrons was passed through the hole in the metal divider, it was diverted when the tube was positively and negatively charged; the beam of electrons would divert towards the positive charge.  This allowed Thomson to conclude that electrons are negatively charged  From this information he got the charge to mass ratio of an electron 8 o Charge:Mass= -1.76 X 10 C/Kg 4. *R.A. Millikan (1909) determined charge on an electron 2  oil drop experiment  He used an atomizer to spray oil drop particles in a drum with air and then calculate the charge on an electron by calculating the mass of one, by timing how long it takes to fall calculated with gravity, then using the charge to mass ratio  When the metal was charged using a power source (as seen in picture) the air between the plates was ionized using x-rays. When the oil drops passed through the hole in the positively charged plate to the ionized air below, the previously neutral oil drops now carried a negative charge  he was looking for the electrical force to = the gravitational force, therefore suspending the oil drops in midair  q (charge on electron) X E () = mass (of oil drop) X gravity o q= -1.60 X 10 -1C 5. *Rutherford (1911) nucleus contains protons (small) 3  conclusion: the nucleus is very small, positively charged, center with electrons surrounding  radium was emitting an alpha particle stream through a thin gold foil at a coated screen that lights up when struck by electrons. Most flashes occurred in the path of the particle stream but occasionally a flash would occur somewhere else on the screen and even some particles bounced straight back from the thin gold foil.  He realized that for them to pass through mostly unaffected that the particle must be mostly empty space but occasionally, when they were being diverted or bouncing back, something larger was striking.  mystery left was why the He:H mass ratio was 4:1? Because He has 2 protons and H has 1, so what else is there? 6. Bohr (1922) electrons in shells 7. Schrodinger (1930) electron orbitals that have different shapes and can be mathematically expressed 8. Chadwick (1932) discovered the neutron HW (9/7): Calculate the mass of an electron using ratio and q - -28 Mass of e = 9.10 X 10 g 4


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