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History of Atomic Structure Chemistry. Democritus (460-370 BC) Goal of Greek philosophers was to explain the natural world  Believed that all materials.

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Presentation on theme: "History of Atomic Structure Chemistry. Democritus (460-370 BC) Goal of Greek philosophers was to explain the natural world  Believed that all materials."— Presentation transcript:

1 History of Atomic Structure Chemistry

2 Democritus ( BC) Goal of Greek philosophers was to explain the natural world  Believed that all materials could be broken down smaller and smaller parts until you reach a point  This point was what Democritus called “atomos” which in Greek meant “indivisble” or “uncuttable” This particulate view of nature was not too popular among many Greek thinkers due to its rejection by Aristotle  Matter is made up of 4 elements: fire, air, water, and earth  Matter was continuous all of one piece Unfortunately for mankind, the ideas of Democritus would not come back into the public domain for nearly 2,000 years!

3 John Dalton: Atomic Theory 1.) All matter is made up of atoms. Atoms are indivisible 2.) All atoms of an element are identical in every respect (have the same masses and same properties) 3.) Atoms of different elements are different (have different masses) 4.) Compounds are formed by a combination of 2 or more different kinds of atoms (same ratio of atoms) 5.) A chemical reaction is a rearrangement of atoms

4 Dalton’s Atomic Theory Dalton’s theory seemed to be quite similar to Democritus don’t you think? One part of Dalton’s atomic theory has been rejected  Make a prediction of which statement you think is incorrect and a hypothesis as to why you think that  We will discuss this more later

5 J.J. Thomson: 3 experiments, 1 big idea Do atoms have parts?  Thomson suggested that they do  Advanced the idea that cathode rays are really streams of very small pieces of atoms  3 experiments led him to this

6 Thomson’s 1 st experiment Electrometer: device for measuring electrical charge Had already been found that cathode rays deposited an electrical charge What Thomson was interested in was whether or not he could separate the charge from rays by bending them with a magnet Found that when rays entered slit and into electrometer, it measured a large amount of negative charge Electrometer did not register much charge if rays were bent so they would not enter slit Thomson concluded that the negative charge and the cathode rays must somehow be stuck together (you cannot separate the charge from the rays

7 Thomson’s 2 nd experiment Prior to Thomson’s 2 nd experiment, all attempts had failed when trying to bend cathode rays with an electric field A charged particle will normally curve as it moves through an electric field, but not if it’s surrounded by a conductor Thomson realized that others had failed probably because traces of gas remaining in the tube were being turned into electrical conductors To test this idea, he took great pains to extract nearly all the gas from a tube, and found that now the cathode rays did bend in an electric field after all Notice which way the cathode ray is being bent after passing through electric field. Why is this?

8 After 2 experiments…. “I can see no escape from the conclusion that cathode rays are charges of negative electricity carried by particles of matter. But, what are these particles? Are they atoms, or molecules, or matter in a still finer state of subdivision?”—J.J. Thomson

9 Thomson’s 3 rd Experiment Sought to determine the basic properties of the particles  Thomson calculated the ratio of the mass of a particle to its electric charge (m/e) What he concluded was astounding:  The mass to charge ratio for cathode rays turned out to be far smaller than that of a charged hydrogen atom—more than 1,000 times smaller 2 possibilities: 1.) the cathode rays carried an enormous charge 2.) they were amazingly light relative to their charge 2 nd possibility was eventually proven

10 Summary of Thomson’s Findings Discovered the electron (negatively charged subatomic particle) in 1897 Developed the “plum pudding” model in 1904  Probably easier to call it the “watermelon” model Seeds: negative particles Fruit: positively charged low density material

11 Ernest Rutherford Student of J.J. Thomson (1894) Early studies based on radioactivity Discovered that some materials emit radiation  Alpha particle: Eventually was confirmed to be helium nuclei, which meant an alpha particle was simply 2 protons and 2 neutrons  Like Thomson, Rutherford was interested in the deflection patterns of alpha particles when exposed to electric & magnetic fields  Found the deflection patterns by measuring alpha particles position on photographic film  By accident, he noticed that if the alpha particles passed through a thin sheet of mica, the images on the film were blurred

12 Rutherford and alpha particles The images were sharp if the mica was not present  Something about the mica sheet was causing the alpha particles to scatter at seemingly random small angles, resulting in blurred images

13 Rutherford & Geiger Wanted to study the effects of alpha particles with matter  Had to develop a way to count individual alpha particles when they hit the screen  Found that a screen coated with zinc sulfide emitted a flash of light each time it was hit by an alpha particle  Rutherford & Geiger had to sit in a dark room for hours and individually count the flashes of light Rutherford asked Geiger to measure the angles of deflection when the alpha particles were passed through a thin ( cm) sheet of gold foil  When the gold sheet was bombarded with alpha particle, Geiger found that the scattering small  Results were consistent with Rutherford’s expectations. Because he knew that alpha particles had a considerable mass and moved quite rapidly, he anticipated that virtually all of the alpha particles would go through the metal foil without much disruption

14 Rutherford, Geiger, & Marsden Marsden was a graduate student working in Rutherford’s lab and Geiger had suggested that a research project should be given to Marsden  Rutherford: “why not let him see whether any alpha particles can be scattered through a large angle?”  Marsden had found that a small fraction (perhaps 1 in 20,000) of the alpha particles were scattered through angles larger than 90 degrees Rutherford was in awe  “It was quite the most incredible event that has ever happened to me in my life. It was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you!”

15 What could the results from this gold foil experiment mean? The + charge in an atom of gold was known, but if this charge were collected on a sphere the size of an atom (like Thomson suggested), the repulsion would be far too weak  To explain the force experienced by the alpha particle, the charge and much of the mass would have to be collected in a much smaller sphere  Published results in 1911 and proposed a model for the atom that is still accepted today  All of the positive charge and essentially of the mass of the atom is concentrated in an incredibly small fraction of the total volume of the atom, which he called the nucleus (Latin for “little nut”)

16 Further findings from Rutherford’s experiment Most of the alpha particles were able to pass right through A small fraction came close to the nucleus of a gold atom as they passed through and were slightly deflected from the positive-positive repulsion of the alpha particle and nucleus But occasionally, an alpha particle would run directly into the nucleus and would result in a great repulsion that deflected the alpha particle through an angle of 90 degrees or more  By carefully measuring the fraction of the alpha particles deflected through large angles, Rutherford was able to estimate the size of the nucleus (JUST LIKE WE DID!)  Found that the radius of the nucleus is at least 10,000 times smaller than the radius of the atom  The vast majority of the atom is therefore empty space!!

17 Someone throw out the plum pudding!! Rutherford revised Thomson’s plum pudding model, showing how electrons could orbit a positively charged nucleus, like planets orbiting a sun Because the majority of the “plum-pudding” atom would be electrically neutral (no charge), the alpha particles would have no problem shooting through Rutherford’s atom

18 Along comes Niels Bohr Also student of Thomson and worked with Rutherford Bohr, and many other, knew that Rutherford’s model made no sense based on one specific reason….  We knew that any charged body (electron) that was in a state of motion other than at rest or in uniform motion in a straight line, will emit energy  Thus the electrons in this “solar system” model would be constantly emitting energy  IF that were the case, the electrons would eventually run out of energy and spiral down into the nucleus and the entire atom would collapse!

19 Bohr Had trouble making sense of line spectra data with Rutherford’s atom

20 Bohr’s Model Similar nucleus to Rutherford’s with both protons and neutrons inside Negative particles (electrons) are in specific orbits around the nucleus Main problem with Bohr model: Only accounted for hydrogen atom Couldn’t explain multi-electron atoms (anything other than hydrogen)

21 Subatomic Particles Based on evidence from Rutherford and Thomson

22 Location of Subatomic Particles

23 Atomic Numbers

24 What is atomic mass? Also referred to as:  Atomic weight  Average atomic mass  Relative atomic mass Important characteristic for elements:  Each element has its own atomic mass Atomic mass Atomic mass is an average Average of the masses of a number of different atoms Special kind of average called a weighted average Different than the usual average you’re probably familiar with in math

25 Understanding Weighted Averages Even though these are different models and have different features, they are both lemonas due to their distinct lemon-like shape Using this analogy, the models of the lemona are similar to the isotopes of an element 29 protons makes both atoms copper—even though they differ in their numbers of neutrons Just like the “lemon-like” shape of a car makes it a lemona— even though they differ in their features

26 Average vs. Weighted Average What would happen if we added in extra information?

27 Weighted Average Because there are so many more GXs than GXLs, the weighted average is much closer to the actual weight of a lemona GX

28 Using Weighted Average with Different Atoms Atomic mass: a weighted average of the masses for all the isotopes of a certain element Mass = 63 amu Mass = 65 amu If we pulled out a random sample of 100 copper atoms, we would find that 69% of them would be Cu-63 and 31% of them would be Cu-65 69% : Cu-63 31% : Cu-65

29 Atomic Mass of Copper 69% : Cu-63 31% : Cu-65 Why are the 2 atomic masses different then?

30 Difference between mass number and atomic mass Mass number: protons + neutrons 1 proton or 1 neutron = 1 amu Therefore, if you have 6 protons and 6 neutrons, your atom is going to weigh 12 amu If you have 6 protons and 7 neutrons, your atom is going to weight 13 amu

31 Practice Gallium has 2 stable isotopes, and the masses of Gallium-69 (60.11% abundant) and Gallium-71 (39.89% abundant) are amu and amu, respectively. Calculate the average atomic mass of Gallium

32 Practice Rubidium has 2 isotopes: Rubidium-85 (atomic mass of amu) and Rubidium-87 (atomic mass of amu). The atomic mass of Rubidium reported on the periodic table is amu. Based on this information, which of the isotopes of Rubidium is more abundant? How do you know? This is more of a thought problem. No real calculations are really necessary

33 Practice Magnesium has 3 stable isotopes. Calculate its average atomic mass, using the information in the chart below. IsotopeMassAbundance Mg amu78.99 % Mg amu10.00 % Mg amu11.01 %


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