Atoms and their structure
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1 Atoms and their structure
2 History of atomic theory Not the history of atom, but the idea of the atom Original idea Ancient Greece (400 B.C..) Democritus and Leucippus Greek philosophers
3 Another Greek Aristotle - Famous philosopher All substances are made of 4 elements Fire - Hot Air - light Earth - cool, heavy Water - wet Blend these in different proportions to get all substances
4 Dalton s Atomic Theory (~1800) 1 All matter is made of tiny indivisible particles called atoms. 2 Atoms of the same element are identical, different elements have different atoms. 3 Atoms of different elements combine in whole number ratios to form compounds 4 Chemical reactions involve the rearrangement of atoms. No new atoms are created or destroyed.
5 Law of Conservation of Matter Matter can not be created or destroyed, just changed from one form to another. --Implies that if you start with g of matter and a chemical reaction occurs, you will have g of matter after the reaction is finished. The law extends to energy as well.
6 Law of Definite Proportions Each compound has a specific ratio of elements It is a ratio by mass Water is always 8 grams of oxygen for each gram of hydrogen
7 Law of Multiple Proportions If two elements form more than one compound, the ratio of the second element that combines with 1 gram of the first element in each is a simple whole number. Ex. CO and CO 2
8 Parts of Atoms J. J. Thomson - English physicist cathode ray tube --vacuum tube
9 Thomson s Experiment - Voltage source + Vacuum tube Cathode Anode
10 Thomson s Experiment - Voltage source +
11 Thomson s Experiment - Voltage source +
12 Thomson s Experiment - Voltage source +
13 Thomson s Experiment - Voltage source + npassing an electric current makes a beam appear to move from the negative to the positive end
14 Thomson s Experiment - Voltage source + npassing an electric current makes a beam appear to move from the negative to the positive end
15 Thomson s Experiment - Voltage source + npassing an electric current makes a beam appear to move from the negative to the positive end
16 Thomson s Experiment - Voltage source + npassing an electric current makes a beam appear to move from the negative to the positive end
17 Thomson s Experiment Voltage source + - nby adding an electric field
18 Thomson s Experiment Voltage source + - nby adding an electric field
19 Thomson s Experiment Voltage source + - nby adding an electric field
20 Thomson s Experiment Voltage source + - nby adding an electric field
21 Thomson s Experiment Voltage source + - nby adding an electric field
22 Thomson s Experiment Voltage source + - nby adding an electric field he found that the moving pieces were negative
23 Thomson s Model Discovered the electron plum pudding model Positive matter, with the electrons able to be removed
24 Robert Milliken American physicist (1909) Famous oil drop experiment Measured the charge on the electron Was able to calculate the mass of the electron (9.1 x kg) That s small! This is 1/1837 the mass of the H atom
25 Charge on the electron Charge on the electron
26 Conclusions from this? Electrons are smaller particles than atoms Since atoms are neutral, positive particles must also exist. Since electrons are so small, the other particles must carry most of the mass of the atom.
27 Other pieces Proton - positively charged pieces 1840 times heavier than the electron Neutron - no charge but the same mass as a proton(approximately). Where are the pieces?
28 Rutherford s experiment Ernest Rutherford English physicist. (1910) Believed the plum pudding model Used radioactivity (to measure size of atom) Alpha particles - positively charged pieces given off by uranium (very tiny) Shot at gold foil
29 Lead block Uranium Florescent Screen Gold Foil
30 What he expected
31 Because, he thought the mass was evenly distributed in the atom
32 What he got
33 Atom is mostly empty Small dense, positive center Alpha particles are deflected by it if they get close enough +
34 +
35 Density and the Atom Since most of the particles went through, it was mostly empty. Because a few pieces turned so much, the positive pieces were heavy. Small volume, big mass = big density This small dense positive area is the nucleus
36 The atom is mostly empty space Two regions Nucleus- protons and neutrons -all the mass positive charge Electron cloudregion where you might find an electron - most all the volume Modern View
37 Name Electron Subatomic particles Symbol Charge e - -1 Relative mass 0 (or 1/1840) Actual mass (g) 9.11 x Proton p x Neutron n x 10-24
38 Size of an atom Atoms are small. ( pm) Measured in picometers, meters Hydrogen atom, 32 pm radius Nucleus tiny compared to atom IF the atom was the size of a stadium, the nucleus would be the size of a marble. Radius of the nucleus near m. Density near 2 x 10 8 metric tons/cm 3.
39 Counting the Pieces Atomic Number = number of protons # of protons determines kind of atom the same as the number of electrons in the neutral atom Mass Number = the number of protons + neutrons
40 Symbols Contain the symbol of the element, the mass number and the atomic number Mass number Atomic number
41 Example 12 6 C Carbon atom Mass # = 12 At. # = 6 Protons = 6 Neutrons = F Fluorine atom Mass #= 19 At. # = 9 Protons =9 Neutrons = 10
42 How come the nucleus doesn t fly apart? Extreme amounts of repulsion when the positive charges (protons) are so close. Neutrons help a little. Requires a very strong force to hold the nucleus together. Called the STRONG NUCLEAR FORCE (most powerful force in nature)
43 Isotopes Atoms of the same element having different numbers of neutrons
44 Naming Isotopes Put the mass number after the name: Carbon-12 Carbon-14 Uranium-235 This format is common in news articles.
45 Measuring Atomic Mass Unit is the Atomic Mass Unit (amu) One amu = one-twelfth the mass of a carbon-12 atom. (standard) Each isotope has its own atomic mass we use the average mass from percent abundance. Periodic table shows the weighted average (decimal value)
46 Calculating averages You have five rocks, four with a mass of 50 g, and one with a mass of 60 g. What is the average mass of the rocks? Total mass = 4 x x 60 = 260 g Average mass = 4 x x 60 = 260 g 5 5 Average mass = 4 x x 60 = 260 g Average mass = 52g
47 The Mole Physically counting atoms is impossible. We must be able to relate measured mass to numbers of atoms. buying nails by the pound. using atoms by the gram
48 The Mole Defined as the number of particles in exactly g of C-12 atoms. (the standard) The specific number of particles is counted as 6.02 x10 23 This number of particles is called 1 mole or Avogadro s number.
49 Molar Mass This is defined as the mass of one mole of a substance. (units are g/mol) For elements: it is the mass value from the periodic table. Used frequently in calculations. Make it your friend
50 Using Molar Mass in Calculations Plan: Molar mass Mass X / Moles
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