Bangs, Flashes, and Explosions

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2 Mechanical Universe S Lesson 49: The Atom This program explores the history of the atom, from the ancient Greeks to the early 20th century, when discoveries by J.J. Thomson and Ernest Rutherford created a new crisis for the world of physics. Text Assignment: Chapter 49 Instructional Objectives Be able to summarize the kinetic theory and discuss the size of atoms. Be able to compare Thomson's model of an atom with Rutherford's planetary model of an atom. Be able to discuss why Rutherford's model of an atom conflicted with Maxwell's theory of charged particles. Be able to discuss the significance of Brownian motion in providing evidence for the existence of atoms.

3 Bangs, Flashes, and Explosions Section 4: Atomic Theory and Periodic Trends Rutherford Experiment Simulation...43 Colored Flames...45 Throwing Flame(test) Flames...47 Bangs, Flashes, and Explosions Illustrated Guide to Chemistry Demonstrations 2005 Chris Schrempp and ExploScience.com. All Rights Reserved Bengal Flames...48 Fiery Tornado...50 The Incredible Glowing Pickle (and grapefruit, and onion, and )...52 Reactions of Lithium, Sodium, and Potassium in Water...54

4 John Dalton, British Physicist ( ) Main contribution: All matter-whether element, compound, or mixture-is composed of small particles called atoms Postulates of Dalton s Atomic Theory 1.Matter consists of indivisible atoms. 2.All of the atoms of a given chemical element are identical in mass and in all other properties. 3.Different chemical elements have different kinds of atoms; in particular their atoms have different masses. 4.Atoms are indestructible and retain their identities in chemical reactions.

5 5.A compound forms from its elements through the combination of atoms of unlike elements in small wholenumber ratios. Dalton s Law of Definite Proportions (one compound always having the same mass relationship)- Suppose a compound is made of element H and element O, giving the formula H2O. Since the weight of H is constant and the weight of O is constant, the H:O weight ratio will always be the same. Hence, the Law of Definite Proportions. If H weighs 1 gram and O weighs 16g the mass relationship in water (H2O)will always be 2g H/ 16 g O = 2:16 =1:8. Dalton s Law of Multiple Proportions (for 2 or more compounds containing the same elements) - Two elements which form more than one compound between them will have compounds whose mass ratio between the two elements are related to each other as a ratio of small whole numbers. H2O2 or H2O but not H2.7O2.2 Example: CO g O = 2.66 g O = g of C 1.0g C CO g O = 1.33 g O = g of C 1.0g C The Structure of the Atom

6 Joseph John Thomson ( ) British Physicist Main contribution: the discovery of the electron! How the electron was discovered Used a Cathode Ray tube 2 electrodes from a high voltage source are sealed in a glass tube in which the air has been evacuated When electricity is applied, a lavender beam of light emits from the electrodes Same lavender beam of light produced regardless of electrode material Since same lavender light is emitted from all types of material, all materials have same type of particles ELECTRONS! Plum Pudding Model

7 The Charge on an Electron Robert Millikan ( ) US Physicist Main contribution: determined the charge on a single electron Famous Oil Drop Experiment Charge on a single electron x Coulombs Along with JJ Thomson s work, Millikan s work helped determine mass of the electron: x kg

8 Ernst Rutherford ( ) British Physicist Main contribution: developed a new model of the atom where there is a small but massive, positively charged nucleus in the atom. The Nuclear Model of the atom Rutherford s model based on his famous gold foil experiment Used alpha particles as bullets (Helium atoms without electrons, have + charge) Fired these particles into very thin gold foil Most went through without deflection Some particles were deflected at large angles Rutherford s Gold Foil Experiment Explanation of Experiment Most of the atom is empty space!

9 There must be very small, dense areas where the alpha particles hit, are deflected, and bounce back These areas must be positively charged! Other Discoveries These scattering experiments led to the determination of the + charge on the nucleus called the Atomic Number, Z Rutherford discovered protons by striking lighter elements, like nitrogen, with alpha particles Led to new definition of element a substance whose atoms contain all the same atomic number (proton #)

10 The Discovery of the Neutron James Chadwick ( ) British Physicist Main Contribution: discovered atom was made of neutral particles that were as massive as protons - these he called neutrons! Isotopes Some elements have 2 or more different versions of the same atom! These different versions all have the same number of protons but differing numbers of neutrons; these are called isotopes!

11 NIELS BOHR Depicted the atom as a small, positively charged nucleus surrounded by waves of electrons in orbit similar in structure to the solar system, but with electrostatic forces providing attraction, rather than gravity, and with waves spread over entire orbits instead of localized planets. The theory that electrons travel in discrete orbits around the atom's nucleus, with the chemical properties of the element being largely determined by the number of electrons in each of the outer orbits. The idea that an electron could drop from a higher-energy orbit to a lower one, emitting a photon (light quantum) of discrete energy (this became the basis for quantum theory).

12 DRAW THIS

13 # of Protons determines what element it will be - a nuclide is a particular nucleus characterized by a definite atomic number and mass number Mass Number Atomic Number, Z One nuclide of sulfur One nuclide of phosphorus Atomic Number = Number of Protons Mass Number = Number of Protons + Neutrons Example: If you look up potassium (K) in the periodic table, it has an atomic number of 19, meaning that all potassium atoms and ions contain 19 protons. Example: A particle with 6 protons and an atomic mass number of 14 has 8 neutrons.

14 REVIEW Determine the number of electrons, protons and neutrons for: B Ni 22 Ni Xe 64 Na 14 F REVIEW 29 Cu 14 N 6 C ISOTOPES Isotopes are atoms of the same element that have a different number of neutrons. Therefore, isotopes have the following characteristics:

15 ISOTOPES 1.Isotopes have the same atomic number (same number of protons), but a different atomic mass number (a different number of neutrons). 2.Isotopes behave the same chemically, because they are the same element. The only difference is that one is heavier than the other, because of the additional neutrons For example, carbon-12 and carbon-14 are both isotopes of carbon. Carbon-12 has 6 neutrons; carbon-14 has 8 neutrons. Isotope 32 S 33 S 34 S 36 S Relative atomic mass Abundance (%) The average weight of Sulfur is a.m.u. and is the accepted mass for Sulfur, (but not mass of all isotopes.)

16 LAB PARTNER EXERCISE ISOTOPES 1) Rubidium has two common isotopes, 85Rb and 87Rb. If the abundance of 85Rb is 72.2% and the abundance of 87Rb is 27.8%, what is the average atomic mass of rubidium? amu 2) Uranium has three common isotopes. If the abundance of 234U is 0.01%, the abundance of 235U is 0.71%, and the abundance of 238U is 99.28%, what is the average atomic mass of uranium? amu 3) Titanium has five common isotopes: 46Ti (8.0%), 47Ti (7.8%), 48Ti (73.4%), 49Ti (5.5%), 50Ti (5.3%). What is the average atomic mass of titanium? amu 4) Explain why atoms have different isotopes. In other words, how is it that helium can exist in three different forms? Neutrons exist to stabilize the nucleus without them, the nucleus would consist of nothing but positivelycharged protons in close proximity to one another. Because there are different ways of stabilizing the protons, there are different isotopes

17 ATOM Subatomic Particles Mass Charge Proton 1 a.m.u +1 Neutron 1 a.m.u 0 Electron 0 a.m.u -1 a.m.u = atomic mass unit 1/12 of a carbon atom How Many: (Mass # in a.m.u) Protons 1- Hydrogen: (1.008) Helium: (4.002) 2 3 Lithium: (6.941) 3 4 Beryllium: (9.012) 4 5 Boron: (10.811) 5 6 Carbon(12.017) 6 Neutrons

18 I MISS CONVERSIONS

19 2H + O -> H2O Two atoms of hydrogen and one atoms of oxygen react together to make two molecules of water. If the atom is so small, so how are we going to pick up and mix 2 atoms of each reactant? How do we slap together 2 trillion atoms of hydrogen (H) and 1 trillion atoms of oxygen (O)? Is the relative weight of these two huge masses any different than the simple atoms? No nting.' 1 Million atoms H = 1 atom H 1 Million atoms O = 1 atom O

20 THE MOLE The mole is the weight scale at the chemical fruit stand: it allows different substances to be measured in a comparable way Moles make it easier to interpret chemical equations. Thus the equation: 2H + O = H2O is understood as "two moles of hydrogen plus one mole of oxygen yields one moles of water."

21 RUN BACK TO BIOLOGY WHILE YOU STILL HAVE THE CHANCE! And at this point God looked down from heaven and said to all confused chemists, LET THERE BE ZOOLOGY

22 SO HOW MUCH IS THIS MOLE? IT EQUATES TO SAND IN THE SAHARA. 602,200,000,000,000,000,000 grains of sand

23 SO LETS BREAK IT DOWN A LITTLE A MOLE LETS US KEEP OUR NUMBERS SIMPLE WITHOUT LOSING TRACK OF ALL THOSE ATOMS. The term dozen represents the number - 12 eggs A deck of cards represents the number - 52 cards The term ton represents the number 2000 pounds The term mole represents the number: x 10 particles/mol A MOLE OF ANYTHING EQUALS A MOLE OF ANTHING ELSE (in # not weight) Pass around vials of one mole of each substance.

24 A dozen eggs will make a nice sized omelet, but a mole of eggs will fill all of the oceans on earth more than 30 million times over. It would take 10 billion chickens laying 10 eggs per day more than 10 billion years to lay that mole of eggs. If I asked you how much does a dozen weigh, You would say that depends on what the dozen is. The same goes for Moles: it simply depends on what it is. All moles weigh differently (depends on the type of atom) So why would need to use such a big number? Using absolute numbers of atoms would take a

25 MOLAR MASS Mass of One Mole (g/mol) ON WHAT IS BASED? A mole is the amount of substance that is the same number of atoms as atoms in 12 grams of Carbon, the number of carbon defined as Avogadro's number. Moles are useful in chemical calculations, because they enable the calculation of yields (and other values)when confronted with different masses (weights) of elements. one oxygen atom weighs almost 16 times as much as a hydrogen atom. For example, helium has an atomic weight of 4.00 AMU. Therefore, 4.00 grams of helium will contain one mole of helium atoms.

26 You can also work with fractions (or multiples) of moles: Not on lecture notes. Mole/Weight Relationship Examples Using Helium Moles Helium # Helium Grams Helium Atoms 1/ x g 1/ x g x g x g x g Molecular Weight (not atomic wt.) When atoms form molecules, the atoms bond together, and the molecule's weight is the combined weight of all of its parts. For example, every water molecule (H2O) has two atoms of hydrogen and one atom of oxygen. One mole of water molecules will contain two moles of hydrogen and one mole of oxygen.

27 A bottle filled with exactly g water will contain 6.02 x 1023 water molecules (not atoms). The concept of fractions and multiples described above also applies to molecules: 9.01 g of water would contain 1/2 mole, or 3.01 x 1023 molecules. You can calculate the molecular weight of any compound simply by summing the weights of atoms that make up that compound. Mass to Mole Conversions: Grams to moles: ( grams) x (1 mole / grams ) Moles to grams: (Number of moles) x ( grams / mole) Moles to atoms: (maintain scientific notation and sig figs) (Number of moles) x (6.022 x 1023 atoms / 1 mole)

28 1. How many moles in 125 g of Hydrogen? _124_mol_ 2. How many moles in g of Copper? _1.000 mol 3. How many moles in 34.7 g of Lithium? 5.00_ mol 4. How many moles in g of Nitrogen? _6.000 mol 5. How many moles in 95.7 g of Carbon? _7.97_ mol 6. How many moles in g of Hydrogen? _ mol 7. How many moles in 98 g of Chlorine? 2.8 mol 8. How many moles in 298 g of Potassium? 7.62_ mol 9. How many moles in 410 g of Sulfur? 13 mol 10. How many grams in 0.5 moles of Hydrogen? _.5g 11. How many grams in 1.5 moles of Magnesium? 36g 12. How many grams in 5.70 moles of Silver? 615g_ 13. How many grams in moles of Uranium? 1004g 14. How many grams in.15 moles of Bromine? 12g

29 15. How many grams in 8.8 moles of Oxygen? 140g 16. How many grams in 9.2 moles of Argon? 370g 17. How many grams in moles of Aluminum?272.5g 18. How many grams in 101 moles of Helium? 19. How many atoms in 0.12 mol of cadmium 7.2 x How many atoms in 2.2mol of xenon 1.3 x How many atoms in 8.0 mol of carbon 4.8 x How many atoms in 5.5 mol of cadmium 3.3 x How many atoms in.33 mol of magnesium 1.9 x How many atoms in 1.59 mol of Americium 9.57 x How many atoms in mol of cadmium 7.23 x g

30 A. How many grams of carbon are needed to give 3 moles of carbon? B. What is the molar weight of H2O2.? C. How many moles of sulfur dioxide, SO2 (g), are in 2000 grams of the gas. D. Determine the number of grams in 4 moles of H2O. E. Determine the number of molecules of H2O in 3 moles H2O. F. Determine the number of moles of CO2 in 454 grams.

31 G. Determine the mass in grams of 3.60 mol of H2SO4.

32 A. How many grams of carbon are needed to give 3 moles of carbon? 1. Calculate the molar mass for carbon. Look up the atomic weight/mass in the periodic table. The molar mass for carbon is 1 mole C = 12.0 grams 2. Determine the mass needed to provide 3 moles of hydrogen. 1 mole C = 12.0 grams 2 mole C = 24.0 grams 3 mole C = 36.0 grams The practical way is to multiply the molar mass by the number of moles.

33 B. What is the molar weight of H2O2? The formula weight for H2O2 = 2(weight from hydrogen) + 2(weight from oxygen) = formula weight for H2O2 2 H atoms x amu + 2 O atom x amu = = amu The molar mass for H2O2 = grams IF I TELL YOU, I NEED A MOLE OF HYDRODGEN PEROIXIDE YOU WOULD BRING ME HYDROGEN PEROXIDE ATOMS IN THE FORM OF 34.02g.

34 C. How many moles of sulfur dioxide, SO2 (g), are in 2000 grams of the gas? 1. Look up the atomic weights in the periodic table for S and O. The atomic weight for sulfur is amu The atomic weight for oxygen is amu 2. Calculate the formula weight for SO2. Add up the masses from all the atoms in the formula The formula weight for sulfur dioxide is: amu S + 2 x (16.00 amu O) = amu SO2 3. Determine the molar mass. Molar mass is a mass in grams that is numerically the same as the formula weight. 1 mole SO2 = grams SO2 = 64. grams SO2 (to the nearest gram ) 4. Convert 2000 grams of SO2 to moles. The "conversion factor" is the molar mass. (2000 grams SO2 )(1 mole SO2 /64. grams SO2 ) = moles SO2 = 31 moles SO2 (rounded to 2 sf)

35 D. Determine the number of grams in 4 moles of H2O? Formula mass H2O = (2 x 1.0) + (1 x 16.0) = mole H2O = formula mass H2O = 18.0 grams H2O 4 moles H2O x.0 grams / 1 mole = 72.0 grams H2O

36 E) Determine the number of moles in 88 grams of CO2 Formula Mass CO2 = (1 x 12) + (2 x 16) = 44 1 mole of CO2 = formula mass CO2 = 44 grams CO2 88 grams CO2 x 1 mole CO2 / 44 grams CO2 = 2 moles CO2

37 E. Determine the number of molecules of H2O in 3 moles H2O 1 mole H2O = X 1023 molecules H2O 3 moles H2O x X 1023 molecules H2O / 1 mole H2O = X 1023 molecules H2O = X 1024 molecules H2O

38 F. Determine the number of moles of CO2 in 454 grams? The atomic mass of C is and the atomic mass of O is The formula mass of CO2 is: (16.00) = AMU Thus, one mole of CO2 weights grams. This relation provides a conversion factor to go from grams to moles. Using the factor 1 mol/44.01 g: moles CO2 = 454 g x 1 mol/44.01 g = 10.3 moles CO2

39 G. Determine the mass in grams of 3.60 mol of H2SO4. The atomic mass is for H; for S; for O. Formula mass H2SO4: 2(1.008) (16.00)=98.08AMU Thus, one mole of H2SO4 weights grams. Using the factor g / 1 mol: 3.60 mol x g / 1 mol = 353 g H2SO4

40 Problems Calculate the molar mass (g/mol) 1) HNO3 2) HC2H3O2 3) C12H22O11 4)

41 1) 63.0 g/mol HNO3 2) 60.0 g/mol HC2H3O2 3) 342 g/mol

42 Atoms Monoatomic (He, Li, Ag, Hg) 7 Diatomic Molecules (O, N, H, F, Cl, Br, I) BrIFN HONCl Multiatomic Molecules(S8, P4) Compounds CO, CH4, H2SO4 Oxygen 16.0 amu Sulfur 32.1 amu Sulfur atom is twice as heavy Helium (4 amu) is one fourth as heavy as Oxygen (16 amu) Phosphorous (31.0) Nitrogen (14.0) P=2.21 Relative Weights 1 Atomic Mass Unit : 1/12 of an atom of Carbon (12 atoms C = 12 amu) Water Molecule 18 amu (1.5 times heavier than C of 12 amu) Sulfuric Acid Molecule 98.1 amu HSO4 (8.175 times heavier than C of 12 amu)

43 2H2 O2 2H2 O *2.02g 32.00g + = 2*18.02g