Unit 6

1. Unit 6

2. W.K. Roentgen studied fluorescence – a phenomenon where certain materials emit light when struck by radiant energy, like UV light. • Found that materials fluoresced when exposed to cathode rays (beam emitted when a cathode has electricity passed through it)

3. Roentgen noticed that there was a mysterious radiation, he called them X-rays, X represented the “unknown”. • He noticed this radiation that passed through black cardboard. • X-rays have a high frequency and can pass through materials like human flesh.

4. Becquerel tried to find X-rays given off by fluoresce material like Uranium. • Instead he discovered the uranium gave off energy which is known as radioactivity – spontaneous emission of particles and energy from atomic nuclei.

5. Crash course chem. • https://www.youtube.com/watch?v=KWAsz59F8gA

6. Inside Atom – subatomic particles • Atoms are made of protons, neutrons, and electrons. • The nucleus (center of atom) is made of protons &neutrons • The electrons orbit the nucleus (far out from it) in shells

7. Isotopes and Atomic Number • JJ Thomson observed two kinds of neon atoms.   • exactly alike chemically, but different in mass. •  He called them isotopes,     • Isotopes - Atoms of the same element that differ in mass.   • have the same number of protons but different numbers of neutron. • mass difference is due to the number of neutrons in isotope • Isotope names:  element – mass egNeon-20

8. Atomic Number • The atomic number (on periodic table) shows the number of protons. • Its symbol is Z.   • determines the identity of the element. • Atoms are electrically neutral so atomic number also shows number of electrons

9. Atomic Mass • Atomic Mass Number - Number of protons + neutrons • Symbol is A.   • Mass only includes protons and neutrons because electrons are basically massless • Unit is the “atomic mass unit” (amu) • Rule of Thumb:  The number of protons determines the identity of the element and the number of neutrons determine the particular isotope of the element     (do isotope sheet)

10. ATOMIC THEORY OF MATTER • Dalton: all matter is composed of small particles called atoms. • Law of conservation of mass – mass is neither created nor destroyed during ordinary chemical reactions or physical changes. • Law of definite proportions – chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound.

11. POSTULATES OF DALTON’S ATOMIC THEORY All matter is composed of atoms. Atom: extremely small particle of matter that retains its identity during chemical rxns. Element: type of matter composed of only one kind of atom. Atoms of a given element have a characteristic mass. Compound: type of matter composed of atoms of two or more elements chemically combined in fixed proportions. Ex. Water, 2 H, 1 O Chemical rxn.: rearrangement of the atoms present in the reacting substances to give new chemical combinations present in the substances formed by the rxn. Atoms are not created, destroyed, or broken into smaller particles by any chemical rxn.

12. ATOMIC SYMBOLS AND MODELS • Atomic symbol: one or two letter notation used to represent an atom corresponding to a particular element. Chlorine: Cl Sodium: Na (latin word natrium)

13. DEDUCTIONS FROM DALTON’S ATOMIC THEORY - Law of multiple proportions: when two elements form more than one compound, the masses of one element in these compounds for a fixed mass of the other element are in ratios of small whole numbers.

14. THE STRUCTURE OF THE ATOM • Nucleus (positively charged, most mass) • Electron (negatively charged, light)

15. DISCOVERY OF THE ELECTRON - J.J. Thomson conducted a series of experiments that showed that atoms were not indivisible particles.

16. How it works. • glass tube with no air, cathode (-), anode (+). • When voltage is turned on the glass tube emits a greenish light. • Greenish light caused by the interaction of the glass with cathode rays (originates for the cathode). • Cathode rays move toward the anode, pass through hole to form beam - Beams bends away from the negatively charged plate and toward the positively charged plate. Concluded that a cathode ray consists of a beam of negatively charged particles (electrons).

17. Youtube video of cathode ray http://www.youtube.com/watch?v=O9Goyscbazk

18. From this experiment Thomson could also calculate the ratio of the electron’s mass, me. Large charge to mass ratio. • Ex. Plum pudding model (Thomson) • Millikan exp. • Observed how a charged drop of oil falls in the presence and in the absence of an electric field. • The charge on the electron is found to be 1.602 x 10^-19 coulombs (C) • Found mass of an electron to be 9.109 x 10^-31 kg.

19. THE NUCLEAR MODEL OF THE ATOM • Rutherford: idea of the nuclear model of the atom • Geiger and Marsden: observed the effect of bombarding thin gold foil with alpha radiation from radioactive substances such as uranium. • Found that most of the alpha particles passed through a metal foil as though nothing were there, but a few (1 in 8000) were scattered at large angles and sometimes almost backward.

20. https://www.youtube.com/watch?v=XBqHkraf8iE

21. Isotopes Atoms of the same element that have different masses. Different numbers of neutrons. Most elements are found naturally with a mixture of isotopes. Ex. Uranium-235 (mass is shown after a hyphen)

22. Atomic mass unit amu 1 amu is 1/12 the mass of a carbon-12 atom The atomic mass of any other atom is determined by comparing it with the mass of the carbon-12 atom. Average atomic mass Weighted average of the atomic masses of the naturally occurring isotopes of an element. percent found mass Hydrogen-1 99.9885% 1.007825 Hydrogen-2 1.115% 2.014102 0.999885 x 1.007825 amu + .0115 x 2.014102 amu = 1.00794 average atomic mass of element

23. Data for an Isotope • Calculations:   • Mass Number =  protons + neutrons  or get from isotope name or symbol • Atomic # = # of P = # of e:  bottom number on symbol; or off of periodic table •     # of neutrons = mass number - atomic number

24. Atomic Mass • Atomic Mass is measured in atomic mass units (u). • Electron = 9.10953 X 10-28 gram  or 0.000549 u      (charge = -1) • Proton    = 1.67265 X 10-24 gram    or    1.0073 u     (charge = +1) • Neutron = 1.67495 X 10-24 gram    or    1.0087 u      (charge = 0) • Atomic mass is based on carbon.   • Carbon has a mass number of 12  :  6 protons and 6 neutrons.  One carbon-12 atom has a mass of 12 u. • All of the other elements are compared to this.  E.g.  Hydrogen was only 1/12th as massive as a carbon therefore it must have an atomic mass of 1 u.

25. Finding Atomic Mass • mass spectrometer is used to find atomic mass • It’s a cathode ray tube with a big magnet around it. • The magnet makes the ray split and the paths of the heavier particles are bent less than the paths of the lighter particles. • atomic masses on the periodic table are based on the “average” atom - an average of the various isotopic forms of the element.

26. Calculating Average Atomic Mass • With isotopes present, you need to use a weighted average to calculate average atomic mass. • (Weight 1 X  No. of 1) + (Weight 2 X No. of 2) + etc. = Total number of isotopes • See Sample Problem on page 492

27. Gamma rays technology • https://www.youtube.com/watch?v=6BddlG1uMXs&feature=relmfu

28. Exposure to Ionizing Radiation • Of the 2000 known isotopes, majority are radioactive. • Radioactive materials release alpha, beta, or gamma radiation (more on that tomorrow!) • Background radiation is always present • Natural sources:  particles from outer space; isotopes found in rocks, soil, and water like uranium; ones in atmosphere like radon; ones in food like K-40 and C-14.

29. cont'd • Manmade sources of radiation - fallout from nuclear weapons testing; airplane flights; fossil fuel and nuclear power use; those released from mining. • Units to monitor radiation exposure: clip • Gray, Gy, SI unit:  quantity of ionizing radiation absorbed by human tissue; 1 Gy = 1 joule of energy absorbed by 1 kg of tissue • Sievert, Sv, - SI unit also,  shows ability of radiation to ionize human tissue

30. Units, cont'd Rad, US unit - absorbed dose of radiation (like the gray) Rem, US unit - ionizing effects on living tissue (like the sievert) Both Rad and Rem are 1/100 of their corrospondingSI unit; That dose if still high so often use millirem (1 mrem = 0.001 rem)

32. Annual exposure • Exp from within body  360 mrem • Natural sources:  300 mrem • Background radiation limit:  500 mrem per yr. • US average;  360 mrem per yr • Limit for workplace eposure:  5000 mrem or 5 rem

33. B3 Health effects of radiation • Ionizing radiation destroys cells and proteins in body • Small doses only destroy a few cells and body can repair itself • Large doses (high radiation density) – too much damage and body can’t repair itself • DNA can be damaged and cause mutations • Can cause leukemia and birth defects

34. 3 Types of Radioactive Decay • 1) Alpha particles – releases a helium nucleus (2 p & 2 n) • Decreases atomic # by 2 and mass # by 4 • can cause burns but can shield by a sheet of paper. • 2)  Beta Decay – a neutron breaks into a proton and beta particle. • beta particle is an electron flying off the atom. • increases atomic number by 1, but no change to mass number. • Can damage body cells, •       blocked  by a sheet of •        aluminum metal

35. 3 Types of Radioactive Decay, cont’d • 3)  Gamma Decay - high energy waves similar to X-rays. • does not change atomic number or mass number. • severe damage to human cells • most penetrating.  Block by thick concrete or lead.

36. Alpha emission reactions • Atom ejects a helium nucleus (alpha particle).  The resulting nucleus contains 2 fewer protons and 2 fewer neutrons. •     238 U    →    234 Th      +           4He •     92              90                     2

37. Beta Particle emissions: • Lose an electron but that mass change is negligible , but nucleus then has 1 more proton •         234 Th       →    0 e    +      234Pa   •           90                   -1               91

38. Nuclear balancing act • In nuclear equations mass number and atomic number (charge) is conserved • Example– write the equation showing beta decay of Cobalt-60: • 60Co → 0e  +   ??? • 27-1            do B5 supplement