2. CP Chemistry – Chapter 3 Mrs. Albertson Spring 2001

3. Lavoisier Father of Modern Chemistry The first to use truly quantitative research Law of Conservation of Mass Identified components of water as hydrogen and oxygen http://encarta.msn.com/encnet/refpages/RefA rticle.aspx?refid=761571807 http://encarta.msn.com/encnet/refpages/RefA rticle.aspx?refid=761571807 http://encarta.msn.com/encnet/refpages/RefM edia.aspx?refid=461541281 http://encarta.msn.com/encnet/refpages/RefM edia.aspx?refid=461541281

4. The Structure of the Atom http://www.watertown.k12.wi.us/hs/teachers/bue scher/atomtime.asp http://www.watertown.k12.wi.us/hs/teachers/bue scher/atomtime.asp http://www.sci.tamucc.edu/pals/morvant/genche m/atomic/index.htm http://www.sci.tamucc.edu/pals/morvant/genche m/atomic/index.htm Atom: Smallest particle of an element that still retains properties of that element 4 th Century B.C. – Democritus first suggested the idea of atoms Nucleus – contains protons and neutrons Electrons are found outside the nucleus Smaller subatomic particles are not discussed here

5. Discovery of Electrons John Dalton Noticed that % of each element in a compound is always the same: Law of Definite Proportions Carbon Dioxide Always 27.3% carbon and 72.9 % Oxygen

6. Dalton’s Atomic Theory All elements are composed of indivisible particles called atoms Atoms of the same element are identical; atoms of different elements are different Atoms of different elements combine in small whole number ratios to form compounds Chemical reactions occur when atoms are separated, joined, or rearranged Atoms of one element are not changed into atoms of another element, subdivided, or destroyed

7. Crooke’s Experiment – 1870’s Gas Tubes w/2 electrodes (conductors) Anode – positive Cathode – negative Cathode ray tube – applied voltage & beam of light composed of particles was deflected by a magnet – able to determine they were charged particles

8. The Discovery of Electrons J.J. Thomson – Investigating the relationship between matter & electricity CRT w/fluorescent screen allowed him to measure deflection when a magnet was used Measure ratio of charge to mass and determined particles were identical regardless of the gas used/subatomic particles

9. Millikan - 1909 Approximated the mass of an electron to be 1/2000 the mass of an H atom Current 1/1837 9.109 x 10 -31 kg

10. Protons Atoms are neutral so a positive charge must exist Thomson – designed exp to test/H+ moved toward negative end of CRT Protons identified by 1920 Deflection of + particles varied w/different gases. Hydrogen had the greatest deflection and smallest mass Mass of Proton 1.673 x 10 -27 kg

11. Thomson’s “Plum Pudding” Model Nobel Prize 1907 Pudding was + charge and most of the mass of the atom Plums: - charged electrons spread throughout to make the atom neutral Ions: + / - charged atoms: result from the loss or gain of electrons Cations – positive charge / lost electrons Anions – negative charge / gain electrons http://www.sci.tamucc.edu/pals/morvant/genchem/ato mic/page6.htm http://www.sci.tamucc.edu/pals/morvant/genchem/ato mic/page6.htm

12. Radioactivity Discovered 1896 – Radioactivity discovered in Uranium by Becquerel Radiation: energy that is emitted from a source and travels through space Radioactivity: spontaneous radiation from the nucleus of an atom Marie/Pierre Curie – radium & polonium

13. Radioactivity By 1900 3 types of radiation identified Alpha – He ions w no elctrons; 1/10 th the speed of light; stopped by paper or clothing Beta – electrons at high speeds / stopped by a few mm of Al Gamma – form of electromagnetic radiation; more energetic than x-rays; stopped by several cm of Pb or more concrete/ no mass or charge

14. Rutherford’s Gold Foil Resulted in a new model of the atom Atoms contain a small dense nucleus Electrons move around like bees in a hive Diameter of nucleus 1/100,000 the size of the atom 1920 Rutherford proposed neutral particles with the same mass as protons http://micro.magnet.fsu.edu/electromag/java/ruth erford/ http://micro.magnet.fsu.edu/electromag/java/ruth erford/ http://www.brainpop.com/science/matter/atomic model/index.weml?&tried_cookie=true http://www.brainpop.com/science/matter/atomic model/index.weml?&tried_cookie=true

15. Chadwick Credited with the discovery of neutrons Nobel Prize – 1935 Neutron Mass – 1.675 x 10 -27 kg

16. Forces in the Nucleus Like charges normally repel Protons are strongly attracted to one another in the nucleus Also neutron/neutron and neutron/proton attractions These are the result of NUCLEAR FORCES

17. Atomic Number & Mass Number Atomic number – the number of protons in the nucleus; defines what element an atom is Mass number Protons + Neutrons = Mass Number Amu – atomic mass units – 1/12 the mass of a carbon-12 atom 1 proton = 1.007276 amu 1 neutron = 1.008665 amu http://www.sci.tamucc.edu/pals/morvant/genche m/atomic/page8.htm http://www.sci.tamucc.edu/pals/morvant/genche m/atomic/page8.htm

18. Isotopes Atoms of the same element with different numbers of neutrons http://www.sci.tamucc.edu/pals/morvant/genchem/atomic/page9.htm Nuclide – general term for any isotope of any element Each isotope has a % abundance in nature Symbols for isotopes: Lithium – 6 / Lithium – 7 Isotopes differ by Number of neutrons Mass number Atomic mass

19. Isotopes Cont. Of 1500 known isotopes, only 264 are stable; others are radioactive Radioactive decay – alpha or beta particles are emitted and the nucleus changes to form a new element or isotope – continues until a stable form is reached

20. Isotopes Cont Atomic mass – average of the mass of an elements isotopes based on % abundance Carbon – 12.011 amu Carbon 13 = 1.11 % Carbon 12 = 98.89 % Example Problem – Find average atomic mass of carbon:.0111 x 13amu =.1443 amu.9889 x 12 amu = 11.8668 amu 12.011 amu Take relative abundance x mass of isotope and add together

21. Example Problem Using the following information, determine the atomic mass of chlorine: Two isotopes are know: chlorine-35 (mass=35.0 amu) and chlorine-37 (mass=37.0 amu). Their relative abundances are 75.4% and 24.6% respectively.

22. Sample Problems ElementPNEMass # *Avg. Atomic mass Nitrogen-15 67 815 Mn(+2)30

23. Relating Mass to Number of Atoms The MOLE SI unit for amount of substance Amt. of substance that contains as many particles as there are atoms in exactly 12 g of carbon-12 Counting unit – just like a dozen Avogadro’s Number Experimentally determined: 6.022 x 10 23

24. Particles & the MOLE 4 Types of Particles Atoms Ions Molecules Formula Units There are 6.022 x 10 23 particles in 1 mole of any pure substance What is a pure substance?

25. Classification of Matter

26. Molar Mass The mass of 1 mole of a pure substance in a unit of g/mol Equal to atomic mass A molar mass of an element contains 1 mol of atoms Examples: Iron (Fe) = 55.85 g/mol 55.85 g of iron contains 6.022 x 10 23 atoms Water (HOH) = 2 mol H atoms x 1.01 g/mol = 2.02 g + 1 mol O atoms x 16.00 g/mol = 16.00 g 18.02 g

27. Gram/Mole Conversions What is the mass of 5.00 mol of Ni (in grams)? Examine Plan Organize Evaluate How many moles are in 70 g of carbon?

28. Conversions w/Avogadro’s Number How many atoms are in 3.0 mol of He? How many atoms are in 52.63 g of Na? How many H atoms are in 3 mol of water? REMEMBER IF GOING FROM A COUNTING UNIT TO GRAMS OR GRAMS TO COUNTING UNIT – GO TO MOLES FIRST g  mol  atoms atoms  mol  g