1.  Atomic Theory & Nuclear Changes Sarah Fleck 2015 – 2016 Chemistry

2. What is an Atom?  Believing in something that you cannot see!  Tiny fundamental particles that make up matter.  First suggested by Democritus more than 2000 years ago.  If 100,000,000 Cu atoms were placed side by side, they would form a line 1cm long.

3. What is an Atom?  Atoms of the same element are identical  The atoms from different elements are different.  Have no net electrical charge; they are electrically neutral  Atoms can combine to form mixtures and/or compounds.

4. What is an Atom made of?  Electrons (e - )  Negatively charged subatomic particles  Discovered by J.J. Thomson  Protons (p + )  Positively charged subatomic particles  Discovered by E. Goldstein  Neutrons (n 0 )  Subatomic particles that have no charge  Discovered by James Chadwick

5. Properties of Subatomic Particles

6. What does this have to do with anything??!! Protons (p + ) = Atomic # Electrons (e - ) = Atomic # Electrons (e - ) = Protons *Atoms have a neutral charge Neutrons = Atomic Mass* – Atomic # *Round Atomic Mass to a whole number

7. The Nucleus  The central core of an atom  Composed of Protons and Neutrons  Tiny compared with the atom as a whole

8. Practice

11. Concept Map Periodic Table Elements Atom ElectronNucleus ProtonNeutron Atomic Mass & Number Isotope

12. Isotopes  Can occur naturally but also can be man-made  Atoms that have the same number of protons but different number of neutrons  They will also have different atomic mass number  Isotopes are chemically alike; due to the p + and e -  These are the subatomic particles responsible for chemical behavior Nitrogen – 14 Nitrogen – 15

13. Potassium Isotopes

14. Isotope Practice  Two isotopes of carbon are carbon – 12 and carbon – 13. Write the symbol for each isotope using superscripts and subscripts to represent the mass number and the atomic number.

15. Isotope Practice  Three isotopes of oxygen are oxygen – 16, oxygen – 17, and oxygen – 18. Write the complete symbol for each, including the atomic number and mass number.

16. Isotope Practice  Three isotopes of chromium are chromium – 50, chromium – 52, and chromium – 53. How many neutrons are in each isotope, given chromium always has an atomic number of 24?

17. Atomic Mass  Electron = 9.11 x 10 -28 g  Proton or Neutron = 1.67 x 10 -24 g  Fluorine Atom = 3.155 x 10 -23 g  Arsenic Atom = 1.244 x 10 -22 g  These masses are small and impractical to work with.  It is more useful to compare the relative masses of atoms using a reference isotope as a standard.

18. Atomic Mass  Carbon – 12 was chosen as the reference isotope and was assigned a mass of 12 atomic mass units (amu).  Atomic Mass Unit (amu) = one-twelfth the mass of a carbon – 12 atom  Helium – 4 atom, with a mass of 4.0026 amu, has about 1/3 the mass of a carbon – 12 atom.  How many carbon – 12 atoms would have about the same mass as a nickel – 60 atom?

19. Atomic Mass  Carbon – 12  6 protons & 6 neutrons in it’s nucleus, its atomic mass is 12 amu.  12 protons and neutrons account for nearly all of it’s mass  1 proton = 1 amu  1 neutron = 1 amu  Why is the atomic mass of most elements not a whole number????  Chlorine (Cl) = 35.453 amu

20. Relative Abundance  Naturally – most elements occur as a mixture of two or more isotopes  Each has a fixed mass and a natural % abundance  Hydrogen has 3 isotopes; hydrogen – 1, hydrogen – 2, and hydrogen – 3  99.985% of H, occurs as hydrogen – 1  The other two are in trace amounts

21. % Abundance NameNatural % Abundance Mass (amu)“Average” atomic mass Hydrogen – 199.9851.00781.0079 Hydrogen – 20.0152.0141 Hydrogen – 3Negligible3.0160 Chlorine – 3575.7734.96935.453 Chlorine – 3724.2336.966 Atomic Mass = the weighted average mass of the atoms in a naturally occurring sample of the element.

22. % Abundance Practice  Which isotope is more abundant: copper – 63 or copper – 65? (The atomic mass of copper is 63.546 amu)

23. % Abundance Practice  Element X has two natural isotopes. The isotope with a mass of 10.012 amu has a relative abundance of 19.91%. The isotope with as mass of 11.09 amu has a relative abundance of 80.09%. Calculate the atomic mass of this element.

24. % Abundance Practice  The element copper (Cu) has naturally occuring isotopes with mass numbers of 63 and 65. The relative abundance and atomic masses are 69.2% for mass = 62.93 amu, and 30.8% for mass = 64.93 amu. Calculate the average atomic mass of copper.

25. % Abundance Practice  Calculate the atomic mass of bromine. The two isotopes of bromine have atomic masses and relative abundance of 78.92 amu (50.69%) and 80.92 amu (49.31%).

27. Nuclear Chemistry  Radioisotopes  Unstable isotopes  Nuclear Reactions  A reaction that unstable isotopes undergo in order to gain stability  Not affected by temperature, pressure, or catalysts  Cannot be slowed down, sped up, or stopped  Emit large amounts of energy

28. Alpha Particles  Helium Nuclei that is emitted from a radioactive source  The double positive charge is not always represented.

29. Beta Particles  Fast moving electrons  Formed by the decomposition of a neutron in an atom  Neutron decomposes to a proton which remains in the nucleas, and an electron is released

30. Gamma Particles  Photons are emitted from the nuclei of an atom  Often emitted along with Alpha & Beta radiation  No mass & no electrical charge  Does not alter the atomic mass or number of an atom  Behavior the same as X-rays but they do not come from the same source λ

31. Alpha, Beta, & Gamma Characteristics  Alpha Particles  Heavy  Slow moving  Less penetrating  Beta Particles  0 mass  Fast Moving  More penetrating  Gamma  Very dangerous  Cannot be completely blocked

32. Fission  Fission  The splitting of a nucleus into smaller fragments  Nuclei of certain isotopes are bombarded with neutrons  Creates a chain reaction  Large amount of energy is released  1 kg of Uranium – 235 = energy generated by 20,000 tons of dynamite  Reaction is instantaneous & uncontrolled

33. Uranium - 235

34. Fusion  Nuclei combine to create a nucleus of greater mass.  Hydrogen nuclei (protons) fuse to make helium nuclei  Reaction also requires 2 beta particles  Release more energy than fission reactions  Take place at temperatures over 40,000,000 °C

35. Half-Life (t 1/2 )  The time required for one-half of the nuclei of a radioisotope sample to decay to products.  After one half-life, half of the original radioactive atoms have decayed into atoms of a new element. The other half remain unchanged.  Parent – Original Element  Daughter – New Element

36. Half-Life

37. Half-Life Formula Number of Half-Lives Amount Remaining Exponential Form 0A0A0 A 0 x (½) 0 1A 0 x ½A 0 x (½) 1 1 2A 0 x ½ x ½A 0 x (½) 2

38. Half-Life Practice#1  Nitrogen – 13 emits beta radiation and decays to carbon – 13 with a half-life of 10 minutes. Assume a starting mass of 2.00 g of nitrogen – 13.  How long is three half-lives?  How many grams of the isotope will still be present at the end of three half-lives?

39. Half-Life Practice #2  Manganese – 56 is a beta emitter with a half-life of 2.6 hours.  What is the mass of the manganese – 56 in a 1.0 mg sample of the isotope at the end of 10.4 hours?

40. Half-Life Practice #3  The mass of cobalt – 60 in a sample is found to have decreased from 0.800 g to 0.200 g in a period of 10.5 years. From this information, calculate the half-life of cobalt – 60.

41. Half-Life Practice #4  A sample of thorium – 234 has a half-life of 25 days. Will all the thorium undergo radioactive decay in 50 days? Explain.