1. Distinguishing Among Atoms

2. Objectives Define isotope and nuclide Use atomic number, mass number, and charge to determine the number of protons, neutrons, and electrons in an isotope or ion Identify carbon-12 as the standard atom for measuring relative atomic mass Determine the process to calculate atomic mass of an element

3. Discovery of the Nucleus

4. Results of foil experiment if Plum Pudding model had been correct.

5. Actual results

6. The Nucleus Positively charged –Contains protons and neutrons Very dense Central portion of the atom –Contains most of the atom’s mass, but... –Occupies only a tiny fraction of its volume

7. Particles in the Atom ParticleMassCharge Proton1 amu+1 Neutron1 amu0 Electron1/1800 amu–1

8. Rutherford’s Model of the Atom The nucleus is positively charged and very dense Most of the atom is empty space –Electrons move in the empty space

9. Models of the Atom DALTON Indivisible - - - - THOMSON “Plum Pudding”

10. Models of the Atom RUTHERFORD Nucleus + + + +

11. Distinguishing Among Atoms The atomic number (Z) identifies the atom –protons in the nucleus The mass number (A) gives number of protons and neutrons in the nucleus Neutrons = A  Z In a neutral atom, electrons equal protons

12. Important Point!! The number of protons identifies the atom, but... The electrons have the most influence on chemical and physical properties

13. The Hydrogen Atom Hydrogen exists in three forms –Protium One p+, one e – Most abundant form –Deuterium One p +, one n 0, one e – –Tritium One p +, two n 0, one e – Exists only in “trace” quantities

14. Isotopes Atoms of the same element with different masses –Same number of protons –Different number of neutrons

15. Atomic Numbers and Mass Numbers Cu 65 29 Atomic Number (Z) Mass Number (A)

16. Cu 65 29 Protons Protons + Neutrons Remember!! Atoms are neutral!! 36 Neutrons 29 Electrons

17. Cu 2+ 65 29 Protons Protons + Neutrons Remember!! Ions are charged!! 36 Neutrons 27 Electrons

18. Calculating Electrons e e – = Z – charge electrons = atomic number  charge

19. Neutral Atoms e e – = Z – charge e – = Z – (0) electrons equal protons e – = Z

20. Positive Ions e e – = Z – charge e – = Z – (+) more protons than electrons e – < Z

21. Negative Ions e e – = Z – charge e – = Z – (  ) more electrons than protons e – > Z

22. NUCLIDE A general term for any isotope of any element Chlorine-37 Uranium-235 Carbon-14

23. Identifying nuclides copper-65 nuclear notation hyphen notation

24. Atomic Mass

25. Relative Atomic Mass The carbon-12 atom is the standard Arbitrarily assigned mass of 12.000  amu –All other atoms are compared to the carbon-12 atom

26. Atomic Mass and Atomic Mass Units One atomic mass unit (amu) equals exactly 1/12 the mass of a carbon-12 atom

27. Atomic Mass and Atomic Mass Units Atomic mass is the mass of an atom expressed in atomic mass units (amu)

28. Atomic Mass of an Isotope –Found by comparing it with the carbon-12 atom –A relative atomic mass –Relative to carbon-12

29. C 12 6 X X = 1.5 times the mass of carbon-12 Relative Mass of X = ??? 18.000 

30. Y Y = 0.5 times the mass of carbon-12 Relative Mass of Y = ??? 6.000  C 12 6

31. Atomic Mass Most elements are mixtures of isotopes –The % composition of the mixture is constant

32. Accepted Atomic Mass A weighted average of the relative masses of the naturally- occurring isotopes –This is the mass shown on the Periodic Table

33. Weighted Averages Tests40% Labs30% Quizzes20% Homework10%

34. Tests 85 Labs 90 Quiz 80 HW 100 Weighted Averages Mean = 88.8% Total Points = Score x Weight (Pct.) x.40 = 34 points x.30 = 27 points x.20 = 16 points Total = 87 points x.10 = 10 points

35. A chemistry teacher counts tests as 35% of a student’s grade, labs as 25%, quizzes as 20%, homework as 10%, and participation as 10%. A student has category averages of 80%, 80%, 90%, 100%, and 100%, respectively. What is the student’s grade?

36. Weighted Averages Grades: Grade = (Category x weight factor) + (Category x weight factor) + … Atomic Mass (AM): AM = (isotope relative mass x weight factor) + (isotope relative mass x weight factor) +... An isotope’s “weight factor” is called its percent abundance in nature

37. Atomic Mass Hydrogen-199.985 % Hydrogen-20.015 % Average Atomic Mass = 1.007 9  1.007 8  x.999 85 = 1.007 6  2.014 0  x.000 15 =.000 30  

38. Quick Check What important information does the atomic number tell us? What is the difference between relative mass and atomic mass?

39. Practice Write the nuclear symbol for bromine-80 –How many protons, electrons, and neutrons are in this atom? Write the hyphen notation for the element that has 6 protons and 7 neutrons –How many electrons does this atom have? –Write the nuclear symbol for this atom

40. Atomic Mass A natural sample of copper is 69.17% copper-63 and 30.83% copper-65. Calculate the average atomic mass of copper to 2 decimal places.