1. Video 3-1 Foundations of Atomic Theory Development of Atomic Models Forces in the Nucleus

2. Chapter 3 Atoms: The Building Blocks of Matter

3. I. Foundations of Atomic Theory Several basic laws were introduced after the 1790’s (emphasis on quantitative analysis):  Law of conservation of mass: mass is neither created nor destroyed during ordinary chemical or physical processes.

4. I. Foundations of Atomic Theory  Law of definite proportions: chemical compounds contain the same elements in exactly the same proportions by mass regardless of the size of the sample.  Ex. NaCl always is composed of 39.34% sodium and 60.66% chlorine by mass.

5. I. Foundations of Atomic Theory  Law of multiple proportions: if two or more different compounds are composed of the same 2 elements, the ratio of mass of the second element combined with a certain mass of the first is always a ratio of small whole numbers.

6. I. Foundations of Atomic Theory  Ex. CO and CO 2 : For the same mass of carbon, the mass of the O in CO to the mass of O in CO 2 will be 1:2  If you had 28 g of CO and 44 g of CO 2, both would contain 12 g of C. The CO would contain 16 g of O and the CO 2 would contain 32 g of O.

7. I. Foundations of Atomic Theory Masses in COMasses in CO 2 Ratios 12 g C 16 g O 12 g C 32 g O C= 12:12 = 1:1 O = 16: 32 = 1:2

8. II. Development of Atomic Models: John Dalton (1808): 1.All matter is composed of extremely small particles called atoms (cannot be subdivided, created, nor destroyed) 2.Atoms of the same element are identical; atoms of different elements are different

9. II. Development of Atomic Models: John Dalton (1808): 3.Atoms combine in simple whole number ratios to form compounds 4.In chemical reactions, atoms combine, separate, or are rearranged.

10. II. Development of Atomic Models: John Dalton (1808): Which of these were later proven wrong and why?  1) Atoms can be subdivided  2) Atoms of the same element do NOT have to be identical

11. II. Development of Atomic Models: J.J. Thomson (1897) and Robert Millikan (1909):  Used cathode rays to determine that atoms contained small negatively charged particles called electrons.  Atoms must also contain positive charges to balance the negative electrons

12. II. Development of Atomic Models:

13.  Other particles must account for most of the mass of the atom  Millikan determine the size of the charge on the electron (oil drop experiment)

14. II. Development of Atomic Models: Ernest Rutherford (1911):  What was the structure of the atom?  Gold Foil Experiment Thomson assumed mass and charged particles were evenly distributed throughout the atom (“plum-pudding” model)

15. II. Development of Atomic Models:

16. Ernest Rutherford (1911): Expected most of the particles to pass with only slight deflection Most particles did, but some showed wide-angle deflections (some almost came back to the source).

17. II. Development of Atomic Models:

18. Ernest Rutherford (1911):  discovery of the NUCLEUS of the atom small, dense, positively charged center of the atom number of PROTONS in the nucleus determines the atom’s identity

19. II. Development of Atomic Models:  Rutherford Atomic Model (solar system model)

20. III. Forces in the Nucleus Repulsive forces should exist between protons in the nucleus (like charges repel).  Why doesn’t the nucleus “fly apart” due to the repulsive electromagnetic force?

21. III. Forces in the Nucleus Strong (nuclear) force:  attractive force that acts over very small distances in the nucleus  causes proton-proton, proton- neutron, neutron-neutron attractions Note: gravitational force is present, but negligible. Why?

23. Video 3-2 Atomic Dimensions Properties of Atoms and Ions Designating Isotopes Elements on the Periodic Table Average Atomic Mass

24. IV. Atomic Dimensions How “big” are subatomic particles? ParticleSymbol Relative Charge Mass Number Relative Mass (amu) Actual Mass (kg) electron e-e- 0 0.0005486 9.109 x 10 -31 protonp+p+ +11 1.0072761.673 x 10 -27 neutron n0n0 01 1.0086651.675 x 10 -27

25. IV. Atomic Dimensions How “big” are subatomic particles?  Atomic radii: 40 to 270 pm  Nuclear radii: about 0.001 pm  Nuclear density: about 2 x 10 8 metric tons/cm 3  1 amu (atomic mass unit) = 1.660540 x 10 -27 kg  Why?

26. V. Properties of Atoms and Ions  atomic number (Z): number of protons in an atom  mass number: number of protons + neutrons in an atom (number of nucleons— particles in the nucleus)  isotopes (nuclides): atoms of the same element that have different masses (different number of NEUTRONS)

27. V. Properties of Atoms and Ions  ions: atoms with a charge (protons  electrons) charge = protons – electrons atoms can only turn into ions by gaining or losing ELECTRONS cation: positive ion anion: negative ion

28. VI. Designating Isotopes  There are two ways to write symbols for an isotope 1.name-(mass number) 2. symbol

29. VI. Designating Isotopes Examples:  Hydrogen has 3 isotopes: Protium Deuterium Tritium  How many neutrons in each isotope?  Note: mass number – atomic number = number of neutrons Hydrogen-1 Hydrogen-2 Hydrogen-3

30. VII. Elements on the Periodic Table  every periodic table will give you at least 3 pieces of information about elements: Li 3 6.941 Atomic Number Symbol Atomic mass (amu)

31. VII. Elements on the Periodic Table  What is the basis for the atomic mass unit (amu)?  1 amu = exactly 1/12 the mass of a carbon-12 atom (6 protons, 6 neutrons, 6 electrons)  All other atomic masses are based on comparisons to C-12 (exactly 12 amu).

32. VII. Elements on the Periodic Table  Example: C-13 has a mass that is 1.083613 times heavier than C-12. The mass of C-13 is (1.083613) x 12 amu = 13.003356 amu

33. VIII. Average Atomic Mass  Carbon has 3 isotopes (nuclides): C-12 (12 amu) C-13 (13.003 amu) C-14 (14.003 amu)  Their average mass should be (12 + 13.003 + 14.003) / 3 = 13.002 amu

34. VIII. Average Atomic Mass  The atomic mass of carbon (periodic table) is 12.011 amu.  WHY?

35. VIII. Average Atomic Mass  Atomic Mass of an element: weighted average of all the atoms in a naturally occurring sample of that element (NOTE: not every atom in that sample has the same mass)  Ex. How would you determine the average age of the students in this class?

36. VIII. Average Atomic Mass  atomic mass = sum of (mass of each isotope x percent abundance)

37. VIII. Average Atomic Mass  Example: C-12 (12 amu) 98.90% C-13 (13.003 amu) 1.10% C-14 (14.003 amu) trace atomic mass of C = (12 amu)(0.9890) + (13.003 amu)(0.0110) + (14.003 amu)(0) = 12.011 amu

38. VIII. Average Atomic Mass  NOTE: the atomic mass of most elements will usually give you an idea of the most common isotope of that element (mass number that is closest to the atomic mass)

40. Video 3-3 Counting Atoms and Stuff

41. IX. Counting Atoms and Stuff  1 mole = 6.0221367 x 10 23 things  Ex. 1 mole of eggs contains 6.0221367 x 10 23 eggs  6.0221367 x 10 23 = Avogadro’s number (N) 1 mol = 6.022 x 10 23 particles

42. IX. Counting Atoms and Stuff  1 mole = 6.0221367 x 10 23 things  1 amu = 1.660540 x 10 -24 g  Suppose you had a sample of one mole of particles. Each particle weighed exactly 1 amu. How many amu would the sample weigh? How many grams would the sample weigh?

43. IX. Counting Atoms and Stuff  1 mole of amu = (6.0221367 x 10 23 )(1.660540 x 10 -24 g) = 1.00000 gram (exactly)  What is the significance of this?  How much will 1 mole of C-12 atoms weigh (in grams)?  12 grams (exactly)

44. IX. Counting Atoms and Stuff Molar Mass:  mass of one mole of a substance  units: grams/mole  equal to the ATOMIC MASS of the element  for compounds, equal to the SUM of the masses of all the elements in the compound (multiply each elements’ atomic mass by the subscript)

45. IX. Counting Atoms and Stuff Example:  Find the molar masses of the following:  NaCl  CO 2  Ca(C 2 H 3 O 2 ) 2 Na = 22.99Cl = 35.45C = 12.01 O = 16.00H = 1.008Ca = 40.08 58.44 grams/mol 44.01 grams/mol 158.168 grams/mol

46. IX. Counting Atoms and Stuff 1 mole of X = atomic mass of X (grams) 1 mole X = 6.022 x 10 23 atoms  These equalities will let you do DIMENSIONAL ANALYSIS Convert grams to moles and moles to grams Convert moles to atoms and atoms to moles  NEVER EVER put 6.022 x 10 23 in front of GRAMS

47. IX. Counting Atoms and Stuff 1 molecule C 6 H 12 O 6 has 6 C atoms, 12 H atoms, and 6 O atoms. 1 mole of C 6 H 12 O 6 has ___ moles of C atoms, ____ moles of H atoms, and ___ moles of O atoms. 6 12 6

48. IX. Counting Atoms and Stuff 1 mole X a Y b = a moles X = b moles Y  Ex. 1 mole Na 2 C 2 O 4 = 2 moles Na = 2 moles C = 4 moles O

49. IX. Counting Atoms and Stuff  1 mole X= 6.0221367 x 10 23 atoms of X  1 mole of X = atomic mass of X (grams)  a moles X = 1 mole X a Y b = b moles Y

50. IX. Counting Atoms and Stuff Moles XMoles X a Y b atoms X grams X grams X a Y b molecules X a Y b 1 mole of X = atomic mass of X 1 mole X= 6.0221367 x 10 23 atoms

51. IX. Counting Atoms and Stuff  Problem-solving:  MAP the problem first (determine what you are starting with, where you want to end up, and the path to follow).  Example: What is the mass of 3.60 moles of Cu?

52. IX. Counting Atoms and Stuff Moles XMoles X a Y b atoms X grams X grams X a Y b molecules X a Y b 1 mole of X = atomic mass of X 1 mole X= 6.0221367 x 10 23 atoms START End

53. IX. Counting Atoms and Stuff  Map out the problem first (determine what you are starting with, where you want to end up, and the path to follow).  Example: What is the mass of 3.60 moles of Cu? Moles of Cu  grams of Cu 3.60 moles Cu x moles Cu grams Cu 1 63.546 = 229 grams Cu

54. IX. Counting Atoms and Stuff Map the following problems FIRST, then solve:  How many moles are in 11.9 grams?  How many atoms are in 3.60 x 10 -10 grams of gold?  How many grams of sodium are in 2.34 moles of Na 2 CO 3 ?  How many grams of Fe are in 13.86 grams of Fe 2 O 3 ?

55. IX. Counting Atoms and Stuff Moles XMoles X a Y b atoms X grams X grams X a Y b molecules X a Y b 1 mole of X = atomic mass of X 1 mole X= 6.0221367 x 10 23 atoms