1. Atomic Mass and the Mole

2. Relative Atomic Mass Units of grams are TOO LARGE for atoms! Relative atomic mass – compare to small particles – amu – “atomic mass units”

3. Average atomic mass Average atomic mass: weighted average mass of all isotopes of an element Weighted by % abundance of isotopes Example: 69.17% Copper-63 30.83% Copper-65 Find average atomic mass on the periodic table Molecular mass of individual compounds, or formula mass of individual formula units found by adding individual atomic masses. Average atomic mass = 63.55 amu (NOT 64 amu – because it’s weighted)

4. Units of amu are used for measuring single atoms, molecules, or formula units

6. The Mole 1 mole = 6.02 × 10 23 particles It is the number of atoms in exactly 12.0 g of carbon-12 Avogadro’s number

7. What is a mole? A counting number (like a dozen) – If you have a dozen cookies, how many do you have? – If you have a mole of cookies, how many do you have? – If you have 3.5 dozen cookies, how many do you have? – If you have 3.5 moles of cookies, how many do you have?

8. How big is a mole? It is HUGE. HUGE. HUGE. A mole of basketballs would fill four bags the size of the earth. If every person on earth counted out atoms during his or her entire life, counting an atom every second, it would take 3 million years to count out a mole of atoms.

9. Molar Mass Mass of one mole of a substance Units are g/mol Find molar mass on periodic table Molar mass of compounds found by adding individual molar masses

10. Units of g/mol (grams per mole) are used for measuring a mole of particles

11. Practice What is the molar mass of each of the following elements? – Sodium (Na) – Nickel (Ni) – Xenon (Xe)

12. Practice What is the molar mass of each of the following compounds? – H 2 O – NH 3 – C 2 H 6 O

13. REMEMBER: Atomic mass and molar mass are the same number, just different units.

15. Late 1700s – Study of reactions led to new ideas Law of conservation of mass: Mass can’t be created or destroyed by ordinary chemical or physical reactions Law of definite proportions: A compound contains the same proportions of mass regardless of sample size Ex. NaCl will always be 39.9% Na and 60.1% Cl Law of multiple proportions: If two or more compounds of the same two elements exist, then the ratio of the masses are ratios of whole numbers Ex. C and O can combine to form both CO and CO 2

16. John Dalton  Billiard ball model (1803)  John Dalton viewed the atom as a small solid sphere.

17. John Dalton’s Atomic Theory: (Early 1800s) 1. All matter is made of atoms 2. The same elements have exactly the same atoms. In other words, all atoms of an element are identical 3. Atoms cannot be divided, created, or destroyed 4. Atoms of different elements combine in whole-number ratios to form compounds 5. Atoms are combined, separated, or rearranged in chemical reactions

18. Modern Atomic Theory (slight changes to Dalton’s theory) 2. Atoms of the same element CAN be different (isotopes and ions) 3. Atoms CAN be divided – protons, neutrons, and electrons – nuclear reactions can split an atom

19. Models of the Atom  Atom: The smallest particle of an element that retains the chemical properties of that element (reactivity, etc.)

20. We now understand that the atom is composed of:  Nucleus Protons Neutrons  electrons

21. Discovery of the Electron

22. J. J. Thompson

23. J. J. Thomson Experiment: Cathode ray tube (1897)  Discovered electron!!!  Measured charge-to-mass ratio of electron  Electron is negatively charged

24. Thomson’s Plum Pudding Model

26. Robert Millikan

27. Experiment: Oil drop experiment (1909)  Measured charge of electron  Both charge of electron and charge- to-mass ratio were used to determine the mass of an electron  Mass of electron is 1/1837 mass of hydrogen atom

28. Discovery of the nucleus

29. Ernest Rutherford

30. Experiment: Gold foil experiment  Discovered nucleus!!!!  All mass of an atom is in the nucleus  Nucleus is VERY massive

32. Rutherford’s model of the atom

34. Subatomic Particles SymbolChargeMass Electrone-e- 0 Protonp+p+ +11 Neutronn0n0 01

35. Atomic Basics  Atomic number: Number of protons Identifies an element  Mass number Mass number = protons + neutrons ALWAYS a whole number NOT on periodic table

36.  A proton and neutron have about the same mass.  An electron is about 1/2000 the mass of a proton  A neutral atom has the same number of electrons as it has protons  Nuclei are held together by nuclear forces  The electrons determine the size of the atom  Electrons move so fast in such a tiny area they make the atom seem solid (like a moving fan blade)

37. Isotopes  Atoms of the same element with different masses. In other words, atoms with the same number of protons but different number of neutrons

38. Hyphen notation: Nuclear symbol: helium-4

39. Quiz 3A 1. What is the atomic mass of fluorine (F)? 2. What is the molar mass of Tin (Sn)? 3. What is the molar mass of magnesium chloride (MgCl 2 )? 4. What is the molecular mass (mass of one molecule) of propane (C 3 H 8 )? Show all work for all calculations!

40. Quiz 3B

41. Quiz 3C

42. Quiz 3D  How many grams are in 3.50 mol Au?  How many molecules are in 1.21 g CO 2 ?