Summarize the five essential points of Dalton’s atomic theory

1. Chapter 3-1: The Atom • Summarize the five essential points of Dalton’s atomic theory • Explain the relationship between Dalton’s Atomic Theory and the laws of conservation of mass and definite composition • Explain the law of multiple proportions.

2. Early Ideas of the Atom Democritus • Greek Philosopher • 460-370 B.C. • Stated – Matter could be divided into smaller & smaller particles until it could no longer be divided. • Called these Particles – Atomos (indivisible) • Not based on any physical evidence, only thought

3. Dalton’s Atomic Theory • All matter is composed of extremely small particles called atoms. • Atoms of a given element are identical in size, mass, and other properties. (Isotopes - atoms same element with different mass.) • Atoms cannot be subdivide, created, or destroyed. • Atoms of different elements can combine in simple, whole-number ratios to form chemical compounds. • In chemical reactions, atoms are combined, separated, or rearranged. Late 1700’s – John Dalton – School Teacher(england) Dalton’s Atomic Theory – Summarized

4. +  Law of Conservation Law of Definite Composition Law of Constant Composition A compound always contains the same elements in the same proportions by mass. The mass ratio of A to B will always be the same. In this case 1:3 or 25% to 75% Atoms and the conservation of mass If atoms are indivisible • mass must be conserved A + B  AB 1 a.u. + 3 a.u.  4 a.u.

5. Law of Multiple Proportions • If two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers. • CO2; 1g to 2.66g CO; 1g to 1.33g • Ratio of oxygen would 2:1.

6. Law of Definite ProportionsCalculations • Hydrogen and Oxygen have a mass ratio of 1:16. What is the mass of oxygen needed to form with 14g of Hydrogen? What is the total mass of this compound?

7. Law of Definite ProportionsCalculations • Magnesium and Oxygen have a mass ratio of 3:2. What is the mass of oxygen needed to form with 20g of Magnesium? What is the total mass of this compound?

8. 3.2 Discovering Atomic Structure • Particle arrangement of the atom. • Protons (+) • Electrons (-) • Neutrons (0) • Michael Faraday • Made the connection between the atom and electricity.

9. J.J Thomson • Used a Cathode Ray Tube (CRT) • Discovered electrons • Charge and mass • Electric current passed through low pressure gas in a glass tube.

10. Robert Millikan • Millikan’s Oil-Drop Experiment • Proved the mass and charge of all electrons are identical. • Used an electric plate created a resistant force acting on falling particles • Voltage Controlled • By varying the charge on different drops, he noticed that the charge was always a multiple of -1.6 x 10 -19 C.

11. Ernest Rutherford • Used the Gold Foil Experiment • Marsden and Gieger proved the existence of the nucleus • Bombarded thin sheets of metal with (+) charged particles. • Results of the Particle deflections • Most traveled through • Some Deflected in Both Directions • Very Few Deflected Back

13. Rutherford vs. Thomson

14. Rutherford results • Atom • Nucleus • Central part of the atom. • Very small and very massive = Very dense. • All the mass of an atom. • Positively charged, location of protons. • Electrons • Surround the nucleus, gives size to the atom. • Negatively charged, atom itself is electrically neutral.

15. E. Goldstein • Used canal rays in a cathode ray tube to prove the existence of protons. • Positively charged particles that had significant mass moved towards the cathode.

16. Sir James Chadwick • Using kinematics, Chadwick was able to determine the velocity of the protons. Then through conservation of momentum techniques, he was able to determine that the mass of the neutral radiation was almost exactly the same as that of a proton. . • Neutrons, were neutrally charged particles found in the nucleus to add mass to the atom and to act as “nuclear glue.”

17. Summary of the Atom • Nucleus - very dense, (+)charged, center of an atom. • Protons • Positively charged particle found in the nucleus • Gives mass to the atom. • Neutrons • Neutrally charged particle found in the nucleus. • Gives mass to the atom and acts as “nuclear glue” • Electrons • Negatively charged particles that give the atom its size.

18. 3.3 Modern Atomic Theory Properties of Subatomic Particles

19. Atomic Mass Units (amu) • the equivalent mass of an atom. • Protons = 1amu • Neutrons = 1amu • Electrons = 0 amu

20. Subatomic calculations • Atomic Number (Z) –the number of protons within an atom. (Atomic # = p+) • Every element has its own unique atomic number. • p+ = e- , in an atom • Mass # = p+ + n0

21. Isotope Abbreviations • Atomic Symbol Notation: • X = atomic symbol • Mass Number form: • name of element – mass#

22. Isotope Form Practice • Write the element uranium with a mass # of 235 in atomic and mass form: • Mass # = • Atomic form =

23. Subatomic Particle Practice • How many protons, neutrons and electrons do the following contain? • Carbon – 13 • Carbon • p+ = • e- = • no = Gold p+ = e- = n0 =

24. Symbolic Form Mass Form Atomic # Mass # p+ n0 e- Zinc-63 30 63 30 33 30 Osmium-191 76 191 76 115 76 Tungstun-180 74 180 74 106 74 Barium-139 56 139 56 83 56 Manganese-53 25 53 25 28 25

25. Isotopes • Atoms of the same element with different mass, due to the number of neutrons. • The number of protons determines the identity of an atom !!!!!!! • The number of neutrons can be different for the same type of element. • Hydrogen

26. Hydrogen Isotope Names • 3 isotopes of hydrogen • Protium, Hydrogen –1 • Deuterium, Hydrogen –2 • Tritium, Hydrogen –3

27. Relative Atomic Masses • Although we know actual masses of atoms it is more practical to use their relative masses • The arbitrary point – Carbon-12 • 1 atomic unit (a.u.) = Exactly 1/12 of C-12 • All other atoms, weighed based upon Carbon-12

28. Atomic Mass Average mass of all naturally occurring isotopes of an element. • Calculation • Atomic mass = (mass of each isotope)(% abundance)/100 • Add the answers to each of the isotopes to get the average mass.

29. Modern Atomic Theory • Determine the atomic mass of Oxygen.

30. Modern Atomic Theory

31. Modern Atomic Theory • Determine the atomic mass of Uranium. Uranium =

32. Relating Mass to Particles • The Mole • The number of atoms of an element equal to the number of atoms in exactly 12.0 g of carbon-12. • Referred as a counting number. • 1 mole of any element = 6.02 x 10 23 atoms • Avogadro’s number, there are exactly 6.02 x 1023 atoms of any element in 1 mole of that element. 1mol C = 1mol H = 1mol S =

33. Molar Mass • Mass in grams of 1 mole of any substance. • Equivalent to the formula mass of a compound and to the atomic mass of an element. • 1 mol of S = 32g S • 1 mol of C = 12g C

34. Mole Conversions • 1 mol = Formula mass (or Atomic Mass) • Or the Molar Mass • 1 mol = 6.02 x 10 23 particles • The mole is the central unit in converting the amount of substances in chemistry.

35. Mass – Mole Conversions • Use 1 mol = Formula Mass conversions. • Mass to Moles Conversion • How many moles are in 250.g of Na? • Moles to Mass Conversion • How many grams are in .55 moles of C?

36. Practice Mass-Mole • Determine the number of moles in the following substances: • 6.60g N • 100.g Fe • Determine the mass for the following substances: 3) 6.25 mol Cu 4) .650 mol P

37. Mole – Particle Conversion • Use 1 mol = 6.02 x 10 23 particles • Particle to Mole Conversion • How many moles are equivalent to 550. atoms of S? • Mole to Particle Conversion • How many atoms are in .525 moles of Ca?