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
Late 1700’s – John Dalton – School Teacher(england)
Dalton’s Atomic Theory – Summarized
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.

4. Law of Definite Composition
Law of Conservation
Law of Definite Composition
Atoms and the conservation of mass
If atoms are indivisible
mass must be conserved
A + B  AB
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% + 
1 a.u 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 Proportions Calculations
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 Proportions Calculations
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 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 Electrons 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 Electron e-
Symbols
Relative Charge
Mass Number
Relative Mass
Actual Mass
Electron e- -1 u
9.11 x 10-28g
Proton p+ +1 1 u
1.67 x 10-24g
Neutron n0 u
1.68 x 10-24g

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. Isotope
Mass (amu)
% abundance
Oxygen – 16
99.762
Oxygen – 17
.038
Oxygen – 18
.200

30. Modern Atomic Theory

31. Modern Atomic Theory Determine the atomic mass of Uranium. Isotope
Mass (amu)
% abundance
Uranium-235
.720
Uranium - 238
99.280
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 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 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) mol Cu
4) .650 mol P

37. Mole – Particle Conversion
Use 1 mol = 6.02 x 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?