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. Today’s Definition of an Atom
Atoms are the smallest particle of matter that retains its unique properties.
We use a scanning tunneling microscope to see the surface of atoms.

5. 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.

6. Lavoisier and Proust Antoine Lavoisier (1743-1794) Proust (1754-1826)
Father of Modern Chemistry
Redefined the compound
2 or more elements
Performed Experiments to test ideas.
Proved the law of conservation
Proust ( )
Discovered that in the formation of water – the proportions of H and O were constant.
8:1 Hydrogen to Oxygen
Law of Definite Proportions
Elements combine in definite mass ratios to form compounds

7. 3.2 Discovering Atomic Structure
Particle arrangement of the atom.
Protons (+)
Electrons (-)
Neutrons (0)
Michael Faraday
Made the connection between the atom and electricity.

8. 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.
Plum-Pudding model or Chocolate Chip Cookie Dough

9. J.J Thomson CRT with an electric field.
Moved away from neg. charged cathode proving the particles were negatively charged and determined the mass to be 1/1800 of a hydrogen atom.

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

12. Rutherford Concluded Nuclear Model Nucleus Central part of the atom.
Very small and very massive = Very dense.
All the mass of an atom.
Positively charged.

13. Atoms are Empty Space
C:\Documents and Settings\murban\Desktop\space within atoms

14. Henry Becquerel Radioactivity
Spontaneous emission of radiation from an element.
Caused by unstable atoms, which release subatomic particles.
Alpha, Beta and Gamma radiations.
Discovered the radioactivity of elements, like uranium, using unexposed photo film.
Rutherford’s gold foil exp

15. Summary of the Atom Nucleus Electrons Protons Neutrons
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.

16. Atomic Structure
Development of the atomic Model

17. 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

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

19. 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

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

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

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

23. Symbolic Form
Mass Form
Atomic #
Mass # p+ n0 e-
Zinc-63 74
106
139 56 25 28

24. Ions atoms that have lost or gained one or more electrons.
Octet rule – atoms try to maintain a full outer electron shell.
8 for most elements, 2 for hydrogen
Ions charge indicates the # of electrons gained or lost.
(+) charge = loss of electrons
(-) charge = gain of electrons
Charge of ion = # protons - # electrons

25. Ion Practice
Write the symbol for the ion with 9 protons and 10 electrons.
F has 9 protons
9-10 = -1
F 1-

26. Subatomic Particles of Ions
How many protons and electrons are present in the following?
Cl 1- Na 1+ Fe 3+
Cl 1- :
Na 1+ :
Fe 3+ :

27. Ion Practice
What is the formula for the ion of Nitrogen, which has 10 electrons?

28. 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

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

30. 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

31. 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.

32. Modern Atomic Theory Determine the atomic mass of Oxygen. Isotope
Mass (amu)
% abundance
Oxygen – 16
99.762
Oxygen – 17
.038
Oxygen – 18
.200

33. Modern Atomic Theory Determine the atomic mass of Uranium. Isotope
Mass (amu)
% abundance
Uranium-235
.720
Uranium - 238
99.280
Uranium =

34. 3.4 Changes in the Nucleus Nuclear Reactions Strong Nuclear Forces
Changes that occur within an atom’s nucleus.
Change the composition of the atom’s nucleus by altering the number of protons or neutrons.
Strong Nuclear Forces
Hold the nucleus together, neutrons help increase or add to this nuclear force.

35. Radioactive Decay
Occurs when an atom emits one of the different forms of radiation.
The original nucleus decomposes,”decays,” to form a new nucleus.
The new nucleus will continue to decay until it reaches a stable configuration.

36. Types of Radioactive Decay
Alpha Decay (Not a serious health hazard)
Nucleus ejects a helium nucleus, called an alpha particle, and becomes a smaller nucleus with less positive charge.
What would alpha decay of U-238 look like?

37. Alpha Decay

38. 3.4 Beta Decay 100 times more penetrating than alpha
The effect of turning a neutron in the nucleus into a proton, causing a beta particle (electron) to eject from the atom and the formation of a neutrino.
Write the equation for the beta decay of sodium – 24.

39. Beta Decay

40. 3.4 Gamma Decay
High energy non-particle, with no charge or particle mass called a photon.
Caused by the nucleus changing to a lower energy level, releasing photons in the form of x-ray or other high end radiation.
Only stopped by thick concrete or lead.
Extremely dangerous.

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42. Fusion vs. Fission Fusion Fission
Process of combining nuclei of two atoms.
Occurs in the stars.
Fission
Process of splitting the nucleus of an atom.
Nuclear reactors or bombs.

43. Gamma Decay