1. Atomic Structure

2. Dalton’s Atomic Theory
John Dalton ( ) had four theories
All elements are composed of submicroscopic indivisible particles called atoms
Atoms of the same element are identical. The atoms of anyone element are different from those of any other element
Atoms of different elements can physically mix together or can chemically combine w/ one another in simple whole-number ratios to form compounds
Chemical reactions occur when atoms are separated, joined, or rearranged. However, atoms of one element are never changed into atoms of another elements as a result of a chemical reaction

3. Atom
Atom- smallest particle of an element that retains the properties of that element
Dalton’s Model of the Atom:
Billiard Ball Model

4. Basic Laws Conservation of Mass
Law of Definite Proportion- compounds have a constant composition by mass.
They react in specific ratios by mass.
Multiple Proportions- When two elements form more than one compound, the ratios of the masses of the second element that combine with one gram of the first can be reduced to small whole numbers.

5. What?! Water has 8 g of oxygen per 1 g of hydrogen.
Hydrogen peroxide has 16 g of oxygen per 1 g of hydrogen.
16/8 = 2/1
Small whole number ratios

6. Example of Law Of Multiple Proportions
Mercury has two oxides. One is 96.2 % mercury by mass, the other is 92.6 % mercury by mass.
Show that these compounds follow the law of multiple proportion.
Speculate on the formula of the two oxides.

7. Assume 100 g sample 96.2/3.8 = 25.3 g Hg/1 g O
25.3/12.5= 2/1of Hg Hg2O and HgO
Amount Hg
Amount O
Compound A
96.2 g
3.8 g
Compound B
92.6g
7.4 g

8. A Helpful Observation
Gay-Lussac- under the same conditions of temperature and pressure, determined that gas compounds always react in whole number ratios by volume.
Avogadro- interpreted that to mean
at the same temperature and pressure, equal
volumes of gas contain the same number of
particles (called Avogadro’s hypothesis)

9. Dalton’s Atomic Theory Revised after New Evidence
Part 2: Discovery of Isotopes: Atoms of the same element can have different masses
Part 3: Atoms are Divisible made up of subatomic particles
Part 3: Atoms cannot be created nor destroyed in ordinary chemical reactions They can in NUCLEAR REACTIONS

10. Electron J.J Thomson (1856-1940) – discovered the electron in 1897
Electron is the negative charged subatomic particle
An electron carries exactly one unit of negative charge & its mass is 1/1840 the mass of a hydrogen atom

11. Cathode Ray
The Cathode Ray tubes pass electricity through a gas that is contained at a very low pressure

12. The Electron
Figure 2.4
Streams of negatively charged particles were found to emanate from cathode tubes.
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13. The Electron
Figure 2.4
Thompson measured the charge/mass ratio of the electron to be 1.76  108 coulombs/g.
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14. Thomson’s Atomic Model
Plum Pudding Model
Thomson thought electrons were like plums embedded in a positively charged “pudding”, so his model was called the “plum pudding” model

15. Mass of Electron
In 1909 Robert Millikan determined the mass of an electron with his Oil Drop Experiment
He determined the mass to be x kg
The oil drop apparatus

16. Millikan’s Experiment
Atomizer - +
Oil
Microscope

17. Millikan’s Experiment
Atomizer
Oil droplets - +
Oil
Microscope

18. Millikan’s Experiment
X-rays
X-rays give some drops a charge by knocking off electrons

19. Millikan’s Experiment+

20. Millikan’s Experiment- - + +
They put an electric charge on the plates

21. Millikan’s Experiment- - + +
Some drops would hover

22. Millikan’s Experiment- - - - - - - + + + + + + + +

23. Millikan’s Experiment- - + +
Measure the drop and find volume from 4/3πr3
Find mass from M = D x V

24. Millikan’s Experiment- - + +
From the mass of the drop and the charge on
the plates, he calculated the charge on an electron

25. Radioactivity Discovered by accident Bequerel Three types
alpha- helium nucleus (+2 charge, large mass)
beta- high speed electron
gamma- high energy light

26. Ernest Rutherford
Rutherford ( ) proposed that all mass and all positive charges are in a small concentrated region at the center of the atom
He used the Gold-Foil Experiment to prove his theory & in 1911 he discovered the nucleus
Nucleus- central core of an atom, composed of protons and neutrons
The nucleus is a positively charged region and it is surrounded by electrons which occupy most of the volume of the atom

27. Rutherford’s Experiment
Used uranium to produce alpha particles
Aimed alpha particles at gold foil by drilling hole in lead block
Since the mass is evenly distributed in gold atoms alpha particles should go straight through.
Used gold foil because it could be made atoms thin

28. Florescent
Screen
Lead block
Uranium
Gold Foil

29. What he expected

30. Because

31. Because, he thought the mass was evenly distributed in the atom

32. What he got

33. How he explained it
Atom is mostly empty space
Small dense, positive piece at center
Alpha particles are deflected by it if they get close enough +

34. +

35. Modern View The atom is mostly empty space Two regions
Nucleus- protons and neutrons
Electron cloud- region where you have a chance of finding an electron

36. Neutron James Chadwick (1891-1974) – discovered the neutron in 1932
Neutron is a subatomic particle with no charge but their mass is nearly equal to that of a proton

37. Quark
Protons & Neutrons can still be broken down into a smaller particle called the Quark
The Quark is held together by Gluons

38. Subatomic particles Relative mass Actual mass (g) Name Symbol Charge
Electron e- -1
1/1840
9.11 x 10-28
Proton p+ +1 1
1.67 x 10-24
Neutron n0 1
1.67 x 10-24

39. Sub-atomic Particles
Z - atomic number = number of protons determines type of atom
A - mass number = number of protons + neutrons in the nucleus
Number of protons = number of electrons if neutral

40. Rules for Atomic Structure
Atomic # = # of Protons
# of Protons = # of Electrons if neutral
Protons – Electrons = charge if not neutral
Mass # = # of Protons + # of Neutrons
M.A.N :
mass number – atomic number = # neutrons

41. Isotopes
Isotope- atoms that have the same number of protons but different number of neutrons
Since isotopes have a different number of neutrons the isotope has a different mass number.
Isotopes are still chemically alike because they have the same number of protons and electrons

42. Examples of Isotopes

43. Isotopes
Isotopes are atoms of the same element with different masses.
Isotopes have different numbers of neutrons. 11 6 C 12 6 C 13 6 C 14 6 C
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44. X X Na Nuclear Symbols Mass number Atomic number
Contain the symbol of the element, the mass number and the atomic number X
Mass
number X A Z
Atomic
number 23 Na 11

45. Nuclear Symbols Proton = p+ Electron = e- Neutron = n0 Charge
Mass # - Superscript
Atomic # - Subscript

46. Naming Isotopes Put the mass number after the name of the element
carbon- 12
carbon -14
uranium-235

47. Na Electrical Charges Neutral atoms have no net electrical charges
Given the number of negative charges combines with the number of positive charges = Electrically Neutral
11 protons
10 electrons
12 neutrons 23 1+ Na 11

48. Example 1
Write the isotopic symbols for three isotopes of Silicon in which there are 14, 15, and 16 neutrons 28 Si 14 29 Si 14 30 Si 14

49. Au Example 2 197 79 Find the number of protons number of neutrons
number of electrons
Atomic number
Mass Number 79
118 79 79
197

50. Example 3
How many neutrons are in the 3 isotopes of magnesium-24, magnesium-25, and magnesium-26? 24 Mg 12
12 n 25 Mg
13 n 12 26 Mg
14 n 12

51. Example 4
Write the nuclear symbol for the atom that has an atomic number of 9 and a mass number of 19. How many electrons and neutrons does this atom have? 19 F 9
9 electrons
10 neutrons

52. Example 5
Write the formula for the fluoride ion. Its mass number is still 19. How many protons, neutrons and electrons are there? 1- 19 F 9
9 protons
10 electrons
10 neutrons

53. Example 6
How many protons, neutrons & electrons are found in the calcium-40 ion? 2+ 40 Ca 20
20 protons
18 electrons
20 neutrons

54. Measuring Atomic Mass Unit is the Atomic Mass Unit (amu)
One twelfth the mass of a carbon-12 atom
Each isotope has its own atomic mass. We need the average from the percent abundance
Each isotope of an element has fixed mass and a natural % abundance
You need both of these values to find the Atomic Weight

55. Calculating Atomic Weight
Cl amu and 75.77% abundance
Cl amu and 24.23% abundance
To solve for Cl-35
AMU x Abundance
x .7577 =
You solve for Cl-37

56. Atomic Weight Cont. Cl-37 AMU x Abundance 36.966 x .2423 = 8.957
Now you combine your two answers =
35.453
Look at Cl on the table. What is the Atomic Weight?

57. Example 7
There are 2 isotopes of lithium, one with a mass of amu and an abundance of 7.42%. The other isotope has a mass of amu and an abundance of 92.58%. Calculate the atomic weight (average atomic mass) of lithium.

58. Example 8
Calculate the atomic weight of copper. Copper has two isotopes. One has 69.1% and has a mass of amu. The other has a mass of amu. What is the atomic weight???

59. Example 9
There are two isotopes of nitrogen, one with an atomic mass of amu and one with a mass of amu. What is the percent abundance of each? Nitrogen’s average atomic mass is amu.
X+ Y = 1 so st isotope = X
2nd isotope (Y) = 1-X

60. Atomic Weight & Decimals
Atomic Weight- of an element is a weighted average mass of the atoms in a naturally occurring sample of an element
Atomic Weights use decimal points because it is an average of an element

61. Experimental Determination of Atomic Mass
Atomic and molecular masses can be measured with great accuracy with a mass spectrometer.
Figure 2.13
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