1. Unit 2: Atomic Structure & Nuclear Chemistry

2. This chapter will introduce
Section 1: Counting Subatomic Particles
Section 2: Isotopes
Section 3: Nuclear Decay
Section 4: Half-Life Calculations
Section 5: Fission & Fusion

3. Section 1 Atomic Structure
The atom is defined as the smallest particle of an element that retains the properties of that element.
It consists of 2 regions & 3 sub-atomic particles

4. What’s in an atom? An atom is made of three sub-atomic particles
Location
Mass
Charge: Fix
Proton
Nucleus
1 amu = 1.6710-27 kg +1
Neutron
Nucleus
1 amu = 1.6710-27 kg
Electron
Outside the nucleus
amu
9.1010-31 kg -1
1 amu (“atomic mass unit”) = 1.66  kg

5. The TWO Regions of the Atom
: small, dense center
containing the protons &
neutrons
Electron Cloud: contains electrons and mostly empty space

6. Counting Subatomic Particles: Atomic Number
Every atom has a different number of protons which determines the identity of the atom
The atomic number shows the number of protons. Atomic number = protons
To get the Atomic Number: Find the whole number of the element from the periodic table
Nitrogen’s Atomic # = 7

7. Calculating Electrons in a Neutral Atom (APE)
A neutral atom has the same number of protons and electrons
Protons = electrons

8. Mass Number To calculate mass number: 2 options
The total number of protons + neutrons in the nucleus of a specific atom is called the mass number.
To calculate mass number: 2 options
1. Add the # of given protons and # neutrons together
Mass # = protons + neutrons
When no info has been given:
2. Round atomic mass to a whole number to get an
element’s most likely common mass number
Nitrogen’s atomic mass = amu mass # = 14

9. Calculating Neutrons (MAN)
To calculate the number of neutrons: 2 options
1.Subtract the atomic number from the
mass number
2. Subtract the proton # from the mass #
# Neutrons = Mass # - Atomic #
14 -7 = 7 neutrons

10. Calculating Electrons in an ION
An ion has a charge.
The number of protons is NOT EQUAL to the number of electrons
Can be a positively charged ion called a cation or a negatively charged ion called an anion
How to calculate number of electrons:
Subtract electron number from proton number.
Overall Charge = protons - electrons

11. Determining the Number of Electrons
Charge = # of protons – # of electrons
Atomic number = # of protons
Charge = -1
Example:
How many electrons does Br-1 have?
Atomic number for Br = 35 = # of protons
-1 = 35 - electrons
Electrons = 36

12. Determining the Number of Electrons
Charge = # of protons – # of electrons
Atomic number = # of protons
Example:
How many electrons does Al+3 have?
Charge = +3
Atomic number for Al = 13 = # of protons
+3 = 13 - electrons
Electrons = 10

13. Displaying Information of atoms: Nuclear Symbol
Element Symbol
1 or 2 letters, found on the periodic table X A C Z
Charge
# protons - # electrons
(assumed to be “0” if blank)
Mass number
# protons + # neutrons
Atomic number # of protons

14. Example: Nuclear symbol
Element Symbol
O = Oxygen O 16 -2 8
Charge -2
Mass number 16
8 p
10 e
8 n
Atomic number 8

15. Self Check
What are the number of protons, neutrons & electrons in each atom?
19F Fe 204Hg
9 p
10 e
10 n
26 p
26 e
31 n
80 p
80 e
124 n

16. Displaying Information of atoms: Hyphen notation
Element Name
copper – 65
mass number
____ p _____e ______n
Find the Missing Values
NuclearSymbol
Hyphen Notation
Atomic #
Mass #
Charge
Proton
Neutron
Electron
Magnesium-25 +2 82
126

17. Let’s Practice Remember: Atomic number is the identity
Atomic number = protons
Charge = proton - electrons
Mass # = protons + neutrons
Nuclear Symbol
Hyphen Notation
Atomic #
Mass #
Charge
Proton
Neutron
Electron
Magnesium -25 +2 82
126 12 25 12 13 10
Lead-208
208 82

18. SECTION 2 How do Atoms Differ?
Isotopes– are atoms of the same element with a different number of neutrons or mass number
Most elements contain a mixture of 2 or more isotopes. Each one having its own
mass and abundance.
Some isotopes are radioactive—but not all…many are quite stable!

19. C C Carbon-12 Carbon-13 Identifying Isotopes
Isotopes can be differentiated by their different mass numbers in the element symbol 12 C 13 C
Carbon-12
Carbon-13
Or by the mass number following their name.

20. Isotopes

21. You Try!
What the number of protons, neutrons and electrons in each isotope? 19F 18F 203Hg 194Hg P= E= N= P= E= N= P= E= N= P= E= N=

22. Mass Number versus Atomic Mass
Average Atomic Mass
# of protons + # of neutrons
Average of actual masses
Always a whole number
Not a whole number
For one specific isotope only
Weighted average of all isotopes
Is not found on the periodic table
Is found on the periodic table

23. Calculating a Weighted Average!
Practice Problems:
(1) Mrs. Soto’s chemistry semester grades are calculated using a weighted average of three category scores:
Major Grades= 60% of your grade
Minor Grades= 30% of your grade
Semester Exam=10% of your grade
If a student had the following scores, what would they receive for the semester?
Major= 80
Minor= 60
Semester Exam=65

24. Weighted Average
Step (1): Multiply each score by the % that it is weighted.
Step (2): Add these products up, and that is the weighted average!
60% x 80 = 48.0
30% x 60 = 18.0
10% x 65 =
Add them up!!
A “normal average” would be calculated by simply adding the raw scores together and dividing by 3…
= 205 ÷ 3 = 68.3 = D +
72.5

25. Calculating Average Atomic Mass
Average atomic mass is a weighted average of the masses of all naturally occurring isotopes.
Actual mass (not mass number)
“Sum of”
Average
atomic
mass  ( )  =
Abundance of isotope
Mass of isotope
What fraction of the time is that isotope present?

26. Example of Finding Average Atomic Mass
Find the atomic mass of chlorine
Chlorine-35 has a mass of amu.Chlorine-37 has a mass of amu & is 24.22% abundant.
Remember that percents add up to 100.
So they said the second isotope is present 24.22% of the time.
This means that the first isotope is present = 75.78% of the time
Isotope
Mass
Percent
Decimal 1
amu
75.78
0.7578 2
amu
24.22
0.2422
This chart summarizes the information in the problem:
= amu
(this is what’s on the periodic table for Cl!)

27. SELF CHECK! Isotope Mass Percent Decimal 1 10.012 amu 19.91 .1991 2
Example:
Element X has 2 natural isotopes. Calculate the average atomic mass. 1st isotope has a mass of a.m.u with 19.91% abundance % of the 2nd element has a mass of a.m.u.
This chart summarizes the information in the problem:
Isotope
Mass
Percent
Decimal 1
amu
19.91
.1991 2
amu
80.09
.8009
= amu
(Element X is Boron!)

28. SELF CHECK ! Isotope Mass Percent Decimal 1 62.93 amu 69.11 .6911 2
Example:
Calculate the average atomic mass of copper if it has 2 isotopes % has a mass of a.m.u and the rest has a mass of a.m.u.
Remember that percents add up to 100.
So they said the first isotope is present 69.11% of the time.
This means that the second isotope is present = 30.89% of the time
Isotope
Mass
Percent
Decimal 1
62.93 amu
69.11
.6911 2
64.93 amu
30.89
.3089
This chart summarizes the information in the problem:
= amu

29. Section 3: Nuclear Chemistry
Nuclear chemistry is the study of the changes of the nucleus of atoms.
Nuclear Reactions involve changes within the nucleus where as chemical reactions involve the loss, gain or sharing of electrons.

30. The Nucleus
Remember: The nucleus is made up of protons and neutrons. The are collectively called nucleons.

31. Radioactivity
A stable nucleus holds together well. An unstable nucleus will decay or break down, releasing particles and/or energy in order to become stable.
An atom with an unstable nuclei is considered “radioactive”.

32. There are several ways radioactive atoms can decay into different atoms!
Transmutation:
Type of nuclear reaction that will change the number of protons and thus will create a different element.
Atoms with an atomic number larger than 92 are created through this process

33. He α U Th He Alpha Decay Loss of an -particle (a helium nucleus)
Atomic number decreases by 2 and mass number decreases by 4
Penetrating Power: LOW: Can be blocked by clothing or thin paper
Example OR He 4 2 α 4 2 U
238 92
 Th
234 90 He 4 2 +

34. Alpha Decay

35. Alpha Decay
Uranium Thorium

36.  e I Xe e Beta Decay Loss of a -particle (a high energy electron)
Atomic number increases by 1 and mass number stays the same. A neutron becomes a proton and a high speed electron that is discharged from the nucleus.
Penetrating Power: Medium: Can be blocked by thin metal or wood
Example  −1 e or I
131 53
 Xe 54 e −1 +

37. Beta Decay

38. Beta Decay
Thorium Protactinium

39. Gamma Emission
Loss of a -ray (high-energy radiation that almost always accompanies the loss of a nuclear particle)
Atomic number and mass number stays the same
Penetrating Power: High: Can only be blocked by thick metal or thick concrete
Example  I
131 53
 e +

41. Writing Balanced Nuclear Equations
The sum of the atomic number & mass number must balance on both sides of the equation
Often problems will have 1 particle missing & you will need to identify it
Alpha decay of thorium-230
Beta decay of cesium-137

42. Np Pu ____ You Try! Beta decay of zircomium-97
Alpha decay of americium-241
Complete this: Np
235 93
 Pu
239 94
____ +

43. Making New Elements & Isotopes: Bombarding the Nucleus
All Transuranium elements (atomic numbers >92) have been made by bombarding the nucleus with neutrons & other atoms in accelerators

44. Nuclear Transformations
Nuclear transformations can be induced by accelerating a particle and colliding it with the nuclide.
These particle accelerators are enormous, having circular tracks with radii that are miles long.

45. Section 4 Radioactivity & Half Life
Radioactive isotopes decay at a characteristic rate measured in half life.
A half life is the time required for half of the amount of radioactive atoms to decay. The time ranges from seconds to millions of years

46. Common Radioactive Isotopes
Isotope Half-Life Radiation Emitted
Carbon ,730 years b, g
Radon days a
Uranium x 108 years a, g
Uranium x 109 years a

47. Radioactive Half-Life
After one half life there is 1/2 of original sample left.
After two half-lives, there will be
1/2 of the 1/2 = 1/4 the original sample.

48. Graph of Amount of Remaining Nuclei vs Time
A=Aoe-lt A

49. Half Life Calculations
HOW TO’s
1. To calculate the number of half lives, divide the half life (T1/2) into the total time (T). 
T/T1/2 = # of half lives 
2. Use the equation to calculate remaining amount left over after a certain number of half lives have passed.
Amt remaining = (initial amt) (.5)n (# of half lives)

50. Example
You have 100 g of radioactive C-14. The half-life of C-14 is 5730 years.
How many grams are left after one half-life?
How many grams are left after two half-lives?

51. Examples
Suppose you have 20 grams of sodium-24. Its half-life is 15 hours. How much is left over after 60 hours.

52. Examples
Uranium-238 has a half life of 4.46 x 109 years. How long will it take for 7/8th of the sample to decay?

53. You Try 1.
1.5 grams of a 12.0 g sample are left after 114 s. What is the half life of radium-222 ?

54. You Try 2. A sample of 3x107 Radon atoms are trapped
in a basement that is sealed. The half-life of
Radon is 3.83 days. How many radon atoms
are left after 31 days?
answer:1.2x105 atoms

55. Section 5 Nuclear Fission v Nuclear Fusion Fission: How does one tap all that energy?
Large atoms split into smaller atoms that generate huge amounts of energy.
Carried out in nuclear reactors.
Could result in a chain reaction of fission like the atomic bomb

56. Nuclear Fission
Bombardment of the radioactive nuclide with a neutron starts the process.
Neutrons released in the transmutation strike other nuclei, causing their decay and the production of more neutrons.
This process continues in what we call a nuclear chain reaction.

57. Nuclear Fission
If there are not enough radioactive nuclides in the path of the ejected neutrons, the chain reaction will die out.
Therefore, there must be a certain minimum amount of fissionable material present for the chain reaction to be sustained: Critical Mass.

58. Nuclear Reactors
In nuclear reactors the heat generated by the reaction is used to produce steam that turns a turbine connected to a generator.

59. Nuclear Reactors
The reaction is kept in check by the use of control rods.
These block the paths of some neutrons, keeping the system from reaching a dangerous supercritical mass.

60. Nuclear Fusion Smaller atoms are combine to form a large atom.
Occurs in the sun and stars
Generates huge amounts of energy

61. Nuclear Fusion Fusion would be a superior method of generating power.
The good news is that the
products of the reaction are
not radioactive.
The bad news is that in order to achieve fusion, the material must be in the plasma state at several million kelvins.
Tokamak apparati like the one shown at the right show promise for carrying out these reactions.
They use magnetic fields to heat the material.