1. Unit 3 – Atoms: The Building Blocks of Matter - Nuclear Chemistry
Learning Objectives and Learning Target Packet
-Text support Ch. 3 and 21

2. Atomic Theory Continues…
Six Easy Pieces Read Aloud
We have come so far in our understanding of atoms. Centuries of researching and countless scientists devoting their lives to create the understanding of the atom today (textbook concepts).
However it is NOT over. The more we understand about atoms (how they work, their make up, etc…) the greater our ability to advance science and technology in all aspects of our lives (i.e. medicine).

3. Check for Understanding
What is the charge of a proton?
Where would you find a proton in an atom?
What is the charge of an electron?
Where would you find an electron in an atom?
What is the charge of a neutron?
Where would you find a neutron in an atom?
How big is an electron compared to a proton?
How big is a neutron compared to a proton?

4. The Atom
The smallest particle of an element that retains the chemical properties of that element.
Consists of two regions:
Nucleus
1. Small region located at the center of an atom
2. Made up of at least one positively charged particle (proton)
3. Made up of usually one or more neutral particles (neutrons)
Region surrounding the nucleus – Electron Cloud
1. Very large compared to size of nucleus
2. Contains the negatively charged particles (electrons)

5. Refresh: What are Atoms?
Atoms are tiny particles that determine properties of all matter.
Atoms are the building blocks for molecules.
Atoms form elements.
Element: A substance that cannot be broken down into simpler substances by chemical means.

6. Parts of an Atom
Proton: A subatomic particle that has a positive charge and is found in the nucleus of the atom.
Neutron: A subatomic particle that has NO charge and is found in the nucleus of the atom.
Electron: A subatomic particle that has a negative charge and moves around the outside of the nucleus.

7. Label the Atom

8. Subatomic Particles Particle Charge Mass (kg) Location Proton +1
1.67 x 10-27
(1 AMU)
nucleus
Neutron
Electron -1
9.11 x 10-31
(~0 AMU)
Outside nucleus

9. Electron Orbital
-Electrons orbit the nucleus in orbital clouds. -Electrons with different amounts of energy exist in different energy levels.

10. The Electron Cloud Model

11. Electrons in each energy level
Each energy level can hold a limited number of electrons.
The lowest energy level is the smallest and the closest to the nucleus.

12. Atomic Number
The atomic number of an element is the number of protons in the nucleus of an atom of that element.

13. Charge of Atoms
Atoms are not charged even though they have particles that contain charges.
Atoms are neutral because they have EQUAL numbers of protons and electrons.
Ex: Helium Atom
Charge of 2 protons:
Charge of 2 neutrons:
Charge of 2 electrons:
Total charge of He atom: +2 -2

14. Mass Number
The sum of the protons and neutrons in the nucleus is the mass number of that particular atom. (mass # = p + n)

15. Atomic Number vs. Mass Number
Atomic Number: Equal to the number of protons in the nucleus of the atom.
(number of electrons = the number of protons)
Mass Number: Equal to the number of protons AND neutrons in an atom’s nucleus.
Average Atomic Weight (below symbol on PT):
Weighted average of the atomic masses of the naturally occurring isotopes of an element
We will round this number to the nearest hundredth (TWO decimal places)
Example Oxygen’s average atomic mass is = 16.00

16. Application- TOGETHER
Iodine (I) Iron (Fe)
Atomic number ____ Atomic Number ____
Mass Number Mass Number ____
Number of Protons ____ Number of Protons____
Number of Neutrons ____ Number of Neutrons 33
Number of electrons ____ Number of Electrons ____

17. Application- ON YOUR OWN
Nickel (Ni) Radon (Rn)
Atomic number ____ Atomic Number ____
Mass Number ____ Mass Number 222
Number of Protons ____ Number of Protons____
Number of Neutrons 31 Number of Neutrons____
Number of electrons ____ Number of Electrons ____

18. Isotopes
Isotopes of an element have different mass numbers because they have different numbers of neutrons, but they all have the same atomic number.

19. Isotope Examples Carbon – 12 Carbon – 13 Carbon 14 6 protons
6 electrons
6 neutrons
Carbon – 13
7 neutrons
Carbon 14
6 protons
6 electrons
8 neutrons

20. Li Example Lithium 6,7 3 Mass Numbers Symbol Atomic Number
Element Name

21. One More Example Ag
Silver 47
107

22. Isotope Tables Breaking it down…
To find the symbol – determine the atomic number of the element. This is the number of protons
To find the protons- determine the atomic number of the element.
To find the electrons – equal to the number of protons of a neutral atom
To find neutrons: Mass Number – Atomic Number = Number of Neutrons
To find Mass Number: Atomic Number + Number of Neutrons = Mass Number

23. Isotopes

24. Practice

25. Molar Mass Molar Mass: mass in grams of one mole of a pure substance.
Units = g/mol
To calculate molar mass you must find the atomic mass (Units = g).
Ex: Find the atomic mass of Al.
Ex: Find the atomic mass of O2.
Ex: Find the atomic mass of CH4.
* To change atomic mass to molar mass just put the answer (g) over moles.

26. Calculate the molar mass of the following … show ALL your work!
1) NaBr
2) PbSO4
3) Ca(OH)2 
4) Na3PO4
5) (NH4)2CO3
6) C6H12O6 
7) Fe3(PO4)2 
8) (NH4)2S
9) Zn(C2H3O2)2 
10) AgF

27. Warm-up Identify the element: 65 neutrons 48 protons 48 electrons
How many protons, neutrons, electrons in the following isotopes
Mn-56 S-31
Determine the molar mass of the following
H2O
Na2CO3

28. Average Atomic Mass
Why is the mass on the periodic table not a whole number?
Can we have parts of a proton or neutron?

29. Average Atomic Mass - POGIL
Group Roles and Members

30. What is the average atomic mass of Copper?
There are two naturally occurring isotopes of copper.
a.m.u. means atomic mass unit and is the mass of one atom of that particular isotope.
Average atomic mass = Σ (a.m.u. x % abundance in decimal form)
Σ means the “sum of”
=( x ) + ( x )
=63.55 amu and also grams/mol….check the periodic table…..the mass number is the average atomic mass for copper on your periodic table!
Isotope
a.m.u. (atomic mass unit)
% Abundance
Copper-63
69.17
Copper-65
30.83

31. Partner Practice: What is the average atomic mass of Oxygen?
Isotope
a.m.u.
% Abundance
Oxygen-16
99.757
Oxygen-17
0.038
Oxygen -18
0.205
a.m.u. or grams/mol … check the periodic table!

32. Write it Out!!
Describe in your OWN words how you can calculate the average atomic mass of an element.
Be sure to describe what information you need to know and how you use it.
**Be prepared to share out your response!**

33. What isotopes are most common?
Percent Abundance
What isotopes are most common?

34. Lithium has two isotopes… Li-6 and Li-7
Lithium has two isotopes… Li-6 and Li-7. What is the % abundance of each?
Isotope
a.m.u.
% Abundance
Lithium-6 ?
Lithium-7
Look up Lithium’s average mass on the periodic table: 6.94
Now it is time for some ALGEBRA!
6.94 = (x) (1-x). Solve for x …
6.94 = x x (apply distributive property here!)
Combine like terms
6.94 – = x
= x and therefore x= and 1-x =

35. Continued… Isotope a.m.u. % abundance Lithium-6 6.015122 7.59%
92.41%

36. Partner Practice: Calculate the average atomic mass of argon to two decimal places.
Isotope
a.m.u.
% Abundance
Argon-36
35.97
0.337%
Argon-38
37.96
0.063%
Argon-40
39.96
99.600%
39.95 g/mol

37. Application… Skittlium!
Lab Reflection
Question #5:
Determine the average atomic mass for the following sample of skittlium: (Please show your work for full credit!)
Type of Skittlium
Number of Atoms
Skittlium-58 3
Skittlium-59 9
Skittlium-60 1
Skittlium-61 6
Skittlium-62

38. The Mole

39. Mole Music Video

40. The Mole & Avogadro’s Number
NA = 6.02x1023

41. Counting Particles

42. Scientists Counting Unit
Since atoms are so small, it is hard to count individual atoms.
To solve this problem, chemists count by moles.
Moles: SI unit for measuring the amount of a substance.
One mole of Carbon

43. Mole Continued Avogadro’s Number: 6.02x1023 = 1 mole
One mole of anything contains
6.02 x 1023 “particles”
Particles can be atoms, ions,
molecules, electrons, formula units, etc.
Avogadro’s Number: 6.02x1023 = 1 mole
602,200,000,000,000,000,000,000

44. How Big is Avogadro’s Number?
One mole of rice grains would cover the land masses of Earth in a depth of 75 meters.
If you had a mole of pennies and gave away 1 million dollars of it (100,000,000 pennies) a day to everyone in the world it would take you more than years to distribute all your money.
One mole of sheets of paper stacked one on top of the other would reach beyond the solar system.
One mole of basketballs could create a new planet the size of the earth.

45. Conversions with Avogadro’s Number
Start with what you know.
Use the following conversion factor:
6.02 x 1023 particles
1 mole
Cancel units.
Solve (preform the math).

46. “Mole Map”

47. Practice Conversions: Moles to Particles
Determine how many particles of sucrose are in 3.50 moles of sucrose.
Determine the number of atoms in 2.50 mol of Zn.

48. Practice Conversions: Moles to Particles
3. Given 3.25 mol AgNO3, determine the number of formula units.
4. Calculate the number of molecules in 1.15 mol of water.

49. Warm-Up
Calculate the number of molecules in 1.15 mol of water.

50. “Mole Map”

51. Practice Conversions: Particles to Moles
How many moles are in 5.75x1024 atoms Al?

52. Practice Conversions: Particles to Moles
How many ATOMS are in 3.75x1024 molecules of CO2?
How many moles are in 2.50x1020 atoms of Fe?

53. Practice, Practice, Practice
Mole Conversion Worksheet
Practice, Practice, Practice

54. Check for Understanding
1.5 moles Cu = ________ atoms Cu
8.51 x SO3 = ________ moles SO3
Solutions:
9.0 x 1023atoms Cu
28.3g SO3

55. Molar Mass and Conversions
Adding Mass-Mole and Mole-Mass
Finally, put it all together and what do you get?!

56. Molar Mass Molar Mass: mass in grams of one mole of a pure substance.
Units = g/mol
To calculate molar mass you must find the atomic mass (Units = g).
Ex: Find the atomic mass of Al.
Ex: Find the atomic mass of O2.
Ex: Find the atomic mass of CH4.
* To change atomic mass to molar mass just put the answer (g) over moles.

57. Mole-Mass Conversions
Conversion Factor: Ex: While working in a Chemistry lab, Gary needs 3.00 moles of Mn for a chemical reaction. How much Mn does Gary need to mass?
Moles 1
Mass (g)
P. Table

58. Practice with Mole-Mass Conversions
Determine the moles in each of the following:
g Ag
g Zn
kg Fe
Determine the mass in grams of each of the following:
3.57 mol Al
42.6 mol Si
3.45 mol Co

59. Mass-Atoms Conversions
This involves two conversions.
Conversion #1: Mass to moles
Conversion #2: Moles to particles
Ex: How many atoms of gold are in a pure gold nugget having a mass of 25.0 g?
Moles 1
“Particles”
6.02x1023
Mass
(g)

60. Practice Mass-Atom Conversions
How many atoms are in the following samples?
55.2 g Li
0.230 g Pb
45.6 g Si

61. Reverse It… Atoms-Mass Conversions
This involves two conversions.
Conversion #1: particles to moles
Conversion #2: moles to grams.
Moles 1
“Particles”
6.02x1023
Mass
(g)

62. REVERSE IT…. Practice Atom-Mass Conversions
How many grams are in the following samples? x 1024 atoms Bi x 1022 atoms He x 1015 atoms U

63. Tie up your shoes its…
PRACTICE TIME!!!
Mole Conversion Worksheet

64. Mole Music Video

65. Unit 3 – Atoms: The Building Blocks of Matter & Nuclear Chemistry
Review
Unit 3 – Atoms: The Building Blocks of Matter & Nuclear Chemistry

66. Review Questions What is the charge of a proton?
Where would you find a proton in an atom?
What is the charge of an electron?
Where would you find an electron in an atom?
What is the charge of a neutron?
Where would you find a neutron in an atom?
How big is an electron compared to a proton?
How big is a neutron compared to a proton?
What part of the atom takes up the most space?
What part of the atom contains the most mass?

67. Complete the following table
Particle
Charge
Mass (amu)
Location +1 1
Neutron -1

68. Review Questions Complete the following table. Isotope Atomic Number
Mass Number
Number of Protons
Number of Neutrons
Number of Electrons
Zn-64 9 10 24 20

69. Individual Review Time
Review Ch. 3 Worksheets, Practice Packet, and Notes!
Atomic Structure
Isotopes
Mole Conversions
Molar Mass
Average Atomic Mass
% Abundance

70. Explore Pages 649-671 10 minutes Unit Packet Questions and Practice
Nuclear Chemistry
Explore Pages
10 minutes
Unit Packet Questions and Practice

71. Nuclear Chemistry – Pages 649-671
Nuclear Reactions = Changes in the # of protons or neutrons in an atom’s nucleus
Nuclide - The nucleus of an atom
Stability of nuclide depends on the ratio of neutrons to protons
Too few or too many neutrons make the nuclide unstable and nuclide decays to form more stable nuclides.
The most stable nuclei cluster over a range of neutron-proton ratios called the band of stability (page 683 in book)
For lower atomic numbers, typically 1:1 ratios of protons to neutrons, for higher atomic numbers typically 1:1.5
Nuclides are most stable when they reach a magic numbers (natural occurrences that are more stable).

72. Types of Radiation
alpha = rapidly moving helium nuclei (2 protons and 2 neutrons)
beta negative (-) = extremely fast electrons, little mass
gamma – electromagnetic radiation – no mass, no charge, energy only
positron or + = particle with same mass as an electron, but a positive charge
Other types of radiation do exist – microwaves, radar, light, etc.
Radiation is classified as ionizing or non-ionizing
Ionizing radiation = enough energy to change atoms and molecules into ions (alpha, beta, X-rays, gamma)/Can cause changes in living cells
Non-ionizing radiation = cannot ionize matter (radio, light)

73. What is nuclear decay?
Nuclear decay is an exothermic process discovered by Henri Becquerel
Each decay results in a nucleus that contains less energy
Stability of nuclide depends on the number of neutrons to protons
Too few or too many neutrons makes the nuclide unstable
Nuclide decays to form more stable nuclides
Describing Nuclear Decay = Equations are used to represent
Initial nucleus = parent nucleus, new nucleus = daughter nucleus
After any decay, the daughter nuclide is less energetic than the parent
The daughter is also more stable – better neutron/proton ratio

74. Alpha Emission
Occurs with heavy nuclei – decreases neutrons and protons to reach stability.
In alpha emission, the parent nuclei decays into the daughter nuclide by emitting alpha particles (helium nucleus, 42He)
Example – Uranium-238
Nuclide loses 2 protons and 2 neutrons

75. Beta Emission 1. Occurs if too many neutrons to protons
2. Either electron or positron is released from the parent nuclide
3. -, an electron is given off, but is it not one of the original electrons
a. The electron is formed from the change of a neutron into 1 proton and 1 electron
b. The daughter nuclide thus has a new atomic number because the number of protons increases by 1; mass stays the same
c. New atomic number = new element symbol
4. Example: Thorium Th  0-1 Pa
5. Positron or + emission occurs when the parent nuclide has too few neutrons for the number of protons
a. Results in conversion of a proton into a neutron – Releases positron
b. Decreases the atomic number by one, mass stays the same
6. Example: Potassium K  0+1 Ar

76. Electron Capture
1. Occurs with too few neutrons to protons, like positron
2. An inner orbital electron is captured by the nucleus of its own atom.
3. This electron combines with a proton and a neutron is formed.
4. Decreases the number of protons and increases neutrons.
5. The atomic number is decreased by one, but the mass stays the same.
6. Electron capture is shown on the left side of the equation.
10647Ag + 0-1e  10646Pd

77. Neutron and Gamma Decay
Equation results in no change in parent nuclide.
Neutrons can also be released from nuclear reactions
25399Es + 42He  10n Md

78. Types of Radiation Compared
Type of Radiation
Symbol
Change in Mass #
Change in Atomic #
Charge
Alpha
Decreases by 4
Decreases by 2 2+
Beta
No Change
Increases by 1 1-
Gamma

79. Nuclear Fission
Nuclear Fission = a very heavy nucleus splits into more stable nuclei or intermediate mass and releases large amounts of energy
Can occur spontaneously or when nuclei are bombarded with particles.
A chain reaction = a reaction in which the material that starts the reaction is also one of the products and can start another reaction.
For example, when fission of uranium=235 occurs, the 2 or 3 neutrons given off cause the fission of other uranium-235 nuclei.
Continues until all atoms are split or neutrons fail to hit uranium nuclei. (mass hits a certain level.
The minimum amount of nuclide that provides the number of neutrons to sustain a chain reaction = critical mass.
Uncontrolled chain reactions = nuclear bombs
Controlled chain reactions = nuclear reactors (used to produce energy or radioactive nuclides.)
Nuclear Power Plants use heat from reactors to produce electrical energy.

80. Nuclear Fusion
Light mass nuclei combine to form heavier, more stable nucleus.
Like Nuclear Fission, releases high amounts of energy - Releases even more energy per gram of fuel than nuclear fission.
Occurs on the sun and in other stars.
Like Nuclear Fission, uncontrolled fusion reactions are used to make bombs – uncontrolled fusion reactions of hydrogen are the source of energy for the hydrogen bomb.
Fission reactions are used to provide heat and pressure to trigger the fusion.
We are investigating ways to use this power in a controlled way.

81. Calculating Half-lives
A. Half-life (t1/2) = time required for half of the atoms of a radioactive nuclide to decay.
B. Solve half-live problems, first determine the rate constant (k)
k = ln(0.5)/half life years
C. Use the k value to determine how long it takes for decay to occur
t1/2 = ln(percent left as decimal)/k

82. Half-life Examples
The radioactive isotope cobalt-60 has a half-life of 5.27 years. What is the rate constant?
K = ln(0.5)/5.27 =
How long will it take 20% of the cobalt-60 to change to nickel- 60?
t1/2 = ln(.80)/ = 1.69 years
How long will it take 70% of the cobalt-60 to decay to nickel-60?
t1/2 = ln(.30)/ = 9.12 years

83. Be prepared for the Spiral Assessment!