1. Final Exam
Study Schedule Spring 2018

2. Final Exam Study Schedule
Monday May 21
Tuesday May 22
Rounding Significant Figures, Scientific Notation
Measurement, Accuracy, Precision
Atomic Nuclear
Periodic Table
Electrons, Electron Configurations
Balancing Equations

3. Accuracy & Precision
The degree to which the result of a measurement or calculation conforms to the correct value or a standard.
How close you are to the actual answer.
A measurement or calculation as represented by the consistency of results.
How close together your results are to one another regardless of the actual answer.

4. Taking Measurements
When taking measurements, you always measure one digit past what you can see.
This is called the “limit of precision.”

5. Examples:
How would you read the two graduated cylinders to the right?

6. REMEMBER UNITS!!! All numbers must be followed by a unit Liters = L
Milliliters = mL
Grams = g
Naked numbers have no meaning in science! Thus will receive no credit!

7. Accuracy Precision and Error

8. Accuracy Precision and Error

9. Scientific Notation Why Scientific Notation?
Scientists sometimes have to deal in very large or very small quantities, and scientific notation is a way of expressing that number.
Remember the Power of 10?
The Power of 10 Video

10. How To: Scientific Notation
Rules to Follow:
A positive exponent indicates that the number is greater than 1.
Example:
1,500, = 1.5 x 106
A negative exponent indicates that the number is less than 1.
= 2.5 x 10-5
Scientific Notation: Coefficient must be between 1 and 9
Standard Notation

11. What Is Scientific Notation?

21. Scientific Notation Practice!
Convert the following to scientific notation:
0.005
5,050
0.0008
1,000.
1,000,000
5 x 10-3
5.05 x 103
8 x 10-4
1.000 x 103
1 x 106
5 x 10^-3
5.05 x 10^3
8 x 10^-3
1 x 10^3
1 x 10^6

22. Do all numbers matter? How many mL are in this graduated cylinder?
Devices for measurement have a ‘limit of precision Not all numbers in a value are “significant.” Significant figures show the accuracy and precision of a measurement
How many
mL are in this graduated cylinder?
meniscus

23. Counting Significant Figures
4 Rules to Remember:
Non-zero Numbers are ALWAYS Significant
Sandwiched Zeros are ALWAYS Significant
Leading Zeros are NEVER Significant
Trailing Zeros are SOMETIMES Significant

24. NON-ZERO NUMBERS
Value Sig Figs
8 mm
42 lbs ?
678 mm ?
9 lbs ?

25. SANDWICHED ZEROS
Value Sig Figs 50.8 mm min lb ? m ?

26. LEADING ZEROS Value Sig Figs 0.008 mm 1 0.0156 oz 3 0.0042 lb ?
mL ?

27. TRAILING ZEROS
When a decimal is PRESENT: Trailing Zeros ARE Significant When a decimal is ABSENT: Trailing Zeros are NOT Signficiant Value Sig Figs 25,000 in yr 3 48,600 gal ? 25,005,000. g ?

28. Calculations Using Sig Figs
There are different rules for different mathematical operations: RULE: When multiplying and dividing the answer has the same number of significant digits as the value with the fewest.
Example:
23.0 cm x 432 cm x 19 cm =
The calculated answer is 188,784 cm3
The reported answer is rounded to two sig figs because 19 cm has the least amount in the problem.
The answer is 190,000 cm3

29. Calculations Using Sig Figs
There are different rules for different mathematical operations: RULE: When adding and subtracting the answer is limited not by the significant digits, but the limit of precision!
Example:
mL mL mL =
The calculated answer is mL
The reported answer is rounded to the tenths place because 46.0 mL has the lowest limit of precision.
The answer is mL

30. Subatomic Particles
The subatomic particles are distinguished by their charge, mass, and location.
Each has a different “job” to do with respect to its structure.
Particle
Location
Charge
Mass
(AMU)
“Job”
Proton
Nucleus +1
~1amu
Determines identity of element
Neutrons
Supplies proper mass to hold nucleus together
Electrons
Electron Cloud -1
~0 amu
Determines bonding and reactivity

31. What Is this AMU? normally measure using units of g or kg
Protons, neutrons, and electrons are all matter and part of the definition is that all matter has mass
normally measure using units of g or kg
These numbers are so very very small, that they are hard to grasp, so we came up with the Atomic Mass Unit
The AMU is simply a comparison unit to show that protons and neutrons have about the same mass, and that the electron is still very small in comparison.

32. Average Atomic Mass
The average mass of all the atoms that make up this element.

33. #Protons + #Neutrons = Mass Number
The mass number equals the total number of subatomic particles in the nucleus.
#Protons + #Neutrons = Mass Number

34. Atomic Number
The Atomic Number is found at the top of each periodic square. Each element has its own Atomic Number
- atomic number: the number of protons in the nucleus of an atom.
The Protons identify the atom. Change the atomic number and you have a new element!

35. Atomic Number and Mass Number, continued

36. You Try! 1. List the Atomic Number for the following elements!
-B -F -K
2. List the Mass Numbers for the following!
-Li -O -S

37. Bohr Diagrams How can we draw an atom on our paper?
Start with a circle in the center – this will represent the nucleus….What goes in the nucleus??
Protons and Neutrons!
Draw another circle around it – this will represent the first energy level for the electrons.

38. Bohr Diagram, cont. The 1st energy level can only hold 2 electrons
The 2nd can hold up to 8
The 3rd can hold a maximum of 18, but we won’t draw Bohr Diagrams for elements that big

39. Bohr Diagram vs. Current Atomic Theory
Bohr diagrams are a vast simplification of what we now understand the atom to be!
The atom actually has a bunch of different shapes of clouds within the energy levels which all overlap
But more on that later…

40. Let’s Build An Atom: 1. Hydrogen 2. Oxygen 2. Carbon
BELL WORK How many protons, neutrons and electrons are part of each stable atom? Does the proton count resemble the electron count? What happens if you add more electrons? What does the atom become? Which component (Proton, Neutron, Electron) is the single identifier of the atom? What is the number of protons called? What happens when you add too many neutrons? What do you get? What is it called?
Standards: H.C.2.A.1, H.C.2.A.2

41. Elements Atoms and Nuclear
Discovery of Atom & Atom Structure
Periodic table
Electrons
Isotopes and Nuclear Decay
Fission and Fusion

42. The Periodic Table Dmitri Mendeleev & Periodic Table (CCC) Periodicity
Groups (Columns) and Periods (Rows)
[TED TALKS PERIODIC TABLE VIDEOS]

43. Periodic Table of Elements
[Go to ptable.com]

44. 1s2 An example The “s” tells you the electron’s cloud shape.
In this case it’s a spherically shaped cloud. It is called the s sublevel.
1s2
The “2” simply tells you how many electrons are in this cloud.
The “1” tells you how far from the nucleus the electrons are.
In this case 2 electrons are creating the cloud.
In this case, its in the 1st energy level, which is closest level to the nucleus.
Note: Each orbital can only hold 2 electrons. They must have opposite spins.

45. Electron Configuration

46. Emission vs Absorption

47. Why we study electrons
Knowing the position of electrons helps us with many topics:
Why chemical reactions occur
Why some atoms are more stable than others
Why some elements react with only certain atoms
We want to know how many electrons an atom has, and where they’re located.

48. The Electron Cloud
Bohr’s model of the atom only showed us that electrons have different energy levels
Our current “quantum mechanical” model of the atom goes even further and talks about the different orbital shapes inside these energy levels.
Quantum numbers link
The world of quantum mechanics 1 2 3 4 5

49. 4 QUANTUM NUMBERS n l m s N - Principal Energy Level
Higher values of n mean more energy for the electron and the corresponding radius of the electron cloud or orbital is further away from the nucleus. Values of n start at 1 and go up by integer amounts. The higher the value of n, the closer the corresponding energy levels are to each other.

50. QUANTUM NUMBERS n l m s
L - Angular Quantum Number
In chemistry, there are names for each values of ℓ. The first value, ℓ = 0 called an s orbital. s orbitals are spherical, centered on the nucleus. The second, ℓ = 1 is called a p orbital. p orbitals are usually polar and form a teardrop petal shape with the point towards the nucleus. ℓ = 2 orbital is called a d orbital. These orbitals are similar to the p orbital shape, but with more 'petals' like a clover leaf. They can also have ring shapes around the base of the petals. The next orbital, ℓ=3 is called an f orbital. These orbitals tend to look similar to d orbitals, but with even more 'petals'.

51. QUANTUM NUMBERS n l m s M - Magnetic Quantum Number
The third quantum number is the magnetic quantum number, m. These numbers were first discovered in spectroscopy when the gaseous elements were exposed to a magnetic field. This relationship shows for every value of ℓ, a corresponding set of values of m ranging from -ℓ to ℓ is found. This number determines the orbital's orientation in space.
For example, p orbitals correspond to ℓ=1, can have m values of - 1,0,1. This would represent three different orientations in space for the twin petals of the p orbital shape. They are usually defined to be px, py, pz to represent the axes they align with.

52. QUANTUM NUMBERS n l m s S - Spin Quantum Number
The fourth quantum number is the spin quantum number, s. There are only two values for s, +½ and -½. These are also referred to as 'spin up' and 'spin down'. This number is used to explain behavior of individual electrons as if they were spinning in a clockwise or counterclockwise. The important part to orbitals is the fact that each value of m has two electrons and needed a way to distinguish them from one another..

53. QUANTUM NUMBERS

54. QUANTUM NUMBERS SUMMARY
Principal
Angular Momentum
Magnetic
Spin
Symbol N l Ml Ms
Values
N-1
-l to l
What it describes
How far the electron is f rom nucleus
Describes the shape or type of orbital
Describes the electron’s position in the orbital
Describes the spin
N=1, 2, 3, 4, 5, 6, 7
l=0, 1, 2, 3
s, p, d, f
Ml=
s 1 position
p 3 positions -1, 0, 1
D 5 positions -2, -1, 0, 1, 2
F 7 positions -3, -2, -2, 0, 1, 2, 3
+1/2 counter-clockwise
-1/2 clockwise
Maximum # e-
2N2
2e for each position
S=2, p=6, d=10, f=14
Maximum # sub levels
Eg. e level 2
l=0 or l=1 s, p
Pauli Exclusion Principle:
NO 2 ELECTRONS in the same atom have the same 4 quantum numbers

55. Orbital – region where electrons of a specific energy are likely to be.
-up to 2 electrons per orbital

56. Summary of Principal Energy Levels, Sublevels, Orbitals
Principal E Level N
Number of Sublevels =N
Types of Sublevels
Orbitals
=N2
# Electrons in Energy Level
=2N2
N=1 1
1s (1 orbital) 2
N=2
2s (1 orbital), 2p (3 orbitals) 4 8
N=3 3
3s (1 orbital)
3p (3 orbitals)
3d(5 orbitals) 9 18
N=4
4s (1 orbital)
4p (3 orbitals)
4d(5 orbitals)
4f (7 orbitals) 16 32
N=5
N=6
N=7

57. Orbital Diagrams
The order in which will fill a sublevel all comes down to energy!
Electrons fill the lowest energy sublevels first
The periodic table helps to show this order!

58. C. Periodic Patterns s p d (n-1) f (n-2) 1 2 3 4 5 6 7 6 7
© 1998 by Harcourt Brace & Company

59. Help a brother out….
When working out electron configuration, it’s important to understand orbital diagrams.
Here are 3 rules that will help us understand and interpret orbital diagrams.

60. A. General Rules Electrons fill the lowest energy sublevel available.
Aufbau Principle
Electrons fill the lowest energy sublevel available.
You cannot skip any sublevels

61. A. General Rules Pauli Exclusion Principle
Each orbital can hold only TWO electrons with opposite spins.
Represents an electron rotating clockwise
Represents an electron rotating counter - clockwise

62. A. General Rules WRONG RIGHT Hund’s Rule
Within a sublevel, place one e- per orbital with the same spin before pairing them.
“Empty Seat Rule”
WRONG
RIGHT

63. A Few more things….
Each arrow you draw represents an electron. (remember the three rules!)
Each orbital holds 2 electrons and has different shapes
The S level has only 1 orbital= total of 2 e-
The P level has 3 orbitals = total of 6 e-
The D level has 5 orbitals = total of 10 e-
The F level has 7 orbitals = total of 14 e-

64. Practice - Fill in the orbital diagram for Magnesium
*Use your PT to help you out!

65. Electron Configuration
An atom’s electron configuration…
Describes the location of electrons within the atom
Identifies the shape of the electron clouds
- regions where the electrons are held.
Uses numbers and letters to describe electrons’ location and which electron cloud it is part of.
Essentially, summarizes the orbital diagram!

66. 1s2 An example The “s” tells you the electron’s cloud shape.
In this case it’s a spherically shaped cloud. It is called the s sublevel. Sublevel = Shape
1s2
The “2” simply tells you how many electrons are in this cloud.
The “1” tells you how far from the nucleus the electrons can go.
In this case 2 electrons are creating the cloud.
In this case, its in the 1st energy level, which is the closest level to the nucleus.
Note: Every orbital can only hold 2 electrons. They must have opposite spins.

67. 1s2 2s2 2p4 O B. Notation 8e- 1s 2s 2p Electron Configuration
Orbital Diagram O
8e- 1s 2s 2p
Electron Configuration
1s2 2s2 2p4

68. C. Periodic Patterns Shorthand Configuration
This is a shorter way to do e- configurations
Steps…
1) Write the symbol of the element your doing
2) Find the noble gas (group 18) that comes before the element you are doing.
3) To begin, put that noble gas in [brackets]
4) Starting with that noble gas, finish the electron configuration

69. S 16e- 1s2 2s2 2p6 3s2 3p4 S 16e- [Ne] 3s2 3p4 B. Notation
Longhand Configuration
S 16e-
1s2
2s2
2p6
3s2
3p4
Core Electrons
Valence Electrons
Shorthand Configuration
S 16e- [Ne] 3s2 3p4

70. C. Shorthand practice
Example - Germanium
[Ar]
4s2
3d10
4p2

71. Practice these in shorthand
Li N Mg C

72. Draw this! This will help It is a way to remember
the order in which electrons
fill their orbitals.
When you reach the far
Left side, you reach a
“wall” and must go back to
The right hand side.