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Unit 3 1
Name: ______________________________ Period:_________
Unit 3: The Periodic Table and Atomic Theory
Day Page # Description IC/HW Due Date
Completed?
1 2-3 Periodic Table and Quantum Model Notes IC
1 4-5 Orbital Diagrams Notes IC
1 14 3-A: Orbital Diagrams Worksheet HW
2 VISION Electron Configuration Homework HW
2 6-8 Electron Configurations and Oxidation States IC
2 15-16 3-B:Electron Configurations and
Oxidation Sheets Worksheet HW
2 17 3-C Oxidation States Worksheet HW
3 9 Electron Configuration Exceptions and
Excited Electrons IC
3 10 Quantum Numbers Notes IC
3 18 3-D: Quantum Number Practice IC
3 19 3-E: Quantum Numbers Homework HW
4 11-12 Periodic Trends Notes IC
4 13 Periodic Table: What is the Trend? IC/HW
4 21 My First Periodic Table IC
4 20 3-F: Periodic Trends Worksheet HW
22-23 Unit 3 Review HW
X Unit 3 Test IC
Unit 3 2
Periodic Table Notes
John Newlands proposed an organization system based on increasing atomic mass in
1864.
He noticed that both the chemical and physical properties repeated every 8 elements
and called this the __________________________________.
In 1869 both Lothar Meyer and Dmitri Mendeleev showed a connection between atomic
mass and an element’s properties.
Mendeleev published first, and is given credit for this.
He also noticed a periodic pattern when elements were ordered by increasing
__________________________________.
By arranging elements in order of increasing atomic mass into columns, Mendeleev
created the first Periodic Table.
This table also predicted the existence and properties of undiscovered elements.
After many new elements were discovered, it appeared that a number of elements were
out of order based on their ________________________.
In 1913 Henry Mosley discovered that each element contains a unique number of
___________________________.
By rearranging the elements based on ____________________________, the
problems with the Periodic Table were corrected.
This new arrangement creates a periodic repetition of both physical and chemical
properties known as the _______________________________.
Periods are the __________________
Groups/Families are the ____________
Valence electrons across a period are in the same energy level
There are equal numbers of valence electrons in a group.
Valence electrons are ______________ energy electrons in the
___________ energy level of the atom (The outside)
Unit 3 3
Quantum Model Notes
Heisenberg's Uncertainty Principle- Can determine either the __________ ___________________________ of an electron, cannot determine both.
The Quantum Model- atomic orbitals are used to describe the possible position of an electron.
The location of an electron in an atom is described with 4 terms.
o Energy Level- Described by _____________. The higher the level, the more energy an
electron has to have in order to exist in that region.
o Sublevels- energy levels are divided into sublevels. The # of sublevels contained within an energy level is equal to the integer of the __________________________________.
o Orbitals- Each sublevel is subdivided into orbitals. Each orbital can hold ________ electrons.
o Spin- Electrons can be spinning clockwise (+) or counterclockwise (-) within the orbital.
Energy Level
Indicates relative sizes and energies of atomic orbitals. Whole numbers, ranging from _________________________.
The energy level is represented by the letter _________.
Sublevels
Number of sublevels present in each energy level is equal to the n.
Sublevels are represented by the letter __________.
In order of increasing energy:
Orbitals –ml
- s Sublevel- Only __________ orbital in this sublevel level.
- p Sublevel- _____________ orbitals present in this sublevel.
- d Sublevel- ______________ orbitals present in this sublevel.
- f Sublevel- _______________ orbitals present in this sublevel.
Unit 3 4
Energy Level Sublevels Present # of Orbitals Total # of Orbitals
in Energy Level
Total # of Electrons in Energy Level
1 s
2 s, p
3 s, p, d
4 s, p, d, f
Drawing Orbital Diagrams
An orbital diagram shows the _________________ of electrons in an atom.
The electrons are arranged in energy levels, then sublevels, then orbitals. Each orbital can only contain 2 electrons.
Three rules must be followed when making an orbital diagram. o Aufbau Principle- An electron will occupy the ___________________ energy
orbital that can receive it. To determine which orbital will have the lowest energy, look to the periodic
table. o Hund’s Rule- Orbitals of equal energy must each contain __________
_____________________ before electrons begin pairing. o Pauli Exclusion Principle- If two electrons are to occupy the same orbital, they
must be spinning _______________________________.
Energy Levels (n) determined by the ROWS
Sub Levels (s,p,d,f)- determined by the sections
Orbitals - determined by the # of columns per sublevel
Unit 3 5
Orbital Diagrams
S
As
Mn
N
Sc
Unit 3 6
Electron Configurations and Oxidation States Electron configurations are shorthand for orbital diagrams. The electrons are not shown in
specific orbitals nor are they shown with their specific spins.
Draw the orbital diagram of oxygen:
The electron configuration should be:
Manganese
Arsenic
Promethium
The Noble Gas shortcut can be used to represent the electron configuration for atoms with many electrons. Noble gases have a full s and p and therefore can be used to represent the inner shell electrons of larger atoms.
For example: Write the electron configuration for Lead.
Write the electron configuration for Xenon.
Substitution can be used:
Manganese
Arsenic
Promethium
Unit 3 7
Valence electrons, or outer shell electrons, can be designated by the s and p sublevels in the highest energy levels
Write the noble gas shortcut for Bromine
Br = [Ar]4s23d104p5 Write only the s and p to represent the valence level.
Br = 4s24p5 This is the Valence Configuration. Bromine has 7 valence electrons.
Silicon
Uranium
Lead Octet Rule and Oxidation States
The octet rule states the electrons need ______________ valence electrons in order to achieve maximum stability. In order to do this, elements will gain, lose or share electrons.
Write the Valence configuration for oxygen O = 2s22p4- 6 valence electrons
Oxygen will gain 2 electrons to achieve maximum stability O = 2s22p6- 8 valence electrons
o Now, oxygen has 2 more electrons than protons and the resulting charge of the atom will be -2
o The symbol of the _____________ formed is now O-2.
Elements want to be like the Noble Gas family, so they will gain or lose electrons to get the same configuration as a noble gas.
When an element gains or losses an electron, it is called an ___________.
An ion with a positive charge is a ____________________.
An ion with a negative charge is an ___________________.
Unit 3 8
Oxidation States:
Element Total # of electrons
Valance Configuration
Gain or
Lose e-
How many?
New Valence Configuration
Total # of e-
Cs
Cl
Rb
Ca
Na
Element or Ion
Symbol Number
of Protons
Number of Electrons
Number of
Neutrons
Mass Number
Valence Configuration
Calcium (+2)
40
Aluminum (+3)
26
Barium (0)
139
Sulfur (-2)
34
Potassium (+1)
39
Unit 3 9
Electron Exceptions: - Some electrons configurations for transition metals have exceptions to the pattern due to potential
lower energy configurations
- Inner transition metals also have exceptions, but pattern is harder to recognize.
- Copper
o Configuration following all rules:
o Actual configuration:
- Chromium o Configuration following all rules:
o Actual configuration:
Electrons in Excited States Electrons in their lowest energy electron configuration in an atom are said to be in the ground state. When electrons are in the ground state they usually obey the three rules governing electron configurations (don’t forget the exceptions we covered earlier).
- Aufbau Principle – electrons enter orbitals of lowest energy first - Hund’s Rule – electrons enter orbitals of equal energy one at a time with parallel spins - Pauli Exclusion Principle – no two electrons in the same atom can have the same four quantum
numbers When electrons in the ground state absorb energy they change their configuration to an excited state. This means that either the Aufbau principle or Hund’s rule is violated. However, the Pauli Exclusion principle can NEVER be violated. Any violation of the Pauli Exclusion Principle is impossible. For each of the following, decide if the electrons are in a ground state, excited state, or impossible
state. If a violation
Unit 3 10
Quantum Number Notes The quantum mechanical model uses three quantum numbers, n, l and ml to describe an
orbital in an atom. A fourth quantum number, ms, describes an individual electron in an orbital.
n = _______________________________________ o This describes the energy level and can be described as an integer from 1-7. The
larger the number, the larger the orbital. As the numbers increase, the electron will have greater energy and will be less tightly bound by the nucleus.
l = _________________________________________ o This describes the shape of the orbital level and can be described as an integer from
0 to n-1. o 0 is used to describe s orbitals; 1 is used to describe p orbitals. o 2 is used to describe d orbitals; 3 is used to describe f orbitals.
ml = ______________________________________ o This describes the orientation of the orbital in space and can be described as an
integer from - l to l o s sublevels have one orbital, therefore possible values of ml include 0 only. o p sublevels have three orbitals, therefore possible values of ml include -1, 0, 1. o d sublevels have five orbitals, therefore possible values of ml include -2, -1, 0, 1, 2. o What about f?
ms= __________________________________ o This describes the spin of the electron in the orbital. o The possible values for ms are +½ and -½. o The positive spin reflects the first electron in a specific orbital and negative spin
reflects the second electron in an orbital.
Examples: o Give the four quantum numbers for the 8th electron in Argon:
o What is the maximum number of electrons that can have the following quantum numbers: n = 2 and ms = -½
o Which of the following quantum numbers would NOT be allowed in an atom n = 2, l = 2 and ml = -1 n = 4, l = 2 and ml = -1 n = 3, l = 1 and ml = 0 n = 5, l = 0 and ml = 1
Unit 3 11
Periodic Trends- Notes
Shielding: As you go down the periodic table, the number of shells increases which results in greater electron-electron repulsion.
o The more shells there are, the further from the nucleus the valence
electrons are.
o Therefore, more shielding means the electrons are ______________
attracted to the nucleus of the atom.
Atomic Radius is defined as half the distance between adjacent nuclei of
the same element.
o As you move DOWN a group an entire energy level is added with
each new row, therefore the atomic radius ___________________.
o As you move LEFT-TO-RIGHT across a period, a proton is added, so
the nucleus more strongly attracts the electrons of a atom, and
atomic radius ___________________________.
Ionic Radius is defined as half the distance between adjacent nuclei of the
same ion.
o For _________________ an electron was lost and therefore the ionic
radius is smaller than the atomic radius.
o For ___________________ an electron was gained and therefore the
ionic radius is larger than the atomic radius.
o As you move down a group an entire energy level is added, therefore
the ionic radius increases.
o As you move left-to-right across a period, a proton is added, so the
nucleus more strongly attracts the electrons of a atom, and ionic
radius _____________________.
However! This occurs in 2 sections. The cations form the first group, and the anions form the second group.
Unit 3 12
Isoelectronic Ions: Ions of different elements that contain the same number of electrons.
Ionization energy is defined as the energy required to _______________ the first electron from an atom.
o As you move down a group atomic size increases, allowing electrons to be further from the nucleus, therefore the ionization energy ______________________.
o As you move left-to-right across a period, the nuclear charge increases, making it harder to remove an electron, thus the ionization energy ______________________.
Electronegativity is defined as the relative ability of an atom to attract electrons in a ____________________________________.
o As you move down a group atomic size increases, causing available electrons to be further from the nucleus, therefore the electronegativity ______________________.
o As you move left-to-right across a period, the nuclear charge increases, making it easier to gain an electron, thus the electronegativity _________________________.
Reactivity is defined as the ability for an atom to react/combine with other atoms. o With reactivity we must look at the metals and non-metals as two separate groups.
Metal Reactivity- metals want to lose electrons and become cations o As you move down a group atomic size increases, causing valence electrons to be
further from the nucleus, therefore these electron are more easily lost and reactivity ____________________.
o As you move left-to-right across a period, the nuclear charge increases, making it harder to lose electrons, thus the reactivity _______________.
Non-metal Reactivity- non-metals want to gain electrons and become anions o As you move down a group atomic size increases, making it more difficult to attract
electrons, therefore reactivity ___________________. o As you move left-to-right across a period, the nuclear charge increases, making it
easier to attract electrons, thus the reactivity _____________.
Unit 3 13
Periodic Table : What is the Trend?
Definition Trend
Atomic Size (Atomic Radius)
Electronegativity
Ionization Energy
Metal
Non-Metal
Shielding
Unit 3 14
3- A: Orbital Diagram Worksheet Give the COMPLETE orbital diagram for the following elements:
1. Ar
2. P
3. Fe
4. Ca
5. Br
6. Mn
7. U
Unit 3 15
3- B: Electron Configuration Practice Give the COMPLETE electron configuration for the following elements:
1. Ar
2. P
3. Fe
4. Ca
5. Br
6. Mn
7. U
Unit 3 16
3-C: Electron Configuration and Oxidation States Worksheet
Give the noble gas shortcut configuration for the following elements: 1. Pb
2. Eu
3. Sn
4. As
Give ONLY the valence (outer shell) configuration for the following
elements:
1. Ba
2. Po
3. S
4. F
Unit 3 17
3-D: Oxidation States Worksheet Complete the following tables:
Element
Total # of
electrons
Valence Configuratio
n
Gain or
Lose e-
How Many
?
Ion Symb
ol
New Valence
Configuration
Total # of e-
Cesium 55 6s1 Lose 1 Cs+ 5s25p6 54
Oxygen
Fluorine
Barium
Aluminum
Element or Ion
Symbol Number
of Protons
Number of
Electrons
Number of
Neutrons
Mass Number
Valence Configuration
Calcium (+2)
Ca2+ 20 18 20 40 3s23p6
Chlorine (-1)
37
Magnesium (0)
23
Oxygen (-2)
16
Lithium (+)
8
Unit 3 18
3-E: Quantum Number Practice: Write the four quantum numbers which describe the location of the highest energy electron of the following:
1. N #7 ____________________________________________________
2. Ni #28___________________________________________________
3. Xe #54___________________________________________________
4. Re #75___________________________________________________
5. Pu #94___________________________________________________
6. Br #35___________________________________________________
Give the four quantum numbers which describe the location of each of the following:
7. The 4th electron in carbon_____________________________________________
8. The 25th electron in Hg_______________________________________________
9. The 57th electron in Ho_______________________________________________
10. The 49th electron in Xe______________________________________________
Identify the element whose highest energy electron would have the following four quantum numbers:
11. 3, 1, -1, +1/2___________________________________________________
12. 4, 2, +1, +1/2___________________________________________________
13. 6, 1, 0, -1/2____________________________________________________
14. 4, 3, +3, -1/2___________________________________________________
15. 2, 1, +1, -1/2___________________________________________________
Which of the following represents a permissible set of quantum numbers? (answer “yes” if permissible and “no” if no permissible)
16. 2, 2, +1, -1/2___________________________________________________
17. 5, 1, 0, +1/2____________________________________________________
18. 6, 3,-2, +1/2____________________________________________________
19. 7, 0, 0, -1/2____________________________________________________
20. 4, 1, 8, +1/2____________________________________________________
Unit 3 19
3-F: Quantum Numbers Worksheet Rules for assigning quantum numbers: n: can be 1, 2, 3, 4, … Any positive whole number
l: can be 0, 1, 2, … (n-1) Any positive whole number, up to (n-1)
ml : can be (-l), (-l +1), (-l +2), … 0, 1, … (l -1), l Any integer, from – l to l . ms: can be +½ or - ½
1) If n=2, what possible values does l have?
2) When l is 3, how many possible values of ml are there?
3) What are the quantum numbers for the 17th electron of Argon?
4) What are the quantum numbers for the 20th electron of Chromium?
5) What are the quantum numbers for the 47th electron of Iodine?
6) Give the quantum numbers for ALL of the electrons in Nitrogen.
7) Determine what the error is in the following quantum numbers, and explain. (HINT: Some do
not have an error)
a. (1,1,0, - ½ )
b. (3,2,-1,+ ½ )
c. (2, -1, 1, - ½ )
d. (3, 3, 2 , + ½ )
e. (2, 0, 1, - ½ )
f. (-1, 0, 0, + ½ )
g. (2, 1, 0, + ½ )
Unit 3 20
3-G: Periodic Trends Worksheet
1. Explain why a magnesium atom is smaller than both sodium AND calcium. 2. Would you expect a Cl- ion to be larger or smaller than a Mg2+ ion? Explain
3. Explain why the sulfide ions (S2-) is larger than a chloride ion (Cl-). 4. Compare the ionization energy of sodium to that of potassium and EXPLAIN.
5. Explain the difference in ionization energy between lithium and beryllium. 6. Order the following ions from largest to smallest: Ca2+, S2-, K+, Cl-. Explain your order.
7. Rank the following atoms/ions in each group in order of decreasing radii and explain
your ranking for each.
a. I, I- b. K, K+ c. Al, Al3+
8. Which element would have the greatest electronegativity: B or O? Explain. Hint: a positive electron affinity means that the element wants to form a negative charge.
Unit 3 21
My First Periodic Table
Use the following clues to place these “alien” elements on the periodic table below. Assume they follow the same trends as our periodic table. Clues: 1. The following are groups of elements: E & F P, Q, & R K & L S & T U & V X, Y & Z M & N A, B, C & D 2. A is the lightest known element. 3. D is the first element in the third period. 4. S has nine protons. 5. P is inert and has the highest ionization energy in its group. 6. T has one less electron than Q. 7. R has eight more electrons than P. 8. K has one more proton than M and one less than E. 9. X has the lowest electronegativity in its group. 10. Y has the smallest atomic radius of its group. 11. Z has the electron structure s2. 12. U has 3 more protons than P. 13. M has an atomic radius smaller than N. 14. E has 6 valence electrons. 15. V has a smaller electronegativity than U. 16. B has the lowest ionization energy of all the elements. Periodic Table:
Unit 3 22
Unit 3 Review: Give the Orbital Diagram for the following elements:
1. Chromium
2. Nitrogen Give the COMPLETE electron configuration for the following elements:
3. Argon
4. Phosphorous Give the Noble Gas electron configuration for the following elements:
5. Plutonium
6. Mercury
7. Complete the table.
Element Total # of electrons
Valence Configuration
Gain or Lose e-
How Many?
Ion Symbol
New Valence Configuration
Total # of e-
Phosphorous
Chlorine
Cesium
Lithium
Give the 4 quantum numbers for the last electron of the following elements:
8. Phosphorous
9. Manganese
10. Silver
11. Promethium
12. Iodine
Unit 3 23
Determine if the following sets of quantum numbers would be allowed in an atom. If not, explain why and if so, identify the corresponding atom.
13. n = 2, l = 1, ml = 0, ms = +2
1
14. n = 4, l = 0, ml = 2, ms = -2
1
15. n = 1, l = 1, ml = 0, ms = +2
1
Give the element with the LARGER radius, ionization energy, electronegativity and reactivity.
ELEMENTS ATOMIC RADIUS IONIZATION ENERGY
ELECTRONEGATIVITY
Sodium and Aluminum
Chlorine and Iodine
Oxygen and Fluorine
Magnesium and Calcium
Circle the element / ion with the larger radius.
16. Mg or Mg 2+ 18. S or S2- 20. N3- or F-
17. Sr2+ or Br- 19. Cl- or Mg2+ 21. B or F
For each of the following families, give their relative reactivity, the number of valence electrons, and at least one additional piece of information (such as how they are found in nature or what other group the generally react with).
22. Alkaline Earth Metals
23. Alkali Metals
24. Halogens
25. Noble Gases