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