Chemistry Annotation Guide If you are NOT using the following annotation, put in your key to the left of each item. Mr. T s Items to be annotated

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Chemistry Annotation Guide

If you are NOT using the following annotation, put in your key to the left of each item.

Your Key

Mr. T’s

Key Items to be annotated

Mr. T’s √ or ×

headings and subheadings.

key content vocabulary

Write DEF next to definitions

(sometimes #2 and #3 will be the same and in that case I expect to see a box AND DEF) Underline important ideas (include captions for visuals).

Write “Eq” in the margin next to an important equation.

Underline key steps to solving an example problem, or write those steps in the margin.

Write any questions in the margin.

Write a learning objective next to each subheading.

Using this handout effectively

The studying and annotations and other homework assignments are designed to require a maximum of 45 minutes to an hour of your time. This time limit, however, does not include things like test make-up points or extra credit activities. Students are expected to plan their time. If students wait until the last minute to complete assignments that are assigned multiple days in advance, then the time required to complete the assignment may require more than the 1 hour limit.

This document is a “work in progress.” I have made an attempt to put in as much about the knowledge you need for this unit of study as I can. This document, however, DOES NOT contain everything that you need to know to make 100% on the Unit 4 test. You must also rely on the knowledge that you should have acquired in previous units of study, classroom explanations, your notes, the textbook sections that were assigned for study, and—in some cases—creative thinking and problem-solving skills.

Atomic Structure

This document is a “work in progress.” I have made an attempt to put in as much about the knowledge you need for this unit of study as I can. This document, however, DOES NOT contain everything that you need to know to make 100% on the Unit 2 test. You must also rely on the knowledge that you should have acquired in previous units of study, classroom explanations, your notes, the textbook sections that were assigned for study, and—in some cases—creative thinking and problem-solving skills.

Unit 2: Chemistry I Honors The Standards in This Unit

According to the South Carolina Science Standards, standard C-2, students will demonstrate an

understanding of atomic structure and nuclear processes. Most of this unit will deal with these concepts.

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Key questions for this next section:

1) What are the representative elements? 2) What are groups or families of elements?

3) What elements make up the alkali metals and how are they different from the group IA elements?

4) What elements make up the alkaline-earth metals? 5) What elements make up the halogens?

6) What elements make up the noble gases? 7) What is a quantum of energy?

8) What do quantum numbers tell you about electrons? 9) What are ground state electrons?

10) What are primary electron energy levels?

11) What would a primary electron energy level look like if we could see one? 12) What is a quantum number?

13) What is the 1st quantum number? What does the 1st quantum number tell you about an electron?

14) What symbol is used to designate the 1st quantum number?

15) What is the 2nd quantum number and what symbol is used to designate the 2nd quantum number? 16) What does the 2nd quantum number tell you about an electron?

17) How is the 2nd quantum number represented in orbital notation and electron configurations? 18) What is the shape of an s orbital?

19) What is the shape of a p orbital?

20) How are the s and p orbitals arranged (how do they fit together)?

Indicator C-2.1

According to Indicator C-2.1 in the South Carolina Science Standards students should be able to illustrate electron configurations by using orbital notation for representative elements.

Students should be able to:

• Understand that the representative elements are those elements within the first two groups (groups IA and IIA on the far left of the periodic table in your class references) and the last six groups on the right of the Periodic Table (groups IIIA through VIIIA)..

Group or family of elements: A group of elements are those that in the same column. They are also called a family if they are in the same column.

Alkali metals: All of the elements in group or family IA, except for hydrogen, are also called the alkali metals.

Alkaline-earth metals: All of the elements in group or family IIA are also called the alkaline-earth metals.

Halogens: All of the elements in group or family VIIA are also called the halogens. Noble gases: All of the elements in group or family VIIIA are also called the noble gases.

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Figure 1. Out line of the periodic table showing the representative elements in darker outline. Quantum numbers – Fully understanding quantum numbers requires an understanding that could take up an entire course… or more! In this class you will only be expected to understand just enough to see how quantum numbers relate to the location of electrons in the electron cloud of atoms.

• Understand the first two quantum numbers and use them to describe the location of electrons in representative elements in the ground state

Definition:

A quantum of energy is the minimum quantity of energy that can be lost or gained by an atom. Atoms do not gain or loose energy in a continual way, like the temperature measured in a

thermometer going up or down appear to do so smoothly. The electrons in an atom gain and lose energy in distinct or discrete steps called quanta. Quantum numbers, however, do not describe energy. Quantum numbers describe the location or properties of electrons. Those locations are all that we are concerned about in this course but those locations have an effect on the energy.

Explanation:

Electrons have potential energy due to their position within the electron cloud and that potential energy can be thought of like a shelf on which the electron sits. The electron absorbs potential energy and you can think this potential energy as like being put on this make-believe shelf and it looses energy when it falls off, but the electrons energy is never going to be somewhere in midair between the shelf and the floor. Electrons are assigned quantum numbers to designate the region in which they can be found in the structure of the electron cloud and they have that region because of the amount of potential energy they have. A set of 4 quantum numbers for an electron is unique

Periodic Table of the Elements Outline

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Figure 2. An illustration of how an atom’s electron cloud might look if sliced in half to reveal the primary electron energy levels. Fifth primary energy level

Fourth primary energy level

First primary energy level Nucleus

Third primary energy level Second primary energy level

to that electron. This is different from orbital notation and electron configurations which are designed to show the energy structure of the entire electron cloud of an atom.

Ground state electrons are those that have not been excited. Electrons exist in many places or regions within the electron cloud. There are primary energy levels, sublevels within those primary levels, and orbitals within those sublevels.

Note that each of these primary electron energy levels illustrated here correspond to rows of the periodic table. The term corresponds to does NOT mean that they are the same thing. In this context the term means that there is a relationship between primary energy levels and the elements in each row—or period—of the periodic table.

The primary levels can be thought of as concentric spheres, with each higher level encompassing the lower levels… like balls one inside of another. Each higher level or region, then, has a spherical shape but which does not include the lower levels. For every primary energy level after the 1st one, it’s as if you have a solid sphere with an inner sphere cut out of it. In other words, after the 1st primary energy level, each higher level is kind of like a very thick-walled, hollow ball. The sublevel regions within those concentric primary level spheres have different shapes and arrangements depending on the sublevel. s sublevels are spherical, while p sublevels look like two mushroom heads, and still others are even more complex. When electrons are excited they jump from one energy level to another, the sublevel shapes can become even more complex still. In this unit students will only be expected to know about the spherical (s orbitals) and mushroom head shapes (p orbitals) in the unexcited state. Chemists and physicists call that unexcited state the “ground state.”

o 1st quantum number

 Understand the aspect of electron location described by the first quantum number. (Energy level; sometimes called a shell)

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The first quantum number is a term used to describe the primary energy level in which a ground state electron is usually found.

 Understand that the 1st

quantum number is designated by numbers 1 through 7 and understand the meaning of each of those numbers in reference to the location of the electron.

Explanation:

The principal quantum number for ground state electrons can be primary energy level 1 (the spherical region or level closest to the nucleus) to level 7 (a spherical region whose outer edge is further from the nucleus than all the other 6 levels). There are 7 rows on the periodic table and the rows on the periodic table correspond to those levels. Elements in each row have electrons in the primary energy level corresponding to that row.

Additionally, elements in each row also have electrons in each of the lower primary energy levels. The symbol for a principal quantum number is the letter n. If n = 2, then the

electron in is the 2nd primary energy level.

These same numbers (1 through 7) will also be used in electron configurations and orbital notation.

Note: The farther the primary energy level extends from the nucleus, the greater the potential energy of the electrons in that level.

o 2nd quantum number

 This means that students should understand the type of orbital, its shape and location within the primary energy level.

Explanation:

The orbital quantum number tells you in which sublevel the electron can be found. Unless the sublevel is s, it won’t tell you which orbital the electron is in. The only reason the 2nd Understand the aspect of electron location described by the 2nd quantum number.

quantum number can tell you which orbital the electron is in is if it’s an s orbital is because s sublevels only have one orbital.

s sublevels have a single a orbital with a spherical shape. There is only one s orbital within each primary energy level. The size and shape of an s orbital is the same as that of the primary energy level. So, the s orbital in any sublevel is the same size and shape as the primary energy level in which the s orbital is located. Go back and look at figure 2. p sublevels have 3 p orbitals. p orbitals have the shape of a double mushroom head with each head separate from the other . There are 3 p orbitals within each primary energy level.

Students will not need to know the shape of the d and f orbitals.

The 2nd quantum number tells you in which set of orbitals the electron is found, but does not tell you which specific orbital within the set the electron is found. So, the orbital quantum number might tell you that the electron is in the p sublevel but it won’t tell you in which of the 3 p orbitals the electron can be found.

 Understand that—in orbital notation and electron configurations—the 2nd

quantum number is

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Explanation:

We’ll learn how to use the designations s, p, d, and f in electron configurations and orbital notation later in this unit. For now, we’ll simply relate the letters to the orbital quantum number.

The orbital quantum number might be better understood as a sublevel quantum number. Sublevel s always has a quantum number of 0, sublevel p always has a quantum number of 1, sublevel d always has a quantum number of 2, and sublevel f always has a quantum number of 3.

Note: Within the same primary level, electrons in sublevel p have more potential energy than electrons in sublevel s; electrons in sublevel d have more potential energy than electrons in either sublevel s or p; and electrons

in sublevel f have more potential energy than electrons in sublevels s, p, or d. So, electrons in sublevel 5f have more potential energy than electrons in 5d, 5d electrons have more potential energy than electrons in 5p, and 5p electrons have more potential energy than electrons in 5s.

Students will need to know the shape of the s orbital and the fact that it is basically the same as the shape of the primary energy level. Students will also need to know the shape and orientation of the p orbitals. p sublevels have the shape of a double mushroom head with the top of the mushroom head facing away from the nucleus and the open end of the mushroom head facing the nucleus. Each of these mushroom head shapes is called a lobe.

Figure 3. A single p orbital is illustrated here. Note the 2 lobes on such a p orbital and the relationship (orientation) of the lobes to the nucleus of the atom.

Open end of lobe of p orbital Lobe 1 of

a p orbital

Lobe 2 of a p orbital

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Figure 4. All 3 p orbitals in a p sublevel are illustrated here. Note that each p orbital has 2 lobes oriented so that the p orbitals are as far apart as possible and surround the nucleus of the atom. The 2nd number (or orbital quantum number) for the p sublevel is 1.

Figure 5. All 3 p orbitals in a p sublevel are illustrated here surrounded by a s orbital. All of the orbitals are in the same primary energy level but the p orbitals have a higher potential energy than does the s orbital.

There are 3 p orbitals in every primary level beginning with level 2 (level 1 does not have a p sublevel). The 3 p orbitals are oriented (or lined up) so that they are as far away from each other as they can get. If a line were drawn through the center of each mushroom head (or lobe) and through the nucleus, each of the 3 sets of mushroom heads representing the 3

2 lobes of one of the 3 p orbitals. We can designate this orbital as being ml = 1. A line that

would go through the middle of the 2 lobes and

through the nucleus would be considered the

x axis.

1 lobe of one of the 3 p orbitals. The opposite

lobe cannot be seen because it is hidden behind the orbitals. We can designate this orbital

as being ml = 2. This

would be considered as being on the y axis.

2 lobes of one of the 3 p orbitals. We can designate this orbital as being ml = 0.

A line though the middle of the 2 lobes and the nucleus would be considered as being

on the z axis.

An s orbital sounding 3 p orbitals. We can designate this orbital as

ml = 0. The 2

nd

quantum number (or orbital quantum number) for the s

sublevel is 0. Lower primary energy

levels would be located within the orbitals

pictured here. Remember that these

regions are not solid objects, but clouds of

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orbitals would be oriented along an x, y, or z axis. The 3 axes represent the 3 space dimensions (up-and-down, side-to-side, and front-to-back –or- height, width, and depth). Note: Very often textbooks will represent p orbitals as a “dumbbell” or “double balloon” shape because this idea is easier for students to grasp.

Figure 6. An illustration of a p orbital (on the z axis) illustrated here as a “dumbbell” or “double balloon” shape. Each “balloon” is a lobe and each orbital has 2 lobes.

Figure 7. An illustration of 3 p orbitals illustrated here as “double balloons” shape. Each “balloon” is a lobe and each orbital has 2 lobes.

Review Questions:

21) Where do we find the representative elements? 22) What is a quantum of energy? What is a quanta?

23) What does the quantum number assigned to an electron tell you about that electron? 24) What is a ground state electron?

25) What is a “primary energy level” in the electron cloud of an atom?

26) Describe the shape of these primary energy levels? Be specific, especially about the higher levels. 27) What is the shape of an s sublevel? What is the shape of a p sublevel?

28) What does the term principal quantum number mean? What part of the periodic table does the principal

quantum number relate to?

29) What numbers represent the principal quantum number?

30) What does the term 2nd quantum number mean? What part of the periodic table does the 2nd quantum number relate to?

31) What numbers represent the 2nd quantum number?

32) What letters relate the orbital notation and electron configuration to the 2nd quantum number? 33) What shape do s and p orbitals have?

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Chemistry I CP students are not expected to know the 3rd or 4th quantum numbers. Key questions for this next section:

34) What is the 3rd quantum number and what symbol is used to

designate the 3rd quantum number?

35) What does the 3rd quantum number tell you about an electron?

36) How is the 3rd quantum number represented in orbital notation and electron configurations?

37) What is the 4th quantum number and what symbol is used to

designate the 4th quantum number?

38) What does the 4th quantum number tell you about an electron?

39) How is the 4th quantum number represented in orbital notation and electron configurations?

Chemistry I CP students are NOT expected to know the 3rd or 4th quantum numbers. o 3rd quantum number

 Understand what aspect of electron location this describes.

Explanation:

The 3rd quantum number indicates in which orbital an electron can be found.

The s sublevel only has one spherical orbital. The p sublevel has 3 double-mushroom-head (or double balloon) shaped orbitals. Students are NOT expected to know the shapes of the d or f orbitals but…

Chemistry I CP and Chemistry I Honors students ARE expected to know that a d

sublevel has 5 d orbitals and

Chemistry I Honors students ARE expected to know that an f sublevel has 7 f

orbitals.

Chemistry I Honors students are expected to know…

The numbers for the 3rd quantum numbers range from:

 0 for an s orbital (Note: s sublevels only have one orbital)

 -1, 0, and 1 for the three p orbitals (Note: p sublevels have 3 orbitals, so there are 3 numbers; one for each orbital)

 -2, -1, 0, 1, and 2 for the d orbitals (Note: d sublevels have 5 orbitals, so there are 5 numbers; one for each orbital)

 -3, -2, -1, 0, 1, 2, and 3 for the f orbitals. (Note: f sublevels have 7 orbitals, so there are 7 numbers; one for each orbital)

Examples:

s sublevels only have one orbital—an s orbital—so the 3rd quantum number for either of the 2 electrons in an s orbital is 0.

p sublevels have 3 p orbitals. Either of the 2 possible electrons in any of the p orbitals might have an orbital number of -1, while either of the 2 possible electrons in another of

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the p orbital might have an orbital number of 0, and the last 2 possible electrons in the last p orbital would have an orbital number of 1.

All that students need to know about 3rd quantum number for d orbitals is that there are 5 d orbitals in a sublevel starting with the 3rd primary energy level and that the numbers range from -2 to 2.

All that students need to know about 3rd quantum number for f orbitals is that there are 7 f orbitals in a sublevel starting with the 4th primary energy level and that the numbers range from -3 to 3.

Chemistry I Honors students are expected to…  Understand that 3rd

quantum number is designated by one of 7 numbers and understand what each of those numbers mean in reference to the location of the electron.

Explanation:

An electron in an s sublevel only has one a 3rd quantum number of 0 (so, ml = 0).

An electron in a p sublevel can have one of 3 3rd quantum numbers: -1, 0, or 1. (so, ml = -1,

ml = 0, or ml = 1).

In this class students will only be expected to know the shapes of orbitals in the s and p sublevels in any of the 7 primary energy levels. See figures 2 through 6 above. If a

student is told that an electron has a set of quantum numbers in which n = 3, l = 1, and ml =

-1, they should know that the electron is in the 3rd primary energy level, in a p-sublevel, and in the p orbital designated as -1. They should know that this corresponds to the electron in the orbital notation that has been highlighted below:

Students should also know that there is a direct correspondence between the 1st 2 quantum numbers and an electron configuration. n = 3 and l = 1 should tell them that the electron has is found in the 3p sublevel as shaded below:

1s2 2s2 2p6 3s2 3p6

There is, however, NO direct correspondence between the 3rd quantum number (ml = -1)

and an electron configuration. In the configuration above all that you can determine about the 3rd quantum number of electron n = 3, l = 1, ml = -1, and ms = -½ is that it is one of the

6 electrons in the 3p sublevel. That superscript 6 is shaded below: 1s2 2s2 2p6 3s2 3p6.

Chemistry I Honors students are expected to know and explain details about the… o 4th quantum number

 Understand what aspect of electron location this describes.

Explanation:

An electron can have one of 2 spins: ms = +½ (which is represented in an orbital notation as ↑) and ms = -½ (which is represented in an orbital notation as ↓).

There is no direct correspondence between these 4th quantum numbers and the electrons represented in an electron configuration.

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 Understand that it is sometimes designated by a positive (+) or a negative (-) instead of +½ and -½.

 Understand that two electrons occupying the same orbital must have opposite spins Explanation:

In physics students are taught that particles with the same charge are pushed apart by opposing electromagnetic fields. So why can 2 electrons with the same charge occupy the same orbital? The explanation that scientists have come up with is that the electrons have opposite spins that cancel their opposing electromagnetic fields (also called magnetic moments). There may be no real electron spin. It’s simply a term scientists use to describe a behavior that they don’t fully understand. In other words, there is evidence that 2

electrons occupy the same orbital but we cannot really say that we understand why it happens. We just know that it does happen, and we call this property “spin” because it seems similar to other electromagnetic properties that are caused by an actual spin.

There is NO direct correspondence between the 4th quantum number (ms = −½ or +½) and

an electron configuration. The electron configuration just doesn’t tell us anything about electron spin.

There IS, however, a pretty direct correspondence between the 4th quantum number and direction of arrows in orbital notation. An arrow pointing up is pretty much the same thing as ms = +½ and an arrow pointing down is pretty much the same thing as ms = −½ . So,

for example, the shaded arrow in the illustration below has an ms quantum number of +½.

So, for another example, the shaded arrow in the illustration below has an ms quantum

number of −½.

Chemistry I CP and Chemistry I Honors students are expected to… o Understand that no two electrons in an atom can have

same set of quantum numbers (this is called the “Pauli exclusion principal”).

o Understand how many of each type of orbital are possible in each of the 1st 5 primary energy levels.

Explanation:

Obtain a copy of the periodic table and look at the rows of the table as you read the following explanation. Each level corresponds to a row or period on the periodic table.

Row 1 has only 1 sublevel and one type of orbital: an s sublevel and an s orbital.

Row 2 has 2 sublevels and 2 types of orbitals on top of the level 1 sublevels and orbitals: an s sublevel with 1 s orbital and a p sublevel with 3 p orbitals.

Row 3 has 2 sublevels and 2 types of orbitals on top of the level 1 and 2 sublevels and orbitals: an s sublevel with 1 s orbital and a p sublevel with 3 p orbitals.

↑↓ 1s ↑↓ 2s ↑↓ ↑↓ ↑↓ 2p ↑↓ 3s ↑↓ ↑↓ ↑↓ 3p ↑↓ 1s ↑↓ 2s ↑↓ ↑↓ ↑↓ 2p ↑↓ 3s ↑↓ ↑↓ ↑↓ 3p

A whimsical application of the Pauli

exclusion principal is the school bus rule. Read about the school bus rule

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Row 4 has 3 sublevels and 3 types of orbitals on top of the lower sublevels and orbitals: an s sublevel with 1 s orbital and a p sublevel with 3 p orbitals. Additionally, beginning with the 3rd element in the 4th row of the periodic table (scandium, atomic number 21) a 3rd sublevel is introduced into the atomic electron cloud structure. This is the d sublevel and it falls into the 3rd primary energy level. By the way, that’s not a typo… the d sublevel is in the 3rd primary level and not the 4th. There are 5 d orbitals in the 3d sublevel.

Row 5 has 3 sublevels and 3 types of orbitals on top of the lower sublevels and orbitals: an s sublevel with 1 s orbital, a p sublevel with 3 p orbitals, and a d sublevel with 5 d orbitals. Chemistry I Honors students are expected to know that…

Row 6 has 4 sublevels and 4 types of orbitals: an s sublevel with 1 s orbital, a p sublevel with 3 p orbitals, and a d sublevel with 5 d orbitals. Beginning with the 4th element with a 6th primary energy level (cerium, atomic number 58) a 4th sublevel is introduced into the atomic electron cloud structure. This is the f sublevel and it falls into the 4th primary energy level. There are 7 f orbitals in the f sublevel.

Row 7 also has 4 sublevels and 4 types of orbitals: an s sublevel with 1 s orbital and a p sublevel with 3 p orbitals. Beginning with the 3rd element with a 7th primary energy level (actinium, atomic number 89) a 3rd sublevel is introduced into the 6th primary level. This is the d sublevel. There are 5 d orbitals in the d sublevel. Additionally, beginning with the 4th element with a 7th primary energy level (thorium, atomic number 90) a 4th sublevel is introduced into the 5th primary level. This is the f sublevel. There are 7 f orbitals in the f sublevel.

40) How many orbitals are found in an s sublevel? What quantum numbers can be assigned to electrons in these orbitals?

41) How many orbitals are found in a p sublevel? What quantum numbers can be assigned to electrons in these orbitals?

42) How many orbitals are found in a d sublevel? What quantum numbers can be assigned to electrons in these orbitals?

43) How many orbitals are found in an f sublevel? What quantum numbers can be assigned to electrons in these orbitals?

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Key questions for this next section:

47) What 2nd (ml) and 3rd (ms) quantum numbers do we assigned

to electrons in an s orbital?

to electrons in a p orbital?

to electrons in a d orbital?

to electrons in a f orbital?

51) What regions of the periodic table correspond to the s, p, d, and f orbitals?

 Understand that two electrons can occupy each orbital

Explanation:

Since an s sublevel (quantum number l = 0) has only one s orbital (quantum number ml =

0), and each orbital can have 2 electrons, the s sublevel can have only 2 electrons (quantum number ms = +½ and –½).

Since a p sublevel (quantum number l = 1) has 3 p orbitals (quantum numbers ml = -1, 0,

and 1), and each orbital can have 2 electrons (each with a quantum number ms = +½ and –

½), the p sublevel can have 6 electrons.

Since a d sublevel (quantum number l = 2) has 5 d orbitals (quantum numbers ml = -2, -1,

0, 1, and 2), and each orbital can have 2 electrons (each with a quantum number ms = +½

and –½), the d sublevel can have 10 electrons.

Since a f sublevel (quantum number l = 3) has 7 d orbitals (quantum numbers ml = 3, 2,

-1, 0, -1, 2, and 3), and each orbital can have 2 electrons (each with a quantum number ms =

+½ and –½), the f sublevel can have 14 electrons.

Each of the electron energy sublevels represent different regions of the periodic table.

s s u b le v e l r e g iln p sublevel region d sublevel region f sublevel region Periodic Table of the Elements

Sublevel Regions

He He

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Figure 7. This illustration outlines the regions of the periodic table where the last electrons in the electron configuration of elements are found (see also figure 8). This illustration can help a high school chemistry student to write orbital notations, and electron configurations. Note that helium belongs in the s sublevel region.

Table 1. Quantum numbers, symbols, meaning, possible values, and their relationship to the periodic table, electron configurations, and orbital notation.

name

sym

bol orbital

meaning range of values value example

1st quantum number, principal quantum number

n

Primary

energy

level

(sometimes called a shell).

Whole positive numbers beginning with 1 and going up to 7 for ground state electrons

n = 1, 2, 3 , …

The 1st quantum number (the principal quantum number) is the same for quantum numbers, orbital notation, and electron configurations. They also correspond to the rows on the periodic table. In other words, the electrons with the highest 1st quantum number in any element will be the same as the number of the row in which the element is found.

2nd quantum number

l

Energy

sublevel

(sometimes called a subshell).

Whole positive numbers beginning with 0 and going up to 6 for ground state electrons, and the largest number that it can be is one less than the principal quantum number. However, ground state electrons never have an azimuthal quantum or angular momentum number larger than 4. (0 ≤ l ≤ n-1)

For n = 3, the 2nd quantum number (the orbital quantum number, azimuthal quantum number, or angular momentum number) can be 0, 1, or 2.

If the he 2nd quantum number is 0 then the sublevel is symbolized as s in both orbital notation and electron configurations.

If the 2nd quantum number is 1 then the sublevel is symbolized as p in both orbital notation and electron configurations.

If the 2nd quantum number is 3 then the sublevel is symbolized as d in both orbital notation and electron configurations.

If the 2nd_{quantum number is 4 then the sublevel is symbolized as}

f in both orbital notation and electron configurations.

3rd quantum number

m

l

Orbital

number

(sometimes called the energy shift).

Whole numbers from the negative value of l to the positive value of l . -l < ml < l

If the orbital quantum number (l) is 2, the value of the 3rd quantum number (the magnetic or projection of angular

momentum number) can be any whole number from -2 to positive 2.

So, for l = 2, ml = -2, -1, 0, 1, or 2.

There is no corresponding number for the 3rd quantum number in electron configurations.

In orbital notation, the box in which the electron symbol is found (either an up or down arrow) corresponds to the 3rd quantum number. The 1st box would correspond to the lowest number (-2 in the example above) and the last box would be the highest number (2 in the example above).

4th quantum

number

m

s

Spin

-½ or +½

(sometimes written simply as – or +)

The 4th quantum number for an electron (spin projection quantum number or simply, “spin number”), is either -½ or +½.

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s sublevel region

l = 0

ml = 0

Periodic Table of the Elements

with sublevel regions highlighted and quantum numbers listed

p sublevel region l = 0 ms = +½ ms = –½

Notice that helium has been moved from the far right to a position next to hydrogen. 1 IA 2 IIA 13 IIIA 14 IVA 15 VA 16 VIA 17 VIIA 18 VIIIA n = 1 1 H Hydrogen 1.00794 2 He Helium 4.002602 d sublevel region l = 2 ms = +½ ms = –½ ml = –1 ml = 0 ml = 1 ml = –1 ml = 0 ml = 1 n = 1 3 Li Lithium 6.941 4 Be Beryllium 9.01218 3 IIIB 4 IVB 5 VB 6 VIB 7 VIIB 8 9 VIIIB 10 11 IB 12 IIB 5 B Boron 10.811 6 C Carbon 12.0107 7 N Nitrogen 14.0067 8 O Oxygen 15.9994 9 F Fluorine 18.9984032 10 Ne Neon 20.1797 n = 1 11 Na Sodium 22.989769 12 Mg Magnesium 24.3050 ms = +½ ms = –½ 13 Al Aluminum 26.981538 14 Si Silicon 28.0855 15 P Phosphorous 30.973762 16 S Sulfur 32.065 17 Cl Chlorine 35.453 18 Ar Argon 39.948 ml = –2 ml = –1 ml = 0 ml = 1 ml = 2 ml = –2 ml = –1 ml = 0 ml = 1 ml = 2 n = 1 19 K Potassium 39.0983 20 Ca Calcium 40.078 21 Sc Scandium 44.955912 22 Ti Titanium 47.867 23 V Vanadium 50.9415 24 Cr Chromium 51.9961 25 Mn Manganese 54.938045 26 Fe Iron 55.845 27 Co Cobalt 58.933195 28 Ni Nickel 58.6934 29 Cu Copper 63.546 30 Zn Zinc 65.409 31 Ga Gallium 69.723 32 Ge Germanium 72.64 33 As Arsenic 74.92160 34 Se Selenium 78.96 35 Br Bromine 79.904 36 Kr Krypton 83.798 n = 1 37 Rb Rubidium 85.4678 38 Sr Strontium 87.62 39 Y Yttrium 88.90585 40 Zr Zirconium 91.224 41 Nb Niobium 92.90638 42 Mo Molybdenum 95.94 43 Tc Technetium (97.9072) 44 Ru Ruthenium 101.07 45 Rh Rhodium 102.90550 46 Pd Palladium 106.42 47 Ag Silver 107.8682 48 Cd Cadmium 112.411 49 In Indium 114.818 50 Sn Tin 118.71 51 Sb Antimony 121.760 52 Te Tellurium 127.60 53 I Iodine 126.90447 54 Xe Xenon 131.293 n = 1 55 Cs Cesium 132.905451 56 Ba Barium 137.327 57 La Lanthanum 138.90547 72 Hf Hafnium 178.49 73 Ta Tantalum 180.94788 74 W Tungsten 183.84 75 Re Rhenium 186.207 76 Os Osmium 190.23 77 Ir Iridium 192.217 78 Pt Platinum 195.084 79 Au Gold 196.966569 80 Hg Mecury 200.59 81 Tl Thallium 204.3833 82 Pb Lead 207.2 83 Bi Bismuth 208.98040 84 Po Polonium (208.9824) 85 At Astatine (209.9871) 86 Rn Radon (222.0176) n = 1 87 Fr Francium (223) 88 Ra Radium (226) 89 Ac Actinium (227) 104 Rf Rutherfordium (261) 105 Db Dubnium (262) 106 Sg Seaborgium (266) 107 Bh Bohrium (264) 108 Hs Hassium (277) 109 Mt Meitnerium (268) 110 Ds Darm stadtium (271) 111 Rg Roentgenium (272) 112 Uub Ununbium (277) 114 Uuq Ununquadium (289) (287) 116 Uuh Ununhexium (289) 118 (293)

Figure 8. This periodic table also shows the regions of the periodic table where the last electrons in the electron configuration of elements are found (see also figure 7). This also shows how where the quantum numbers for electrons correspond to these orbital notations and electron configurations. This illustration can help a high school chemistry student to write quantum numbers for individual electrons.

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s sublevel

region Periodic Table of the Elements _{with sublevel regions highlighted}

1 IA

2

IIA p sublevel region

1 H Hydrogen 1.00794 2 He Helium 4.002602

Notice that helium has been moved from the

far right to a position next to hydrogen. IIIA13 14 IVA 15 VA 16 VIA 17 VIIA He 18 VIIIA 3 Li Lithium 6.941 4 Be Beryllium 9.01218 d sublevel region 5 B Boron 10.811 6 C Carbon 12.0107 7 N Nitrogen 14.0067 8 O Oxygen 15.9994 9 F Fluorine 18.9984032 10 Ne Neon 20.1797 11 Na Sodium 22.989769 12 Mg Magnesium 24.3050 3 IIIB 4 IVB 5 VB 6 VIB 7 VIIB 8 9 VIIIB 10 11 IB 12 IIB 13 Al Aluminum 26.981538 14 Si Silicon 28.0855 15 P Phosphorous 30.973762 16 S Sulfur 32.065 17 Cl Chlorine 35.453 18 Ar Argon 39.948 19 K Potassium 39.0983 20 Ca Calcium 40.078 21 Sc Scandium 44.955912 22 Ti Titanium 47.867 23 V Vanadium 50.9415 24 Cr Chromium 51.9961 25 Mn Manganese 54.938045 26 Fe Iron 55.845 27 Co Cobalt 58.933195 28 Ni Nickel 58.6934 29 Cu Copper 63.546 30 Zn Zinc 65.409 31 Ga Gallium 69.723 32 Ge Germanium 72.64 33 As Arsenic 74.92160 34 Se Selenium 78.96 35 Br Bromine 79.904 36 Kr Krypton 83.798 37 Rb Rubidium 85.4678 38 Sr Strontium 87.62 39 Y Yttrium 88.90585 40 Zr Zirconium 91.224 41 Nb Niobium 92.90638 42 Mo Molybdenum 95.94 43 Tc Technetium (97.9072) 44 Ru Ruthenium 101.07 45 Rh Rhodium 102.90550 46 Pd Palladium 106.42 47 Ag Silver 107.8682 48 Cd Cadmium 112.411 49 In Indium 114.818 50 Sn Tin 118.71 51 Sb Antimony 121.760 52 Te Tellurium 127.60 53 I Iodine 126.90447 54 Xe Xenon 131.293 55 Cs Cesium 132.905451 56 Ba Barium 137.327 57 La Lanthanum 138.90547 72 Hf Hafnium 178.49 73 Ta Tantalum 180.94788 74 W Tungsten 183.84 75 Re Rhenium 186.207 76 Os Osmium 190.23 77 Ir Iridium 192.217 78 Pt Platinum 195.084 79 Au Gold 196.966569 80 Hg Mecury 200.59 81 Tl Thallium 204.3833 82 Pb Lead 207.2 83 Bi Bismuth 208.98040 84 Po Polonium (208.9824) 85 At Astatine (209.9871) 86 Rn Radon (222.0176) 87 Fr Francium (223) 88 Ra Radium (226) 89 Ac Actinium (227) 104 Rf Rutherfordium (261) 105 Db Dubnium (262) 106 Sg Seaborgium (266) 107 Bh Bohrium (264) 108 Hs Hassium (277) 109 Mt Meitnerium (268) 110 Ds Darm stadtium (271) 111 Rg Roentgenium (272) 112 Uub Ununbium (277) 114 Uuq Ununquadium (289) (287) 116 Uuh Ununhexium (289) 118 (293)

Figure 9. This periodic table also shows the regions of the periodic table where the last electrons in the electron configuration of elements are found (see also figure 7). This illustration can help a high school chemistry student to write orbital notations, and electron configurations.

Key questions and goals for next section:

52) Write orbital notation for all the elements of the periodic table.

53) Write electron configurations for all the elements of the periodic table. 54) Write electron dot notation for all the elements of the periodic table.

Use standard orbital notation to illustrate the electron configuration of any element in any of the periods based on the element’s position on the periodic table.

Orbital notation for any element includes the following:

• Write the symbol for the element (1st

letter is always capitalized and any subsequent letters must be lower case). • Write a colon following that symbol.

• Below the line and to the right write the primary energy level number followed by the sublevel letter.

• Above that primary level number and sublevel letter place squares for every orbital in that sublevel (squares are required even if the square remains empty).

• Use arrows to represent electrons in those squares. The 1st

arrow always points up and the 2nd arrow always points down. Only 2 electrons can be in an orbital and therefore only 2 arrows (1 up and 1 down) can be in an

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orbital/square. One upward pointing arrow is placed in each square of a sublevel before a 2nd, downward pointing arrow is placed in each orbital/square.

Step-by-step instructions to writing orbital notations (refer to figure 7 and figure 8 to see the sublevel regions s, p, d, and f in these orbital notations):

H:

1s

Write the number 1 below the line (as illustrated) to show that hydrogen’s only electron is in the 1st primary energy level. Write the letter s below the line (as illustrated) to show that hydrogen’s only electron is also in an s sublevel (note hydrogen’s location in figure 8). Draw a square above the 1 and the s to illustrate the single orbital in the s sublevel. Draw an upward pointing arrow to represent the 1st electron in that orbital.

C:

1s 2s 2p

Start with the 1st element, hydrogen. Write its symbol, followed by a colon (:).

N:

1s 2s 2p

Now that the 2 p orbitals have one electron in each, you can go back and pair them up as needed. See oxygen below.

He:

1s

Helium’s orbital notation is almost the same as hydrogen’s except that it has a 2nd electron in the 1s orbital. The 2nd electron in that orbital must be represented by a downward pointing arrow.

O:

1s 2s 2p

Li:

1s 2s

1s 2s 2p 3s

When we get to sodium, we enter a new row on the periodic table and therefore we are putting electrons in a new primary energy level; level 3. Recall that we always put electrons into an s sublevel and an s orbital first.

Mg:

1s 2s 2p 3s

The rule for writing the 2nd electron arrow in an orbital is an application of the Pauli exclusion principal, but a more whimsical name for the rule is the school bus

rule.

Once again, we need squares for the three p orbitals. Only we

This rule about putting the 1st arrow in each orbital pointing up and the 2nd electron arrow pointing down is an application of the Pauli exclusion principal. A more whimsical version that applies only to orbital notation is called the school bus rule. Here’s how it works:

Imagine that you are getting on a school on which none of the students know each other. Each student takes a seat by themselves until all the seats have one student on it. Finally, students start pairing up when there are no more empty seats.

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Al:

1s 2s 2p 3s 3p

are now in the 3rd primary energy level.

Si:

1s 2s 2p 3s 3p

Even though some of orbital squares are empty, you must never-the-less draw the squares. This is because—once you put an electron into a new

1s 2s 2p 3s 3p 4s

When you write orbital notation for electrons in the d sublevel region (see figure 8), you must drop to one lower primary energy level. Notice below how the highest primary electron energy level in scandium (Sc) is 4, but the d sublevel in scandium is in 3d.

Sc:

1s 2s 2p 3s 3p 4s 3d

Ti:

1s 2s 2p 3s 3p 4s 3d

V:

1s 2s 2p 3s 3p 4s 3d

When you get to chromium (Cr), an exception to the school bus rule occurs. The d sublevel is much more stable when it has an electron in each orbital. To create this extra stability, an electron from 4s drops down into the last orbital/square, leaving 4s with only 1 electron/arrow. The exception to the school bus rule also occurs when we get to molybdenum (Mo).

Cr:

1s 2s 2p 3s 3p 4s 3d

Please note that the there are no exceptions to the school bus rule in the f sublevel! Unlike the d sublevels, having one electron in each of the orbitals of the f sublevel does not provide enough additional stability for an electron to be pulled from an upper energy level.

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Er: 1s 2s 2p 3s 3p 4s 3d 4d 5s 4d 5p 6s 5d 4f Tm: 1s 2s 2p 3s 3p 4s 3d 4d 5s 4d 5p 6s 5d 4f Yb: 1s 2s 2p 3s 3p 4s 3d 4d 5s 4d 5p 6s 5d 4f Lu: 1s 2s 2p 3s 3p 4s 3d 4d 5s 4d 5p 6s 5d 4f Hf: 1s 2s 2p 3s 3p 4s 3d 4d 5s 4d 5p 6s 5d 4f Ta: 1s 2s 2p 3s 3p 4s 3d 4d 5s 4d 5p 6s 5d 4f

Please note that tungsten (W) does not follow exceptions to the school bus rule that chromium and molybdenum followed. See the orbital notations for chromium and molybdenum and figure 8. Tungsten’s 5d sublevel does not offer enough additional stability for an electron to be pulled from an upper energy level. The same thing happens for seaborgium (Sg), which is immediately below tungsten on the periodic table.

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When you write orbital notation for chromium (Cr) and molybdenum (Mo), you must remove one electron arrow from the next highest s orbital/sublevel and place that in the last orbital square (putting one electron arrow in each of the orbital squares in the d sublevel). When you write orbital notation for copper (Cu), silver (Ag), and gold (Au), you must remove one electron arrow from the next highest s orbital/sublevel and place that in the last orbital square (putting two electron arrows-one up and one down- in each of the orbital squares in the d sublevel).

When you write orbital notation for electrons in the f sublevel, you must drop to two lower primary energy levels (one more below the d sublevel primary level).

When you write orbital notation for electrons in the f sublevel, you must put one electron arrow in the 1st box of the d sublevel, and then fill the f sublevel. After filling the f sublevel, you return to fill the d sublevel.

Note: the symbol for the element MUST be shown followed by a colon (:) and then followed by the orbital notation.

o Additional methods of illustrating electron configuration include Practice Problem:

1. Using ONLY your copy of the periodic table, write the orbital notations for magnesium, molybdenum, and polonium. AFTER you have written these orbital notations check your work against the list above. If your work is not correct, study again the previous section, pick more elements and try again.

• Students must be able to illustrate the electron configurations for all elements on the periodic table  Electron configuration notation

For oxygen, the orbital notation for oxygen

O:

1s 2s 2p

becomes the electron configuration for oxygen: O: 1s22s22p4 .

All that you have to do is count the number of electron arrows in a sublevel in the orbital notation and that number becomes a superscript on the sublevel letter in the electron configuration.

Note: the symbol for the element MUST be shown followed by a colon (:) and then followed by the electron configuration.

Electron configuration for any element includes the following:

• Write the symbol for the element (1st

letter is always capitalized and any subsequent letters must be lower case). • Write a colon following that symbol.

• Write the primary energy level number followed by the sublevel letter.

• Count the number of electrons in that sublevel and enter that number as a superscript on the sublevel letter. "The use of the letters s, p, d, and f has a curious origin; they are the initial letters of the adjectives sharp, principal, diffuse, and fundamental, which happened to be used by spectroscopists in the period around 1890 to describe different observed spectral series of the alkali metals. These letters accordingly are not abbreviations of words that describe the orbitals in a meaningful way."

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• Continue this process until all the electrons in the atom are accounted for.

 There is a “shorthand” for longer electron configurations which we will call noble gas notation. In noble gas notation, the part of the electron configuration that matches the next lowest noble gas is removed from the configuration and replaced with the atomic symbol of the noble gas in brackets.

Example:

The electron configuration for chromium is: Cr: 1s2 2s2 2p6 3s2 3p6 4s1 3d5.

The next lowest noble gas is argon and its electron configuration is: Ar: 1s2 2s2 2p6 3s2 3p6.

Notice that argon and chromium have the same electron configuration up to the 3p sublevel:

Cr: 1s2 2s2 2p6 3s2 3p64s1 3d5 Ar: 1s2 2s2 2p6 3s2 3p6.

You can replace the part of chromium’s electron configuration that matches that or argon with the symbol for argon in brackets. Therefore the noble gas electron configuration notation is:

Cr: [Ar] 4s1 3d5.

Electron configurations for the elements (refer to figure 7 and figure 8 to see the sublevel regions s, p, d, and f in these orbital notations and to the orbital notations for each of the elements). The full electron configuration looks like this: Mg: 1s2 2s2 2p6 3s2, while the noble gas electron configuration looks like this: Mg: [Ne] 3s2. All that you do is to replace the section of the configuration that corresponds to the next lowest noble gas.

H: 1s1 He: 1s2 Li: 1s2 2s1 Li: [He] 2s1 Be: 1s2 2s2 Be: [He] 2s2 B: 1s2 2s2 2p1 B: [He] 2s2 2p1 C: 1s2 2s2 2p2 C: [He] 2s2 2p2 N: 1s2 2s2 2p3 N: [He] 2s2 2p3 O: 1s2 2s2 2p4 O: [He] 2s2 2p4 F: 1s2 2s2 2p5 F: [He] 2s2 2p5 Ne: 1s2 2s2 2p6 Ne: [He] 2s2 2p6 Na: 1s2 2s2 2p6 3s1 Na: [Ne] 3s1 Mg: 1s2 2s2 2p6 3s2 Mg: [Ne] 3s2 Al: 1s2 2s2 2p6 3s2 3p1 Al: [Ne] 3s2 3p1 Si: 1s2 2s2 2p6 3s2 3p2 Si: [Ne] 3s2 3p2 P: 1s2 2s2 2p6 3s2 3p3 P: [Ne] 3s2 3p3 S: 1s2 2s2 2p6 3s2 3p4 S: [Ne] 3s2 3p4 Cl: 1s2 2s2 2p6 3s2 3p5 Cl: [Ne] 3s2 3p5 Ar: 1s2 2s2 2p6 3s2 3p6 Ar: [Ne] 3s2 3p6 K: 1s2 2s2 2p6 3s2 3p6 4s1 K: [Ar] 4s1 Ca: 1s2 2s2 2p6 3s2 3p6 4s2 Ca: [Ar] 4s2 Sc: 1s2 2s2 2p6 3s2 3p6 4s2 3d1 Sc: [Ar] 4s2 3d1 Ti: 1s2 2s2 2p6 3s2 3p6 4s2 3d2 Ti: [Ar] 4s2 3d2

The same exceptions in the orbital filling order still apply for chromium, copper, molybdenum, silver, and gold. Notice that these elements have just 1 electron in their highest s sublevel.

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Ga: [Ar] 4s2 3d10 4p1 Ge: [Ar] 4s2 3d10 4p2 As: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3 As: [Ar] 4s2 3d10 4p3 Se: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p4 Se: [Ar] 4s2 3d10 4p4 Br: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5 Br: [Ar] 4s2 3d10 4p5 Kr: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 Kr: [Ar] 4s2 3d10 4p6 Rb: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1 Rb: [Kr] 5s1 Sr: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 Sr: [Kr] 5s2 Y: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d1 Y: [Kr] 5s2 4d1 Zr: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d2 Zr: [Kr] 5s2 4d2 Nb: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d3 Nb: [Kr] 5s2 4d3 Mo: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1 4d5 Mo: [Kr] 5s1 4d5 Tc: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d5 Tc: [Kr] 5s2 4d5 Ru: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d6 Ru: [Kr] 5s2 4d6 Rh: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d7 Rh: [Kr] 5s2 4d7 Pd: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d8 Pd: [Kr] 5s2 4d8 Ag: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1 4d10 Ag: [Kr] 5s1 4d10 Cd: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 Cd: [Kr] 5s2 4d10 In: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p1 In: [Kr] 5s2 4d10 5p1 Sn: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p2 Sn: [Kr] 5s2 4d10 5p2 Sb: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p3 Sb: [Kr] 5s2 4d10 5p3 Te: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p4 Te: [Kr] 5s2 4d10 5p4 I: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p5 I: [Kr] 5s2 4d10 5p5 Xe: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 Xe: [Kr] 5s2 4d10 5p6 Ce: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s1 Ce: [Xe] 6s1 Ba: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 Ba: [Xe] 6s2 La: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 5d1 La: [Xe] 6s2 4d1

Notice that—unlike the orbital notation—the f sublevel is listed before the d sublevel in the electron configuration.

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Practice Problem:

2. Using ONLY your copy of the periodic table, write the both kinds of electron configurations for phosphorous, silver, and mercury. AFTER you have written these electron configurations check your work against the list above. If your work is not correct, study again the previous section, pick more elements and try again.

 Electron Dot notation (to show valence electrons)

Most students learned about electron dot notation in physical science. The orbital notation for selenium is:

Se: 1s22s22p63s23p64s23d104p4.

Selenium’s highest primary energy level is 4. Notice that there are 6 electrons in the highest primary level of selenium.

The electron dot notation for selenium is or or It doesn’t matter on which side you put

the single or paired electrons so long as you follow the rules.

• • • • •

S

• • • • • • •

Cl

• • • • • • •

Ar

• • • • •

In the 4th row of the periodic table you will 1st enter the d sublevel region. When writing orbital notation, electrons in the d sublevel region drop to one primary energy level lower, so those electrons are not included in the electron dot notation of these elements: scandium through zinc. When writing electron dot notation, this same pattern continues through the remainder of the periodic table. d and f sublevel electrons are not included in the electron dot notation.

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Practice Problem:

3. Using ONLY your copy of the periodic table, write the electron dot notation for oxygen, rubidium, arsenic, and iodine. AFTER you have written these orbital notations check your work against the list above. If your work is not correct, study again the previous section, pick more elements and try again.

• Students must be able to illustrate the electron configurations for all elements on the periodic table

• Students must understand that the order in which electrons fill orbitals reflect the most stable electron arrangement for the given number of electrons.

 Students should be able to make general statements such as those below concerning stable electron arrangements

o All “d” orbitals are less stable than the “s” orbitals in the next-highest energy level

o All “f” orbitals are less stable than the “s” and the “p” orbitals which are two energy levels higher, and less stable than the “d” orbitals which are one energy level higher

• Understand exceptions to the normal orbital filling order (Cr, Mo, Cu, Ag, Au), what the exceptions are, and why they are exceptions.

Explanation:

All the elements in the list above (Cr, Mo, Cu, Ag, and Au), take an electron from the s sublevel in the highest primary level and add an electron to a d orbital in order to have all d orbitals half filled (with one electron) or completely full (with 2 electrons). These are exceptions to the normal filling order. They occur because have one electron in each of the d orbitals or 2 electrons in each of the d orbitals is much more stable than not having them. Key questions for next section:

55) What is a “bright line spectrum?” 56) Describe the Bohr model of the atom. 57) How are electrons promoted in energy? 58) How do atoms give off light energy?

59) Why do atoms give off only specific wavelengths of light?

• Students should be able to use a Bohr model of the atom to explain the bright line spectrum in terms of electrons moving between energy levels

Explanation:

Electron configurations and orbital notation are used as a kind of bookkeeping system for electrons in their ground state. The ground state is the lowest energy state of an electron. In this unit, students should be able to use a Bohr model of the atom to explain the bright line spectrum in terms of electrons moving between energy levels. This means that students will know more about the behavior of excited electrons, which means electrons that are NOT in their ground (or lowest) energy state.

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that high school chemistry teachers use the Bohr model to explain is the quantum mechanical behavior of electrons.

Using the Bohr model it’s easy to explain how electrons are excited through exposure to an energy source with a sympathetic frequency. These excited electrons are then promoted to a higher orbit in the Bohr model. When the electron falls back to the original orbit or to some other quantum mechanical orbit between the original orbit and the orbit to which it was promoted, photons of light are emitted. The photons are of a specific energy

frequency corresponding to the length of the drop to a lower orbit. Lots and lots of these photons of specific energy create the “bright line spectrum” or emission spectrum of an element. Check out the following web pages for help in understanding the connection between electrons absorbing energy and the light that is emitted:

http://www.colorado.edu/physics/2000/quantumzone/index.html

http://www.colorado.edu/physics/2000/quantumzone/bohr.html

http://www.colorado.edu/physics/2000/quantumzone/bohr2.html

Also, check out this video on YouTube:

http://www.youtube.com/watch?v=-YYBCNQnYNM

It is well-known that electrons do not travel in orbits, but if we transpose the word energy state for orbit, we get a pretty good idea of how electrons behave. If we use the Bohr model as a kind of circular graph in which the concentric circles no longer represent orbits but represent the quantum levels of energy, then the Bohr model becomes a reasonable representation of what really occurs even if it’s a more abstract idea. When an electron is promoted to a higher energy state by absorbing energy with a sympathetic frequency, it then releases very specific photons of light corresponding to the very specific drops in energy that electrons can go through. These very specific photon frequencies can be used to identify elements and compounds through spectrophotometry. Spectrophotometry is the study of light frequencies emitted or absorbed by substances.

The result of emitted photons from excited electrons is a bright line spectrum. A bright line spectrum is like a segment of a rainbow with most of the colors removed. Only certain bright lines of color remain.

The Bohr model is too simplistic to explain all the behaviors of electrons and the photons they absorb and emit, but it does approximate the behavior of some electrons in an atoms. Since it is simple, it is the model that students are expected to understand and be able to explain in this class.

In the Bohr model, electrons absorb energy (often in the form of photons) that match the electron’s wave energy. The electron is promoted to a higher energy state and then drops back to a lower state while—almost at the same instant—it releases (emits) photons of light of very specific frequencies that match the wave energy of the electron. Electrons can emit different wavelengths of light because the electron can be promoted to different quantum levels of energy and emits photons that match those quantum levels of wave energy.

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Indicator C-2.2

According to Indicator C-2.2 in the South Carolina Science Standards students should be able to summarize atomic properties (including electron configuration, ionization energy, electron affinity, atomic size, and ionic size).

It is essential for all students to

• Understand the following atomic characteristics and properties (in terms of atomic structure) and understand what variables influence the magnitude of the characteristics or properties for a given element.

o Electron configuration

Explanation:

Electron configurations were explained in the last section. In atomic physics and quantum chemistry, electron configuration is the arrangement of electrons in an atom, molecule, or other physical structure. In this class we will only be looking the arrangement of electrons in ground state electrons and using that knowledge to help us understand how electrons are shared in covalent bonding or given up and gained in ionic bonding. Knowledge of the electron

configuration of different atoms is useful in understanding the structure of the periodic table of elements.

Physicists and chemists use a standard notation to describe atomic electron configurations. In this notation, an electron energy sublevel is written in the form nxy, where n is the primary energy level number, x is the subshell label and y is the number of electrons in the subshell. An atom's sublevels are written in order of increasing energy – in other words, the sequence in which they are filled. Note: in this class orbital notations are written in the order in which they are 1st encountered.

For instance, ground-state hydrogen has one electron in the s orbital of the first primary electron energy level, so its configuration is written H: 1s1. Lithium has two electrons in the 1s electron energy sublevel and one in the (higher-energy) 2s sublevel, so its ground-state configuration is written Li: 1s2 2s1. Phosphorus (atomic number 15), is as follows: P: 1s2 2s2 2p6 3s2 3p3.

For atoms with many electrons, this notation can become lengthy and so the noble gas notation is used. It is often abbreviated by noting that the first few sublevels are identical to those of one or another noble gas. Phosphorus, for instance, differs from neon (1s2 2s2 2p6) only by the presence of a third primary electron energy level. Thus, the electron configuration of neon is pulled out, and phosphorus is written as follows: P: [Ne] 3s2 3p3.

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Key questions for next section:

62) What is “first ionization energy?”

63) Explain the trend in 1st ionization energy as it relates to the position of each atom on the periodic table. 64) Explain the trend in each subsequent ionization energy as it relates to the 1st, 2nd, 3rd and all subsequent

electrons being removed from an atom.

o Ionization energy

Explanation:

The ionization energy (sometimes called the ionization potential or symbolized EI) of an atom is the energy required to remove an electron from the isolated atom or ion. It is typically measured in kilojoules per

mole kJ

mol

 

 

 . Isolated means, “By itself,” or not bonded to other atoms. In this class students will not be required to memorize these but will only be required to understand 2 ionization energy patterns or trends. Students must understand that ionization energy refers to the removal of any electron whether it’s the 1st one, the 2nd, of any subsequent electron. First ionization energy, however, refers only to the energy required to remove one (the 1st) electron from an atom.

The first of the 2 trends deals with 1st ionization energy. The 1st ionization energy of elements on the periodic table is highest for helium (which is in the upper right corner) and lowest for francium (which is in the lower left hand corner). In between these 2 elements, there is a general pattern of 1st ionization energy that grows larger as you approach helium.

The second of the 2 trends deals with all ionization energies for any given (single) atom. The 1st ionization energy is always lowest. The 2nd ionization energy (the energy required to remove a 2nd electron from that atom) is always higher than the 1st. the 3rd ionization energy is higher still, and so forth.

Periodic table of the elements with 1st ionization energies. H

13.5

(in eV)

Note: No values are shown for those elements whose 1st ionization energies are not known.