Mr. Joanis’ Science Class! ~ Chemistry ~ WEEK 13 Name: ______Period: ____ Outermost Energy Levels A Basic Periodic Trend

In our last lesson, you learned how to determine the electron configuration for any given element. To complete that assignment, you likely had to make use of a periodic table. For our next few lessons, we’ll be focusing more on how the periodic table is arranged, what the arrangement of the elements in the periodic table can tell us, and what properties of elements exist that we can determine from their placement on the periodic table. These goals connect directly with one of the physical science standards we aim to meet in this class:

PS1-1: Use the periodic table as a model to predict the relative properties of elements based on the patterns of electrons in the outermost energy levels of atoms.

There’s one vocabulary term in this standard that we have already learned. Knowing this vocabulary term is key to our understanding of the standard as a whole, and to our ability to meet this standard.

The outermost energy level of an atom ties directly to the concepts we’ve been discussing in our previous lessons. As you should already know, the quantum model of the atom separates “energy levels” into two more specific concepts: the main energy level, and the sublevel. Main energy levels are represented with numbers (i.e. “1, 2, 3, 4, etc…”), and sublevels are represented with one of four letters (i.e. “s, p, d, or f”). By combining a main energy level and a sublevel together, we arrive at our new and more correct understanding of an energy level. For example: 2p is an energy level where the electrons in an atom may reside. 2s is another energy level where electrons may reside. 4f, 3d, 1s, and 5p are also all examples of energy levels where electrons may reside in an atom.

Considering all of this, the outermost energy level of an atom is the energy level that is the highest- energy energy level where some number of electrons reside. The diagram below, which is also present in your textbook, shows the energy levels in order of ascending (increasing) energy. The 2s orbital is, for example, higher energy than 1s and lower energy than 2p. This means that if there are some electrons in the 2s orbital but none in any orbital that is higher in energy, 2s is the outermost energy level of whatever atom you are describing. If the 2p orbital is the highest energy level with any electrons in it, then 2p would be the outermost energy level in that case.

Recall what you learned in previous lessons as well: Higher energy levels will not be filled with any electrons until all lower energy levels δ have been completely filled (i.e. 2p can’t have any electrons in it unless 2s and 1s are already full).

Getting back to the standard on the first page of this lesson: What is meant by “patterns of electrons” in the outermost energy levels of atoms? Lets take a look at a few elements on the periodic table in order to find out what patterns we can see.

Consider the orbital notation electron configurations for carbon, silicon, and germanium...

As you may have already noticed, there’s something similar about the outermost energy level in each of these elements. Carbon’s, silicon’s, and germanium’s outermost energy levels are all in the p- sublevel, and they all have exactly 2 electrons in them.

Something worth noting about this collection of three elements is that they also all appear within the game group on the periodic table.

The vertical columns of elements on the periodic table are called groups. They are also numbered on most periodic tables. The column that contains hydrogen, lithium, and sodium is “Group 1” of the periodic table. It is called “Group 1” because it is the first of the columns, if you look at the periodic table from left to right. Each column after that is numbered in a similar way. The second column of the periodic table – the column that begins with beryllium, is “Group 2.” The column that carbon, silicon and germanium lie on is called “Group 14.”

Similar to groups, the horizontal rows on the periodic table are called periods. Hydrogen and helium exist on “period 1.” Lithium, carbon, and a number of other elements exists on “period 2.”

So, is it coincidence that these three Group 14 elements all have similar looking outermost energy levels? Let’s look at another set of elements’ electron configurations to see if they have a similar pattern. If construct the orbital notation electron configurations for three of the Group 1 elements (hydrogen, lithium and sodium), they should look something like this…

As you can see, a similar trend appears here. Again, all of the outermost energy levels of hydrogen, lithium, and sodium atoms are part of the same sublevel (s-sublevel) and contain the same number of electrons (1).

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If you look at the electron configurations of any two elements in the same group of the periodic table, you would notice that this relationship is always true (with a few exceptions that we don’t have to worry about yet).

As we will learn in our lesson on chemical reactions, the number of electrons an atom has in its outermost energy level has a great effect on how that atom can react with other atoms.

All of the elements in Group 1 of the periodic table, as we have demonstrated, have the same number of electrons in their outermost energy level, and the sublevel of that outermost energy level is also the same for all Group 1 elements. Therefore, we can expect all of the Group 1 elements to have similar properties. This is how the modern periodic table was arranged. Today, we call the Group 1 elements of the periodic table the alkali metals, and the Group 2 elements alkaline-earth metals. These two groups make up what we call the “s-Block” of the periodic table, because all elements in that block have an outermost energy level with an s-sublevel.

Hydrogen is a special exception to this rule. As we can see, it has a similar electron configuration to its Group 1 family members, but hydrogen is a gas at room temperature and not a solid metal. Although its properties as a “standalone” element don’t really match up with any group in the periodic table, the arrangement of electrons in its outermost energy level does mean that is reacts in similar ways to the other Group 1 elements (something that we will explore in much more detail when we learn about chemical reactions).

The “d-Block” elements are known as the transition metals. They are the metals with typical metallic properties such as high luster, good electrical conductivity, and malleability.

The “p-Block” elements are known as the main-group elements. Each group within the p-Block has its own specific name that we will discuss later, as we delve into their reactive properties. These six groups in this block all have an outermost energy level with a p-sublevel.

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Using the information in the packet up to this point, you should be able to determine certain pieces of information about an element based just on its electron configuration. For example, we can solve the following problem:

Without looking at the periodic table, identify the group, period, and block in which the element that has the electron configuration [Xe]6s2 is located.

Because this is an electron configuration in noble gas configuration, and because of the rules of noble gas configuration that we learned in a past lesson, we know that the element being represented here has “6s” as its outermost energy level. The main energy level of all s-block elements is equal to that element’s period, so we know that our element exists on Period 6.

Because our element’s outermost energy level has 2 electrons in it, we know it exists in the 2nd group of the s-Block (Group 2).

Because the element’s outermost energy level has an s-sublevel, we know that it exists in the s-Block.

Therefore, our answer to the question above is: Group 2, Period 6, s-Block. We were able to find this answer without even looking at the periodic table, but you can check your answer by using the periodic table to find the element with a matching electron configuration. Here, we were talking about cesium (Cs).

Here’s one last piece of information that you will need to answer questions in this worksheet. In our next lesson, we will explain how and why this statement is true: The further down an element is within its group, the more reactive that element is.

Or, stated in a different way… Elements of a greater period are more reactive.

Answer the following questions:

1. Circle all electron configurations which are impossible. • [Ne]3s3 • [Ar]3d10 • [Ar]4s23d1 • [Xe] 6s26p3 • [He]2s22p6 • 1s1 2. Is lead (Pb) a transition metal? Why or why not.

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3. How many groups are there in the d-Block of the periodic table?

a. What is true about the electrons in the outermost energy level of d-Block elements?

4. List one transition metal

a. Give its electron configuration, in noble gas configuration.

b. List another transition metal in the same group

c. Give the electron configuration of that second transition metal, in noble gas configuration.

d. What is similar about the electron configuration of atoms of these two elements?

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5. Without looking at the periodic table, identify the group, period and block in which the element that has an electron configuration of [Kr]5s1 is located.

6. Without looking at the periodic table, identify the group, period and block in which the element that has an electron configuration of [Ar]4s23d8 is located.

7. Without looking at the periodic table, identify the group, period and block in which the element that has an electron configuration of [Ne]3s23p5 is located.

a. Without looking at the periodic table, write the electron configuration for the Group 17 element in the second period. Is this element likely to be more reactive or less reactive than the element with an electron configuration of [Ne]3s23p5.

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8. Write the orbital notation electron configurations for helium (He), neon (Ne) and argon (Ar) below.

a. What similarities exist in the outermost energy level of all three of these elements?

b. Does it make sense to you that helium be placed in Group 18 of the periodic table? Why or why not?