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Electron Configuration Lab

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Introduction: Electron configurations are the summary of where the electrons are around a nucleus (FSU, 2020). As we learned earlier, each neutral atom has a number of electrons equal to its number of protons. What we will do now is place those electrons into an arrangement around the nucleus that indicates their energy and the shape of the orbital in which they are located. Below is a summary of the types of subshells, with its orbitals (region of probability that describe the likely position of electrons in an atom) in Figure 1. So, based on what we know about the quantum numbers, you need 2 electrons to fill an s subshell, 6 electrons to fill a p subshell, 10 electrons to fill a d subshell and 14 electrons to fill the f subshell. What we have not discussed is how these subshells get filled...the order of fill.

Figure 1: Orbitals and Electron Capacity

Figure 2: Electron Energy Filling Diagram

Order of Fill The order in which electrons are placed into the orbitals is based on the order of their energy. This is referred to as the Aufbau principle. The lowest energy orbitals fill first. This order was determined by calculation and is summarized in Figure 2 above. How to Write an Electron Configuration The symbols used for writing the electron configuration start with the shell number (n) followed by the type of orbital (s, p, d, f) and finally the superscript indicates how many electrons are in the orbital.

Figure 3: Symbols Used for Electron Configuration 1

Example: Looking at the periodic table, you can see that oxygen has 8 electrons. Based on the order of fill in Figure 2, these 8 electrons would fill in the following order 1s, 2s and then 2p. So, oxygen's electron configuration would be 1s22s22p4. Special Cases Configurations of ions present a special case of electron configuration and also demonstrate the reason for the formation of those ions in the first place. If you need to write the full electron configuration for an anion , then you are just adding additional electrons and the configuration is simply continued. Example: We know that oxygen always forms 2- ions when it makes an ion. This would add 2 electrons to its normal configuration making the new configuration: 1s22s22p6. With 10 electrons you should note that oxygen's electron configuration is now exactly the same as neon's. We talked about the fact that ions form because they can become more stable with the gain or loss of electrons to become like the noble gases and now you can see how they become the same. The electron configurations for Cations are also made based on the number of electrons but there is a slight difference in the way they are configured. First you should write their normal electron configuration and then when you remove electrons you have to take them from the outermost shell (valence shell). Note that this is not always the same way they were added. Example: Iron has 26 electrons so its normal electron configuration would be: 1s 22s22p63s23p64s23d6 When we make a 3+ ion for iron, we need to take the electrons from the outermost shell first so that would be the 4s shell NOT the 3d shell: 1s22s22p63s23p63d5 Noble gas abbreviation One other note on writing electron configurations: A short cut. When writing some of the lower table configurations the total configuration can be long. In these cases, you can use the previous noble gas to abbreviate the configuration as shown below in Figure 4. You just have to finish the configuration from where the noble gas leaves it: Element Magnesium, Mg Bromine, Br Strontium, Sr

Number of Electrons 12 35 38

Electron Configuration

Noble Gas Abbreviation

1s22s22p63s2 [Ne] 3s2 2 2 6 2 6 2 10 5 1s 2s 2p 3s 3p 4s 3d 4p [Ar] 4s23d104p5 2 2 6 2 6 2 10 6 2 1s 2s 2p 3s 3p 4s 3d 4p 5s [Kr] 5s2

Figure 4: Noble Gas Abbreviation Exceptions There are 19 exceptions to the expected electron configuration including Cr, Cu, Nb, Mo, Ru, Rh, Pd, Ag, La, Ce, Gd, Pt, Au, Ac, Th, Pa, U, Np, and Cm. This is mainly due to the fact that totally-filled and exactly half-filled subshells have an unusual stability. Notice that in the expected electron configuration for chromium (Figure 5) there is a totally filled 4s2 but a partially filled 3d4. In chromium, one of the 4s electrons goes up to the 3d so that both are exactly half-filled 4s13d5. 2

Element Chromium, Cr Copper, Cu Silver, Ag

Number of Electrons 24 29 47

Expected Electron Configuration

Actual Electron Configuration

1s22s22p63s23p64s23d4 1s22s22p63s23p64s23d9 1s22s22p63s23p64s23d104p65s24d9

1s22s22p63s23p64s13d5 1s22s22p63s23p64s13d10 1s22s22p63s23p64s23d104p65s14d10

Figure 5: Examples of Exceptions to Expected Electron Configurations Orbital Diagrams Another way to represent the order of fill for an atom or ion is by using an orbital diagram often referred to as “electron configuration diagram”. Examples of this can be seen in Figure 6. The boxes are used to represent the orbitals and to show the electrons placed in them. The order of fill is the same but as you can see from above the electrons are placed singly into the boxes of degenerate orbitals (orbitals of equal energy) before filling them with both electrons. This is called Hund's Rule: "Half fill before you Full fill" and again this rule was established based on energy calculations that indicated that this was the way atoms actually distributed their electrons into the orbitals. Periodic Properties One of the cool things about electron configurations is their relationship to the periodic table. Basically, the periodic table was constructed so that elements with similar electron configurations would be aligned into the same groups, Figure 7. The periodic table shown above demonstrates how the configuration of each element was aligned so that the last orbital filled is the same except for the shell number. The reason this was done is that the configuration of an element gives the element its properties and similar configurations yield similar properties. Let's go through some of the Periodic Properties that are influenced directly by the electron configuration:

Figure 6: Electron Configuration Diagram

Figure 7: Periodic Table Showing Last Orbital Filled

3

Atomic Size The size of atoms increases going down in the periodic table (Figure 8). This should be intuitive since with each row of the table you are adding a shell (n). What is not as intuitive is why the size decreases from left to right. But again, the construction of the electron configuration gives us the answer. What are you doing as you go across the periodic table? Answer, adding protons to the nucleus and adding electrons to the valence shell of the element. What is not changing as you cross a period? Answer, the inner shell electrons. So, think of it this way, the inner shell electrons are a shield against the pull of the nucleus. As you cross a period and increase the number of protons in the nucleus you increase its pull but since you are only adding electrons to the new shell the shield is not increasing but remains the same all the way across. This means the pull on the electrons being added to the valence shell is increasing steadily all the way across. What happens if you pull harder on the electrons? Well, they come closer to the nucleus and the size of the atom decreases. The effect of the nucleus pulling on the electrons being added across a period is called the effective nuclear charge and is calculated as ZEff = #protons - Core # Electrons. So, for example the pull felt by Sulfur would be ZEff = 16 - 10 = +6

Figure 8: Atomic Radius

Electronegativity Electronegativity may be the most important of the periodic properties you can learn and understand since so many other properties are depend on its value. Electronegativity is an atom’s ability to pull electrons towards itself. Electronegativity is generally expressed by the Pauling Scale and the values were determined experimentally. The table below shows the scale values for the elements. The electronegativity values increase from left to right and bottom to top in the periodic table excluding the Noble gases. The most electronegative element is Fluorine. From these electronegativity values we can derive the patterns of two other periodic properties: Ionization Energy and Electron Affinity.

Figure 9: Electronegativity 4

Ionization Energy Ionization energy is the amount of energy required to remove an electron from an atom. All ionization energies are positive values because all of these removals (even those for elements that form positive ions) require input of energy. The more electronegative the element, the higher the ionization energy. Examples: 1. Which has a higher ionization energy, S or Ca? It would be more difficult to take an electron away from sulfur. Sulfur is smaller and more electronegative. Sulfur has the higher ionization energy. 2. Which has a higher ionization energy, K or K+? It would be more difficult to take away an electron from K + because it has already lost one electron and is electron poor. The potassium ion has the higher ionization energy.

Figure 10: Ionization Energy

Electron Affinity

Electron affinity is the change in energy of an atom when an electron is added to the atom to form a negative ion. In other words, the neutral atom's likelihood of gaining an electron. Energy is typically released when an atom gains an electron and becomes an anion. For this reason, the higher more negative number indicates a large electron affinity. The electronegativity and Electron Affinity increases in the same pattern in the periodic table. Left to right and bottom to top. Example: Which has the largest most negative electron affinity, Cl or Br? While both halogens desire to gain an electron, chlorine with it’s smaller atomic radius has a greater electron affinity. Chlorine has the largest negative electron affinity.

Figure 11: Electron Affinity

The Four Quantum Numbers A total of four quantum numbers are used to describe completely the movement and trajectories of each electron within an atom (CHEM, 2019). Each electron in an atom has a unique set of quantum numbers; according to the Pauli Exclusion Principle, no two electrons can share the same combination of four quantum numbers. Quantum numbers are important because they can be used to determine the electron configuration of an atom and the probable location of the atom's electrons. Quantum numbers are also used to understand other characteristics of atoms, such as ionization energy and the atomic radius. 5

In atoms, there are a total of four quantum numbers: the principal quantum number (n), the orbital angular momentum quantum number (l), the magnetic quantum number (ml), and the spin quantum number (ms). The principal quantum number, n, describes the energy of an electron and the most probable distance of the electron from the nucleus. In other words, it refers to the size of the orbital and the energy level an electron is placed in. The number of subshells, or l, describes the shape of the orbital. It can also be used to determine the number of angular nodes. The magnetic quantum number, ml describes the energy levels in a subshell, and ms refers to the spin on the electron, which can either be up or down. The Principal Quantum Number (n) The principal quantum number, n, designates the principal electron shell. Because n describes the most probable distance of the electrons from the nucleus, the larger the number n is, the farther the electron is from the nucleus, the larger the size of the orbital, and the larger the atom is. n can be any positive integer starting at 1. When n=1 it designates the first principal shell (the innermost shell). n = 1, 2, 3, 4…. The Angular Momentum Quantum Number (l) The angular momentum quantum number l determines the shape of an orbital, and therefore the angular distribution (Figure 12). The number of angular nodes is equal to the value of the angular momentum quantum number l. Each value of l indicates a specific s, p, d, f subshell (each unique in shape.) The value of l is dependent on the principal quantum number n. Unlike n, the value of l can be zero. It can also be a positive integer, but it cannot be larger than one less than the principal quantum number (n-1): l = 0, 1, 2, 3, 4…(n−1) Angular Momentum Quantum Number, l

Type of Subshell

0

s

1

p

2

d

3

f

Shapes of Subshells

Figure 12: Angular Momentum Quantum Number 6

The Magnetic Quantum Number (ml) The magnetic quantum number ml determines the number of orbitals and their orientation within a subshell. Consequently, its value depends on the orbital angular momentum quantum number l. Given a certain l, ml is an interval ranging from -l to +l so it can be zero, a negative integer, or a positive integer. As seen below, this signifies the possibilities for ml, but to find the correct possibility the electron diagram must be drawn (Figure 6). s orbital 0

p orbital -1

0

+1

-2

-1

0

+1

+2

-3

-2

-1

0

+1

d orbital f orbital +2

+3

Figure 13: Magnetic Quantum Number Values Based on Subshells The Electron Spin Quantum Number (ms) Unlike n, l, and ml, the electron spin quantum number, ms, does not depend on another quantum number. It designates the direction of the electron spin and may have a value of +1/2, represented by ↑, or –1/2, represented by ↓. This means that when ms is positive the electron has an upward spin, which can be referred to as "spin up." When it is negative, the electron has a downward spin, so it is "spin down." The significance of the electron spin quantum number is its determination of an atom's ability to generate a magnetic field or not. Labeling the Four Quantum Numbers When you are asked to label the 4 quantum numbers for an atom or ion, the first step is to find the number of electrons. After that, write the complete ground state electron configuration. Next, draw the electron configuration diagram for the subshell that the last electron went into. The 4 quantum numbers are only based on the last electron that was added. 2p Now, we are ready to label the 4 quantum numbers. 1s 2s Example: Find the 4 quantum numbers for carbon n = 2, because the last electron is in the 2nd shell. l = 1, because the last electron is in a p orbital. ml = 0, because the last electron is in the middle position. ms = + ½ , because the last electron is in the up position

-1 0 +1

References FSU. (2020). Electron Configurations. Retrieved from https://www.chem.fsu.edu/chemlab/chm1045/e_config.html. 7

Pre-lab Questions: Complete these questions BEFORE you attempt the lab. (12 pts) 1. Define the following. a. Electron configuration b. Aufbau principle c. Noble gas abbreviation for electron configuration d. Electron configuration diagram e. The principal quantum number, n f. The angular momentum quantum number, l g. The magnetic quantum number, ml h. The spin quantum number, ms

2. Explain the periodic trend for atomic size

3. Explain the periodic trend for electronegativity

4. Explain the periodic trend for ionization energy

5. Explain the periodic trend for electron affinity 8

The Lab: For each element or ion give the full electron configuration, the noble gas abbreviation, the electron configuration diagram and the 4 quantum numbers. (70 pts) Element or Ion

Electron Configuration Whole:

Quantum Numbers Diagram:

n l

N

ml Abbreviated: Whole:

ms Diagram:

n l

N3-

ml ms

Abbreviated: Whole:

Diagram:

n l

Ca

ml ms Abbreviated: Whole:

Diagram:

n l

Ca2+

ml ms

Abbreviated: Whole:

Diagram:

n l

O

ml Abbreviated:

ms 9

Element or Ion

Electron Configuration Whole:

Quantum Numbers Diagram:

n l

O2-

ml ms Abbreviated: Whole:

Diagram:

n l

Cr

ml ms Abbreviated: Whole:

Diagram:

n l

Fe

ml ms Abbreviated: Whole:

Diagram:

n l

Fe3+

ml ms Abbreviated: Whole:

Diagram:

n l

Cl-

ml ms Abbreviated: 10

Post-lab Questions (7 pts): 1. With respect to size, rank Fe, Fe2+, Fe3+ from smallest to largest and explain.

2. What is the reason for chromium’s unexpected electron configuration?

3. Why can no 2 electrons in the same element or ion have the same 4 quantum numbers?

4. When iron is oxidized, loses electrons, why does the electron come from the valence shell (4s) instead of the highest energy shell (3d)?

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