Chem 115 Activity 5(2)

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Dec 6, 2023

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Chem 115 Activity 5 Introduction: This activity covers sections 2.6-2.8 from the text including objectives 2.6-2.8. This is an in-class activity! You will find on BlackBoard an assignment that will ask you to enter some of your answers. Some of these questions will be duplicated there. You will get five attempts at each question, so this will allow you to check your answers, as well as receive a grade. If you are unsure of your answer after four attempts, you should ask for help. Not all questions will be graded, but you will still be asked to enter a response. When discussing answers with your group, make sure you not only agree on answers, but also why they are correct. Can you verbally defend why your answer is correct? Electron Orbital Energy Levels Learning Objective 2.6 : Describe the distribution of electrons into shells, subshells, and orbitals around the nucleus of an atom. Learning Objective 2.7 : Write the electron configuration for an atom to describe how electrons are distributed into specific orbitals. We have learned that the atom is composed of two main areas: the nucleus and the electron cloud. The nucleus leads to the identity of the atom, as the number of protons provides the atomic number. The electron cloud is just as important and is responsible for the properties that the atom shows. Describing the location of electrons as a “cloud” is somewhat correct, as the movement of electrons is wavelike, which results in the electrons moving around a very large space around the nucleus in a somewhat unpredictable way. However, despite this unpredictability there are some rules that are observed for electrons. This is largely determined by the quantization of energy for an electron. Classical thinking would believe that electrons can speed up or slow down any way they want and therefore change their energy in any way they want. However, electrons are limited to specific speeds and therefore energies around the nucleus. A way to think about how the energy is limited relating this to spare change. Coins are an example of “quantized” amounts of money. For example, a penny is worth 1¢, a nickel is worth 5¢, a dime is worth 10¢, and a quarter is worth 25¢. There is no coin for 3¢ or for 8¢. There is an additional concept with electrons with respect to occupation of orbitals. This has to do with how many electrons are allowed at a certain energy level. This is sort of like saying you are only allowed 2 pennies in your pocket and 5 dimes. Addition rules would say you can only put nickels in your pocket after you have 2 pennies in there first. Your text goes through this in detail, but a quick summery of the rules are as follows: Rule 1: Electrons occupy the lowest-energy orbitals available, beginning with 1s. (Aufbau Principle) Rule 2: Each Orbital can hold only two electrons, which must be opposite spin. Rule 3: Two or more orbitals with the same energy are each half-filled by one electron before any one orbital is completely filled by the addition of a second electron. (Hund’s rule) In order to apply the first rule, we need to know how the energy levels come together. We can reference Model 1 to show the order of the energy levels. 1

Model 1: Orbital Energy Diagram This doesn’t show the total picture as each subshell consist of a single orbital or multiple orbitals. Each orbital can hold two electrons (one spin up and one spin down). The label of the subshell will indicate how many orbitals that subshell should have. Orbitals in each subshell Subshell # of orbitals # of electrons s 1 2 p 3 6 d 5 10 f 7 14 We use a simplified method to describe the distribution of electrons in an atom. This is called electron configuration . In an electron configuration we list each subshell with respect to the shell it is in and number of electrons in it. So an atom with an electron configuration of 1 s 2 2s 2 2p 5 would have 2 electrons in the s subshell that is found in shell 1 , 2 electrons in the s subshell and 5 electrons in the p subshell found in shell 2. To further determine how we come up with electron configurations complete the following questions. 1. Draw boxes representing orbitals next to each subshell in Model 1 up to 3d. One box for each orbital. The first three subshells have been done for you. 2. How many subshells does shell number 1 have? __1___ How many total orbitals? _1____ 3. How many subshells does shell number 2 have? _2___ How many total orbitals? __4___ 4. How many subshells does shell number 3 have? ___3__ How many total orbitals? ___9__ 5. How many total electrons can the first shell hold? 2 6. How many total electrons can the second shell hold? 2

8 How many total electrons can the third shell hold? 18 7. Discuss with your team members, are there more electrons in the lower energy levels or higher energy levels? Lower because they fill up faster 8. How many electrons does a carbon atom have? 6 electrons a. How many orbitals are required to hold all the electrons in a carbon atom? 4 b. Starting from the first shell in and proceeding in order that orbitals are listed in your diagram, work together as a team to determine which orbitals would hold the electrons in a carbon atom. Write the electron configuration for a neutral carbon atom. 1s^2 2s^2 2p^2 c. Did you follow the three rules discussed earlier to determine this electron configuration? Yep 9. Using this same method, fill in the boxes with electrons for an element that has 15 electrons. a. What is the electron configuration for an atom with 15 electrons? 1s^2 2s^2 2p^6 3s^2 3p^3 Information I Energy levels start close to the nucleus and then increase in energy as they get farther away from the nucleus. So, the higher the energy, the farther the orbital is from the nucleus. We can also say that the farther you get from the nucleus, the more space there will be for electrons to move without “bumping into each other”. So, having more electrons in higher energy levels seems very reasonable. Quantum mechanics can show why these energy levels are distributed the way they are and have the population they have. We are not going to worry about that, and just observe the final product. Valence Electrons Learning Objective 2.8 : Identify the valence shell electrons for an atom, and which subshell of electrons (s, p, d, f) correlate with which groups in the periodic table. Earlier in this activity it was mentioned that the electrons provide the properties of a given atom. We can be even mores specific and say that the outermost shell of electrons provides the properties that result in how an atom behaves in nature. Please note, that it is the outermost shell, not subshell that we need to worry about. This outermost shell is called the valence shell and the electrons present there are called the valence electrons . These electrons can interact with other valence electrons in other atoms to result in bonds being formed. 3

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