Chapter 4 Arrangement of Electrons in Atoms Section 3: Electron Configuration Objectives: Be able to define: Aufbau Principle, Pauli Exclusion Principle, Hund s rule. Be able to list the number of electrons it takes to fill each energy level. Be able to describe the electron configurations for the atoms of any element using orbital notation, electron configuration notation, and, when appropriate, noble-gas notation. Electrons are added to the orbitals one by one according to three basic rules. 1. principle 2. exclusion principle 3. rule First rule shows the in which electrons occupy orbitals. According to the Aufbau principle, an electron occupies the -energy orbital that can receive it. (Lower energy orbitals are to the of the atom). Aufbau Principle fill from the up. The second rule reflects the importance of the quantum number. The Pauli Exclusion Principle, no electrons in the same atom can have the set of quantum numbers. The third rule requires placing as many electrons as possible in orbitals in the same sublevel to minimize electron-electron repulsion. Hund s rule states that orbitals of energy are each occupied by electron before any orbital is occupied by a electron, and all electrons in singly occupied orbitals must have the spin state. If multiple orbitals have the energy, one electron goes into of them before they start to up. Aufbau House Story Where are the s, p, d, f orbitals located on the periodic table? Identify them on the diagram below.
Diagonal Rule Steps: 1. Write the energy levels to bottom. 2. Write the orbitals in s, p, d, f order. Write the same number of as the energy level. 3. Draw lines from the top right to the bottom left. To get the correct order, follow the. 1s 2s 2 3s 3 4s 4 5s 5 6 6 7 7 Why are d and f orbitals always in lower energy levels? The and orbitals require amounts of. It is better (lower in energy) to skip a that requires a large amount of energy (d and f orbtials) for one in a level but lower energy. This is the reason for the diagonal rule! BE SURE TO FOLLOW THE ARROWS IN ORDER! How many electrons can be in a sublevel? Remember: A maximum of electrons can be placed in an. Number of orbitals s orbitals p orbitals d orbitals f orbitals Number of electrons The arrangement of electrons in an atom is known as the atom s configuration. Since atoms of different elements have numbers of electrons, a electron configuration exists for the of element. Electrons in atoms tend to assume arrangements that have the possible energies. The lowest energy arrangement of the electron for each element is call the electron s - state configuration.
Electron Configurations A list of the electrons in an atom (or ion) Must go in (Aufbau principle) 2 electrons per orbital, We need electron configurations so that we can determine the number of electrons in the energy level. These are called electrons. The number of electrons determines how and this atom (or ion) can to in order to make a molecule. Example: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 etc. Electron Configurations 2p 4 Electron configuration practice: Complete an electron orbital diagram for: 1. 10 electrons 2. 24 electrons 3. 13 electrons 4. 3 electrons 5. 5 electrons 6. 26 electrons 7. 22 electrons 8. 17 electrons 9. H 10. Li 11. N 12. Ne 13. K 14. Zn 15. Pb
Noble Gas Configuration and Notation Noble Gases are in group of the Periodic Table. They contain valence shells. The Noble Gas Notation is when a Noble gas is used to represent the core (inner) and the shell is shown. Example: Bromine (Br) has an atomic number of 35. Br electron configuration: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 Noble gas notation: [Ar] 4s 2 3d 10 4p 5 The [Ar] represents the core electrons and only the valence electrons are shown Which Noble Gas Do You Choose? Think of Price is Right. You want to get as close as possible without going over the atomic number you are wanting. Noble Gas # of electrons He Ne Ar Kr Xe Example: Chlorine Noble Gas Notation Example 1. Determine the number of electrons to place 2. Determine which noble gas to use 3. Start where the noble gas left off and write electron configuration notation for the valence electrons Example: As
Practice Shorthand Notation Write the shorthand notation for each of the following atoms: 1. Cl 2. K 3. Ca 4. I 5. Bi Valence Electrons Electrons are divided between and electrons. B Core = [ ], valence = Br Core = [ ] valence = No. of valence electrons of a main group atom = Group number (for A groups) Atoms like to either or their outermost level. Since the outer level contains s electrons and p electrons (d & f are always in lower levels), the optimum number of electrons is. This is called the rule. Valence electrons: Electrons in an atom s highest shell (s and p electrons) and in partially filled subshells of lower shells d or f electrons). Example 1: K (potassium) has the electron configuration 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 Its highest occupied shell is and it has one electron in the 4s orbital so it has valence electron. Example 2: S (sulfur) has the electron configuration 1s 2 2s 2 2p 6 3s 2 3p 4. Its highest occupied shell is and it has electrons in the 3s orbital and electrons in the 3p orbital. Sulfur has total valence electrons. Example 3: Co (cobalt) has the electron configuration 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7. Its highest occupied shell is 4 and it has two electrons in the 4s orbital. Cobalt also has electrons in a partially filled 3d orbital so it has a total of valence electrons.
Example 4: Se (selenium) has the electron configuration 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 4. Its highest occupied shell is 4. Selenium has two electrons in the 4s orbital and four electrons in the 4p orbital. Selenium has a full 3d orbital so these ten electrons are valence electrons. Selenium has total valence electrons. How many valence electrons do the following elements have? F Sb Nd Mo Rb Ag Keep an Eye On Those Ions! Electrons are lost or gained like they always are with ions ions have electrons, ions have electrons. (Remember electrons have a negative charge.) The electrons that are lost or gained should be added/removed from the energy level (not the highest orbital in energy!) Example 1: Tin (Sn) atomic number Atom: [Kr] 5s 2 4d 10 5p 2 Sn +4 ion: [Kr] 4d 10 Sn +2 ion: [Kr] 5s 2 4d 10 Note that the electrons came out of the highest energy level, not the highest energy orbital! Example 2: Bromine (Br) atomic number Atom: [Ar] 4s 3d 4p Br -1 ion: [Ar] 4s 3d 4p Note that the electrons went into the highest energy level, not the highest energy orbital!
Practice: Write the longhand notation for these: F - Li + Mg +2 Write the shorthand notation for these: Br - Ba +2 Al +3 Two exceptions to the Aufbau principle remember Remember d and f orbitals require LARGE amounts of energy. If we can t fill these sublevels, then the next best thing is to be full (one electron in each orbital in the sublevel). The most common exceptions are d 4 and d 9. Copper and Chromium 3d 4. Move 1 4s to 4d. 3d likes to be half full or completely filled. d 4 is one electron short of being HALF full. In order to become more (require less energy), one of the closest s electrons will actually go into the d, making it d 5 instead of d 4. For example: Cr would be [Ar] 4s 2 3d 4, but since this ends with a d 4 it is an exception to the rule. Thus, Cr should be [Ar] 4s 1 3d 5. Procedure: Find the closest orbital. Steal electron from it, and add it to the. OK, so this helps the d, but what about the poor s orbital that loses an electron? Remember, half full is good and when an s loses 1, it too becomes full! So having the s half full and the d half full is usually in energy than having the s full and the d to have one empty. Write the shorthand notation for: Cu
Orbital Diagrams Orbital diagrams are representation of an electron configuration. One arrow represents one and shows and which within a sublevel. Rules: 1. Same rules as before (Aufbau principle, d 4 and d 9 exceptions, two electrons in each orbital, etc. etc.) 2. One additional rule: Hund s Rule Monopoly Rule In orbitals of EQUAL ENERGY (p, d, and f), place electron in each orbital before making any pairs All single electrons must spin the way Example 1: Lithium atomic number electron configuration total electrons 3s 2s 3p 2p Example 2: 1s Carbon atomic number electron configuration total electrons. Here we see for the first time HUND S RULE. When placing electrons in a set of orbitals having the energy, we place them as long as possible. 3s 2s _ 3p 2p 1s
Draw these orbital diagrams 1. Oxygen (O) 2. Chromium (Cr) Ion Configurations To form from elements, add 1 or more e- from the highest sublevel. P [Ne] 3s 2 3p 3 + 3e- ---> P 3- [Ne] 3s 2 3p 6 or [Ar] 3s 3p 3s 3p 2s 2p 2s 2p 1s 1s Practice: Chlorine (Cl) 1s 2s 2p 3s 3p Fluorine ion (F-) 1s 2s 2p 3s 3p
Summarizer Which of the rules do each of the following electron energy diagrams violate? (Circle the violation.) a. b. c. Write out the following for Calcium. 1- Electron configuration 2- Shorthand notation (noble gas notation) 3 Orbital diagram