Ch. 9 - Electron Organization. The Bohr Model [9.4] Orbitals [9.5, 9.6] Counting Electrons, configurations [9.7]

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1 Ch. 9 - Electron Organization The Bohr Model [9.4] Orbitals [9.5, 9.6] Counting Electrons, configurations [9.7] Predicting ion charges from electron configurations. CHEM 100 F07 1 Organization of Electrons Early ideas of electron organization pointed to the cloud - nucleus in the center, electrons freely floating around. If this model were true, an excited element would emit a full spectrum of light. CHEM 100 F07 2 1

2 Organization of Electrons We have observed that when elements are excited, they emit discrete lines of light. The emission is caused by electrons moving about. The discrete lines indicate additional organization of electrons - it s not just a cloud! CHEM 100 F07 3 Organization of Electrons - The Bohr Model Developed in 1913 by Niels Bohr Electrons lie in discrete, spherical orbitals (n level) around the central nucleus. Postulates: Energy-level: An electron can only have specific energy values in an atom. Transitions: An electron can change energy only by going from one energy level to another. Over time, modifications to the Bohr Model led to the organization of electrons in orbitals. CHEM 100 F07 4 2

3 The s orbital The Orbitals CHEM 100 F07 5 The p orbitals The Orbitals CHEM 100 F07 6 3

4 The d orbitals The Orbitals For a cool representation of all of the orbitals from n = 0 thru n = 10, check out: The Grand Orbital Table CHEM 100 F07 7 Filling in Electrons The Rules: Aufbau Principle: Electrons fill in electrons by order of energy, from low high. Not all orbitals are available for all energy (n) levels. Pauli Exclusion Principle: Each individual orbital takes 2 electrons only! There is one s orbital = 2 electrons. There are three p orbitals = 6 electrons. There are five d orbitals = 10 electrons. There are seven f orbitals = 14 electrons. Hund s Rule: If there are multiple orbitals at the same energy, they fill singly first, before electrons pair. CHEM 100 F07 8 4

5 n = 1 n = 2 n = 3 Filling in Electrons 1s orbital 2 electrons 2s orbital 2 electrons 2p orbital 6 electrons 3s orbital 2 electrons 3p orbital 6 electrons 3d orbital 10 electrons Lower energy rows have fewer orbitals available, therefore there are fewer elements there! n = 4 4s orbital 2 electrons 4p orbital 6 electrons 4d orbital 10 electrons 4f orbital 14 electrons Beyond n = 4, all levels have s, p, d and f orbitals. CHEM 100 F07 9 Filling in Electrons Electrons get filled into orbitals individually: s: unoccupied orbital orbital with 1 electron orbital with 2 electrons The Pauli Exclusion Principle! p: One electron Two electrons Three electrons Four electrons Five electrons Six electrons Hund s Rule: fill orbitals singly first, then start pairing! CHEM 100 F

6 The Aufbau Principle As electrons get added to elements, the get inserted into the orbitals in order of energy. This is not in numerical order! The diagram at right shows the order that electrons fill. To create the diagram: List the orbitals in order. Then, draw diagonal lines downward from right to left. Once you complete a diagonal, loop back around. 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s CHEM 100 F07 11 Summary of Electron-filling Rules The Pauli Exclusion Principle Electrons are like spinning magnets, and must have opposite alignment. We imagine this as one and one. The Aufbau Principle The orbitals get filled from the lowest energy to the highest energy. Hund s Rule When multiple orbitals are present, each orbital gets filled singly at first, and pairing begins. The single electrons all have the same alignment. CHEM 100 F

7 Writing Electronic Configurations To determine the electron configuration: 1) Find the number of electrons for the element. 2) Fill the electrons in order of the Aufbau Principle. 3) Use Hund s Rule and the Pauli Exclusion Principle for orbital diagrams. Example: Nitrogen - Element #7 7 electrons Orbital Diagram 1s 2s 2p 2p 2p Electronic Configuration The number of electrons within each set of orbitals 1s 2 2s 2 2p 2 Each individual orbital gets a box. Electrons are filled into the boxes until the total is reached. The energy, or n level The orbital CHEM 100 F07 13 Electron Configurations Determine the orbital diagrams and electron configurations for the following elements. He Li C F Mg P CHEM 100 F

8 Electron Configurations Determine the orbital diagrams and electron configurations for the following elements. He Li 6e - C 1s 2 2s 2 2p 2 F Mg P 1s 1s 1s 2 1s 2 2s 1 2s 1s 2s 2p 2p 2p 1s 2s 2p 2p 2p 1s 2s 2p 2p 2p 1s 2 2s 2 2p 5 1s 2s 2p 2p 2p 3s 3p 3p 3p 3s 1s 2 2s 2 2p 6 3s 2 1s 2 2s 2 2p 6 3s 2 3p 3 CHEM 100 F07 15 Electron Configurations and the Periodic Table Electron configurations can be read off of the periodic table CHEM 100 F

9 Electronic Configurations Use the periodic table to determine the electronic configurations for the following elements. Be S 16 e - 1s 2 2s 2 2p 6 3s 2 3p 4 Ca V Ge CHEM 100 F07 17 Electronic Configuration Shorthand Consider the electronic for Argon and Calcium: Ar: 1s 2 2s 2 2p 6 3s 2 3p6 As a noble gas, Argon s orbitals are completely filled. Ca: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 [Ar] 4s 2 We can use the last noble gas as a shorthand in electronic configurations! The core electrons The valence electrons Not only are shorthand configurations easier to write, but they identify the valence electrons, which are the electrons that are available for reaction! CHEM 100 F

10 Electronic Configuration Shorthand Use the shorthand notation to write the electronic configurations for the following elements. Also, indicate the number of valence electrons for each. K Mn As [Ar] 4s 2 3d 10 4p 3 15 or 5 valence electrons Pd In Cs CHEM 100 F07 19 Electronic Configurations for Ions Let s consider calcium: [Ar]4s 2 What is the typical charge on a Calcium ion? Are electrons removed or gained for a cation? How many valence electrons does calcium have? What conclusion can you make from these responses? CHEM 100 F

11 Electronic Configurations for Cations A cation has fewer electrons than the neutral atom. These electrons are removed from the highest n level first! Examples: Al: [Ne] 3s 2 3p 1 Al +3 : Sn: [Kr] 5s 2 4d 10 5p 2 Sn +2 : Sn +4 : What do you notice about the resulting cation electron configurations? CHEM 100 F07 21 Electronic Configurations for Anions An anion has more electrons than the neutral atom. These electrons are added to the atom according to the Aufbau Principle. Examples: P: [Ne] 3s 2 3p 3 P -3 : Br: [Ar] 4s 2 3d 10 4p 5 Br -1 : What do you notice about the resulting anion electron configurations? How does this conclusion compare with the cation configurations? CHEM 100 F

12 Electronic Configurations for Ions Write the electronic configurations for the following elements and their ions: Mg/Mg +2 : Fe/Fe +2 /Fe +3 :Fe [Ar]4s 2 3d 6 Fe +2 [Ar]3d 6 O/O -2 : Fe +3 [Ar]3d 5 Mn/Mn +2 /Mn +7 : CHEM 100 F07 23 Electronic Configuration Summary For neutral atoms, the number of electrons = number of protons = atomic number. Electrons are inserted according to the Pauli Exclusion Principle, the Aufbau Principle and Hund s Rule. Configurations using the shorthand rely on the previous noble gas and the valence electrons. For cations, electrons are removed from at atom, typically to reach a noble gas configuration, or other stable point. For anions, electrons are added to an atom to reach a noble gas configuration. CHEM 100 F

13 CHEM 100 F07 25 Periodic Trends Based on similarities in electronic configurations within a group and a period, some generalizations and trends can be predicted. Atomic Size Ionization Energy CHEM 100 F

14 Atomic Size This trend typically looks at the atomic radius, but also applies to the volume. Each atom is treated as a marble - a hard sphere. The size is loosely based on the valence shell, or the n-level The radius increases as you go down a group. The size is determined by the number of protons available to attract the electrons. The radius decreases as you across a period. CHEM 100 F07 27 Atomic Size CHEM 100 F

15 Atomic Size Choose the larger atom in each pair: C or O Li or K P or Al Br or I CHEM 100 F07 29 Ionization Energy Defined as the amount of energy required to remove an electron from an atom in the gas phase. M (g) + IE M +1 (g) + e- (g) Electrons are always removed from the valence shell. The IE tends to decrease down a group. Electrons get further away from the nucleus! The IE tends to increase across a period. Electrons are closer to the nucleus! CHEM 100 F

16 Ionization Energy Choose the atom with the lower ionization energy in each pair: C or O Li or K P or Al Br or I CHEM 100 F07 31 Electron Affinity Defined as the amount of energy released from an atom in the gas phase when an electron is captured. M (g) + e - M -1 (g) + EA Electrons are always added to the last unfilled orbital. The EA tends to decrease down a group. The electron spends less time near the nucleus. The EA tends to increase across a period. The addition of an electron gets the atom closer to a noble gas. CHEM 100 F

17 Electron Affinity Choose the atom with the higher in each pair: N or F Ca or Mg P or Al Br or I CHEM 100 F07 33 Periodic Trend Summary EA increases IE increases Radius decreases Radius increases IE decreases EA decreases CHEM 100 F

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