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Chem 1314 Chapter 8

by: Morgan Walker

Chem 1314 Chapter 8 Chem 1314

Morgan Walker
OK State
GPA 3.2

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About this Document

These notes cover the material from chapter 8
General Chemistry
Dr. Jimmie Weaver
Class Notes
General Chemistry
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This 2 page Class Notes was uploaded by Morgan Walker on Thursday March 31, 2016. The Class Notes belongs to Chem 1314 at Oklahoma State University taught by Dr. Jimmie Weaver in Winter 2016. Since its upload, it has received 20 views. For similar materials see General Chemistry in Chemistry at Oklahoma State University.


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Date Created: 03/31/16
Chapter 8 Periodic Properties of Elements Periodic Properties  Predicted by how the element is placed on the periodic table  Include the atomic radius, density and electron count as well as its metallic character Periodic Table Development  Developed by Mendeleev, in order of increasing atomic mass from left to right that way elements of the same property fell in the same column  Quantum Mechanics describes it by describing how the electrons fill the orbitals inside the atoms of the elements Electron Configuration  Shows which orbitals contain an electron in the atom 2 o He 1s means that Helium’s two electrons are in the s orbital  each orbital can hold only two electrons that have opposite spins (Pauli exclusion principle) o one spins up designated by a half up arrow o one spins down designated by half down arrow o  Hund’s rule states that the orbitals fill up with one electron before going to pairing o If you have a s orbital, p orbital and f orbital you fill all the way up to the f orbital and then with the leftover electrons fill it up again.  Filling- 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s  EXECPTIONS Cr Cu Ni Mo Ag Pd W Au Pt Sg Rg Ds 1 5 1 10 0 10 4s 3d 4s 3d 4s 3d  These have a different configuration that usual because it would take less energy to be in the d orbital than the s orbital. Electron Configuration and Periodic Table  Since the orbitals fill up by increasing atomic number we can determine an elements electron configuration by where it is positioned on the periodic table  When all the principal energy levels are full the configuration is most stable o The noble gases are the most stable since they have the lowest energy configuration  An element with more than one valence electron it will be the most active metals since they lose their electrons to achieve a noble gas configuration  An element with 6-7 valence electrons are the most active non-metals, they will gain electrons to achieve noble gas configuration  Valence electron- electron in the highest principal energy level 2 5 o Any electron in the last two levels. If an element ends in 4s 3p the electron will have 2 valence electrons because of the 4s 2  Core electron- electron in the lower energy level o All the other sublevels counted up. (you can subtract valence from number of electrons to get core electrons) Nuclear Charge and Trends in Atomic Size  Atom size is effected very heavily on its outermost electrons o the further down the periodic table we go the more orbitals we get therefore the larger the atoms radius is  atomic radius decreases when we move across a row in the periodic table because the net charge of the outermost electrons increases o when we move across the row we get more electrons and they pair up, therefore decreasing the radius  transition elements stay the same since electrons are added to the n(highest)-1 orbital while the total number of electrons stays the same Properties of Ions  if an ion has a charge you can determine the configuration by adding or subtracting the number of electrons to the atoms neutral configuration +3 o Cl would have three less electrons than Cl  The main group of ions you would remove the electrons in the order that you added them in the configuration o Ended with 3p in Cl, end in 3p for Cl+3  For the transition metals the electrons in the s orbital are removed before the d orbital 2 3 0 1 +3 o 4s 3d for V, 4s 3d for V  Cations have a much larger radius while anions have a much smaller radius compared to their neutral atom  Ionization energy will decrease as you move down the column and increases as you move across the row Halogens, Alkali Metals and Noble Gases  Halogens have one fewer electron than noble gases and will want to gain one in order to get the Noble Gas configuration  Alkali Metals have one more electron than Noble Gases and will lose it to be at Noble Gas configuration  Noble gasses are unreactive because their electron configuration is the most stable Info from Pearson Mastering Chemistry


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