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CHEM 110, Week 4 Notes

by: BritneyMoore

CHEM 110, Week 4 Notes CHEM 110

CSU Pomona
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These notes cover quantum mechanics, electronic energy levels and transferring energy, electronic configuration, periodic law, octet rule, cations, anions. These notes also discuss topics such as c...
Chemical Principles I
Dr. Hoda Mirafzal
Class Notes
General Chemistry




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This 6 page Class Notes was uploaded by BritneyMoore on Tuesday April 26, 2016. The Class Notes belongs to CHEM 110 at California State Polytechnic University taught by Dr. Hoda Mirafzal in Spring 2016. Since its upload, it has received 11 views. For similar materials see Chemical Principles I in Chemistry at California State Polytechnic University.

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Date Created: 04/26/16
CHEM 110 04/24/2016 ▯ Electronic Transitions continued  Red has the lowest energy level  Violet has the most energy  When an electron moves from the 2 ndto the 1 energy level, light is emitted st nd  When an electron moves from the 1 to 2 energy level, energy in the form of heat or light is absorbed  If the atom emits green light when an electron moves from the 3 to rd nd 2 energy level, itrds listly to emit blue light when an electron moves from the 3 to 1 energy level  Larger energy gaps, lose more energy and a more energetic light is emitted ▯ Quantum Mechanics  Describes the arrangement and space occupies by electrons in atoms  Boehr model is very simplistic and not all orbitals are circular ▯ Electronic energy levels (shells)  Contain electrons that are similar in energy and distance from the nucleus  Lowest energy electrons are closest to the nucleus  Identified by n=1,2,3,4…  The 1 shell (n=1) is lowest in energy nd  The 2 shell (n=2) is next lowest in energy  The quantum # n tells you about the energy of an electron and its distance from the nucleus  An electron’s energy is determines by the orbital # the are in  Max number of electrons in any electronic energy level= 2n^2 Shell 1 N=1 (2(1^2))= 2e- Shell 2 N=2 (2(2^2))= 8e- Shell 3 N=3 (2(3^2))= 18e- Shell 4 N=4 (2(4^2))= 32e-  All electrons in the same energy level have similar energy st  The order of filling energy levels for the 1 36 electrons: 1 2 3 4 3 4 2e- 8e- 8e- 2e- 10e- 6e- ▯ Electronic Configuration  Lists the shells containing electrons  Written Element 1, (2e-) 2, (8e-) 3, (8e-) 4, (2e-) 3, (10e-) 4, (6e-) He 2 C 2 4 F 2 7 Ne 2 8 Al 2 8 3 Fe 2 8 14 8 2 6  The number of valence electrons is consistent with the group number of the element ▯ Periodic Law  All the elements in a group have the same electronic configuration in their outermost shells  Same number of valence electrons  Elements in a group have similar properties ▯ Octet Rule  An octet in the outer shell makes atoms stable  Octet= 8 valence electrons  Very stable and do not want to participate in chemical reactions  Electrons are lost, gained, or shared to form an octet  Halogens and Alkalai metals are the MOST reactive ▯ Electron Dot Structures  Symbols of atoms with dots to represent the valence electrons ▯ Ion  An atom/ group of atoms that have a charge  Cation is positively charged (atoms lose e-)  Anion is negatively charged (atoms gain e-)  Ionic compounds are a cation bonded to an anion (metal and nonmetal bonded together) ▯ Cations  Metals lose their valence electrons and form cations but not all cations are metals  Often, the cations formed have a completed shell (octet rule) Na= 2, 8, 1 Group 1 metals 1 valence e- Ion + Na+= 2,8 Ca= 2, 8, 8, 2 Group 2 metals 2 valence e- Ion 2+ Ca 2+= 2, 8, 8 Al= 2, 8, 3 Group 13 metals 3 valence e- Ion 3+ Al 3+= 2, 8 ▯ Formation of a Sodium Ion Sodium Atom Sodium Ion Na Na + EC: 2, 8, 1 EC: 2, 8 (=Ne, Octet) 11p+ 11 p+  11e-  10 e- 0 Charge 1 p+ Charge ▯ Anions  Non-metals always gain valence electrons and form anions  The anions formed often have a completed shell  F= 2, 7 Group 17 7 valence e- Ion - F-= 2, 8 S= 2, 8, 6 Group 16 6 valence e- Ion 2- S 2-= 2, 8, 8 N= 2, 5 Group 15 5 valence e- Ion 3- N 3-= 2, 8 ▯ Formation of a Fluoride Ion Fluoride Atom Fluoride Ion F F- EC: 2, 7 EC: 2, 8 (=Ne, Octet) ▯ 9 p+ ▯ 9 p+  9 e-  10 e- 0 Charge 1- Charge ▯ Ionic Compound  Contains a cation bonded to an anion ▯ Formulas of ionic compounds  Ionic compound (metal + non-metal), (cation + anion)  Must be neutral  Have the simplest formula that balances the charges on the cation with the anion ▯ Writing and ionic formula  Write the formula for the ionic compound that will form between Ba 2+ and Cl- o Balance the charges: 1 st 2+ and 2 Cl- o Write the cation 1 and anion 2nd: Ba Cl o Write the number of ions needed as subscripts: BaCl2 ▯ Binary Ionic Compounds  Contains 2 different elements: a cation and an anion  Naming for ionic compounds that contain metals in groups: 1, 2, 3, Ag, Zn, and Cd  Name the metal 1 and the nonmetal as “-ide”  Ex: NaCl Sodium Chloride ▯ Variable Cations  Metals except groups 1, 2, Al, Zn, Cd, Ag usually form several different ions each with a different charge  Write the charge of the cation in Roman Numerals and in parenthesis after the name of the metal  Ex: PbCl2 (Pb 2+): lead (II) chloride ▯ Polyatomic Ions  A group of atoms bonded together with an overall charge.  Know the names and formulas and charges of the ions on the handout. ▯ Ternary Compounds  Contains at least 3 elements  If the subscript is 2 or more, write the polyatomic ion in parentheses in the formula  Ex: NaHCO3 Sodium bicarbonate (baking soda) ▯ Naming Ternary Compounds  Same rules as naming binary compounds (cation and anion) but use the names of the polyatomic ions where appropriate ▯ Chemical Bond  An attraction between 2 or more atoms due to an interaction between valence electrons ▯ Types of bonding  Ionic o Between a cation and an anion. One atom has lost electrons and one has gained electrons.  Covalent o 2 atoms share electrons  Polar covalent o 2 atoms share electrons unevenly. One atom gets more of the electron density and the other atom gets less. ▯ Electronegativity  A measure of how strongly an atom attracts electrons ▯ Bond Polarity  One atom’s electronegativity is much higher than the others  A really large electronegativity difference, about >2) ▯ Covalent Bonds  Formed between 2 atoms with similar electronegativity  No electronegativity difference ▯ Polar Covalent Bonds  Formed between 2 atoms with small differences in electronegativity ▯ Elements that are always found as covalent compounds  Memorize this list: o H2, N2, O2, F2, Cl2, Br2, I2, P4, S8 ▯ Binary Covalent Compounds  2 bonded non-metals  Naming: o Name each element o End the last element in “-ide” o Add prefixes to show the number of atoms but the 1 element st does npt get the mono prefix  Prefixes: Mono: 1 Di: 2 Tri: 3 Tetra: 4 Penta: 5 Hexa: 6 ▯ ▯


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