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CEM 141, Notes, week 3

by: Leah DiCiesare

CEM 141, Notes, week 3 Cem 141

Leah DiCiesare

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About this Document

These are the notes from the third week of lectures, 9/19-9/23, as well as the notes from 1.8 of the textbook.
General Chemistry
J. Hu
Class Notes
interactions, atoms, Molecules, London dispersion forces, covalent bonds
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This 4 page Class Notes was uploaded by Leah DiCiesare on Saturday September 24, 2016. The Class Notes belongs to Cem 141 at Michigan State University taught by J. Hu in Fall 2016. Since its upload, it has received 25 views. For similar materials see General Chemistry in Chemistry at Michigan State University.


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Date Created: 09/24/16
Textbook Notes: week 3  1.8 Interactions between Helium Atoms and Hydrogen Molecules o When He atoms approach each other the LDFs act upon their interaction; these  forces are so weak though that the atoms do not stick together o When H atoms approach each other they form a very stable interaction (~1000x  stronger than He­He LDFs)  H­H interaction, the atoms are held together by attraction of each nucleus  for both electrons   Called a covalent bond  Have to supply energy to break covalent bond  H 2molecules are attracted to each other with LDFs o Strength of LDFs depends on: shape of molecule, surface area, and number of  electrons  The further electrons are from the nucleus, the more polarizable (floppy)  the electron cloud becomes, which makes the LDFs stronger Lecture Notes: 9/19­9/23  Interaction Graph   o The deeper the potential energy well (the potential energy minimum), the more  energy needed to break the interaction; depth tells how strong the interaction is o Δ E = the amount of energy needed to separate the atoms; this is found from: the PE is close to zero at the flattening out part of the graph minus the PE at the potential  energy minimum o The potential energy minimum is where the interaction is the most stable and the  attractive forces equal the repulsive forces o For noble gases, the Van der Waals (LDF) radius = 1/2 the distance between atomic  nuclei at the potential energy well (the x­axis value/the position of the atoms that  lines up with the potential energy minimum) o Electron cloud of Xe is "floppier", more polarizable, than He because it is bigger o London Dispersion Forces Increases with size of particle (number of electrons) Lecture Notes: 9/19­9/23  Increases with surface area  Part of a variety of intermolecular forces (between particles) o Larger atoms have higher melting and boiling points  Formation of Covalent Bonds o When two H atoms approach, they are more strongly attracted than two He atoms o Form a covalent bond o PE curve of H has a deeper well; H­H interaction is stronger o Internuclear distance between H atoms is shorter when interaction is most stable o Electron clouds are overlapping but are still stable o The nuclei share the electrons o H2 ­ hydrogen molecule, has different properties that H; new chemical species o Hydrogen usually exists as H2 ­ diatomic molecule o Total of 7 naturally occurring diatomic molecules: B2 ,2I ,2N ,2Cl 2 H2, O2, F   (BrINClHOF ­ acronym for Bromine, Iodine, Nitrogen, Chlorine, Hydrogen, Oxygen, and Flourine)  IMF's and Bonds o Intermolecular Forces (IMFs) act between neighboring particles; Van der Waals  interaction (catch­all term for weak interactions) ­ between atoms or neutral  molecules o Bonds are more permanent; stronger and harder to break ­ within a molecule o Breaking bonds requires the input of energy (to the system) ­ making it less stable o Forming bonds releases energy (from the system) ­ making it more stable Stoichiometry  Stoichiometry: relationship between macroscopic and molecular level  Allows us to calculate how much "stuff" can be produced in a reaction  Ex. CH  4g) + 2O  2g) ­> CO  2g) + 2H 2 o 1 molecule CH ,42 molecules O ,21 molecule CO ,22 molecules H O 2  Can't measure single molecules so we use the mole to connect molecular level to  macroscopic level which is measurable  Mole: a really big number; Avogardo's Number: 6.022*10  "things"  Stoichiometric coefficients establish the mole ratio between reactants and products o Ex. CH 4(g) + 2O 2(g) ­> CO 2(g) + 2H 2 1 mole CH , 2 moles O , 1 mole CO , 2 moles H O  4 2 2 2 o Mole ratio can be used to convert the mass or moles of one reactant (or product)  to the mass or moles of another reactant (or product) in the reaction  Molar Mass = mass in grams of 1 mole of a substance (g/mol)  Can be calculated by adding up the atomic masses from the periodic mass o Ex. 1 mole H 2 = 18 grams: O = 16 g + H =1 g + H = 1 g = 18 grams o Ex. NaOH: Na = 23 g, O = 16 g, H = 1 g, 23+16+1=40 g/mol o Ex. UF 6 U = 238 g, F = 19 g, (6*19) + 238 = 352 g/mol  Mass is conserved in a chemical reaction Lecture Notes: 9/19­9/23  Balancing reactions: same number of each kind of atom on each side of reactions; have to be integers o Ex. 2C 6 14 19O  2> 12CO  +214H O 2  Mass, Mole, and Molar Mass o Molar mass (g/mole) = mass (g) / mole (mol) o Mass (g) * 1/ molar mass (g/mol) = mole o Mole (mol) * molar mass (g/mol) = mass (g) o Ex. How many moles of calcium nitrate (Ca(NO ) 3 2re in 325 g?  Ca(NO 3 2= 164 g/mol  325g/164g/mol = 1.98 mol


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