CEM 141, Notes, week 3
CEM 141, Notes, week 3 Cem 141
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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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