Chapter 5 Notes: Molecules and Compounds
Chapter 5 Notes: Molecules and Compounds CH 101
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This 5 page Class Notes was uploaded by Rebecca de la O on Monday October 3, 2016. The Class Notes belongs to CH 101 at University of Alabama - Tuscaloosa taught by Jared Allred in Fall 2016. Since its upload, it has received 4 views.
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Date Created: 10/03/16
Chapter 5: Molecules and Compounds - Chemical bond describes the forces that hold atoms together o Strong bonds = lower in energy and more stable (between compounds or molecules) - Ionic bond = held together because 2 atoms with opposite signs are attracted to each other + - o Ex. NaCl Na + Cl o 1 atom needs a large negative electron affinity o The other must have a small ionization energy o Nonmetal + metal - Covalent bond = 2 nonmetals get close, neither wants to give up an electron because both have large -EA and large IE so they share their valence electrons o Ex. H 2, CO ,2C H2 6 Chemical Formula - Describes the ratio of elements in a compound o Ex. H 2 1 oxygen atom 2 hydrogen atoms - Molecule: discrete, covalently bonded with a finite number of atoms o Nonmolecular NaCl - Molecular formula gives the actual number of atoms in a molecule - Empirical formula gives the relative number of atoms in a compound - Ex. C 2 6 molecular CH 3 emipircal C H O molecular CH O emipircal 6 12 6 2 C 3 8 molecular C3H 8 emipircal Structural Formula - Explains arrangement in a molecule - Ex. H 2 2 H -O-O-H CO 2 O = C = O Ionic Formula - Charge must be balanced - Ex. Ca 2+ F- Ca 2+ 2F - CaF 2 o Ca has 2+ because it wants to lose 2 electrons to be stable o F has - because it wants to gain 1 electron to be stable o You need two 1- from the fluorine to balance the 2+ charge Naming Ionic Compunds - Metal name + nonmetal name + “ide” - Ex. NaCl sodium chloride Al2O 3 aluminum oxide ScN scandium nitride - If an element can easily gain or lose different number electrons to be stable (like transition metals), use roman numerals to denote how many electrons are involved - Ex. Ti3+ O Ti 2 3titanium (III) oxide 4+ 2- Ti O Ti 2 4TiO 2itanium (IV) oxide Polyatomic Ions 2- - Ions with multiple atoms (Ex. CO 3 - carbonate) - Most stable form is named with suffix “ate” 2- - Lose 1 oxygen add “ite” (Ex. CO 2 - carbonite) - Lose 2 oxygen add “hypo” as a prefix and “ite” (Ex. CO 2-- hypocarbonite) 2+ - Gain 1 oxygen add “per” as a prefix (Ex. CO 4 - percarbonate) - Learn this table: IONS STABLE LOSE 1 LOSE 2 GAIN 1 OXYGEN OXYGEN OXYGEN 2- 2- 2- 2- 2+ CO 3 CO 3 CO 2 CO CO 4 carbonate Carbonite hypocarbonit percarbonite e - - - - + NO 3 NO 3 NO 2 NO NO 4 nitrate nitrite hyponitrite pernitrate 2- 2- 2- 2- 2+ SO 4 SO 4 SO 3 SO 2 SO 5 sulfate sulfite hyposulfite persulfate - - - - + ClO 3 ClO 3 ClO 2 ClO ClO4 Chlorate chlorite hypochlorite Perchlorate 3- 3- 3- 3- 3+ PO 4 PO 4 PO 3 PO 2 PO5 phosphate Phosphite hypophosphit perphosphate e Also know these: OH - O 22- H2O 22- NH 4+ MnO 4- CN - CrO - 4 hydroxid peroxide Hydroge Ammoniu Permangana Cyanide Dichromat e n m te e peroxide Example Test Question: What is the formula of calcium nitride? a) Ca 3 2correct b) CaN incorrect because Ca has a 2+ charge and N has a 3- charge and they must balance c) Ca(NO )2 2incorrect because there is an oxygen d) Ca(NO )3 2incorrect because there is an oxygen e) Ca(NO )3 incorrect because there is an oxygen Hydrates - When water is part of the ionic compound - To name, add “hydrate” at the end and the number of H O mole2ules as the prefix - Ex. CaCl 2 6H O2 calcium chloride hexadrate Prefixes Naming Molecules - Different from ionic because they don’t have charges to tell you how to balance - Numerical prefix + first element numerical prefix + second element + “ide” o Omit prefix “mono” for the first element - Ex. CO carbon monoxide CO 2 carbon dioxide H 2 hydrogen dioxide S8 octasulfur CCl 4 carbon tetrachloride P2O 5 diphosphorous pentaoxide Using Formulas to Convert to Mass and Moles - Formula mass = the mass per mole of a molecule, compound, or ion - Formula mass = ƩnA W sum of the atomic weight times the number of atoms of each element - Ex. H 2 (2 x 1.008 g/mol) + (1 x 16.00 g/mol) = 18.016 g/mol MgO (1 x 24.31 g/mol) + (1 x 16.00 g/mol) = 40.31 g/mol P 4 x 30.97 g/mol = 123.88 g/mol 4 Percentage by Mass Ex. SF 6 sulfur hexafluoride a) What is the percentage of fluoride by atoms? 6 fluorine atoms + 1 sulfur atom = 7 total atoms part 6 = x100 x100 whole 7 = 85.714 % are F atoms b) What is the percent of fluorine by mass? massof allthe fluorine x100 % = massof thecompound Mass of all the fluorine = 6 x 19.00g/mol (atomic mass) = 114.0 g/mol Mass of SF = (1 x 32.06 g/mol) + (6 x 19.00g/mol) = 146.9 g/mol 6 114g/mol 146.9g/mol x100 = 78.03 % F * Strategy: keep track of the # of atoms and convert to mass only when needed Ex. In Na2CO 3 for every formula unit there are 2 units of Na, 1 unit of C, and 3 units of O (look at subscripts) If we use units of mol: 1 mol Na 2O h3s 2 mol Na, 1 mol C, and 3 mol O 2- + 1 mol CO 3 and 2 mol Na Formula mass of Na CO2= (3 x 22.99 g/mol) + (1 x 12.01 g/mol) + (3 x 16.00 g/mol) = 105.99 g/mol If we have 12.6 g Na C2 , 3ow many grams of O are there? 2 ways to solve this: 1) Molar ratio of Na 2O :3O 1:3 massof Na2CO3 # mol Na C2 = 3 formula mass ¿molof Na2CO3 1 ¿molO = 3 Mass of O = atomic mass x # of mol massof Na2CO3 = atomic mass of O x 3 x formula mass 3xatomicmass of O 2) Mass % of O = formulamassof Na2CO3 3 x atomicmass of O Mass of O = mass of Na CO2x 3 formulamassof Na2CO 3 3x16.00g/mol = 12.6 x 105.99g/mol = 5.71 g O Lewis Structure - Non-quantum mechanical (nothing to do with electrons being in orbitals) - Based on number of valence electrons (look at their group/column number) - Draw the elemental symbol and put dots around it (dots correspond to # of valence electrons) Lithium has 1 valence electron Nitrogen has 5 valence electrons Neon has a full octet - To draw the Lewis structure of ionic compounds: Magnesium oxide MgO Beryllium bromide BeBr 2
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