Chemistry 1210, Week 3 Notes
Chemistry 1210, Week 3 Notes CHEM 1210
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This 4 page Class Notes was uploaded by Grace Campbell on Friday September 9, 2016. The Class Notes belongs to CHEM 1210 at Ohio State University taught by in Fall 2016. Since its upload, it has received 4 views.
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Date Created: 09/09/16
CHEMISTRY 1210 WEEK 3 NOTES 9/7/16 – 9/9/16 Chapter 3: Stoichiometry: Calculations with Chemical Formulas and Equations A. Chemical Equations-‐ symbolic representation of a chemical reaction a. Coefficients-‐ indicate the number of atoms, molecules, or formula units of each substance that is involved in a reaction b. Each substance always written as: i. (s) = solid 1. most ionic substances are solids at room temperature 2. most elements are solids at room temperature ii. (g) = gas iii. (l) = liquid iv. (aq) = aqueous solution (dissolved in water) c. reactants-‐ chemical formulas to the left of the arrow d. products-‐ chemical formulas to the right of the arrow i. substances that were produced from the reaction e. Balancing Equations i. Final equation should contain the smallest possible whole-‐number coefficients ii. Never change subscripts, instead use a coefficient iii. The number of atoms of each element before the reaction has to equal the number of atoms of each element after the reaction iv. Charges on the left side of the equation must equal the charges on the right side of the equation v. Tips: 1 1. Balance the molecule that has the largest number of atoms of a single element first a. Excluding H, O, and polyatomic ions 2. Balance polyatomic ions as a whole unit (if they did not change) 3. Balance hydrogen and oxygen last a. If they are both in an equation, balance the one that is in an equal number of compounds on both sides i. If they both are, balance the one that is in the least number of compounds on both sides B. Simple Patterns of Chemical Reactivity a. Elements of the same group often react in similar ways b. Combination reaction-‐ 2 or more substances react to form one product i. EX: 2Mg (s) + 2 O (g) à 2MgO (s) ii. Oxides reacting with water: 1. Metal oxides a. Basic oxides produce basic metal hydroxides b. EX: K 2 + H2O à 2KOH c. OH-‐ causes substances to be basic 2. Nonmetal oxides a. Acidic oxides produce acids b. EX: SO 3 + 2 O à 2 H 4O c. Decomposition reaction-‐ one substance undergoes a reaction to produce 2 or more other substances i. EX: CaCO (s3 à CaO (s) 2 CO (g) d. Combustion reaction-‐ a rapid reaction that produces a flame; A reaction with O 2 i. Will always only produce CO and 2 2 ii. Incomplete combustion 1. Occurs when there is not enough oxygen present in the reaction iii. Oxide reaction-‐ a combustion reaction that requires intermediate steps 2 C. Formula Weights a. Molecular weight-‐ the sum of the atomic weights of the atoms in the chemical formula of a molecule b. Formula weight-‐ the sum of the atomic weights of the atoms in the chemical formula of an ionic substance c. Percent composition-‐ the percentage by mass contributed by each element in a substance i. = (number of atoms of element)(atomic weight) x 100 formula weight of substance D. Avogadro’s Number and The Mole a. Avogadro’s number = 6.022 x 10 atoms i. Represented by N A b. 1 amu = 6.022 x 10 grams c. Molar Mass i. 1 mole = 6.022 x 10 particles 1. a mole is always the same number, but different samples of 1 mol can have different masses ii. an element’s atomic weight/molecular weight/formula weight (amu) is equal to the mass (g) of 1 mole of that element d. Interconverting masses, moles, and number of particles i. Grams ( use molar mass) à moles (use Avogadro’s number) à formula units ii. Formula units (use Avogadro’s number) à moles (use molar mass) à grams E. Empirical Formulas from Analyses a. The ratio of the numbers of moles of all elements in a compound gives the subscripts in the compound’s empirical formula b. Calculating the mole ratio of each element in a compound 3 i. Mass % elements à (assume 100g sample) grams of each element à (use molar mass) moles of each element à (calculate the mole ratio) à then you can find empirical formula c. Finding empirical formula i. Divide larger number of moles by the smaller to obtain the mole ratio (also the atom ratio) ii. Multiply by the simplest factor to get whole numbers iii. Write formula d. Molecular Formulas from Empirical Formulas i. The subscripts in the molecular formula of a substance are always whole number multiples of the subscripts in its empirical formula ii. Whole number multiple = molecular weight Empirical formula weight e. Combustion analysis i. Determines empirical formula for compounds principally containing C and H 4
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