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Chemistry 1210 Chapter 3 Notes

by: Emily Notetaker

Chemistry 1210 Chapter 3 Notes Chem 1210

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Emily Notetaker

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These notes cover all of chapter 3, including important concepts from books that were left out in lecture. This chapter will introduce the class to a more calculation-based course. There is a separ...
General Chem 1
Dr. Bartoszek-Loza
Class Notes
General Chemistry, Chemistry
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This 9 page Class Notes was uploaded by Emily Notetaker on Saturday September 10, 2016. The Class Notes belongs to Chem 1210 at Ohio State University taught by Dr. Bartoszek-Loza in Summer 2016. Since its upload, it has received 18 views.


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Date Created: 09/10/16
Lecture #4 Thursday, September 1, 2016 2:20 PM Chemical Equations (3.1) Hydrogen gas and oxygen gas react to produce steam: 2H 2g) + O 2g) 2H 2 Reactants Products There are two types of numbers • Subscripts (substance) - ratio of atoms in a moleculein a formula ○ Water: two hydrogen atoms, one oxygen atom • Stoichiometriccoefficient (equation) - ratio between reactants and products ○ 2 hydrogen molecules,1 oxygen molecule,2 water molecules Conservationof matter (and mass) is required • Atoms are neither created nor destroyedin a chemical reaction CH 4g) + O 2g) --> CO 2g) + 2H 2(g) Patterns of Chemical Reactivity(3.2) Combination:form one product • A + B ---> AB 2Mg(s) + O (2) --> 2MgO Redox N 2g) +3H 2g) --> 2NH 3 Decomposition:form two or more products • AB ---> A + B CaCO (3) ---> CO 2g) + CaO(s) Coral reefs 2H 2 (2q) ---> Combustion:burning reactions in "air" (O ) 2 usually excess oxygen (completecombustion) • Hydrocarbons (C,H) combustion always makes CO + H O 2 2 • Hydrocarbons with oxygen always makes CO + H O 2 2 ○ Mother Eats Peanut Butter  Meth- 1, Eth- 2, Prop- 3, But- 4 (number of carbons) ○ Ex: Propane is burned in air write balance the equation Formula Weights (3.3) • the Periodic table gives weighted average of naturally occurring isotopes ○ Average weight (AW): average atomic mass based upon weighted average of naturally occurring isotopes  The mass of an atom is assigned using relative weights with the standard of Carbon being exactly 12 amu's • Carbon is 12.011amu ( this is what is on the periodic table) 12 13 ○ 98.892% C and 1.07% C • Formula Weight: sum of atomic weights of atoms in a formula (everything has a formula weight) ○ Ionics • Molecular weight: The sum of atomicweight in a molecule ○ nonmetals • Percent composition:what is the mass percentage compositionof copper (II) nitrate 3SF? ○ x 100% ○ Mass % copper(II) nitrate trihydrate? Cu(NO )3 2 FW = 187.5amu x100%= 14.9%Nitrogen atoms Avogadro's Number and the Mole (3.4) • The mole is a convenient unit for an enormouscollection. Convenientfor amounts of atoms handled routinely in the lab. 12 ○ 1 mole is the number of atoms in exactly 12g of C • 1 mole of something = 6.022x10 23 • Mass/ion mole atom ○ How many:  Moles of copper are in 3.0 g of copper? 3.0 g Cu x (1 mol/63.55g) =  Grams in one copper atom?  Atoms of copper are in 3 g of copper?  H atoms are in 2.5 g of C 2 6?  PO 43ions in 2.6x10 mol Mg (PO3) 4 2 2+ ○ What mass of Ca (PO3)2 4s needed for an experiment that requires 0.836 moles Ca ? Lecture #5 Tuesday, September 6, 2016 2:45 PM Empirical Formulas from Analyses (3.5) For an unknown chemical: a qualitative analysis is performed to determine which elementsare present then a quantitative analysis is done to det. The amount of each element.Once the percent compositionis known the empirical formula can be determined. • Empirical formula: The smallest whole number ratio of atoms ○ Ascorbic acid (Vitamin C) has the empirical formula C H O The molecular weight of ascorbic acid is 176 3 4 3 g/mol. What is the molecular formula? ○ Analysis of an unknown compound gave 39.72% C, 1.67% H and 58.61% Cl. The molecular weight was found to be 181.4. Determine the molecular formula. • Combustion Analysis: experimental method used to determine the empirical formula. ○ Combustion produces carbon dioxide and water CH 4g) + 2O (2) --> CO 2g) + 2H 2(g) Combustion of a 0.1000 g sample produced 0.1910 g Co2 and 0.1172 g H2O. What is the empirical formula of this compound? Quantitative Information from Balanced Chemical Equations (3.6) How many grams of KOH were formed in a reaction of 3.0 g of potassium with excess water? Calculation Method: • Step 1: Convert grams to moles • Step 2: Multiply by the mole ratio • Step 3: Convert moles to grams A 25.5 g sample of potassium chlorate was decomposed. How many • Moles of oxygen were produced? • Grams of oxygen were produced? • Grams of KCl were produced? Limiting Reactants (3.7) • Limiting reactant: reactant used up first. • Excess reactant: reactant that remains. • An excess of one reactant is often used to ensure maximum product formation How many moles of NH3 can be formed from 3.0 mol of N2 and 6.0 mol of H2? • Limiting reactant: H 2 • Excess Reactant: N 2 • Amount of NH for3ed: 4.0 mol How much excess reactant remains? g(excess) = g(initial) - g(reacted) What mass of NH3 can be produced from a mixture of 100 g of N2 and 500 g of H2? • Limiting reactant: N 2 • Excess reactant: H 2 • Amount of NH for3ed: 114 g What was the percent yield of ammonia if the actual yield of ammonia is 105 g? • Percent yield = (actual yield/theoreticalyield)x100 • Theoreticalyield: product calculated when all limiting reactant is consumed. Real life: Excess of one reactant often used to ensure maximum product formation Chapter 3 Overview Tuesday, September 6, 2016 8:14 PM Ch. 3| ChemicalReaction and Reaction Stoichiometry Stoichiometry- the area of study that examinesthe quantities of substances consumed and produced in chemical reactions • Built upon Law of Conservationof Mass: atoms are neither created nor destroyed during a chemical reaction 3.1 | Chemical Equations Chemical equations represent chemical reactions. 2H + O ----> 2H O 2 2 2 • Reactants - Chemical formulas to the left of the arrow, they are the starting substances. • Products - the chemical formulas to the right of the arrow, they are the substances that are produced Steps to construct a balanced chemical equation: 1. Start by writing the formulas for the reactants on the left side and products on the right side of the arrow 2. Balance the equation by determining the coefficients that provide equal numbers of each atom on both sides. 3. All equations should contain the smallest whole-number coefficients Indicating the state of reactants and products • (g) - gas • (l) - liquid • (s) - solid • (aq) - aqueous 3.2 | Simple Patterns of Chemical Reactivity The three types of reactions that we will see frequently: • Combination reactions - two or more substances react to form one product ○ A combinationreaction between a metal and nonmetal reaction forms an ionic solid • Decompositionreactions - one substance undergoes a reaction to produce two or more other substances • Combustion reactions - rapid reactions that produce a flame ○ Most of these reactions involve O2 from air as a reactant ○ A hydrocarbon is a reactant in a reaction (consists of hydrogen and carbons) ○ These reactions always produce CO2 and H2O 3.3 | Formula Weights Chemical formulas and chemical equations have a quantitative significance in their subscripts and coefficients • For example: In H2O, there are two hydrogens to every oxygen Formula weight - the sum of the atomicweights of the atoms in the chemical formulas of a substance • For example: H S2 fo4mula weight is 98.1 amu with 2 Hydrogens + 1 Sulfur + 4 Oxygens Percent Composition • The percentage compositionof a compound is the percentage by mass contributed by each • The percentage compositionof a compound is the percentage by mass contributed by each element in the substance • Calculation Example: 3.4 | Avogadro's Number and the Mole • Mole - the amount of matter that contains as many objects (atoms, molecules)as the number of atoms in exactly 12 g of Carbo23 • Avogadro's number: 6.022x10 ○ Named after Amedeo Avogadro ○ From experiments scientists have determined that there is this amount in one mole of something ○ This is such a large number • Conversionfrom Moles to atoms Example: • Molar mass - the mass in grams of 1 mole of a substance ○ One mole of something is a different mass for each element  Cl has an atomic weight of 35.5 amu, so the mass of 1 mol is also 35.5 g  Au has an atomic weight of 197 amu, so the mass of 1 mole is 197 g. ○ Calculating Molar Mass Example: ○ Calculating Numbers of Molecules and Atoms from Mass Example: (The rest of the sections were well covered in lecture with exact calculations that will be on the midterm and final) Quiz 3 Tuesday, September 6, 2016 2:25 PM Practice Questions: What is the sum of the coefficients of the products for the combustion ofC3H8O? The mass of copper (II) nitrate is 188 g/mol. It is a hygroscopic material- it absorbs water on exposure to air. It has been determined that 3 water molecules are absorbed. a.) What happens to the molecular weight? b.) What happens to the percent oxygen? Answer: Both increase How many atoms are in 3g of copper? (MW Cu = 64)(Avagadro's number = 6.02x10^23) Answer: A sample contains C, H, and O. On combustion, water and carbon dioxide are formed. The amount of C (g) can be calculated from the amount of carbon dioxide. The amount of H (g) can be calculated from the amount of water. How can you find the amount of O? Answer: grams O = sample weight + (g C + g H) Chapter 3 Page 9


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