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# Chapter 3&4 study guide Chem 142

Marketplace > University of Washington > Chemistry > Chem 142 > Chapter 3 4 study guide
Jessie Yuan
UW

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Basically from lectures, hope this helps. Thank you guys!
COURSE
General chemistry
PROF.
Dr. Li Xiaosong
TYPE
Study Guide
PAGES
14
WORDS
KARMA
50 ?

## Popular in Chemistry

This 14 page Study Guide was uploaded by Jessie Yuan on Tuesday November 17, 2015. The Study Guide belongs to Chem 142 at University of Washington taught by Dr. Li Xiaosong in Fall 2015. Since its upload, it has received 76 views. For similar materials see General chemistry in Chemistry at University of Washington.

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Date Created: 11/17/15
CHAPTER 3&4 REVIEW NOTES (WEEKLY BUNDLE) JESSIE YUAN (U WASHINGTON) STOICHIOMETRY The Atomic Mass Unit (amu):  Defined as 1/12 the mass of a carbon-12 atom  The masses of all other atoms are given relative to this standard  The atomic masses you find on the periodic table are a weighted average of the masses of each isotope of that element. Atomic Mass:The Atomic Mass (aka Atomic Weight or Average Atomic Mass) is the average of the atomic masses of all of the element's isotopes, weighted by isotopic abundance. Don’t confuse “Atomic Mass” with the mass of one atom!! Counting by Weighing:  Chemical reactions occur at the microscopic level, between individual molecules and/or atoms.  In the lab, we measure substances in terms of grams or milliliters…these are macroscopic measurements.  The number of molecules in 1 g of water will be different than the number of molecules in 1 g of glucose, because these molecules have different masses.  We need a way to convert between the microscopic and macroscopic descriptions. Percent Composition of Compounds  Sometimes, it is helpful or necessary to know a compound’s composition in terms of the masses of its elements.  We can also deduce a molecular formula based on a given percent composition.  Mass Fraction and Mass % Mass fraction = mass of one component (element, particle, etc.)/total mass (of molecule, box of marbles, etc.) Empirical and Molecular Formulas  Empirical Formula - The simplest formula for a compound that agrees with the elemental analysis! The smallest set of whole numbers of atoms.  Molecular Formula - The formula of the compound as it really exists. It must be a multiple of the empirical formula. Chemical Equations  Chemistry is the study of the rearrangement of matter due to the flow of energy.  In a chemical reaction, some bonds are broken and others are formed, resulting in a reorganization of the atoms.  Atoms are neither created or destroyed in a chemical reaction - Law of Conservation of Mass (Reactants and products must occur in numbers that give the same number of each type of atom on both sides of the arrow.) Formulas of Elements and Compounds  The name or chemical formula of a compound gives you information about the relative number of atoms in the compound sodium chloride: NaCl copper(II) nitrate: Cu(NO3)2 carbon tetrachloride: CCl4  The formula of most elements (particularly metals) is simply the element symbol tungsten: W calcium: Ca boron: B carbon: C  Elemental forms of nonmetals are often found as molecules H2, N2, O2, F2, Cl2, Br2, I2 S8, C60, P4 How to Balance Equations Mass Balance (or Atom Balance)- same number of each element on each side of the equation: (1) start with largest/most complicated molecule (2) progress to other elements, leaving lone elements for last (3) make all whole numbers (4) re-check atom balance Mole Ratios We can use a balanced chemical equation to predict the number of moles of products that a given number of moles of reactants will produce. Theoretical vs. Actual Yield ¡ The theoretical yield of a reaction is the amount of product that would be formed under ideal reaction conditions in which starting materials are completely consumed (up to the LR). This is a calculated number. ¡ The actual yield is the amount of product that is actually produced in real life (in the lab). ¡ The actual yield is always less than the theoretical yield because… LR starting materials may not be completely consumed side reactions may occur the reverse reaction may occur there may be loss of product during purification steps Percent Yield: A comparison of… How much product we actually produced to how much product we could theoretically produce gives us the percent yield of a chemical reaction. Solutes, Solvents and Solutions ØSolute ¡ Substance being dissolved, mixed, diluted. ¡ Example: compounds extracted from coffee grounds, sugar, milk. ØSolvent ¡ Substance doing the dissolving, mixing, dilution. ¡ Example: water ØSolution ¡ Final combination of dissolution, mixing, and dilution. ¡ Example: morning coffee Water as a Solvent ¡Water is an important solvent – dissolves many substances ¡“Aqueous” means a solution in which water is the solvent ¡Water is a POLAR molecule Polar and Nonpolar Solutes ¡ Water dissolves some non-ionic substances if they are polar (ethanol-water) ¡ Ethanol molecules are polar (contain directional O-H bond) Ionic Solutes ¡ Polar water molecules dissolve ionic compounds (salts) ¡ “Hydration” breaks ionic compounds into anions and cations ¡ Water dissolves different ionic compounds to different degrees (more in Ch. 8) The Role of Water as a Solvent: Dissolution of Ionic Compounds Electrical conductivity: the flow of electricity in a solution indicates the presence of ions in solution. Electrolyte: a substance that conducts a current when dissolved in water. • Ions become solvated/hydrated– they are surrounded by water molecules. • These ions are labeled “aqueous” – they are free to move throughout the solution and conduct electricity. Electrolytes and Non-Electrolytes ¡ If a solution conducts electricity, it contains ions ¡ A solution that contains many ions is a strong electrolyte. ¡ A solution that contains only a few ions is a weak electrolyte. ¡ A solution that contains no ions is a nonelectrolyte. Strong Electrolytes strong electrolytes: substances that are good conductors of electricity ¡ These substances break up to produce many ions in water ¡ many ions present to move electrons/conduct electricity: strong electrolyte Weak Electrolytes weak electrolytes: substances that are weak/poor conductors of electricity ¡ These substances mostly remain intact as compounds, producing very few ions in water ¡ only a few ions present to move electrons/conduct electricity: weak electrolyte Non-electrolytes nonelectrolytes: substances that cannot conduct electricity ¡ These molecules never break down into ions. ¡ They always remain intact as neutral molecules that have no charge: no ions to move electrons/conduct electricity Dissolving compounds in water ¡When an ionic compound is dissolved in water: ¡ Ions are “hydrated” ¡ Separated from their solid crystal ¡ Become individual ions in solution Concentration ¡Many chemical reactions take place “in solution” ¡ Still need to know amounts of reactants and products ¡ How do we make solutions of known concentrations? ¡We measure concentration in terms of moles per volume… Molarity = M = moles of solute/liters of solution Dilution ¡ The number of moles (n) of solute stays the same…only the volume of solution (V) changes. n = MV ¡We can formalize this relationship: M1V1 = M2V2 = n where… M1, V1 = molarity and volume of concentrated solution M2, V2 = molarity and volume of diluted solution Types of Chemical Reactions ¡ Precipitation Reactions ¡ Acid-Base Neutralization Reaction ¡ Oxidation-Reduction (Redox) Reactions ¡ Further classified as: Combination Decomposition Combustion Single-replacement reactions Solubility Rules Total and Net Ionic Equations Conventional (molecular) equation: a bookkeeping of all species present, and arranged for charge neutrality. Total Ionic equation: all aqueous species are split up into their component ions. Net Ionic equation: indicates exactly the chemical change that occurs, and nothing more. Selective Precipitation • Some ionic compounds are soluble, while others are not. • We can use this behavior to remove species selectively. • Example: separating Ag+ from Ba2+ and Fe3+. • Notice that selective precipitation is nothing more than an application of the solubility rules. Acid-Base Reactions Generally, in an acid-base reaction, ¡ H+ from acid reacts with the OH– from base to form water, H2O ¡ The cation (M+) from base combines with anion from acid (X–) to form a salt. Acid-Base Rxns Bronsted-Lowry Theory: acid/base reactions are protontransfer processes. ¡ acid is proton-donor (H+ ion donor). ¡ base is proton- acceptor (H+ ion acceptor). Strong and Weak Acids ¡ Strong acids (think “strong electrolyte”) undergo complete ionization. ¡ Weak acids (think “weak electrolyte”) undergo incomplete ionization. (Weak acids are like insoluble salts…they don’t like to dissociate very much.) Oxidation-Reduction (Redox) Reactions ¡“Redox” Chemistry: Reduction and Oxidation ¡Oxidation: Loss of electrons ¡Reduction: Gain of electrons (a reduction in oxidation number) Redox Reactions ¡ In a redox reaction, one species loses electrons and another species accepts those electrons. ¡ Electrons are neither created nor destroyed during the reaction…charges are conserved. Oxidation Numbers (or States) ¡ Oxidation number (state) represents the number of electrons required to produce the “effective charge” on a species. ¡ Oxidation numbers (states) are chosen so that: ¡ charges are conserved ¡ in ionic compounds, the sum of the oxidation numbers on the atoms is the same as the charge on the whole ion the oxidation numbers of the atoms/ions in a species must sum to the total charge of the species. Oxidation-Reduction (Redox) Reactions Types of Redox Reactions ¡ Combination Reaction ¡ Decomposition Reaction ¡ Replacement (or Displacement) Reaction ¡ Combustion Reaction Activity Series Activity Series: Relative order of elements arranged by their ability to replace cations in aqueous solution Or, stated another way…their ease of oxidation, or loss of e- to form a cation. Li > K > Ba > Sr > Ca > Na > Mg > Al > Mn > Zn > Fe > Cd > Co > Ni > Sn > Pb > (H) > Cu > Ag > Au Balancing REDOX Equations: The Half-Reaction Method Step 1: Write the half-reactions for the chemical equation. Step 2: For each reaction, balance the atoms other than O and H. Step 3: Add H2O to balance O, then H+ to balance H. Step 4: Balance the charge by adding electrons. The net charge of the reactants should equal the net charge of the products. Step 5: Add the two half-reactions together, making sure e- lost equal e- gained, and canceling any species that appear on both sides of the reaction. The reaction is now balanced in an acidic solution. Step 6: If you need to balance in a basic solution, first balance in acidic solution (!), then add OH- to both sides to neutralize any H+ present. Cancel any species that appear on both sides of the reaction.

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