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Factor each polynomial. See Examples 1 through 10.x4 - 4x3

Intermediate Algebra | 6th Edition | ISBN: 9780321785046 | Authors: Elayn El Martin-Gay ISBN: 9780321785046 180

Solution for problem 38 Chapter 5.5

Intermediate Algebra | 6th Edition

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Intermediate Algebra | 6th Edition | ISBN: 9780321785046 | Authors: Elayn El Martin-Gay

Intermediate Algebra | 6th Edition

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Problem 38

Factor each polynomial. See Examples 1 through 10.x4 - 4x3

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Chemistry Notes Week 12 Chapter 8: Chemical Reactions A chemical equation uses chemical symbols to denote what occurs in a chemical reaction A. NH3 + HCl —> NH4Cl B. Ammonia and Hydrochloride reaction to form ammonium chloride Labels (indicate the physical state) A. Gas (g) B. Liquid (l) C. Solid (s) D. Aqueous (aq)- dissolved in water E. Example: NH3(g)+ HCl(g)—> NH4Cl(g) Balancing Chemical Equations A. Chemical equations must be balanced so that the law od conservation of mass is obeyed B. Balancing is achieved by writing stoichiometric coefficients to the left of the chemical formulas C. Requires a trial-and-error approach D. Generally, it will facilitate the balancing process if you do the following: 1. Write the unbalanced equation 2. Balance each of the atoms, leaving 02 at the end 3. Continue process until it becomes even Determination of Empirical Formula A. In the combustion of 18.8 g of glucose (CxHyOz) 27.6 g of CO2 and 11.3 g of H2O are produced B. It is possible to determine the mass of carbon and hydrogen in the original sample by multiplying the (original mass) x (original molar mass) x (mole ratio) x (new molar mass) = new mass C. Divide by the smallest subscript to find the whole number ratio D. To get the empirical formula, follow these steps: 1. Multiply mass of sample by their molar masses to get moles of the samples, then obtaining the whole number ratio 2. Find the molar mass of the smallest ratio of the substance 3. Divide the molar mass by the empirical formula mass 4. Multiply the subscripts by the ratio Calculations with Balanced Chemical Equations A. Consider the complete reaction of 3.82 moles of CO to form CO2 B. Calculate the number of moles of CO2 produced C. Moles you have x mole ratio = moles you should get Limiting Reactants A. Consider the reaction between 5 moles of CO and 8 moles of H2 to produce methanol B. CO + 2H2 —> CH2OH C. How many moles of H2 are necessary in order for all the CO to react Moles H2 = 5 moles CO x mole ratio = 10 moles H2 D. How many moles of CO are necessary in order for al of the H2 to react Moles CO = 8 moles H2 x mole ratio = 4 moles CO Reaction Yield A. The theoretical yield is the amount of product that arms when all the limiting reactant reacts to form the desired product B. The actual yield is the amount of product actually obtained from a reaction C. The practical yield tell what percentage the actual yield is of the theoretical yield D. % Yield = actual/theoretical x 100% **Begin Chapter 9** General Properties of Aqueous Solutions A. A solution is a homogenous mixture of two or more substances B. The substance present in the largest amount (moles) is referred to as the solvent C. The other substances are the solutes D. An electrolyte is a substance that dissolves in water to yield a solution that conducts electricity Strong Electrolytes and Weak Electrolytes A. An electron that disassociates completely is known as a strong electrolyte B. Water soluble ionic compounds/ strong acids/ strong bases C. Look up chart for strong acids D. A weak electrolyte is a compound that produces ions upon dissolving but exists in solution predominately as molecules that are NOT ionized Precipitation Reactions A. An insoluble product that separates from a solution is called a precipitate B. A chemical reaction in which a precipitate forms is called a precipitation reaction Solubility Guidelines for Ionic Compounds in Water A. Water is a good solvent for ionic compounds because it is a polar molecule B. The polarity of water results from electron distributions within the molecules C. The oxygen atom has an attraction for the hydrogen atom’s electrons and is therefor partially negative compared to hydrogen D. Hydration occurs when water molecules remove the individual ions from an ionic solid surrounding them, preventing them from recombine so the substances dissolve E. Solubility is defined as the maximum amount of solute that will dissolve in a given quantity of solvent at a specific temperature G. There are many exceptions (see below) Molecular Equations A. In a molecular equation, compounds are represented by chemical formulas as though they exist in solution as molecules or formula units B. Na2SO4(aq) + Ba(OH)2(aq) —> 2NaOH(aq) + BaSO4(s) C. Reactions in which cations in two ionic compounds exchange anions are called metathesis or double replacement reactions Ionic Equations A. In an ionic equation, compounds that exist completely or predominately as ions in solution are represented as those ions B. In the reaction between aqueous sodium sulfate and barium hydroxide, the aqueous species are represented as ions which form precipitates Net Ionic Equations A. An equation that includes the species that are actually involved in the reaction is called a net ionic equation B. Ions that appear on both sides of the equation are called spectator ions C. Spectator ions do not participate in the reaction Precipitation Reactions (to determine the molecular, ionic, and net ionic equations) A. Write and balance the molecular equation, predicting the products by assuming that the cations trade anions B. Write the ionic equation by separating strong electrolytes into their constituent ions C. Write the net ionic equation by identifying and cancelling spectator ions on both sides of the equation D. If both the reactants and products are all strong electrolytes, all the ions in solution are spectator ions. In this case, there is no net ionic equation and no reaction takes place Acid-Base Reactions A. Acids can either be strong or weak B. A strong acid is a strong electrolyte Strong Acids and Bases A. Strong bases are strong electrolytes (dissociate completely) B. Strong bases are the hydroxides of Group 1A and heavy Group 2A C. A weak acid is a weak electrolyte, it does not dissociate completely D. Most acids are weak acids Bronsted Acids and Bases A. An Arrhenius acid is one that ionizes in water to produce H+ ions B. An Arrhenius base is one that dissociates in water to produce OH- ions C. A Bronsted acid is a proton donor D. A Bronsted base is a proton acceptor E. In these definitions, a proton refers to a hydrogen atom that has lost its electron, also known as a hydrogen ion (H+) F. Bronsted acids donate protons to water to form the hydronium ion (H3O+) G. A monoprotic acid has one proton to donate H. Hydrochloric acid is an example I. A polyprotic acid has more than one acidic hydrogen atom J. Sulfuric acid, H2SO4, is an example of a diuretic acid, there are two acidic hydrogen atoms K. Polyprotic acids lose protons in a stepwise fashion L. Bases that produce only one mole of hydroxide per mole of compound are called monobasic (sodium hydroxide is an example) K. Some strong bases produce more than one hydroxide per mole of compound (Barium hydroxide per mole of compound (Barium hydroxide is an example is an example of a dibasic base)

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Chapter 5.5, Problem 38 is Solved
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Textbook: Intermediate Algebra
Edition: 6
Author: Elayn El Martin-Gay
ISBN: 9780321785046

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Factor each polynomial. See Examples 1 through 10.x4 - 4x3