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Chapter 3: Problem 6

You know that chemical \(A\) reacts with chemical \(B\). You react \(10.0 \mathrm{~g} A\) with \(10.0 \mathrm{~g} B\). What information do you need to determine the amount of product that will be produced? Explain.

Short Answer

Expert verified
To determine the amount of product produced in the reaction between chemical A and B, you need the balanced chemical equation, the molar masses of A, B, and the product, and the stoichiometry of the reaction. Then, convert the masses of reactants into moles, identify the limiting reactant, calculate the theoretical yield of the product using stoichiometry, and finally convert the moles of the product to grams to determine the amount of product produced.

Step by step solution

01

Write the balanced chemical equation for the reaction between A and B

It's important to know the balanced equation representing the chemical reaction between A and B to understand the stoichiometry or mole ratios of reactants and products involved in the reaction. For the purpose of explanation, let's assume the balanced chemical equation is: A + B -> C
02

Determine the molar masses of A, B, and the product C

To convert the given masses of reactants A and B into moles, we need to know their molar masses. We also need to know the molar mass of the product C to express the amount of product formed in grams. Suppose the molar masses of A, B, and C are: \(M_A\), \(M_B\), and \(M_C\)
03

Calculate the moles of reactants A and B

Using the molar masses of A and B, convert the given mass of reactants into moles using the formula: Moles of A = \(\frac{mass~of~A}{molar~mass~of~A} = \frac{10.0~g}{M_A}\) Moles of B = \(\frac{mass~of~B}{molar~mass~of~B} = \frac{10.0~g}{M_B}\)
04

Identify the limiting reactant

The limiting reactant is the reactant that is completely consumed in a chemical reaction and determines the amount of product that can be formed. Compare the mole ratios of A and B in the balanced equation with the calculated moles of A and B. The reactant with the smaller mole ratio is the limiting reactant.
05

Calculate the theoretical yield of product C using stoichiometry

Since we now know the limiting reactant, we can use stoichiometry to determine the theoretical yield of product C. According to the balanced equation, 1 mole of A reacts with 1 mole of B to form 1 mole of C. Theoretical yield of C (moles) = Moles of limiting reactant (either A or B, whichever is smaller)
06

Convert the moles of product C to grams

In order to find the amount of product formed in grams, use the molar mass of product C: Amount of product C (grams) = Moles of product C * Molar mass of C This final value gives the amount of product C that will be produced from the given amounts of reactants A and B.

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Most popular questions from this chapter

A sample of a hydrocarbon (a compound consisting of only carbon and hydrogen) contains \(2.59 \times 10^{2.3}\) atoms of hydrogen and is \(17.3 \%\) hydrogen by mass. If the molar mass of the hydrocarbon is between 55 and \(65 \mathrm{~g} / \mathrm{mol}\), what amount (moles) of compound is present, and what is the mass of the sample?

What amount (moles) is represented by each of these samples? a. \(20.0 \mathrm{mg}\) caffeine, \(\overline{\mathrm{C}_{8} \mathrm{H}_{10} \mathrm{~N}_{4} \mathrm{O}_{2}}\) b. \(2.72 \times 10^{21}\) molecules of ethanol, \(\mathrm{C}_{2} \mathrm{H}_{3} \mathrm{OH}\) c. \(1.50 \mathrm{~g}\) of dry ice, \(\mathrm{CO}_{2}\)

DDT, an insecticide harmful to fish, birds, and humans, is produced by the following reaction: $$ 2 \mathrm{C}_{6} \mathrm{H}_{3} \mathrm{Cl}+\mathrm{C}_{2} \mathrm{HOCl}_{3} \longrightarrow \mathrm{C}_{14} \mathrm{H}_{4} \mathrm{Cl}_{5}+\mathrm{H}_{2} \mathrm{O} $$ \(\begin{array}{ll}\text { orobenzenc chloral } & \mathrm{D}\end{array}\) In a government lab, \(1142 \mathrm{~g}\) of chlorobenzene is reacted with \(485 \mathrm{~g}\) of chloral. a. What mass of DDT is formed? b. Which reactant is limiting? Which is in excess? c. What mass of the excess reactant is left over? d. If the actual yield of DDT is \(200.0 \mathrm{~g}\), what is the percent yield?

Acrylonitrile \(\left(\mathrm{C}_{3} \mathrm{H}_{3} \mathrm{~N}\right)\) is the starting material for many synthetic carpets and fabrics. It is produced by the following reaction. \(2 \mathrm{C}_{3} \mathrm{H}_{6}(g)+2 \mathrm{NH}_{3}(g)+3 \mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{C}_{3} \mathrm{H}_{3} \mathrm{~N}(g)+6 \mathrm{H}_{2} \mathrm{O}(g)\) If \(15.0 \mathrm{~g} \mathrm{C}_{3} \mathrm{H}_{6}, 10.0 \mathrm{~g} \mathrm{O}_{2}\), and \(5.00 \mathrm{~g} \mathrm{NH}_{3}\) are reacted, what mass of acrylonitrile can be produced, assuming \(100 \%\) yield?

When aluminum metal is heated with an element from Group \(6 \mathrm{~A}\) of the periodic table, an ionic compound forms. When the experiment is performed with an unknown Group \(6 \mathrm{~A}\) element, the product is \(18.56 \%\) Al by mass. What is the formula of the compound?
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