Chemistry Stoichiometry Standard Set 3 Review

Name:______Date:______Period:______CHEMISTRY STOICHIOMETRY STANDARD SET 3 REVIEW

3a. Students know how to describe chemical reactions by writing balanced equations. Description Reactions are described by balanced equations because all the atoms of the reactants must be accounted for in the reaction products. An equation with all correct chemical formulas can be balanced by a number of methods, the simplest being by inspection. Given an unbalanced equation, do an inventory of each element to determine how many of each atom are on each side of the equation. If the result is not equal for all atoms, coefficients (not subscripts) are changed until balance is achieved. Vocabulary CHEMICAL REACTIONS CHEMICAL EQUATIONS REACTANTS PRODUCTS NOMENCLATURE [RULES] CHEMICAL FORMULAS BALANCED EQUATIONS CHEMICAL SYMBOLS COEFFICIENTS SUBSCRIPTS Practice Problems

1. Potassium carbonate (K2CO3) is an important component of fertilizer. The partially balanced equation for the reaction of 6 moles of potassium hydroxide (KOH) and 3 moles of carbon dioxide (CO2) to produce potassium carbonate and water is given below.

6KOH + 3CO2 → __ K2CO3 + 3H2O

When this equation is balanced, what is the coefficient for potassium carbonate? A. 2 B. 3 C. 6 D. 9 2. Aluminum reacts vigorously and exothermically with copper(II) chloride. Which of the following is the balanced equation for this reaction? A. Al + CuCl2 → AlCl3 + Cu B. Al + 3CuCl2 → 2AlCl3 + Cu C. 2Al + 3CuCl2 → 2AlCl3 + 3Cu D. 3Al + 2CuCl2 → 3AlCl3 + 2Cu

3. In the formula X2O5, the symbol X could represent an element in Group- A. 1 B. 2 C. 15 D. 18

Chemistry Standard Set 3 1 3b. Students know the quantity one mole is set by defining one mole of carbon-12 atoms to have a mass of exactly 12 grams. Description The mole concept is often difficult at first, but the concept is convenient in chemistry just as a dozen is a convenient concept, or measurement unit, in the grocery store. The mole is a number 6.02 x 1023. Specifically, a mole is defined as the number of atoms in 12 grams of carbon-12. When atomic masses were assigned to elements, the mass of 12 grams of carbon-12 was selected as a standard reference to which the masses of all other elements are compared. The number of atoms in 12 grams of carbon-12 is defined as one mole, or conversely, if one mole of 12C atoms were weighed, it would weigh exactly 12 grams. (Note that carbon, as found in nature, is a mixture of isotopes, including atoms of carbon-12, carbon-13, and trace amounts of carbon-14.) The definition of the mole refers to pure carbon-12. The atomic mass of an element is the weighted average of the mass of one mole of its atoms based on the abundance of all its naturally occurring isotopes. The atomic mass of carbon is 12.011 grams. If naturally occurring carbon is combined with oxygen to form carbon dioxide, the mass of one mole of naturally occurring oxygen can be determined from the combining mass ratios of the two elements. For example, the weight, or atomic mass, of one mole of oxygen containing mostly oxygen-16 and a small amount of oxygen-18 is 15.999 grams. Vocabulary

ISOTOPES

ATOMIC MASS

CARBON-12

MASS-MOLE RELATIONSHIP

Practice Problems 4. How many moles of carbon-12 are contained in exactly 6 grams of carbon-12? A. 0.5 mole B .20 moles C. 3.01 ×1023 moles D. 6.02 ×1023 moles 3c. Students know one mole equals 6.02 × 1023 particles (atoms or molecules). Description A mole is a very large number. The number of atoms in a mole has been found experimentally to be about 6.02 × 1023. This number, called Avogadro’s number, is known to a high degree of accuracy. Avogadro's Law states that the relationship between the masses of the same volume of different gases (at the same temperature and pressure) corresponds to the relationship between their respective molecular weights. Hence, the relative molecular mass of a gas can be calculated from the mass of sample of known volume. Vocabulary 6.02 x 1023

AVOGARDRO’S NUMBER

Chemistry Standard Set 3 2 Practice Problem 5. How many molecules of water are in sample containing 6.00 moles of the compound? A. 3.61 x 1023 molecules B. 3.61 x 1024 molecules C. 1.00 x 1023 molecules D. 1.00 x 1022 molecules

3d. Students know how to determine the molar mass of a molecule from its chemical formula and a table of atomic masses and how to convert the mass of a molecular substance to moles, number of particles, or volume of gas at standard temperature and pressure. Description The molar mass of a compound, which is also called either the molecular mass or molecular weight, is the sum of the atomic masses of the constituent atoms of each element in the molecule. Molar mass is expressed in units of grams per mole. The periodic table is a useful reference for finding the atomic masses of each element. For example, one mole of carbon dioxide molecules contains one mole of carbon atoms weighing 12.011 grams and two moles of oxygen atoms weighing 2 × 15.999 grams for a total molecular mass of 44.009 grams per mole of carbon dioxide molecules.

The mass of a sample of a compound can be converted to moles by dividing its mass by the molar mass of the compound. This process is similar to the unit conversion. The number of particles in the sample is determined by multiplying the number of moles by Avogadro’s number.

The volume of an ideal or a nearly ideal gas at a fixed temperature and pressure is proportional to the number of moles. Students should be able to calculate the number of moles of a gas from its volume by using the relationship that at standard temperature and pressure (0°C and 1 atmosphere), one mole of gas occupies a volume of 22.4 liters.* Vocabulary MOLAR MASS

CONSTITUENT ELEMENTS

CONVERSIONS [MOLE-MASS-PARTICLES]

Practice Problems 6. How many molecules of water are in sample containing 6.00 moles of the compound? A. 3.61 x 1023 molecules B. 3.61 x 1024 molecules C. 1.00 x 1023 molecules D. 1.00 x 1022 molecules

7. How many moles are in 59.6 grams of BaSO4?

Chemistry Standard Set 3 3 A. 0.256 mole B. 3.91 moles C. 13.9 moles D. 59.6 moles

8. How many molecules are contained in 55.0 g of H2SO4? A. 0.561 molecule B. 3.93 molecules C. 3.38 x 1023 molecules D. 2.37 x 1024 molecules

9. How many moles of chlorine are in 100 g chlorine (Cl)? Element Molar Mass (g/mol) Carbon 12.01 Chlorine 35.45

A. 64.6 B. 2.82 C. 100 D. 0.355

3e. Students know how to calculate the masses of reactants and products in a chemical reaction from the mass of one of the reactants or products and the relevant atomic masses. Description Atoms are neither created nor destroyed in a chemical reaction. When the chemical reaction is written as a balanced expression, it is possible to calculate the mass of any one of the products or of any one of the reactants if the mass of just one reactant or product is known.

Balanced chemical equations can be used to predict the mass of any product or reactant. Teachers should emphasize that the coefficients in the balanced chemical equation are mole quantities, not masses. Here is an example: How many grams of water will be obtained by combining 5.0 grams of hydrogen gas with an excess of oxygen gas, according to the following balanced equation? 2H2 + O2 → 2H2O This calculation is often set up algebraically, for example, as and can be easily completed by direct calculation and unit cancellation (dimensional analysis). Recall that the coefficients in the balanced equations refer to moles rather than to mass. The molar ratio will assume a place of central importance in solving stoichiometry problems. The sources for these ratios are the coefficients of a balanced equation Vocabulary CHEMICAL REACTIONS

REACTANTS

PRODUCTS

LAW OF CONSERVATION OF MASS

BALANCED EQUATION

COEFFICIENTS

Chemistry Standard Set 3 4 MOLE-MASS RELATIONSHIP

DIMENSIONAL ANALYSIS

Practice Problems 10. Which model demonstrates the Law of Conservation of Matter?

Zn + 2HCl → ZnCl2 + H2 11. If 0.600 gram of zinc is used, and the amount of HCl is unlimited, what is the amount of zinc chloride that is produced in the above reaction? A. 0.125 gram C. 1.25 grams B. 12.5 grams D. 0.018 gram

12. Following each equation are two requests for molar ratios from the equation.

a) N2 + 3 H2 ---> 2 NH3 Write the molar ratios for N2 to H2 and NH3 to H2.

b) 2 SO2 + O2 ---> 2 SO3 Write the molar ratios for O2 to SO3 and O2 to SO2.

c) PCl3 + Cl2 ---> PCl5 Write the molar ratios for PCl3 to Cl2 and PCl3 to PCl5.

d) 4 NH3 + 3 O2 ---> 2 N2 + 6 H2O Write the molar ratios for NH3 to N2 and H2O to O2.

Chemistry Standard Set 3 5 e) Fe2O3 + 3 CO ---> 2 Fe + 3 CO2 Write the molar ratios for CO to CO2 and Fe to CO.

Chemistry Standard Set 3 6 Mole-Mole Problems

The solution procedure used below involves making two ratios and setting them equal to each other. This is called a proportion. One ratio will come from the coefficients of the balanced equation and the other will be constructed from the problem. The ratio set up from data in the problem will almost always be the one with an unknown in it.

You will then cross-multiply and divide to get the answer.

What happens if the equation isn't balanced? Answer - balance it. You cannot do these problems correctly without a balanced equation.

How will I know which substances to use in the ratio? Answer - you will have to read the problem and understand the words in it.

Here is the first equation we'll use: ` N2 + 3 H2 ---> 2 NH3

Problem #1: if we have 2.00 mol of N2 reacting with sufficient H2, how many moles of NH3 will be produced?

Solution Comments

1. The ratio from the problem will have N2 and NH3 in it. 2. How do you know which number goes on top or bottom in the ratios? Answer: it does not matter, except that you observe the next point ALL THE TIME. 3. When making the two ratios, be 100% certain that numbers are in the same relative positions. For example, if the value associated with NH3 is in the numerator, then MAKE SURE it is in both numerators. 4. Use the coefficients of the two substances to make the ratio from the equation. 5. Why isn't H2 involved in the problem? Answer: The word "sufficient" removes it from consideration.

Let's use this ratio to set up the proportion:

That means the ratio from the equation is:

The ratio from the data in the problem will be:

The proportion (setting the two ratios equal) is:

Solving by cross-multiplying gives x = 4.00 mol of NH3 produced.

Chemistry Standard Set 3 7 N2 + 3 H2 ---> 2 NH3

Example #2 - Suppose 6.00 mol of H2 reacted with sufficient nitrogen. How many moles of ammonia would be produced?

Solving by cross-multiplying and dividing gives x = 4.00 mol of NH3 produced.

N2 + 3 H2 ---> 2 NH3

Example #3 - We want to produce 2.75 mol of NH3. How many moles of nitrogen would be required?

Before the solution, a brief comment: notice that hydrogen IS NOT mentioned in this problem. If any substance ISN'T mentioned in the problem, then assume there is a sufficient quantity of it on hand. Since that substance isn't part of the problem, then it's not part of the solution.

Solving by cross-multiplying and dividing (plus rounding off to three significant figures) gives x = 1.38 mol of N2 needed.

Chemistry Standard Set 3 8 Practice Problems

Here's the equation to use for all three problems:

2 H2 + O2 ---> 2 H2O

1) How many moles of H2O are produced when 5.00 moles of oxygen are used?

2) If 3.00 moles of H2O are produced, how many moles of oxygen must be consumed?

3) How many moles of hydrogen gas must be used, given the data in problem two?

4) Suppose 4.00 grams of H2 were used? How many grams of water would be produced?

Chemistry Standard Set 3 9