1 Appendix D Reaction Stoichiometry D.1 INTRODUCTION In Appendix A, the stoichiometry of elements and compounds was presented. There, the relationships among grams, moles and number of atoms and molecules were reviewed. A similar relationship exists for chemical reactions, and we will now extend this concept of stoichiometry to reactions. In reaction stoichiometry, we are interested in the quantitative relationships between the amounts of reactants and products in a reaction. We will find, as we did in Appendix A, that the mole is the central character in these calculations. D.2 QUANTITATIVE RELATIONSHIPS IN REACTIONS Chemical reactions surround us. Chemists use shorthand notation to describe them in a sentence called a chemical equation. The chemical equation that describes the combustion of benzene is 2C 6(l) + 15O 2(g) 12CO 2(g) + 6H 2O(g) This equation implies that for every two benzene molecules that react, 15 dioxygen molecules must also react and 12 carbon dioxide molecules and six water molecules will be produced. In other words, it tells us about the stoichiometry, or the amounts of reactants and products involved. The coefficients, or numbers in front of each chemical formula, tell us the relative number of molecules involved in the reaction. They also tell us the relative number of moles involved in the reaction. Two moles of benzene will react with 15 moles of dioxygen to form twelve moles of carbon dioxide and six moles of water. It is important to note that the equation does not give us any direct information about the number of grams of each reactant or product, only the moles. If we want to know about a measurable quantity like grams, we will have to do some conversions. The first requirement for any stoichiometric calculation is a balanced equation. Once we have a balanced equation, the calculations will be performed by following three steps. 1. Convert the given quantitative information to moles. Experiments are always set up (thus chemistry problems are always written) such that number of moles of at least one reactant or product can be determined. 2. Use the balanced equation to convert from the moles of given substance to the moles of desired substance. Remember, we re given quantitative information on one reactant or product, we desire quantitative information on another. This is the heart of all stoichiometry problems. 3. Convert from moles of the desired compound to the appropriate quantity. These three steps are purposely vague. In the three previous appendices, we have discussed the conversion of moles to grams for a solid, moles to pressure, volume and temperature for a gas, and moles to volume and molarity for a solution. As you might guess, there are several variations on this three-step theme. The following examples will show some of the variety. D.3 REACTION STOICHIOMETRY INVOLVING GRAMS For calculations involving grams of reactants and products, our three-step scheme looks like this: grams of A grams of B M m moles of A coef of B x coef of A xm m moles of B In the preceding diagram, A and B are products and/or reactants. The gram to mole conversion is achieved through application of the molar mass, the mole to mole conversion comes from the coefficients in the balanced equation.
2 Example 1 Consider the combustion of benzene: 2C 6(l) + 15O 2(g) 12CO 2(g) + 6H 2O(g) What is the maximum mass of CO 2 (M m = 44.0 g/mol) that can be produced from the combustion of 10.0 g of C 6 (M m = 78.1 g/mol)? The given information is grams of benzene. The desired information is grams of CO 2. The road map above tells us we must convert grams of benzene to moles, use the mole ratio from the equation to give moles of carbon dioxide, and then convert back to grams of carbon dioxide g C 6 1 mol C g C 6 = mol C mol C 6 12 mol CO 2 2 mol C 6 = mol CO mol CO g CO 2 mol CO 2 = 33.8 g CO 2 As we have seen several times in the previous appendices, each individual step is not difficult. Putting the steps together in a logical manner is the challenge. In this example, we have done each of the three steps separately. However, we could have strung together the conversion factors to save ourselves some writing. We will do that in the next example. Example 2 How many grams of O 2 (M m = 32.0 g/mol) are required to completely react with 10.0 g of C 6? The given information is the mass of benzene, the desired information is the mass of dioxygen. Start with the given information and apply conversion factors following the road map g C 6 1 mol C g C 6 15 mol O 2 2 mol C g O 2 1 mol O 2 = 30.7 g O 2 Stoichiometry problems do not always relate reactants to products. Here is a situation where both the given and desired information deals with reactants. In this example, we have done our three step calculation by stringing together the three conversion factors appropriate to the three steps. Note that the order of operation can be determined by using the units because the units of the denominator of each conversion factor must be the same as the units of the previous numerator. Using the units to help is called the factor-label method. D.4 LIMITING REACTANTS We know from experience that the amount of product that is formed depends on the amount of reactant that is consumed. You can drive a car only as long as it has gasoline. The gasoline is the limiting reactant because it dictates how much product (miles) can be achieved. The amount of gasoline determines not only how far you can go, but it also determines how much CO and H 2 2O (the reaction products of the combustion of gasoline) can be made. In any chemical reaction, the amount of products that are made is limited by the amount of reactants. When any one reactant runs out, the reaction stops. The reactant that runs out is called the limiting reactant or limiting reagent. Any reactants that do not run out are said to be in excess. In most chemical reactions, one or more of the reactants is in excess. In the gasoline combustion reaction, there is certainly more oxygen available then there is gasoline in the gas tank, and so the oxygen is in excess. In calculating the amount of product formed in a reaction, we always have to identify the limiting reactant. In some cases it is obvious. In Example 1 above, we read that 10.0 g of C 6 reacts with excess O 2. Clearly, C 6 is the limiting reactant, and O 2 is the excess reactant. But consider the following example.
3 Calcium hydride reacts with water to form calcium hydroxide and hydrogen gas, via the following reaction CaH 2(s) + 2H 2O(l) Ca(OH) 2(s) + 2H 2(g) If 10.0 g of CaH 2 reacts with 9.00 g of H 2 O, what mass of Ca(OH) 2 can be formed? We start by determining the limiting reactant, but we cannot tell which reactant will limit the amount of product just by comparing the grams of each reactant, we must determine how much Ca(OH) 2 we can make from each reactant. Based on the amount of CaH 2 we start with, we can make: 10.0 g CaH 2 1 mol CaH g CaH 2 1 mol Ca(OH) 2 1 mol CaH g Ca(OH) 2 1 mol Ca(OH) 2 = 17.6 g Ca(OH) 2 Based on the amount of H 2 O available, we can make: 9.0 g H 2 O 1 mol H 2 O 18.0 g H 2 O 1 mol Ca(OH) 2 2 mol H 2 O g Ca(OH) 2 1 mol Ca(OH) 2 = 18.5 g Ca(OH) 2 Even though we have enough water to make 18.5 g of Ca(OH) 2, there is only enough calcium hydride to make 17.6 g. In this case, CaH 2 is the limiting reactant, H 2 O is in excess, and 17.6 g of Ca(OH) 2 would be produced. Let s determine how much of the excess reactant remains. CaH 2 is the limiting reactant and all amounts are calculated from it, and so we must now determine how much water reacts with the CaH g CaH 2 1 mol CaH 2 2 mol H 2 O 18.0 g H 2 O g CaH 2 1 mol CaH 2 1 mol H 2 O = 8.56 g H 2 O The above is how much water reacts, we now determine how much remains by subtracting the amount that reacts from the initial amount = 0.44 g H 2O remains Let s summarize limiting reactants with another example. Example g of Na 2 is added to 7.00 g of carbon and allowed to react according to the following equation: Na 2(s) + 4C(s) Na 2S(s) + 4CO(g) a) What is the limiting reactant? b) How many grams of Na 2S can be formed? c) How many grams of the excess reactant will be leftover? Perform the three-step calculation twice, starting from the information given for each reactant g Na 2 1 mol Na g Na 2 1 mol Na 2 S 1 mol Na g Na 2 S 1 mol Na 2 S = 13.7 g Na 2 S 7.00 g C 1 mol C g C 1 mol Na 2 S 4 mol C g Na 2 S 1 mol Na 2 S = 11.4 g Na 2 S Comparing the two calculations leads us to conclude that C is the limiting reactant and Na 2 is in excess g of Na 2S can be produced. Next, find the amount of excess reactant that reacts g C 1 mol C g C 1 mol Na g Na 2 = 20.8 g Na 4 mol C 1 mol Na 2 SO 2 4 The amount remaining is given by the difference 25.0 g initially g consumed = 4.2 g of Na 2 remain
4 D.5 REACTIONS INVOLVING GASES Examples 1 and 2 dealt with the combustion of benzene. In each, a known amount of benzene was burned, and we calculated the mass of CO 2 or O 2 produced. However, the measurable quantities of gases are pressure, volume and temperature, not mass. In this section, we use the treatment presented in Appendix B to introduce these quantities into our stoichiometric calculations. If there is known quantitative information about the pressure, volume and temperature of a gas, we can calculate the moles of that gas, using PV = nrt. We can use this relationship in the first or last step of our calculation. We summarize this through the following road map or flowchart: grams of A grams of B P, V, and T of A PV RT M m moles of A coef of B x coef of A xm m moles of B P, V, or T of B Notice that this flowchart is identical to the one that appeared earlier in this appendix, except that the use of P,V, and T information for a gas has been added as an entry into the scheme at the left and as a result from the scheme at the right. The following examples show how the ideal gas law can be used in reaction stoichiometry calculations. Example 4 Octane combusts via the following chemical reaction: 2C 8H 18(l) + 25O 2(g) 16CO 2(g) + 18H 2O(g) How many liters of CO 2, collected at a pressure of 1.00 atm and a temperature of 25 o C, can be produced by the combustion of 35.0 kg of octane (the amount held by a typical car gasoline tank)? We first must recognize this as a reaction stoichiometry problem. We know the mass and therefore the number of moles of octane. We desire the volume of the carbon dioxide, so PV = nrt will be our third step. It is somewhat complicated to use the ideal gas law as a conversion factor, so we will string together the first two steps, and then do the third step separately, solving for volume. 35 kg C 8 H g kg 1 mol C 8 H g C 8 H mol CO 2 2 mol C 8 H 18 = 2.45x10 3 mol CO 2 Next apply the ideal gas law, PV = nrt. V = nrt P ( mol)( L atm K -1 mol -1 )(298 K) = = 5.99x10 4 L 1.00 atm Burning a tank of gasoline generates enough carbon dioxide at 1.00 atm and 25 o C to fill a 16 ft x 16 ft room with an 8 ft ceiling. Example 5 The following reaction is used to quickly inflate some car airbags: 2NaN 3(s) 2Na(s) + 3N 2(s) How many grams of NaN 3 must be used if you wish to fill a 20.0 L airbag with dinitrogen to a pressure of 1.25 atm at 25 o C? In this problem the known information involves a gaseous product (a pressure, volume and temperature are all given) and information on the reactant is desired. This time, the first of the three steps involves the ideal gas law, with the second and third steps being simple application of conversion factors.
5 n = PV RT = (1.25 atm)(20.0 L) = 1.02 mol N ( L atm K -1 mol -1 2 )(298 K) 1.02 mol N 2 2 mol NaN 3 3 mol N g NaN 3 1 mol NaN 3 = 44.2 g NaN 3 D.6 REACTIONS INVOLVING SOLUTIONS In the previous sections, we have discussed the quantitative relationships between reactants and products in a reaction. In all cases, the heart of the problem was the mole ratio, and the only difference in the problems lies in how we get to the mole ratio and what we do after we have applied it. We will now consider our last conversion to and from moles. If a reactant or product is dissolved in solution, the moles of that compound are related to the molarity and the solution volume, M = n/v. (See Appendix C if you do not remember this relationship.) We can add this route to our road map of reaction stoichiometry. grams of A grams of B P, V, and T of A PV RT M m moles of A MxV coef of B x coef of A xm m moles of B P, V, or T of B MandVofA MorVofB The following examples show how solution data can be manipulated along with masses and data on gases to give information on reaction stoichiometry. Example 6 The Pb 2+ ions from water soluble Pb(NO 3) 2 can be precipitated by the addition of KI, forming insoluble PbI 2: Pb(NO 3) 2(aq) + 2KI(aq) PbI 2(s) + 2KNO 3(aq) If 25.0 ml of M Pb(NO 3) 2 is reacted with excess KI, how many grams of PbI 2 will be produced? The moles of Pb(NO 3) 2 can be found since the volume and molarity of the solution are known. We convert the volume to liters then apply molarity to determine moles. The final two steps of the calculation are the mole to mole conversion from the balanced equation, and the conversion from moles of PbI 2 to grams, using the molar mass of g/mol ml mol Pb(NO 3 ) ml 1 mol PbI g PbI 2 = 4.32 g PbI 1 mol Pb(NO 3 ) 2 1 mol PbI 2 2 Example 7 The silver ions in AgNO 3 solution can be precipitated by the addition of aqueous NaCl: AgNO 3(aq) + NaCl(aq) AgCl(s) + NaNO 3(aq) When 35.0 ml of a AgNO 3 solution of unknown concentration is reacted with excess NaCl solution, 8.53 g of AgCl is formed. What is the concentration of the AgNO 3 solution? In this problem, there is some information about the AgNO 3 solution and some information about the solid AgCl. After rereading the problem, it should become clear that the desired quantity is molarity of the AgNO 3 solution. Thus, we need to start at the other end, with the mass of AgCl. You should also recognize that whenever mass data is presented along with either a molar mass or a chemical formula (from which we can get a molar mass), we have an entry into our road map. Here, we will do the first two steps of the calculation in the usual manner, and then use the M = n/v relationship as our third step.
6 1 mol AgCl 8.53 g AgCl g AgCl 1 mol AgNO 3 1 mol AgCl M = n mol = = 1.70 mol/l = 1.70 M V L = mol AgNO 3 The first two steps are as we have done many times now. The third step, using the molarity relationship, may at first seem a little unusual. Molarity is always the ratio of moles of a substance to the volume in liters. The first two steps tell us that there are moles of AgNO 3 contained in the 35.0 ml of solution. We simply take the ratio of these two numbers, first converting 35.0 ml to L, since molarity is moles per liter. Note that the experiment did not have to be done on a 1 L scale in order to calculate the molarity! Example 8 Potassium permanganate solutions can react with acidic hydrogen peroxide solutions via the following balanced equation in water: 2KMnO 4 + 5H 2O 2 + 6HCl 2MnCl 2 + 5O 2(g) + 2KCl + 8H 2O(l) When 15.0 ml of M KMnO 4 reacts with excess H 2O 2 and HCl, how many liters of O 2, collected at a total pressure of 1.00 atm and a temperature of 27 o C, will be formed? The volume and molarity data on KMnO 4 allow us to enter into the road map. The desired information is volume of O 2. We start by converting ml to L, and then proceed through the first two of the three steps in the calculation to find moles of O 2. We will do the third step, manipulation of PV = nrt, separately ml mol KMnO m L 5 mol O 2 = mol O 2 mol KMnO 2 4 V = nrt P ( mol) ( L atm K -1 mol -1 ) (300 K) = = L 1.00 atm As you can see, there are many possible routes along our road map for solving stoichiometry problems. Example 9 How many ml of a M NaOH solution are required to completely react with 40.0 ml of a M H 2 solution. The overall reaction is: H 2(aq) + 2NaOH(aq) 2H 2O(l) + Na 2(aq) We see that there is enough information to calculate moles of H 2. We desire information on NaOH. Start with the volume of L of H 2 solution and proceed as usual L mol H 2 1 L 3 L 2 mol NaOH 1 L 1 mol H mol NaOH 10 ml = 48.0 ml The ml to L conversion was done at the beginning and the end of the problem since the volume information is in ml, but the concentration (molarity) is, of course, in moles per liter. Both molarities were used as conversion factors; M converted volume of solution to moles of H 2, M was turned upside down to convert moles of NaOH to liters of solution.
7 Example 10 Aqueous HCl and NaOH react in the following manner: HCl(aq) + NaOH(aq) H 2O(l) + NaCl(aq) When 30.0 ml of M HCl are mixed with 20.0 ml of M NaOH, what is concentration of the excess reagent? This is a limiting reactant problem. At first glance, it would appear quite different than the limiting reactant problems we saw earlier. However, upon careful reading of the experiment, we see that we have quantitative information on both reactants, enough to calculate moles of both. The desired quantity the concentration of the excess reactant. In essence, we are reacting an acid and a base, and need to determine which reactant is limiting, and find how much of the excess reactant is leftover, as was done in Example 3 above. In order to determine the limiting reactant, we calculate how much product can be made from each reactant. It doesn t matter which product we choose. Let s pick water. 1 L mol HCl 30 ml 10 3 ml 1 L 1 mol H 2 O 1 mol HCl = mol H 2 O 1 L mol NaOH 20 ml 1 mol H 2 O 10 3 ml 1 L 1 mol NaOH = mol H 2 O Fewer moles of water can be made from the NaOH, so NaOH is the limiting reactant, HCl is the excess reactant. Notice that it was not necessary to go all the way through and calculate the grams of water. Clearly, if we multiply each result by g/mol (the molar mass of H 2O), the conclusion is the same, NaOH is limiting. In order to calculate molarity of HCl, we need the number of moles of HCl which were leftover, and the total solution volume. Moles of HCl at the start: 1 L mol HCl 30 ml = mol HCl 10 3 ml 1 L The number of moles of HCl consumed is based on the amount of limiting reactant consumed: 1 L mol NaOH 1 mol HCl 20 ml = mol HCl react 10 3 ml 1 L 1 mol NaOH Moles of HCl remaining: = moles of HCl remain Concentration of HCl at the end moles of HCl mol HCl = = M total volume L This problem is actually very similar to the limiting reactant problems we did before. The difference is that, instead of finding the grams of the leftover reactant, we had to find the concentration, which involved a calculation of the number of moles of the leftover reactant.
8 D.7 EXERCISES Use the following molar masses to do the following problems: C4H8 : g/mol C4H9OH : g/mol Fe2O3 : g/mol Al2O3 : g/mol V2O5 : g/mol NH4VO3 : g/mol NH3 : g/mol V2O3 : g/mol Cu2S : g/mol CuO : g/mol Cu2O : g/mol AgCl : g/mol 1. In the presence of acids, water can react with alkenes to form alcohols: C4H8 + H2O C4H9OH If 250 g of C4H8 reacts with excess H2O, how many grams of C4H9OH can be produced? 2. Aluminum reacts with iron(iii) oxide in the thermite reaction : 2Al(s) + Fe2O3(s) 2Fe(s) + Al2O3(s) a) If 10.0 g of Al reacts with excess Fe2O3, how many grams of Al2O3 can be produced? b) If 25.0 g of Al reacts with 10.0 g of Fe2O3, how many grams of Al2O3 can be produced? c) In the experiment in part b, what is the mass of the excess reactant remaining after complete reaction? 3. Vanadium(V) oxide reacts with ammonia and water as follows: V2O5 + 2NH3 + H2O 2NH4VO3 a) If 50.0 g of V2O5 is reacted with excess ammonia and water, how many grams of NH4VO3 can be produced? b) How many grams of NH3 are required to completely react with 50.0 g of V2O5? 4. Vanadium(III) oxide can be made by reduction of vanadium(v) oxide with hydrogen: V2O5(s) + 2H2(g) V2O3(s) + 2H2O(l) a) How many liters of H2, measured at 1.00 atm and 30 o C, are required to completely react with 75.0 g of V2O5? b) If 10.0 g of V2O5 reacts with 1.65 L of H2, measured at 1.00 atm and 30 o C, how many grams of V2O3 can be produced? 5. Copper(I) sulfide is prepared by heating copper and sulfur in the absence of air: 2Cu(s) + S(s) Cu2S(s) a) How many grams of Cu2S can be produced from the reaction of 25.0 g of Cu with excess S? b) How many grams of sulfur are required to form 75.0 g of Cu2S? c) If a mixture of 135 g of Cu and 45 g of S is allowed to react, how many grams of Cu2S could be produced? d) How many grams of the excess reactant remain in the experiment in part c? 6. Copper(I) oxide can be prepared by thermal decomposition of copper(ii) oxide: 4CuO(s) 2Cu2O(s) + O2(g) a) How many grams of Cu2O can be produced upon the decomposition of 450 g of CuO? b) How many liters of O2, collected at 1.00 atm and 27 o C, can be produced by the decomposition of 450 g of CuO? 7. The silver ions in aqueous silver sulfate can be precipitated by addition of excess chloride: Ag2SO4(aq) + 2NaCl(aq) 2AgCl(s) + Na2SO4(aq) a) How many grams of silver chloride can be formed when 35.0 ml of a M Ag2SO4 solution is reacted with excess sodium chloride solution? b) If 22.7 ml of a silver sulfate solution of unknown concentration yields g of AgCl upon reaction with excess sodium chloride solution, what is the concentration of the silver sulfate solution? 8. Zn metal reacts with hydrochloric acid to produce hydrogen gas and zinc(ii) chloride: Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g) a) If 15.0 g of Zn are added to excess HCl(aq), how many liters of H2(g), collected at 27 o C and 725 mm Hg, are produced? b) If excess Zn is added to 25.0 ml of M HCl(aq), how many liters H2(g), collected at 27 o C and 725 mm Hg, can be produced?
9 9. Potassium permanganate and iron(ii) chloride undergo an electron transfer reaction in acid solution: KMnO 4 (aq) + 5FeCl 2 (aq)+ 8HCl MnCl 2 (aq)+ 5FeCl 3 (aq) + KCl(aq) + 4 H 2 O(l) How many ml of M FeCl2(aq) are needed to completely react with 13.7 ml of M KMnO4? 10. Citric acid reacts with sodium hydroxide in a proton transfer reaction: H3C6H5O7(aq) + 3NaOH(aq) 3H2O(l) + Na3C6H5O7(aq) a) How many ml of M NaOH(aq) are required to completely react with 25.0 ml of M citric acid? b) If 37.5 ml of 1.25 M NaOH(aq) is needed to completely react with 22.5 ml of a citric acid solution, what is the concentration of the citric acid solution? ANSWERS: g 2. a) 18.9 g b) 6.38 g c) 21.6 g 3. a) 64.3 g b) 9.34 g 4. a) 20.5 L b) 4.97 g 5. a) 31.3 g b) 15.1 g c) 169 g d) 11 g 6. a) 405 g b) 34.8 L 7. a) 1.00 g b) M 8. a) 5.92 L b) 8.07 ml ml 10. a) 41.7 ml b) M
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Moles I. Lab: Rice Counting II. Counting atoms and molecules I. When doing reactions chemists need to count atoms and molecules. The problem of actually counting individual atoms and molecules comes from
EXPERIMENT 7 Reaction Stoichiometry and Percent Yield INTRODUCTION Stoichiometry calculations are about calculating the amounts of substances that react and form in a chemical reaction. The word stoichiometry
CHEMISTRY 123-07 Midterm #1 Answer key October 14, 2010 Statistics: Average: 74 p (74%); Highest: 97 p (95%); Lowest: 33 p (33%) Number of students performing at or above average: 67 (57%) Number of students
Practice questions for Chapter 8 2) How many atoms of nickel equal a mass of 58.69 g? (Refer to the Periodic Table.) A) 1 B) 28 C) 58.69 D) 59 E) 6.02 x 1023 Answer: E Section: 8.1 Avogadro's Number 6)
Chemical Reactions Chemical Equations Chemical reactions describe processes involving chemical change The chemical change involves rearranging matter Converting one or more pure substances into new pure
SCH3U- R.H.KING ACADEMY SOLUTION & ACID/BASE WORKSHEET Name: The importance of water - MAKING CONNECTION READING 1. Read P. 368-375, P. 382-387 & P. 429-436; P. 375 # 1-11 & P. 389 # 1,7,9,12,15; P. 436
Chapter 6 Oxidation-Reduction Reactions 1. Oxidation is defined as a. gain of a proton b. loss of a proton c. gain of an electron! d. loss of an electron e. capture of an electron by a neutron 2. Which
SCH 4C1 Unit 2 Problem Set Questions taken from Frank Mustoe et all, "Chemistry 11", McGraw-Hill Ryerson, 2001 1. A small pin contains 0.0178 mol of iron. How many atoms of iron are in the pin? 2. A sample
Name: Tuesday, May 20, 2008 1. What is the molecular formula of a compound with the empirical formula PO and a gram-molecular mass of 284 grams? 2 5 1. P2O 5 3. P10O4 2. P5O 2 4. P4O10 2. Which substance
Chemical Reactions in Water Ron Robertson r2 f:\files\courses\1110-20\2010 possible slides for web\waterchemtrans.doc Properties of Compounds in Water Electrolytes and nonelectrolytes Water soluble compounds
Mass, Moles, & Molar Mass Relative quantities of isotopes in a natural occurring element (%) E.g. Carbon has 2 isotopes C-12 and C-13. Of Carbon s two isotopes, there is 98.9% C-12 and 11.1% C-13. Find
Stoichiometry Atomic Mass (atomic weight) Atoms are so small, it is difficult to discuss how much they weigh in grams We use atomic mass units an atomic mass unit (AMU) is one twelfth the mass of the catbon-12
Formula Conventions (1 of 24) Superscripts used to show the charges on ions Mg 2+ the 2 means a 2+ charge (lost 2 electrons) Subscripts used to show numbers of atoms in a formula unit H 2 SO 4 two H s,
Chem. I Notes Ch. 12, part 2 Using Moles NOTE: Vocabulary terms are in boldface and underlined. Supporting details are in italics. 1 MOLE = 6.02 x 10 23 representative particles (representative particles
PURPOSE EXPERIMENT 5: CHEMICAL REACTIONS AND EQUATIONS To perform and observe simple chemical reactions. To identify the products of chemical reactions and write balanced equations for those reactions.
Larry Brown Tom Holme www.cengage.com/chemistry/brown Chapter 3 Molecules, Moles, and Chemical Equations Jacqueline Bennett SUNY Oneonta 2 Warning!! These slides contains visual aids for learning BUT they
1. How many iron atoms are present in one mole of iron? Ans. 6.02 1023 atoms 2. How many grams of sulfur are found in 0.150 mol of sulfur? [Use atomic weight: S, 32.06 amu] Ans. 4.81 g 3. How many moles
Chapter 3 Stoichiometry 3-1 Chapter 3 Stoichiometry In This Chapter As you have learned in previous chapters, much of chemistry involves using macroscopic measurements to deduce what happens between atoms
UNIT 6 stoichiometry practice test True/False Indicate whether the statement is true or false. moles F 1. The mole ratio is a comparison of how many grams of one substance are required to participate in
Name: Thursday, March 06, 2008 Solutions Review Questions 1. Compared to pure water, an aqueous solution of calcium chloride has a 1. higher boiling point and higher freezing point 3. lower boiling point
Chemistry Zumdahl, 7th edition CH4 Types of Chemical Reactions and Solution Stoichiometry Yellow lead(ii) iodide is produced when lead(ii) nitrate is mixed with potassium iodide. P.126 Contents 4.1 Water,
Stoichiometry 1 The total number of moles represented by 20 grams of calcium carbonate is (1) 1; (2) 2; (3) 01; (4) 02 2 A 44 gram sample of a hydrate was heated until the water of hydration was driven
Chapter 3. Stoichiometry: Mole-Mass Relationships in Chemical Reactions Concept 1. The meaning and usefulness of the mole The mole (or mol) represents a certain number of objects. SI def.: the amount of
Honors Chemistry: Unit 6 Test Stoichiometry PRACTICE TEST ANSWER KEY Page 1 1. 2. 3. 4. 5. 6. Question What is a symbolic representation of a chemical reaction? What 3 things (values) is a mole of a chemical
Stoichiometry and Aqueous Reactions (Chapter 4) Chemical Equations 1. Balancing Chemical Equations (from Chapter 3) Adjust coefficients to get equal numbers of each kind of element on both sides of arrow.
TEKS REVIEW 8B Calculating Atoms, Ions, or Molecules Using Moles TEKS 8B READINESS Use the mole concept to calculate the number of atoms, ions, or molecules in a sample TEKS_TXT of material. Vocabulary
The Mole Atomic mass units and atoms are not convenient units to work with. The concept of the mole was invented. This was the number of atoms of carbon-12 that were needed to make 12 g of carbon. 1 mole
Chem 31 Fall 2002 Chapter 3 Stoichiometry: Calculations with Chemical Formulas and Equations Writing and Balancing Chemical Equations 1. Write Equation in Words -you cannot write an equation unless you
Worksheet # 11 1. A solution of sodium chloride is mixed with a solution of lead (II) nitrate. A precipitate of lead (II) chloride results, leaving a solution of sodium nitrated. Determine the class of
Symbols used in chemical equations: Symbol Meaning + used to separate one reactant or product from another used to separate the reactants from the products - it is pronounced "yields" or "produces" when
Experiment 5 Chemical Reactions OBJECTIVES 1. To observe the various criteria that are used to indicate that a chemical reaction has occurred. 2. To convert word equations into balanced inorganic chemical