Ch+5+Reading+Guide

__**Reading Guide for Ch 5: Chemical Reactions and Quantities**__ The main point of Chapter 5, Chemical Reactions and Quantities, is to help the reader to relate given amounts of compounds, in moles or in units of mass, to chemical formulas and to reactants and products from balanced equations.

__5.1: Chemical Changes__ ->Summary: Chemical changes are different than physical changes because in chemical changes, reacting substances change into new substances with different chemical formulas and properties. Chemical reactions always cause chemical changes because the atoms in the reactants are rearranged in the molecules of the products.

->Struggling Topic 1: State changes, such as liquid water freezing to a solid, are physical changes because the chemical properties and formulas remain the same before, after and during the state change. I have difficulty teaching this concept to students because the state change does seem to create new properties such as the changes in volume, density, temperature, and other factors that students easily identify as having changed. Reflecting on the abstract concept that the chemical properties and formulas are identical is a challenge to some learners because it is so abstract.

->Struggling Topic 2: Dissolving a solute into a solvent is also a physical change that is easily confused for a chemical change because although the chemical properties and formulas of both substances are still present before, during, and after the dissolving of the solute, the substances take on a changed appearance. This is slightly easier for students to grasp because the properties of both solute and solvent are still present. For example, salt water still tastes salty although it behaves in every other way like ordinary water. I also use the idea that physical changes are typically easier to reverse than chemical changes. It is fairly easy to boil or evaporate the water from the salt in salt water compared to trying to un-rust a nail or un-burn a piece of paper.

__5.2: Chemical Equations__ -> Summary: A chemical equation shows what reactants are present before a chemical reaction and what products are formed as a result of a chemical reaction. A balanced chemical equation also shows how no atoms are created or destroyed in the chemical reaction, but that all the atoms are simply rearranged to form new substances. Chemical equations must be balanced by changing the coefficients for chemical formulas so that the same number of each kind of atom is present on both sides of the equation.

->Struggling Topic 1: Students often struggle with finding the total of each kind of atom on the product and reactant side of the equation when coefficients must be combined with subscripts in chemical formulas. For example, asking a student how many phosphorous atoms are present in 2 Mg3(PO4)2 is quite difficult for students. The author shows how many of each element's atoms are present on the reactant and product side when balancing equations, but the author doesn't explicitly state where these numbers come from. I envision some students needing it spelled out that 2 Mg3(PO4)2 must be "distributed" for lack of a better term to determine how many Oxygen atoms are present. A single phosphate ion (PO4) has 4 Oxygen atoms, but in each magnesium phosphate there are two phosphate ions. Thus we multiply the 4 Oxygen atoms in a single phosphate ion by the number of phosphate ions in this compound and get 8 Oxygen atoms per magnesium phosphate. Lastly, the balanced equation tells us there must be two magnesium phosphates, and we know that every magnesium phosphate has 8 Oxygen atoms. This means that we must multiply our 8 Oxygen atoms by the number of magnesium phosphates present in the reactants of this balanced equation, or in this case, by 2 to get a total of 16 Oxygen atoms in 2 Mg3(PO4)2. I think the author assumes that students are able to do this already. This is not a fair assumption for all of our students. ---> Later in this chapter, in section 5.5 on the mole, the author begins to break this down a little more specifically when she discusses how the subscript, //n//, of an atom in a chemical formula, 2AnB, means //n// times the number of moles of the substance. In this case, 2//n// moles of A while only 2 moles of B.

->Struggling Topic 2: Why must the arrow, -->, be used in chemical equations instead of an equal sign, =? I find that students often think in their heads "product + product = reactant + reactant" even though they see an arrow or yields sign. It is important that the author explain that an equal sign means something different than a yields sign. Equal signs, by definition, can be reversed upon themselves as to say A+B=C; therefore, C=A+B. With chemical equations, this is not always the case. Oftentimes the reverse reaction of a chemical equation is significantly different and cannot be written as an equilibrium chemical equation. Because of this, we distinguish between one way reactions (with a yield sign -->) and two way equilibrium reactions (with a two way yields sign <-->). This topic should be explained to the students as soon as they hear about balancing equations and forming equations so that they do not begin to think "equals" when they see "yields."

__5.3: Types of Reactions__ -> Summary: Most chemical reactions can be classified into general reaction types, combination, decomposition, or replacement reactions. Combination reactions involve multiple reactants that combine to form a single product. Decomposition reactions involve single reactants that break down to form multiple products. Replacement reactions involve multiple reactants and multiple products whose ions or elements separate and bond with different ions or elements.

-> Struggling Topic 1: Can classifying these reactions be simply grouped into by the following observation? An equation with only 1 reactant is decomposition, an equation with only 1 product is combination, and any equation with more than 1 product and reactant is a replacement reaction. I am struggling to find an example that disproves this statement.

This is a good generalization. The problem is that when you get into more complicated system then these generalities do not work. There are way more than only four type reaction categories.

-> Struggling Topic 2: In understanding the difference between single and double replacement reactions, what happens if more than 2 reactants are involved? Say AbCd + EfGh + Ij --> AbEf + GhIj + Cd. This reaction seems to have a single replacement and double replacement reaction going on between all the pieces. Are there reactions that combine or blur the line between single and double replacement? Would these be written as half-reactions?

__5.4: Oxidation-Reduction Reactions__ -> Summary: Oxidation-reduction reactions are chemical reactions with a loss and gain of electrons among the reactants. The oxidized reactant loses at least one electron, and the reduced reactant gains at least one electron. These two half-reactions must occur at the same time and with the same number of electrons.

->Struggling Topic 1: Writing half reactions makes it easy to see which reactant is oxidized and which is reduced. The author may have lost some students by not explaining exactly what a half-reaction is. The example in the textbook puts the full reaction just above the two half-reactions, but a verbal explanation could help as well. A half-reaction includes only 1 reactant from the full reaction and a number of electrons (on either the product or reactant side depending upon if it is the oxidation or reduction half-reaction). The half-reaction basically puts on paper the thought process that examines the reactants that changed charges during the full chemical reaction. By writing only the one reactant at the beginning of the reaction and the end of the reaction, it is easier to see whether that reactant lost or gained electrons to acquire a new charge after the reaction.

->Struggling Topic 2: The author assumes that the student will be able to rewrite an equation with ionic compounds as an equation with all the individual ions separated with charges. Some students will struggle to break these ionic compounds correctly. The example in the textbook included a sulfate ion that some students may try to break into its elements. A quick explanation that the copper (II) sulfate can be rewritten as the copper (II) ion, Cu^2+, and the sulfate ion, SO4^2-, could help some of the students who didn't understand the sections on ions as well the first time review this concept.

__5.5: The Mole__ -> Summary: The mole is a unit of measurement that includes 6.02 x 10^23, Avogadro's number, items. Thus, Avogadro's number can be used as a conversion factors to determine how many particles, molecules, formula units, ions, or other units of a substance are present in a mole of that substance. This is helpful because we are usually dealing with so many particles at once, that looking at a smaller unit, the mole, is easier to understand.

-> Struggling Topic 1: Avogadro's number is important because it allows us to define a mole of items, and moles are very commonly used to describe numbers of particles, ions, and molecules. How did a mole come to include 6.02 x 10^23 items? Is there any particular reason why Avogadro's number couldn't have been a different number? Oftentimes it seems like any other conversion factor doesn't have value except that it allows us to deal with smaller and more reasonable numbers. Is this the case with Avogadro's number?

-> Struggling Topic 2: Formula units are the group of ions that make up an ionic compound much as molecules are groups of covalently bonded atoms. What exactly makes it necessary to have separate titles for formula units and molecules? Is it the way that molecular compounds remain closely joined in the sharing of electrons in their covalent bonds while ions in ionic compounds do not remain closely joined because they gave and received electrons and can therefore easily dissociate in solutions that makes the difference between these two titles? I suppose in earlier chemistry understanding, a compound and molecule were equivalent when truthfully, only covalently bonded compounds are molecules.

__5.6: Molar Mass__ -> Summary: The molar mass of an element is a value for how many grams are in a single mole of that element. The molar mass of a compound is found by adding the molar mass of each atom in the compound. A substance's molar mass can be used as a conversion factor to convert moles of a substance into grams of a substance or grams of a substance into moles of a substance.

-> Struggling Topic 1: I was glad to read that the author will round molar mass to the nearest tenths (0.1) place or giving at least 3 significant figures. When doing my homework problems, I often used the most precise molar mass I could find on the periodic table, and the program usually had rounded differently than I had. Many periodic tables have molar mass written out to the thousandths or even ten-thousandths place for some elements. It will remove some frustrations by students to know how far we ought to round these molar mass values.

-> Struggling Topic 2: Many students struggle to keep the conversion units straight between molar mass, Avogadro's number, and the subscripts in individual chemical formulas. The flow chart on p.197 from grams of a compound to moles of a compound to molecules/formula units or to atoms/ions would be quite useful to help students see the big picture of the many conversions that we can do with these concepts.

__5.7: Mole Relationships in Chemical Equations__ ->Summary: A balanced equation can be used to form proportions between the amounts of each substance present in the reaction. Coefficients from balanced equations form proportions between moles of reactants and products. Then these proportions between products and reactants can be used as conversion factors.

->Struggling Topic 1: Students often struggle to keep their units straight as they perform conversions between one reactant or product and another. Do you have suggestions for how to get students to show their work WITH UNITS. I remember as a high school student simply writing //x// number of grams, and then later wishing I knew grams of what. Now with teacher eyes it's easy to see how writing the //x// g H2O for example can save a headache later. Suggestions?

->Struggling Topic 2: Students struggle with limiting reactants in problems that are quite similar to these. Is there a way to introduce these conversions that will help students grasp limiting reactants with more ease later? I would think that we could mention something about //x// amount of reactant A is needed to get //y// amount of product B. The problems in the practice set are worded so that the author can avoid mentioning limiting reactants yet "How many moles are needed to react with _ and _," but I think the typical student thinks in less particular terms " _ moles of _ get me _ moles of that." Although the author's wording is correct, the student's thinking may leave out the check for a limiting reactant step. Perhaps this topic would be easier if there was mention of it earlier in the sequence of instruction.

__5.8: Mass Calculations for Reactions__ ->Summary: Using mole to mole conversion factors from balanced equations allows us to find how many moles of product can be produced given a certain number of moles of reactant. Combining molar mass to this procedure allows us to begin with mass of reactant, convert to moles of reactant, convert to moles of product, and finally convert to mass of product. This is much more easily applied to everyday science because measuring mass is much easier than measuring moles.

->Struggling Topic 1: Just as in 5.7, I think the most challenging part of these conversions is keeping your units straight. Our book does an excellent job illustrating the proper way to show units. Every number has a unit (moles C2H2, or g CO2). Is there a good way to convince students that this is needed?

Most of our students are going to the nursing field. Let's say the student ask the head nurse how much drug do I need to give to the patient and the head nurse says 500. What does tha mean, most likely the student would assume 500 mg, but what about if it was 500 ug, then the student just killed a patient.

->Struggling Topic 2: Significant figures and molar mass is troublesome for me. Does molar mass count as an exact number and get to be ignored in the factoring of the significant figures that belong in my answer? (yes it is used an exact number. Even though it was determined experimentally, the number of significant figures we are using are highly accurate) I remember a few homework problems where the homework program and I got slightly different answers because of rounding. I think this might have been because of the way I took as many digits as I could get when calculating molar mass. This was before I read that the author says round to the tenths place or to 3 significant figures for molar mass.

__5.9: Percent Yield and Limiting Reactants__ ->Summary: The amount of product that is actually produced is always less than the amount that the calculations say should be produced. The actual and expected amounts are used to calculate the percent yield of the reaction. Also, reactions rarely use up all of their reactants. Oftentimes one reactant is used up and there are leftovers from other reactants. The reactant that is used up is the limiting reactant, and the others are excess reactants. When calculating the expected yield for a reaction, you must take the limiting reactant into account. To do this, you can simply calculate the expected yield of product based upon all the given amounts of reactant, but then be sure to choose the smallest yield of all the possible yields because this one will be from the limiting reactant. More of this limiting reactant would be needed to produce any of the larger possible yields that were calculated.

->Struggling Topic 1: I don't recall any practice problems asking for the mass of excess reactants after the reaction has occurred. Is this a viable question? I was intrigued by the table from page 207 that showed 2 moles of hydrogen gas reacting with 5 moles of chlorine gas to produce 4 moles of hydrogen chloride with 3 moles of chlorine gas left. Could it be fair to ask a question like "Hydrogen and chlorine gas react to form hydrogen chloride gas. If 24 grams of hydrogen gas react with 150 grams of chlorine gas, which reactant is the limiting reactant, and how much is left?" I would assume you could answer this question either by using the conservation of mass (150 g Cl2 + 24 g H2 = initial mass = final mass = //x// g HCl +//y// g H2 + //z// g Cl2). The typical problem would ask you to find //x// g HCl, and either //y// or //z// would be 0, but the other could be found by subtracting the mass of HCl from the initial mass. Or you could determine how many moles were used up from the excess reactant and multiply the leftover moles by the excess reactant's molar mass. These two should get the same resulting mass. I guess my question is if there is a use or a need for this kind of question. I don't recall seeing one like this before.

There are applications for this type of problems. Let’s say that I have a solution with an unknown amount of H2 but I know that H2 reacts with Cl2 to form 2HCl. If I add 100 moles of Cl2 and at the end I only have 25 moles, how many moles of H2 I had in the solution initially? 100 - 25 = 75. This is an example of a titration problem. This type of problems are often use for quantification of substance.

->Struggling Topic 2: The text suggests that the percent yield cannot be greater than 100% because the theoretical yield is always larger than the actual yield. Couldn't some error cause the measured yield to be too great and thus make the percent yield greater than 100%?

As you mentioned, there are errors that can cause the calculated yield to be over 100%. However, those are errors, since matter cannot be created. A common error is not drying samples, moist samples weight more than dry sample (because of the presence of water molecules). In this case a yield higher than 100% just shows there were some problems while performing a technique (such as drying).

__Critiques of this chapter__ --A. How clearly the author communicated individual topics ---> 5.2 = the author clearly showed the process for checking balanced equations by totaling the number of each element's atoms on both the product and reactant side of the equation ---> 5.4 = the author shared a shorthand phrase to remind students about oxidation and reduction "LEO the lion goes GER," stands for "loss of electrons is oxidation and gain of electrons is reduction." ---> 5.5 = the author's table of 1 mole of many samples is the same number, 6.02 x 10^23 particles, atoms, molecules, formula units, or other units no matter what the sample was. It was particularly effective to put this example just after having discussed a dozen, a gross, a ream of paper, etc. ---> 5.6 = the author's flow chart for calculations using molar mass and Avogadro's number is a very succinct description of many of the calculations that can be done with the concepts we have reviewed thus far. ---> 5.7 = table 5.5 is an excellent example of how much information is available in a balanced chemical equation. The proportion between the products and reactions is maintained per mole, and the molar mass can be used to identify the relative masses of each substance in the chemical equation. ---> 5.9 = using the peanut butter sandwich analogy for limiting reactants is a good way to get students to realize that limiting reactants are common, every-day, and universal experiences. Hopefully by relating them to something as common as a sandwich, students won't be intimidated by this topic. ---> 5.9 = the flow chart for limiting reactants on page 207 makes it clear that the easiest way to know for sure which reactant is limiting is to simply perform the calculation to see how much product is formed given the amounts of all the reactants that you have. The chart is very clear that the smallest amount of product that is produced is the actual amount produced because it will have used up the reactant that generated this smallest amount of product.

--B. Specifics about the amount of content (did you need more examples?) No ---> 5.1 = the author didn't discuss state changes or solubility changes as physical changes. These two examples are important for students to grasp that substances may change appearance, temperature, and even mix with other substances without undergoing a chemical change. ---> 5.4 = the author gives problems and examples that have the student write half-reactions, but the student should be able to take two half-reactions and write the full reaction too. This "reverse" problem would be a good way to check for student understanding. ---> 5.7 = conversion factors between two different substances only work using moles in the proportion from the balanced equation. This could be easily illustrated by copying figure 5.8 from page 197 with a second compound reflected below the first. You could easily see how using a mole to mole conversion factor from a balanced equation can move you from atoms/ions, grams, or molecules/formula units of one substance in a reaction to any of those three in another substance from the same reaction. In my high school chemistry course the teacher illustrated these as "mole islands" that can be crossed using the necessary conversion factors (mole to mole ratios, molar mass, Avogadro's number, etc).

--C. Chapter's place in the overall text ---> 5.6 = the conversion factors for molar mass and Avogadro's number come after the students have had experience with conversion factors in different units of measurement from chapter 1. This should allow the students to focus on new units with slightly more complicated conversion factors. ---> 5.7 = when we first start converting between different substances in a chemical reaction using a balanced equation it might be good to introduce a very small part of the limiting reactant information that the students will learn later. This might save students from needing to form a new habit of checking to see which reactant will limit the reaction after they have already become quite familiar with these conversions. Maybe if we mention limiting reactants sooner, then students will struggle less with them later. ---> 5.7 and 5.8 = although these two sections are very similar, it is good that the author spent a bit of time in 5.7 talking only about moles and converting between different substances in a chemical reaction before adding mass into the possible problems in 5.8. This should give students some time to get used to the conversion process before making the problems more rigorous. ---> 5.9 = percent yield seemed to be an addition to section 5.9. It seems reasonable to use as an introduction to limiting reactants since the author began by discussing how the actual amount produced in any chemical reaction is never as much as the exact calculation says it should be. This did seem to make the section significantly more lengthy than the previous few sections.