CHEMICAL REACTIONS. Chemical equations are written in the following standard format:

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1 OBJECTIVE(S): Be able to identify if a reaction takes place and record your observations Be able to classify the type of reaction that takes place Write balanced chemical reactions based on physical observations Write balance molecular, complete ionic, and net ionic equations Identify a cation and anion in an unknown Explore the concept of Green Chemistry INTRODUCTION: Chemical Reactions Chemical reactions occur all of the time in our daily lives. Some examples include digesting our food, treating upset stomach (excess acid) with an antacid (a base) such as TUMS, burning fuel in our car to get to school or work, using cleaning products to remove stains from our clothes and to remove hard water deposits on bathroom and kitchen items. Chemists also can make many products using chemical reactions and invent new products by researching different combinations of compounds under different conditions. The purpose of this experiment is to explore different reaction types and identify the cation and anion in an unknown based on your observations of known reactions. In addition, being able to identify the chemical formulas of the products based on observations and solubility rules is important to writing correct molecular equations. Upon writing the molecular equation for each reaction that occurs, the complete ionic, and net ionic equations will also be written. Chemical equations are written in the following standard format: Reactants (physical state) Products (physical state) The physical state of each reactant and product must also be written in parentheses following the chemical formula. Possible physical states include (s) for solid, (l) for liquid, (g) for gas, or (aq) for aqueous (dissolved in water). The number of each type of atom on both sides of the reaction must be equal or balanced using stoichiometric coefficients or numbers placed in front of the chemical formula of the reactants and products. Observations of Chemical Reactions There are several observations that can help us identify or provide evidence that a chemical reaction has taken place. Some observations that a chemical reaction has taken place include the following: Revised:

2 Observation Formation of bubbles or fizzing sound Formation of a solid (a precipitate) or solution becomes cloudy Temperature increases Solution color changes A solid (precipitate) disappears Evidence for formation of a new substance A gas is produced A solid is produced which is insoluble in solution H2O(l) is formed A soluble compound is formed A soluble compound is formed For example, when two aqueous solutions are mixed, the resulting mixture becomes cloudy and yellow. The solid is insoluble, i.e. it does not dissolve in solution. If a substance is soluble, it means that it dissolves in solution and the solution is transparent but not necessarily colorless. The observation that proves that a reaction has taken place was the formation of the yellow precipitate. Remember that a cloudy solution where you cannot see through it is just small solid particles suspended in solution. If you wait long enough, the particles will settle to the bottom. A precipitate is not defined by the size of the solid, only if solid is present or not (opaque solution versus a transparent solution). Some reactions occur very fast and some occur very slowly. If a reaction does occur, then how fast it occurs depends on the concentrations of reactants as well as other factors. In this experiment, it is very important to pay close attention to the concentration of solutions needed as we will be studying how reaction rates change as concentrations change. The letter M is used to represent a concentration unit called molarity. It represents the amount of moles of solute per liter of solution. The number in front of the M determines how concentrated or dilute the solution is. For example, this experiment uses 0.1 M HCl and 1 M HCl. The solution that is less concentrated (has less solute per liter of solution) is the solution with a smaller value for the concentration. Therefore, 0.1 M HCl is more dilute than 1 M HCl as there are less moles of HCl per liter of solution as compared to the 1 M HCl solution. Three Ways to Write a Balanced Chemical Reaction There are three ways to write a balanced chemical reaction, the balanced molecular, complete ionic and net ionic equations. Recall that the balanced molecular equation involves all species written as neutral compounds. For example, when solutions of lead(ii) nitrate and sodium chloride are mixed, a white precipitate is produced. The balanced molecular equation for this reaction is Pb(NO3)2 (aq) + 2 NaCl(aq) à PbCl2 (s) + 2 NaNO3 (aq) We know that the white precipitate is PbCl2 and that the NaNO3 must be dissolved in solution as predicted by the solubility rules. For example, all group 1A metal and nitrate-containing compounds are soluble in water. All species in the molecular equation are written as neutral species (no charges are explicitly written in the equation). The complete ionic equation shows how all the compounds truly exist in solution. If the species is ionic and soluble, then it exists as individual ions in solution. If the compound is any other physical state, it is written as a neutral compound. Be sure that the entire equation still remains

3 balanced. The molecular equation written above would therefore be written as follows for the complete ionic equation Pb 2+ (aq) + 2 NO3 - (aq) + 2 Na + (aq) + 2 Cl - (aq) à PbCl2 (s) + 2 Na + (aq) + 2 NO3 - (aq) The net ionic equation only shows the species that caused the reaction to take place (formation of water, a precipitate, a gas or exchange of electrons). The net ionic equation for the reaction is Pb 2+ (aq) + 2 Cl - (aq) à PbCl2 (s) The observation that a reaction took place was due to the formation of a white solid, PbCl2. Notice that all spectator ions, NO3 - and Na + in this example (those that appear on both the left and right side of the complete ionic equation that remain unchanged and having the same physical state), are eliminated from the reaction. Spectator ions only watch the reaction take place and do not participate in the reaction and hence remain unchanged in their formula and physical state on both sides of the equation. Types or Classification of Reactions The major types of chemical reactions are 1) Combination Several compounds form one compound 2) Decomposition One compound breaks down into several compounds and/or elements 3) Combustion An example is when a compound reacts with oxygen gas, O2, to produce products. If a carbon/hydrogen only containing compound, CxHy, reacts with gaseous O2, it produces carbon dioxide gas, CO2, and water vapor, H2O. If a carbon/hydrogen/oxygen only containing compound, CxHyOz, reacts with gaseous O2, it produces CO2, and gaseous H2O. 4) Single replacement/displacement When an element reacts with a compound and forms a new element and a new compound, a single replacement reaction occurs. The format in the table below will help you determine the products produced. The possible products depend on if the reacting element is a metal or nonmetal. Another way of looking at is to make sure that the ionic compound formed is written correctly, i.e. the cation written first and the anion written second. Type of Reaction Single replacement (A is metal element, B is cation, C is anion) Single replacement (A is nonmetal element, B is cation, C is anion) Format A + BC AC + B A + BC BA + C

4 5) Double replacement/displacement When two aqueous ionic compounds react, their anions are exchanged (or their cations are exchanged). The format in the table below will help you determine the products produced. Be sure that the ionic products produced are written correctly, with the cation written first and the anion written second. Type of Reaction Double displacement (A and C are cations, B and D are anions) Format AB + CD AD + CB If carbonic acid, H2CO3, forms from the double displacement reaction, it will decompose into further products (a secondary reaction). The balanced equation showing the decomposition of H2CO3, carbonic acid, is as follows: H2CO3 (aq) H2O (l) + CO2 (g) There are several types of double replacement reactions: a. Acid/base (or neutralization) An acid (HA(aq) compound) and a base (MOH compound where M is a metal) react to produce a salt or ionic compound and water. b. Precipitation When aqueous, ionic solutions are mixed, a solid product or precipitate forms. 6) Oxidation-reduction (redox) This type of reaction requires an exchange of electrons. To track an exchange of electrons, all of the oxidation states of each element in compounds and/or in ions must be determined (review the rules for assigning oxidation states in your textbook or instructor s notes). In addition, students must be able to determine which element in the equation is being oxidized, which element is being reduced, which reactant is considered the oxidizing agent, and which reactant is the reducing agent. A simple way to identify a redox reaction is when an element on one side of a reaction (as a reactant or product) becomes a part of a compound on the other side or opposite side of the reaction. Identification of an Unknown Chemists must determine the identity of compounds in a solution on a regular basis. For example, working in water quality, the identity and concentration of different compounds must be determined to evaluate whether or not drinking water is safe to ingest. Chemists work with compounds of known identity and study their chemical behavior then test a solution to determine if it is present in a sample. This experiment will allow you to engage in the same process used by chemists. Once you have performed a series of reactions with known compounds and studied

5 their behavior in solution as well as their solubility, you will determine the identity of an unknown cation and anion (i.e. an unknown compound) in solution. This can be done by adding a reactant of known identity to a small amount of your unknown to see if you get observations consistent with a reaction that you already performed. For example, when you add silver nitrate solution, AgNO3, to sodium chloride solution, NaCl, a white precipitate of AgCl forms. If you add silver nitrate to your unknown and a white precipitate forms, then your unknown contains the Cl - ion as long as no other white precipitate formed with the Ag + ion and a different ion based the other reactions you performed. If no white precipitate forms, then no Cl - is present in your unknown. It is important to do as many tests as possible to positively identify the presence of an ion in your unknown. It is also important to recognize that no observation of a chemical change is also important. It just means that you may have soluble ions in solution. For example, if you do all of the tests to identify the cation and you do not identify a chemical reaction, your ion may be a soluble ion such as Na + or NO3 -. Green Chemistry Finally, the Chemistry Department has made a concerted effort to make changes to the experiments in our courses to support Green Chemistry principles and to be better stewards of our environment. According to the Environmental Protection Agency, Green chemistry consists of chemicals and chemical processes designed to reduce or eliminate negative environmental impacts. The use and production of these chemicals may involve reduced waste products, non-toxic components, and improved efficiency. Green chemistry is a highly effective approach to pollution prevention because it applies innovative scientific solutions to real-world environmental situations. This experiment has been modified to introduce the concept of Green Chemistry to our students as well as replace some of the more hazardous chemicals with more benign chemicals. In addition, quantities used are minimal in order to produce less hazardous waste

6 PRE-LAB QUESTIONS: 1. For the following reactions, identify whether the reaction is combination, decomposition, combustion, single replacement, double displacement, acid/base (neutralization), precipitation and/or redox. More than one answer may be possible: a. KOH(aq) + HCl(aq) à H2O(l) + KCl (aq) b. 2 Na(s) + Cl2(g) à 2 NaCl (s) c. Ca(NO3)2(aq) + Na2CO3 (aq) à 2 NaNO3 (aq) + CaCO3 (s) d. HBr (aq) + Cl2 (g) à 2HCl (aq) + Br2 (l) e. 2 C2H6(g) + 7 O2 (g) à 4 CO2 (g) + 6 H2O (g) 2. Write the complete ionic and net ionic equations for all of the reactions listed in question 1 above. 3. For the following redox reaction, answer the questions below: 2 AgNO3 (aq) + Mg(s) à 2 Ag(s) + Mg(NO3)2(aq) a. Identify the oxidation number of each element or ion by writing it the spaces below each species in the reaction above. b. What species was oxidized? c. What species was reduced? d. What substance is the oxidizing agent? e. What substance is the reducing agent?

7 (PRELAB QUESTIONS continued) f. (Optional) Write the complete ionic equation for this reaction. g. (Optional) Write the net ionic equation for this reaction. h. (Optional) Identify any spectator ion(s). 4. This experiment used Pb(NO3)2 and Ba(NO3)2 in the past. These chemicals have been removed from this experiment and replaced with others such as Na2CO3 to reflect Green Chemistry Principles. Explain why these chemicals have been removed from the experiment. You may need to access other resources (books, internet) to help you answer this question. Please cite any resources used. In addition, identify which green chemistry principles might apply to your answer. See the 12 green chemistry principles below (from 1. Prevention It s better to prevent waste than to treat or clean up waste afterwards. 2. Atom Economy Design synthetic methods to maximize the incorporation of all materials used in the process into the final product. 3. Less Hazardous Chemical Syntheses Design synthetic methods to use and generate substances that minimize toxicity to human health and the environment. 4. Designing Safer Chemicals Design chemical products to affect their desired function while minimizing their toxicity. 5. Safer Solvents and Auxiliaries Minimize the use of auxiliary substances wherever possible make them innocuous when used. 6. Design for Energy Efficiency Minimize the energy requirements of chemical processes and conduct synthetic methods at ambient temperature and pressure if possible. 7. Use of Renewable Feedstocks Use renewable raw material or feedstock rather whenever practicable. 8. Reduce Derivatives Minimize or avoid unnecessary derivatization if possible, which requires additional reagents and generate waste. 9. Catalysis Catalytic reagents are superior to stoichiometric reagents. 10. Design for Degradation Design chemical products so they break down into innocuous products that do not persist in the environment. 11. Real-time Analysis for Pollution Prevention Develop analytical methodologies needed to allow for real-time, in-process monitoring and control prior to the formation of hazardous substances. 12. Inherently Safer Chemistry for Accident Prevention Choose substances and the form of a substance used in a chemical process to minimize the potential for chemical accidents, including releases, explosions, and fires

8 Use these diagrams as templates for your spot plate by cutting them out, labeling them and placing them under your spot plates

9 CHEMICALS AND MATERIALS Pieces of Zn metal Pieces of Cu metal Pieces of Mg metal 0.1 M CuCl2 0.1 M FeCl3 0.1 M AgNO3 0.1 M HCl 1 M HCl 0.1 M NaOH 1 M NaOH 0.1 M Na2CO3 0.1 M NaCl Na2CO3 solid steel wool EQUIPMENT TO BE CHECKED OUT Clear spot plates (deep well; 6 wells by 4 wells) SAFETY Goggles must be worn at all times if you or others are working with chemicals in the lab area. HCl is a strong acid and NaOH is a strong base. If any contact on skin occurs, rinse well with water for 15 minutes. WASTE DISPOSAL Solutions and solids must be placed in a proper waste container or as described in your experimental procedure. EXPERIMENTAL PROCEDURE: Part A: Observations of various chemical reactions and their classification 1. General information a. This experiment is designed for students to work individually. If you have questions, please see your instructor. b. If wastes are to be disposed of in the waste container, keep a large beaker at your bench area to collect all of the waste for the duration of the experiment. At the end of the experiment, dispose of the wastes in the proper waste container located in the fume hood. Be sure to check the label first before pouring. Some wastes can be poured down the drain and this will be clearly written in the procedure when to do this

10 c. Before doing each reaction in the experiment, write a possible chemical reaction that might take place. You might want to consult your solubility rules to help identify the soluble compounds from the insoluble compounds in order to write the physical states. After writing the chemical reaction, do the experiment and compare it to the chemical reaction that you wrote to see if the results make sense. 2. Reactions using a spot plate a. In this part of the experiment, you will be using a spot test plate. You may label each well of the spot plate using small pieces of tape from the tape dispenser on the instructor s bench or place a piece of paper underneath and label the paper. DO NOT write on the spot plates. Some of the depressions will be used only for observations of a single solution as some have color, and some of the depressions will be used for mixing and observing two components to see if a reaction takes place. The identity of the solutions to be used in each depression are listed in the table below (table continues on the next page) and quantities needed are described in part b on the next page: Depression location and/or Reaction Letter First Reactant Observations of First Reactant Second Reactant A 0.1 M CuCl2 None (observations only) B Zn metal 0.1 M CuCl2 C Mg metal 0.1 M CuCl2 D Cu metal 0.1 M CuCl2 E 0.1 M Na2CO3 None (observations only) F 0.1 M Na2CO3 0.1 M CuCl2 G 0.1 M HCl None (observations only) H 1 M HCl None (observations only) I Zn metal 0.1 M HCl J Zn metal 1 M HCl K Mg metal 0.1 M HCl L Mg metal 1 M HCl M Cu metal 0.1 M HCl N Cu metal 1 M HCl O 0.1 M AgNO3 None (observations only) Observations of Second Reactant None None None None None P 0.1 M NaCl None (observations None

11 only) Q 0.1 M AgNO3 0.1 M NaCl R 0.1 M AgNO3 Cu(s) S Na2CO3 (s) 0.1 M HCl T Na2CO3 (s) 1 M HCl U 0.1 M NaOH None (observations only) V 0.1 M FeCl3 None (observations only) W 0.1 M NaOH 0.1 M FeCl3 X** 0.1 M NaOH 0.1 M HCl AND 0.1 M FeCl3 Y 0.1 M AgNO3 0.1 M NaOH Z 0.1 M CuCl2 0.1 MNaOH AA 0.1 M FeCl3 Zn(s) BB 0.1 M FeCl3 0.1 M Na2CO3 None None b. The quantities of each reactant are as follows: i. If a reactant is a solution, just place 5 drops into the depression per reactant. You may add up to 10 drops of each solution if it is difficult to make observations. **NOTE: depression X has three components. Add the first two reactants together first, mix and make observations. Then add the third reactant and mix and make observations. ii. If the reactant is a metal, place just one piece of metal in the depression. You may have to use steel wool to remove any tarnish on the metal first. iii. If the reactant is a solid, place a pea sized portion into the depression using the tip of your spatula. c. Recording observations: i. Record any observations of the individual reactants BEFORE mixing in your notebook. ii. Shake the spot plate gently to encourage mixing and/or knocking off any substances that react on the surface of a metal. Record your observations after mixing. NOTE: Some reactions occur quickly and others slowly, so it is best to keep the solutions in your spot plate for at least 10 or more minutes, checking periodically to see if observations change over time. If no evidence of reaction is observed, then write no change observed or NR for no reaction

12 iii. Write the balanced molecular equation for those reactions that take place in your notebook BEFORE moving to the next set of reactants. 3. Reactions using test tubes (see table below). a. 3 ml or 60 drops per solution is needed. Pay attention to other instructions under NOTES column in table below. The resulting solutions of these two reactions can be poured down the drain. First Reactant Second Reactant NOTES 1.0 M NaOH 1.0 M HCl Measure Temperature before and after addition of HCl and record in table; this solution can be poured down the drain 4. Complete Data Analysis for Part A before proceeding to Part B or per professor s instructions. B. Using chemical reactions and solubility to identify ions in an unknown In this part, an unknown containing one compound will be given to students. Students must correctly identify the cation and anion and write the correct chemical formula of the compound in their unknown. The possible cations in the unknown solution are Fe 3+, Na +, Cu 2+, Ag +. The possible anions in the unknown solution are Cl -, OH -, NO3 -, CO3 2-. Recall that sodium and nitrate ions are soluble in solution. Therefore, if you rule out all other ions besides these, then your cation or anion may likely be sodium or nitrate ion. 1. Obtain an unknown from your instructor. Be sure to record your unknown number in your notebook and the initial appearance of your unknown solution. DO NOT WASTE YOUR UNKNOWN! USE ONLY AS MUCH AS THE DIRECTIONS REQUIRE. 2. Think about what reactant or reactants you could add to your unknown that could give you a clear answer as to whether or not a particular ion is present in your unknown. Next, make a table in your lab notebook similar to the one below (the test for Cl - has been added in the first row to help you understand how to use the table): Test for: Cl - Cu 2+ Reactant(s) to be added and its appearance AgNO3; colorless and transparent Zn metal; silvery solid and shiny Initial appearance of unknown solution Colorless and transparent colorless soln Observation after adding solutions together Is ion present and why?

13 3. Next add the appropriate amounts of solution depending on which reaction you are studying (see prior instructions for knowns). For example, if you are testing for Cl -, the test uses 5 drops (and up to 10 drops if necessary) of each solution. Therefore, you would initially add 5 drops of AgNO3 and 5 drops of your unknown. If a white solid formed, then the results would be recorded in the table as follows: Test for: Cl - Cu 2+ Reactant(s) to be added and its appearance AgNO3; colorless and transparent Zn metal; silvery solid and shiny Initial appearance of unknown solution Colorless and transparent colorless solution Observation after adding solutions together White cloudy precipitate formed Zn metal remains unchanged; solution is colorless Is ion present and why? Cl - is present; AgCl(s) formed as it is insoluble in water and is a white solid Cu 2+ is not present; no observations of a chemical reaction occurred. 4. Use fresh solutions for each new test. 5. Complete enough tests to verify the presence or absence of the cation and anion in your unknown. Recall if there is no observable reaction that this is useful information too. 6. Write down the chemical formula of the cation, anion, and compound in your unknown. 7. Place all wastes used in this section (testing for the unknown) into the waste container located in the fume hood. 8. Clean all glassware and spot test plate with soap, tap water, and several rinses with DI water. Pat dry and give back any checked out items to your instructor

14 DATA ANALYSIS: Part A: Observations of various chemical reactions and their classification 1. Results Summary Table a. Placing a summary table at the end of your notebook pages would be useful to identify only the reactions that took place. A table is provided below as an example. It would be best to turn your notebook sideways so you have plenty of room to fit the table and its contents. For example, when doing reaction A in the spot plate, include a column in your data table for each of the following: reaction letter and if in spot plate or test tube balanced molecular equation including physical states of all species observations of the reactants and products underneath species in molecular equation observations after mixing reactants balanced complete ionic equation including physical states of all species balanced net ionic equation including physical states of all species reaction type(s) Results Summary Table Letter/plate or Molecular test tube equation Observations After mixing Reactants Complete Ionic equation Net ionic equation Type(s) of Reaction(s) 2. For every redox reaction that occurred, fill in the following table: Reaction Letter Net ionic reaction Oxidation number of each element in each reactant Oxidation number of each element in each product Species reduced Species oxidized Reducing agent Oxidizing agent

15 B. Using chemical reactions and solubility to identify ions in an unknown 1. Identify the cation in your unknown. Support your answer with the reactions and observations used to arrive at your conclusion. 2. Identify the anion in your unknown. Support your answer with the reactions and observations used to arrive at your conclusion. 3. Write the correct formula of the compound in you unknown. Ask your instructor if he/she will verify your answer. 4. Consult your instructor if any other materials are required and if some type of report is needed. POST LAB QUESTIONS: 1. Some reactions used the same reactants but in different concentrations. a. Does it matter which concentrations were used? Why or why not? b. Let us say that a procedure asked for you to add 10 drops of reactant A into two different test tubes. It next asked for you to add 10 drops of 0.1 M of a second reactant, B, to one test tube and to add 10 drops of 0.5 M of reactant B to the other test tube. You did not observe a reaction in the first test tube but you did observe a reaction in the second test tube. Does this mean that no reaction took place in the first test tube? Explain your answer. c. Is it important when using chemicals in the laboratory to be careful to use not only the correct formula of chemical, but also its concentration? Why or why not? 2. Which metal was the most reactive, Mg, Zn or Cu? Explain your reasoning including use of observations from the experiment. Which metal was the least reactive? Explain your reasoning including use of observations from the experiment. REFERENCES: Written by Jody Williams Tyler (Norco College) and feedback provided by Todd Clements, Terri Beam, and Janet Truttmann (Mt San Antonio College)

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