Silver ion, Ag+, can be recovered from photographic solutions by adsorbing onto an ion exchange resin (Langmuir adsorption)

Ag+ +Ex↔ Ag-Ex logKL =12.30

(where Ex represents an adsorbing site on the resin). According to the manufacturer, the resin has 1.20 x10-5 moles of adsorbing sites per gram of resin.

a) If you have a liter of solution with [Ag+] = 0.030 mM, how much resin will you need to adsorb all the silver assuming 100% binding?

b) Calculate the predicted equilibrium dissolved [Ag+] if you add 10.00 g of resin to 1.00 liter of solution described in part a.

The ideal gas law PV=nRT relates pressure P, volume V, temperature T, and number of moles of a gas, n. The gas constant R equals 0.08206 L⋅atm/(K⋅mol) or 8.3145 J/(K⋅mol). The equation can be rearranged as follows to solve for n: n=PVRT This equation is useful when dealing with gaseous reactions because stoichiometric calculations involve mole ratios.

Part A

When heated, calcium carbonate decomposes to yield calcium oxide and carbon dioxide gas via the reaction CaCO3(s)→CaO(s)+CO2(g) What is the mass of calcium carbonate needed to produce 53.0 L of carbon dioxide at STP? Express your answer with the appropriate units.

Part B

Butane, C4H10, is a component of natural gas that is used as fuel for cigarette lighters. The balanced equation of the complete combustion of butane is 2C4H10(g)+13O2(g)→8CO2(g)+10H2O(l) At 1.00 atm and 23 ∘C, what is the volume of carbon dioxide formed by the combustion of 3.80 g of butane? Express your answer with the appropriate units.

Calcium hydride, CaH2, reacts with water to form hydrogen gas: CaH2(s)+2H2O(l)→Ca(OH)2(aq)+2H2(g) This reaction is sometimes used to inflate life rafts, weather balloons, and the like, where a simple, compact means of generating H2 is desired.

Part A

How many grams of CaH2 are needed to generate 146 L of H2 gas if the pressure of H2 is 824 torr at 22 ∘C?

The metabolic oxidation of glucose, C6H12O6, in our bodies produces CO2, which is expelled from our lungs as a gas: C6H12O6(aq)+6O2(g)→6CO2(g)+6H2O(l)

Part A

Calculate the volume of dry CO2 produced at body temperature (37 ∘C) and 0.990 atm when 23.5 g of glucose is consumed in this reaction

Part B

Calculate the volume of oxygen you would need, at 1.00 atm and 298 K, to completely oxidize 53 g of glucose.

An organic compound contains C, H, N, and O. Combustion of 0.3069 g of the compound produces 0.8301 g CO2 and 0.2975 g H2O. A sample of 0.5322 g of the compound was analyzed for nitrogen. At STP, 30.54 mL of dry N2 (g) was obtained. In a third experiment, the density of the compound as a gas was found to be 3.68 g/L at 143 deg C and 244 torr. Calculate the empirical formula and the molecular formula of the compound.

calcium in powdered milk is determined by ashing a 1.50g sample and then titrating the calcium with EDTA soluton, 12.1mL being required. The EDTA was standardized by titrating 10.0mL of a zinc solution prepared by dissolving 0.632g zinc metal in acid and diluting to 1L (10.8mL EDTA required for titration). What is the concentration of calcium in the powdered milk in parts per million?

An organic compound contains C, H, N, and O. Combustion of 0.3069 g of the compound produces 0.8301 g CO2 and 0.2975 g H2O. A sample of 0.5322 g of the compound was analyzed for nitrogen. At STP, 30.54 mL of dry N2 (g) was obtained. In a third experiment, the density of the compound as a gas was found to be 3.68 g/L at 143 deg C and 244 torr. Calculate the empirical formula and the molecular formula of the compound.

What is the pH of 0.213 M NaF(aq)?

HF: Ka = 7.2 * 10-4

Answer is 8.24

Rank the polarity of the following compounds: ferrocenecarboxylic acid, acetylferrocene, ferrocenecarboxaldehyde, ferrocene, (dimethylaminomethyl) ferrocene, and ferrocenemethanol and give a detailed explanation why.

A 50/50 blend of engine coolant and water (by volume) is usually used in an automobile\'s engine cooling system. If your car\'s cooling system holds 5.60 gallons, what is the boiling point of the solution? Make the following assumptions in your calculation: at normal filling conditions, the densities of engine coolant and water are 1.11 g/mL and 0.998 g/mL respectively. Assume that the engine coolant is pure ethylene glycol (HOCH2CH2OH), which is non-ionizing and non-volatile, and that the pressure remains constant at 1.00 atm. Also, you\'ll need to look up the boiling-point elevation constant for water.

Engineers in Brazil do not have to worry as much about the CLOUD POINT or the POUR POINT of their biodiesel as Canadian engineers do. Define these terms and suggest two ways problems relating to them can be addressed.

Experiment 1: Determining the Chemical Formula for Copper Gluconate The experiments in this lab use a compound named copper gluconate. This compound can be formed when gluconic acid (C6H12O7) and copper solutions react with copper (II) carbonates or copper hydroxide. Copper gluconate has a variety of uses and applications. For example, copper gluconate is used as a primary ingredient in the breath mint Certs®. It is also used as a source of copper in nutritional supplements. Your task will be to determine the chemical formula of the compound by isolating the copper and determining the molar ratio of copper and gluconate in the compound. Figure 3: Certs® Figure 3: Certs® Materials: 1 g Copper gluconate 10 mL 0.5% Salt, NaCl 10 mL Graduated cylinder Scale (1) 50 mL Beaker 250 mL Glass beaker 2 Aluminum washers Sterno® Ring stand Ring for ring stand Stir rod Matches Cupcake wrapper *Oven *Hot pad or towel *Baking pan *20 mL Distilled Water *Kitchen tongs *You must provide Procedure Put the 250 mL beaker on the scale and tare the scale. Measure 1.0 g of the copper gluconate in the beaker on the scale. Record the exact mass in the Data section below. Use the graduated cylinder to measure and pour 10 mL 0.5% NaCl into the beaker with the copper gluconate. You may need to gently swirl the solution if all the copper gluconate does not immediately suspend into the solution. Add two aluminum washers to the solution. Fasten the small ring approximately 6 - 10 inches up on the ring stand and place the beaker on the ring Place the Sterno® directly beneath the beaker. Remove the inner cap on the Sterno® and ignite the inner contents with the matches. Heat the beaker until solution clears. Your solution may not turn completely clear, but some color change should be evident. Alternatively you can also determine when the reaction is complete by looking for the formation of gas bubbles on the surface of the washers. When the formed gas bubbles are gone, then the reaction is complete. Note: Carefully monitor the set-up while the Sterno® is in use. You may need to adjust the height of the ring/beaker to ensure that the beaker is heated enough; and, to avoid exposing the beaker from high temperatures. **Carefully observe the set-up you choose!! Do not leave the beaker unattended while exposed to the Sterno®. Plastic beakers should never be used with heat.** Carefully remove the beaker from heat, and use forceps to replace the lid on the Sterno®. Decant (pour) the clear liquid into a 50 mL beaker. When all that remains in the original beaker are the copper plated washers, rinse the washers with distilled water and decant the remaining liquid, being careful not to lose any copper, into a container. This water can be disposed of down a sink drain. Repeat this process three times. Remove the first washer and use the stir stick to scrape the copper into the metal cupcake wrapper. Rinse the washer with distilled water to be sure all copper is recovered into the wrapper. Repeat the process for the second washer, scraping the copper into the same wrapper. Place the wrapper on a baking pan and put it in the oven at 115 °C (239 °C) to dry the product. Monitor the wrapper and contents and use a hot pad or towel to carefully remove them from the oven after 45 minutes, or after all of the water has evaporated. After the wrapper has cooled to room temperature, empty the dried copper from the wrapper onto the scale and weigh the final mass.

Why is it important in this experiment to be accurate in all your measurements?

List the measurements you will take in this experiment.

What chemical wastes are produced in this reaction?

Record the mass in the Data section below.

Data: Mass of copper gluconate (initial, see Step 2): _____________________

Mass of copper (final; see Step 13): _____________________

Calculations Mass of copper recovered:

Moles of copper :

Mass of gluconate:

Moles of gluconate:

Chemical formula:

1. What is the chemical formula of copper gluconate?

2. List two sources of error in the experiment and explain the impact they had on the results.

3. Create a pie chart showing the percent composition for each element in the compound copper gluconate, clearly label each element and the percentage.

4. Copper chloride can be used as a source of copper for this experiment, but copper gluconate is preferred due the fact that it is a “green” compound. Discuss the environmental

Lab: Spectroscopy and Qualitative Analysis in the Determination of an Unknown Iron Salt

Procedure

Part A:

Obtain one of the provided kits and inspect for completeness (see available list in the laboratory). Ask your instructor for replacement items if necessary. Select one of the available unknown iron salts. All of the iron salts (which may or may not be hydrated) are in the +2 oxidation state and will have one of the following anions: Cl^-, Br^-, SO4^2- and CH3COO^- (acetate). Finally, obtain ~ 30 mL of the standard iron(II) solution (record the concentration).

2. Using the 10 mL beaker, weigh out ~ 0.0400 to 0.0500 g (±0.0001 g) of your salt. Record the mass dispensed. Transfer the salt to a 100.00 mL volumetric flask using a small amount of water. Rinse the beaker 3 to 4 times; transferring the washings to the flask each time. Allow the iron salt to completely dissolve and then dilute to the mark.

3. Using the pipette from your kit, transfer 1.00 mL of the solution from step 2 into a 100.00 mL volumetric flask. To the flask add 1 mL each of: 1% hydroxylamine hydrochloride solution, 1 M sodium acetate and 1% 1,10-phenanthroline solution from the provided burettes. Dilute to the mark and mix thoroughly. This is the unknown solution you will measure the absorbance of.

4. Rinse and fill a burette with the standard iron solution obtained in step 1. Prepare the standard solutions as follows: accurately dispense – recording initial and final burette readings (±0.02 mL) – approximately 1.00, 2.00, 3.00 and 4.00 ml of the standard Fe^2+ solution into separate volumetric flasks. As in step 3, add 1 mL each of: 1% hydroxylamine hydrochloride solution, 1 M sodium acetate and 1% 1,10-phenanthroline solution from the provided burettes. Dilute to the mark and mix thorougly.

5. Using the final 100.00 mL volumetric flask, prepare the blank solution by adding 1 mL each of: 1% hydroxylamine hydrochloride solution, 1 M sodium acetate and 1% 1,10-phenanthroline solution from the provided burettes. Dilute to the mark and mixthoroughly.

6. To analyze your samples ask the lab demonstrators for instruction on the proper use of the spectrophotometer. Obtain the absorbance of each of the standards and the unknown at 508 nm.

Part B:

Qualitative Analysis

Transfer some of your salt (tip of a scoopula) into a beaker and dissolve it in ~ 30 mL distilled water. If you are finding it difficult to dissolve your salt, add two to three drops of 6M HNO3 and mix.Transfer equal portions (about 2-3 cm in height) of this solution into two test tubes:

1. Chloride and Bromide. Add 3 drops of 6M HNO3 to the solution in the first test tube and shake well. Add 5 drops of 0.1M AgNO3 and mix well. If a precipitate has formed upon the addition of silver nitrate a positive test for chloride or bromide has occurred. If no precipitate formed you do not have either the chloride or bromide salt and can move on to the sulfate test. To differentiate between chloride and bromide, add 10 drops of 6M NH4 OH to the mixture (do not mix). If two distinct layers form immediately, then the chloride anion is present.

2. Sulfate. Add 3 drops of 6M HNO3 to the solution in the second test tube and shake well. Add 5 drops of 0.1M BaCl2. The formation of a white precipitate indicates the presence of SO4^2-.

3. Acetate. Place a small amount of the solid salt (~ the size of a pea) into a test tube. Add 20 drops of concentrated H2SO4 (CARE!!!) and 20 drops of 1-pentanol from the dropper bottles in the fume hood to the test tube and mix with a glass stirring rod. Place the test tube in the water bath provided (making a note of which one is yours) for about 5 minutes. Remove the test tube from the bath and check the odour of the solution by bringing your nose slowly to the test tube and waving your hand over the top of the test tube towards your nose. You can also pour it onto a watch glass to make it easier to pick up the odour. A fruity odour (it often reminds people of fake banana smell) indicates the presence of the acetate ion.

_______My Data Below

Concentration of standard iron(II) solution = 50.0 ppm

Weight of unknown salt: 0.0413 g (+- 0.0001 g)

The positive ion test determined that the unknown salt was sulphate (SO4^2-).

Question: Using the mass of the sample analyzed, the dilutions you performed and the positive anion test determine: the identity of your unknown salt, its molecular weight and the water of hydration.

Calculate the appropriate amount of 28 wt% aqueous NH3 and solid NH4Cl to be mixed together to yield 500 ml of a ~0.2 M pH 10 buffer. Check with pH paper. Store in a flask and adjust pH as needed.

Consider the reaction A+2B⇌C whose rate at 25 ∘C was measured using three different sets of initial concentrations as listed in the following table: Trial [A] (M) [B] (M) Rate (M/s) 1 0.20 0.010 4.8×10−4 2 0.20 0.020 9.6×10−4 3 0.40 0.010 1.9×10−3 What is the rate law for this reaction? Express the rate law symbolically in terms of k, [A], and [B].

Calculate the initial rate for the formation of C at 25 ∘C, if [A]=0.50M and [B]=0.075M.