Procedure Reaction 1: Dissolving the Copper 1. Obtain a clean, dry, glass centrifuge tube. 2. Place a piece of copper wire in a weighing paper, determine the mass of the wire and place it in the centrifuge tube. The copper wire should weigh less than 0.0200 grams. 3. In a fume hood, add seven drops of concentrated nitric acid to the reaction tube so that the copper metal dissolves completely. Describe your observations in the lab report. (Caution, Concentrated nitric acid and nitrogen dioxide are very corrosive. Either will turn your skin yellow on contact. Do not leave any spills on the lab bench or in the fume hood.) 4. When the copper has dissolved, add seven drops of distilled water to the tube. Reaction 2: Preparation of Copper(II) Hydroxide 1. Add 15 drops of 3.0 M aqueous sodium hydroxide to the tube. Make sure that the reactants are well mixed. Shake the tube carefully or gently flick the bottom of the tube with your finger. Remember that the contents of the tube may still be corrosive. 2. Add a second 15 drops of NaOH(aq), mix well, and record your observations. If you have two layers at this point it means that you have not mixed the solution well enough. 3. Centrifuge the reaction mixture. 4. The liquid at the top of the centrifuged mixture is called the supernatant while the solid is called a precipitate. Before separating the supernatant from the precipitate it is necessary to ensure that all of the copper(II) hydroxide has been precipitated. The supernatant should be clear and colorless indicating the absence of any Cu+2 ions in the solution. It should also be basic due to an excess of OH- ions. Using a clean glass stirring rod, transfer a drop of the supernatant onto a piece of red litmus paper. If the litmus paper turns blue then the solution is basic and enough NaOH has been added. If the paper does not turn blue, add more NaOH, mix well, recentrifuge, and repeat the litmus paper test until the paper does turn blue. 5. An efficient separation of supernatants and precipitates is key to obtaining a good final yield of copper. The supernatant liquid can be separated from the precipitate by expelling the air from the bulb of a Pasteur pipet, inserting the tip of the pipet into the supernatant, then gently sucking the supernatant into the pipet. If you expel air or liquid into the precipitate with the pipet, you will stir up the precipitate and will have to repeat the centrifugation step. Remove as much liquid as possible and discard it in the waste container provided on the instructor’s cart. It is better to leave a small amount of supernatant liquid than to remove some of the copper(II) hydroxide precipitate. Reaction 3: Formation of Copper(II) Oxide 1. Set up a hot water bath by placing a beaker of water on a hotplate, placing an iron ring around the beaker, and heating the water to boiling. 2. Place the centrifuge tube containing the copper(II) hydroxide into the boiling water. Carefully hold the tube with a test tube clamp so that it doesn’t get water into it. Record your observations. Reaction 4: Formation of Copper(II) Sulfate 1. Add 20 drops of 3.0 M H2SO4 to the solid in the centrifuge tube. Stir carefully to ensure that the copper(II) oxide dissolves completely. Complete dissolution of the mixture will require thorough mixing and possibly heating of the solution. 2. Obtain the mass of a small, clean, glass test tube as accurately as possible. 3. Transfer the liquid from the centrifuge tube into the test tube. Rinse the centrifuge carefully with 1.0 mL of distilled water and transfer the rinse into the test tube containing your sample. 4. Record your observations on the data sheet. Reaction 5: Formation of Copper Metal 1. Add a small quantity of zinc powder to the sample solution. Continue adding zinc in small quantities until the solution loses the blue copper(II) color. Any excess zinc added will need to be removed so don’t add it too quickly or in large quantities. When the solution has turned colorless, add several drops of 3.0 M H2SO4 to the tube to dissolve any left over zinc. You can tell that the zinc has dissolved when addition of sulfuric acid does not generate bubbles. 2. Allow the copper metal to sink to the bottom of the tube and carefully remove the supernatant liquid using a Pasteur pipet. 3. Wash the red-brown copper metal in the tube with 1.0 mL of water. Allow the copper metal to settle to the bottom and remove the excess water. Repeat this rinsing process two more times. 4. Describe your observations on the data sheet. Drying the copper metal 1. After removing as much of the third rinse water as possible you are ready to dry the metal. This must be done carefully in a cool Bunsen Burner flame. If the tube is heated too quickly there is a risk of ejecting the contents of the tube as the water boils. Also, if the flame is too hot you may convert the copper metal back into black copper(II) oxide. The objective is to drive the water from the tube as steam. Make sure that as water condenses on the walls of the tube that you continue to heat until all of the water if gone from the tube. 2. Once all of the water is removed from the tube, cool the tube and its contents then determine the mass of copper by weighing the tube and subtracting the tube + copper weight from the weight of the empty tube (Reaction 4 step 2). If the mass of copper is higher than the original mass of the copper wire it either contains water or zinc or has been converted to copper(II) oxide. Excess water can be removed by reheating the tube and reweighing to constant mass. Excess zinc requires addition of sulfuric acid followed by re-rinsing with water and re-drying. Chemistry 1215 Experiment 9 Lab Report Name ______________________________ Data Sheet Mass of copper wire _______________ Mass of clean, dry test tube _______________ Mass of test tube plus copper _______________ Mass of final copper sample _______________ Percent recovery of copper. Show all calculations. Observations 1. Describe your observations for Reaction 1 including colors, gases formed, etc. 2. Describe your observations for Reaction 2 including colors, gases formed, etc. 3. Describe your observations for Reaction 3 including colors, gases formed, etc. Estimate the temperature of the decomposition of Copper(II) hydroxide. 4. Describe your observations for Reaction 4 including colors, gases formed, etc. 5. Describe your observations for Reaction 5 including colors, gases formed, etc. Write a brief discussion of your results including a statement of the final percent recovery of copper and a discussion of reasons why the recovery differs from 100%. Chemistry 1215, Experiment #9; Copper and its compounds, Pre-lab Name ____________________________________ 1. Write a balanced chemical equation including phase labels for the reaction between aqueous copper (II) nitrate and aqueous sodium hydroxide. 2. Nitrogen monoxide (NO) and nitrogen dioxide (NO2) are toxic, corrosive gases that significantly lower blood pressure when inhaled. How are these gases produced in today’s experiment? What should you do to protect yourself against their toxicity? 3. Iron reacts with oxygen from the atmosphere to produce iron (III) oxide, also known as rust (Fe2O3). What chemical species is oxidized in this reaction? What is the reducing agent? Jaffrey Zagnut couldn’t find any nitric acid so he tried to dissolve his copper sample in hydrochloric acid instead. Unfortunately his copper wouldn’t dissolve in HCl. Why will copper dissolve in nitric acid but not in hydrochloric acid (after all, HCl is a stronger acid than HNO3). Chemistry 1215, Experiment #9; Copper and its compounds, Post-lab Name ____________________________________ 1. Copper (II) hydroxide is converted into copper (II) oxide by heating the test tube containing Cu(OH)2 in a hot water bath. Is it necessary to use distilled water in this water bath? Why or why not? 2. Copper metal doesn’t “rust” in the presence of oxygen at room temperature. However, it will react with O2 at elevated temperatures. Write a balanced chemical equation describing the formation of copper (II) oxide when copper metal is heated in air. 3. When zinc is dissolved in sulfuric acid a gas is produced. What is the chemical identity of this gas? How is it produced? 4. Jaffrey Zagnut started with a 0.032 g sample of copper which he took through the series of reactions described in this experiment. At the end of the experiment he obtained 0.038 g of a black product. What was his percent yield? What is the most likely source of the error in his experiment? (Hint: consider question 2 above)

Mothballs are composed primarily of the hydrocarbon naphthalene (C10H8). When 1.025 g of naphthalene is burned in a bomb calorimeter, the temperature rises from 24.25 ?C to 32.33 ?C.

Find ?Erxn for the combustion of naphthalene. The heat capacity of the calorimeter, determined in a separate experiment, is 5.11kJ/?C.

Express the change in energy in kilojoules per mole to three significant figures.

Pure copper may be produced by the reaction of copper(I) sulfide with oxygen gas as follows:

Cu₂S(s) + O₂(g) → 2Cu(s) + SO₂(g)

If 0.490 kg of copper(I) sulfide reacts with excess oxygen, what mass of copper metal may be produced?

A 35.00-mL solution of 0.2500 M HF is titrated with a standardized 0.1120 M solution of NaOH at 25 C

1) the pH at 0.50 mL before the equivalence point
2) the pH at the equivalence point
3) the pH at 0.50 ml after the equivalence point.

Assume that all 4.0 mL of methyl salicylate was placed in the beaker and reacted, that most of the NaOH solution is water, and that the total water used is 56 mL. The density of methyl salicylate is 1.18 g/mL and the density of water is 1.00 g/mL. Pay close attention to units and dimensional analysis when completing this problem.

Predict under what temperatures (all temperatures, low temperatures, or high temperatures), if any, the reactions will be spontaneous or nonspontaneous.

Part A) N2(g)+O2(g)→2NO(g) ΔH∘rxn=+182.6kJ

Part B) 2N2(g)+O2(g)→2N2O(g) ΔH∘rxn=+163.2kJ

Part C) 4NH3(g)+5O2(g)→4NO(g)+6H2O(g) ΔH∘rxn=−906kJ

A 5.65 −g sample of a weak acid with Ka=1.3×10−4 was combined with 5.20 mLof 6.10 M NaOH and the resulting solution was diluted to 750 mL. The measured pH of the solution was 4.20.

What is the molar mass of the weak acid?

In a waste-water treatment plant a 1.00x10³ L tank with 3.00x10^-2 M Na₂CO₃(aq) is mixed with a 2.00x10³ L tank with of 6.00x10^-3 M Na₂CO₃(aq). Assume sodium carbonate behaves as a strong electrolyte.

a) How many moles of sodium ions are in the product?

b) What is the concentration (in M) of sodium ions in the product?

The acid dissociation constant of the 1,10-phenanthrolium ion (phen) was meeasured at different temperatures using spectroscopic methods.

The values of pKa measured at 298K and 308K were 4.812 and 4.731, respectively.

Part A: Calculate  ΔrGo at 298K. Express your answer in kJ/mol to three significant figures.

Part B: Calculate  ΔrGo at 308K. Express your answer in kJ/mol to three significant figures.​

Part C: Calculate  ΔrHo. Express your answer in kJ/mol to three significant figures.​

Part D: Calculate  ΔrSo. Express your answer in J/(K.mol) to three significant figures.​

Part E: Calculate the pKa of the acid at 359K. Express your answer to three significant figures​

A student finds that a mixture of acetone and water at 740 mmHg boils at 68 Celcius. Pure acetone has a vapor pressure of 1200 mmHg and water has a vapor pressure of 237 mmHg at the given temperature.

a.) What are the mole fractions of each component in the mixture?

b.) If there are 0.500 moles in the solution, what is the composition by mass, of the mixture?

give us your version of Emil Fischer's proof of the structure of glucose. You do not need to flesh out Fischer's proof to reveal the structures of all the D-aldopentoses and D-aldohexoses, but you do have to show us structures for the four sugars (three D, one L) that come directly from the Fischer proof, the determination of the structure of D-Ribose, and mechanisms for all reactions you (and Emil Fischer) use(d).

For the reaction SO₂(g) + NO₂(g) == NO(g) + SO₃(g), Kc = 85.0 at 460

Questions

A.      Use your results to determine if the forward reaction in the potassium chromate/HCl reaction endothermic or exothermic. Explain your answer, using Table 1 to help construct your thoughts.

B.      Write the equation for the equilibrium constant (K) of the reaction studied in this exercise.

2K2Cr4 + 2HCl ------> Cr2o7 + H2O + 2KCl

Use the information below to answer Questions C, D, and E:

The equilibrium constant (K) of the reaction below is K = 6.0 x 10^-2, with initial concentrations as follows: [H₂] = 1.0 x 10^-2 M, [N₂] = 4.0 M, and [NH₃] = 1.0 x 10^-4M.

N2 + 3H2 ------> 2NH3

If the concentration of the reactant H₂ was increased from 1.0 x 10^-2 M to 2.5 x 10^-1M, calculate the reaction quotient (Q) and determine which way the equilibrium position would shift.

If the concentration of the reactant H₂ was decreased from 1.0 x 10^-2 M to 2.7 x 10^-4M, calculate the reaction quotient (Q) and determine which way the equilibrium position would shift.

If the concentration of the product NH₃ was decreased from 1.0 x 10^-4 M to 5.6 x 10^-3M, calculate the reaction quotient (Q) and determine which way the equilibrium position would shift.

Exercise 2.58. Part A: Determine the number of Kr atoms in a 5.55 −mg sample of krypton.

Part B:Determine the molar mass, M, of an element if the mass of a 2.80×1022−atom sample of the element is 2.09 g.

Part C: Determine the identity of an element if the mass of a 2.80×1022−atom sample of the element is 2.09 g. (Answer choices: titanium. potassium, scandium, calcium.)

Exercise 2.56. Without doing detailed calculations, indicate which of the following quantities contains the greatest number of atoms: 6.022×1023Ni atoms, 25.0 g nitrogen, 52.0 g Cr, 10.0cm3 Fe (d=7.86g/cm3). (answer choices: Ni, Nitrogen, Cr, Fe)

Excercise 2.62. A particular lead–cadmium alloy is 8.0% cadmium by mass. What mass of this alloy, in grams, must you weigh out to obtain a sample containing 6.30×10²³ Cdatoms? Express your answer using two significant figures.

Exercise 3.6 Determine the mass, in grams, of,

Part A 7.32 mol N2O4

Part B 3.22×10²⁴ O2 molecules

Part C 18.8 mol CuSO4⋅5H2O

Part D 4.14×10²⁴ molecules of C2H4(OH)2

Excercise 3.11

Part A moles of N2O4 in a 145 −g sample

Part B N atoms in 43.5 g of Mg(NO3)2

Part C N atoms in a sample of C7H5(NO2)3 that has the same number of O atoms as 12.4 g C6H12O6

thank you in advance ^^

Calculate the free-energy change for ammonia synthesis at 25 ? C(298 K) given the following sets of partial pressures:

a.1.0 atm N2, 3.0 atm H2, 0.020 atm NH3

b. 0.010 atm N2, 0.030 atm H2, 2.0 atm NH3