Which of the 2 reactants is the limiting reagent C. Calculate the mass of urea formed D. How much excess reagent is left at the end of the reaction E. If the actual yield of urea formed was 980g what is its yield % a. Ans. 2NH3 + CO2 ( (NH2)2CO + H2O (17) (44) (60) 637.2 1142 980 b. HN3 c. 1124g d. 318g of CO2 e. 87.2% 1b .Iron is reduced from a reaction between aluminium and iron (III) oxide at a temperature of 3000°C. In this reaction 124g of Al and 601g of iron (III) oxide are reacted. A. Write a balance equation for the above reaction B.
14.1 Multiple-Choice and Bimodal Questions 1) Consider the following reaction: [pic] The average rate of appearance of B is given by[pic]. Comparing the rate of appearance of B and the rate of disappearance of A, we get [pic] A) -2/3 B) +2/3 C) -3/2 D) +1 E) +3/2 Answer: B Diff: 1 Page Ref: Sec. 14.2 2) Nitrogen dioxide decomposes to nitric oxide and oxygen via the reaction: [pic] In a particular experiment at 300 °C, [pic]drops from 0.0100 to 0.00650 M in [pic] The rate of appearance of [pic] for this period is __________ M/s. A) [pic] B) [pic] C) [pic] D) [pic] E) [pic] Answer: A Diff: 1 Page Ref: Sec. 14.2 3) Which substance in the reaction below either appears or disappears the fastest?
1. A solution containing 1.000M acetic acid (CH3COOH) and 1.000M ethyl alcohol (CH3CH2OH) at 150oC produced 0.171 mole/liter of the product ethyl acetate (CH3COOCH2CH3) when equilibrium was established. Determine Kc for the reaction. CH3COOH (aq) + CH3CH2OH (aq) CH3COOCH2CH3 (aq) + H2O (l) 2. Using the equilibrium constant from above, calculate the equilibrium concentrations of all the compounds in the reaction if 1.000M acetic acid is reacted with 2.000M ethyl alcohol.
Who is right, John or Anna? Explain your answer. 3 When 200 g of calcium nitrate, Ca(NO3)2.2H2O is heated at 120ºC the mass decreases by 36 g. (a) Why does the mass decrease? (b) How much calcium nitrate is left after heating? 4 Calculate the molar mass of these compounds: (relative atomic masses: H = 1; N = 14; O = 16; S = 32; Cu = 64; Br = 80; Pb = 207) (a) copper nitrate, Ca(NO3)2 (b) lead bromate, Pb(BrO3)2 (c) ammonium sulfate, (NH4)2SO4 5 The equation for the complete combustion of methane is shown below.
Calculations involving Solutions (1) Answer the following questions on file paper. Remember to set out your working neatly and clearly, showing all necessary steps. Give answers to the appropriate number of significant figures. Use Ar values when required as given on your copy of the Periodic Table. 1) Calculate the number of moles in each of the following solutions:- a) 2.0 dm3 of 0.050 mol.dm-3 hydrochloric acid b) 50 dm3 of 5 mol.dm-3 sulfuric acid c) 10 cm3 of 0.25 mol.dm-3 potassium hydroxide [3] 2) Calculate the volume in cm3 of each solution that contains the following numbers of moles:- a) 0.00500 moles of NaOH with concentration 0.100 mol.dm-3 b) 1.00×10-5 moles of HCl with concentration 0.0100 mol.dm-3 c) 9.25×10-3 moles of KCl with concentration 0.250 mol.dm-3 [3] 3) Find the concentrations of the following solutions in both mol.dm-3 and g.dm-3:- a) 0.400 moles of HCl in 2.00 dm3 of solution b) 12.5 moles of H2SO4 in 5.00 dm3 of solution c) 1.05g of NaOH in 500 cm3 of solution [6] 4) In a titration, 25.0 cm3 of a solution of sodium hydroxide required 18.80 cm3 of hydrochloric acid of concentration 0.0500 mol.dm-3 for neutralisation.
Siddharth Rajendran Chemistry HL Urea Dissolution Lab Raw Data:- (Expected Values) Change in Enthalpy: 14 kJ mol-1 Change in Gibbs free energy: 6.86 kJ mol-1 Change in Entropy: 69.5 J mol-1 Molar Mass of Urea: 60.06 g mol-1 Heat Capacity: 4.184 J g-1º Data Table 1: To calculate the Enthalpy change Mass of Urea Tablet (g) (±0.01g) | Volume of Water(mL) (±0.05mL) | Initial Temperature (Cº) (±0.2 Cº) | Final Temperature(Cº) (±0.2 Cº) | 3.04 | 50.0 | 21.3 | 17.4 | Initial Observations:- * There was a decrease in temperature at a fast rate. * The temperature of the solution was slowing down continuously but the rate started decreasing. * The Urea dissolved and the rate was decreasing continuously. * The temperature gradually started to increase after almost the Urea present had dissolved. Data Table 2: Mass, Volume and Temperature during Dissolution of Urea (To calculate Keq) Mass of Urea(g) (±0.01g) | Initial Temperature(Cº) (±0.2 Cº) | Final TemperatureCº) (±0.2 Cº) | Initial Volume(mL)(±0.05 mL) | Final Volume(mL)(±0.05 mL) | 3.76 | 21.4 | 22.9 | 5.02 | 7.14 | Processing Raw Data * Determining the Final temperature of dissolution of Urea in the Styrofoam cup.
of mixture Metal C 25.605g 24.6mL 25.2°C 100.5°C 28.7°C Calculations: Show your work and write a short explanation with each calculation. Part I: 1. Calculate the energy change (q) of the surroundings (water). We can assume that the specific heat capacity of water is 4.18 J/ (g · °C), and the density of water is 1.00 g/mL. (4 points) q = m × c × Δt Given: q=?
Label each axis, reactants, products, heat of reaction (ΔH), and energy of activation (Ea). Define heat of reaction and energy of activation. [pic][pic] 6. Given the following; [H2] = 0.00537, [HCl] = 0.0354, ΔH = -56.3 kJ, ΔG = +26.6 kJ; Calculate Keq for the following reaction: Al(s) + HCl(g) [pic] H2(g) + AlCl3(s) Balanced equation: 2Al(s) + 6HCl(g) [pic] 3H2(g) + 2AlCl3(s) [pic] a) Are the reactants or products favored at equilibrium? Explain.
The rate law for this reaction is rate = k[acetone]m[H+]n[I2]p where k is the rate constant for the reaction and m, n, and p are the orders of the reaction with respect to acetone, hydrogen ions (acid), and iodine, respectively. Although orders of reaction can be any value, for this lab we will be looking only for integer values for the orders of reaction (0, 1, 2 are acceptable but not 0.5, 1.3, etc.) The rate of the reaction can also be expressed as the change in the concentration of a reactant divided by the time interval: rate = - Δ[ I 2 ] Δt The iodination of acetone is easily investigated because iodine (I2) has a deep yellow/brown color. As the acetone is iodinated and the iodine converted to the iodide anion, this color will disappear, allowing the rate of the reaction to be easily monitored. We can study the rate of this reaction by simply making I2 the limiting reactant in a large excess of acetone and H+ ion.
Questions 1. Calculate the temperature change DeltaT, for each reaction by subtraction the initial temperature, T1, from the final temperature, T2 (DeltaT=T2-T1) (Reaction 1): T2=3.129 T1=23.8 (3.129-23.8=20.6) DeltaT=20.6 (Reaction 2): T2=41.0 T1=24.5 (41.0-24.5-16.5) DeltaT=16.5 2. Tell Which reaction is exothermic. Explain. The first reaction is exothermic because the final temperature is lower than the final temperature from the second reaction showing that it let out more energy to its surroundings.