A quantity of 0.225 g of a metal A molar mass = 27.0gmolliberated 0.303 L of molecular hydrogen (measured at 17 °C and 741. mmHg) from an excess ofhydrochloric acid. Deduce from these data the corresponding equation and write the formulas for an oxide and a sulfate of A.Note: Reference the Fundamental constants and Conversion factors for non-SI units tables for additional information.Part: 0/3Part 1 of 3Using the smallest whole-number coefficients, write the balanced equation for the reaction of the metal, A, with hydrochloric acid. Include the phaseabbreviations.☑→A) + 2HCl(aq) ACI, (e) + H₂(aq)ローロ

Step 1: Calculation of Moles of Metal AMoles of Metal = Mass/Molar MassMoles of Metal = 0.225 g/27.0…

Answered: A quantity of 0.225 g of a metal A…

Chemistry

10th Edition

ISBN:9781305957404

Author:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Publisher:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Chapter1: Chemical Foundations

Section: Chapter Questions

Problem 1RQ: Define and explain the differences between the following terms. a. law and theory b. theory and...

See similar textbooks

Similar questions

A mixture of Oxygen and Chlorine gases contain twice as much Oxygen (by mass - in grams) as Chlorine (2xO2g= Cl2g) If the total pressure of the mixture is 565mmHg, then what is the partial pressure of chlorine ? Using (from Ideal Gas law - Dalton) Partial Pressure Partial Pressure (?) Moles of Chlorine (?) / (divided) by Total Moles O2+Cl2 or written nA / nT (Total Pressure) I've tried to determine what multiple of Oxygen, in grams - by atomic mass, will give me 1/2 the weight of Chlorine - in order to determine Mole (mol) Count for each DIATOMIC. I'm not frustrated, but i'm a little tired of going through this, for 2 hours, I have looked at - similar formula - variations, but this is the challenge. thanks
Using the ideal gas law (PV=nRT), calculate the temperature of a gas in a canister 10.4 L in volume. The pressure is 234,567 Pa and there are 32 moles. (R= 0.0821 L atm/ mol K) (101,325 Pa/1 atm).
Magnesium metal reacts with hydrochloric acid to produce hydrogen gas, H2. Mg(s) + 2HC1(aq)MgCl, (aq) + H2 (g) Calculate the volume (in liters) of hydrogen produced at 29°C and 720. mmHg from 9.40 x 10-2 mol Mg and excess HCI Volume L. %3D
0.374 m3 of hydrogen was observed at −7.0 oC and atmospheric pressure (105 Pa). The relative molecular mass of hydrogen is 2.02 g mol-1 and the gas constant, R, is 8.314 J K−1 mol−1 a) Calculate the number of moles (round to 2 decimal places): b) Calculate the mass of the gas in g (round to 3 significant figures):
The combustion of fuel in your car engine requires oxygen gas, which is supplied as air (21% oxygen molecules) into the engine. Consider a car that is using 100% ethanol, C2H5OH, as fuel. If your engine intakes 2.71 L of air per minute at 1.00 atm and 25ºC, what is the maximum volume of ethanol (0.789 g/mL) that can be burned per minute? Hint: You can ignore the "per minute" information because both the ethanol and air are being quantified per minute. Enter the answer to three significant figures in units of mL.
Use the ideal gas law equation to calculate the following: The pressure exerted by 2.00 moles of oxygen gas confined to a volume of 500mL, at 20.0℃. The volume of hydrogen gas in a steel cylinder if 0.525 mol of the gas exerts a pressure of 3.00 atm at a temperature of 10.0℃.
For the equation below, if you started with 32.6 L of B₂H4 gas at 2.1 atm and 25°C, and your reaction went in 35% yield, how many molecules of product did you isolate? B₂H4 + 2 HF --> 2 BH₂F Please note: If you examine the equation carefully, you will see that it is not balanced correctly (and in fact cannot be balanced). Nevertheless, do NOT correct the coefficients or the chemical substances. Use the equation as it has been given to you. a. 5.9 X 1023 molecules b. 1.2 X 1024 molecules c. 9.6 X 1024 molecules d. 3.8 X 1025 molecules
a.) The percent-by-mass of an unknown compound was found to be 24.3% carbon and 71.7% chlorine, with the remainder being hydrogen. Determine the empirical formula of this compound. b.) In a separate experiment, 5.6 grams of the vaporized compound was found to occupy 1.09 liters at 226°C and 785 mmHg. Determine the molar mass and molecular formula of the compound.
How many grams of NaClO3 were decomposed?
#4
Butane, CAH10, is an easily liquefied gaseous fuel. Calculate the density of butane gas at 0.857 atm and 22 °C. Give the answer in grams per liter. Given that R = 0.0821 L-atm/(mol-K), AW of C = 12.01 amu, and AW of H = 1.01 amu. a. 1.83 g/L. O b. 2.08 g/L. O c. 1.56 g/L. o d. 1.35 g/L. O e. 2.32 g/L.
Consider the reaction shown below: 2 HCl (aq) + CaCO3 (s) ⟶ CO2 (g) + CaCl2 (aq) + H2O (l) How many grams of CaCO3 would need to be consumed to fill a 908.001 mL balloon with CO2 gas, assuming a constant temperature of 21.7 oC and constant pressure of 0.958 atm? Express your answer in units of grams using at least three significant figures.

SEE MORE QUESTIONS

Recommended textbooks for you

Chemistry

ISBN:

9781305957404

Author:

Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Publisher:

Cengage Learning

Chemistry

ISBN:

9781259911156

Author:

Raymond Chang Dr., Jason Overby Professor

Publisher:

McGraw-Hill Education

Principles of Instrumental Analysis

Chemistry

ISBN:

9781305577213

Author:

Douglas A. Skoog, F. James Holler, Stanley R. Crouch

Publisher:

Cengage Learning

Chemistry

ISBN:

9781305957404

Author:

Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Publisher:

Cengage Learning

Chemistry

ISBN:

9781259911156

Author:

Raymond Chang Dr., Jason Overby Professor

Publisher:

McGraw-Hill Education

Principles of Instrumental Analysis

Chemistry

ISBN:

9781305577213

Author:

Douglas A. Skoog, F. James Holler, Stanley R. Crouch

Publisher:

Cengage Learning

Organic Chemistry

Chemistry

ISBN:

9780078021558

Author:

Janice Gorzynski Smith Dr.

Publisher:

McGraw-Hill Education

Chemistry: Principles and Reactions

Chemistry

ISBN:

9781305079373

Author:

William L. Masterton, Cecile N. Hurley

Publisher:

Cengage Learning

Elementary Principles of Chemical Processes, Bind…

Chemistry

ISBN:

9781118431221

Author:

Richard M. Felder, Ronald W. Rousseau, Lisa G. Bullard

Publisher:

WILEY

SEE MORE TEXTBOOKS