Fundamentals, Gases-CH301
The reaction below has a percent yield of 45.0%. H2(g) + Cl2(g) ⟶2HCl(g) How many moles of HCl gas are produced if 15.5 L of Cl2 at STP and excess H2 are reacted?
0.623 mol 1. Balance the equation. 2. Write down all units given. n= x R= 0.08206 T= 273.15 P= 1 atm v= 15.5 L 3. Determine moles needed. 1 mol of H2 and 1 mol of Cl2 yields 2 mol of HCl 4. Use PV=nRT to calculate moles of Cl2. n= PV/RT 1 x 15.5/0.08206 x 273.15= 0.691 mol of Cl2 5. Use moles of Cl2 to calculate how many moles of HCl are created. 0.691 mol Cl2 x 2 mol HCl/1 mol Cl2= 1.382 mol HCl 6. Calculate theoretical yield. 1.382 x .45= 0.6219 mol HCl
How many moles of gaseous carbon dioxide are there in 15 L at STP?
0.67 moles 1. Write out all units given. V= 15 L T= 273.15 P= 1 atm R= 0.08206 n= x? 2. Use PV=nRT to solve for moles. n= PV/RT 3. Plug in and solve. 1 x 15/ 273.15 x 0.08206= 0.66 moles of CO2
The density of a gaseous compound of phosphorous is 0.943 g/L-1 at 461 K when its pressure is 708 Torr with a molar mass of 38.3 g/mol. If the compound remains gaseous, what would be its density at 0.5 atm and 252 K? Please answer in g/L.
0.926 g/L 1. Write out all units given. V=? P= 0.5 atm T= 252 k M= 38.3 g/mol R= 0.0 2. Use density equation to determine density. M x (P/RT)= p 3. Plug in and solve. 38.3 x (0.5/0.08206 x 252)= 0.926 g/L
How do you derive an empirical formula?
1. Assume a mass that is easy to work with- 100 grams. 2. Calculate the number of moles using the formula: mass/molecular weight= moles. 3. Scale your subscripts to make them all integers.
What are the steps to balancing an equation?
1. Identify the most complex molecule. 2. Choose an element that's in only 1 product and 1 reactant. Balance. 3. Balance the remaining elements from most to least complex. 4. Count to check. If and only if one of your coefficients is a fraction: 5. Multiply out to remove the fraction by the dividend.
How do you derive a molecular formula?
1. Take the molecular weight of your unknown formula and the molecular weight of your empirical formula and divide them to determine your ratio. 2. Multiply the quotient by the empirical formula to obtain the subscripts for your molecular formula.
How do you calculate percent yield?
1. Write and balance out equation if necessary. 2. Calculate moles of given element with formula: mass given/mass of elements= moles 3. Calculate theoretical mols of the product from reaction by: dividing moles calculated in step 2/by amount of elements present in balanced equation x the coefficient of the reactant 4. Calculate theoretical mass of product by multiplying theoretical moles from step 3 by mass of the product. 5. Calculate percent yield by dividing grams given in problem of actual yield by theoretical yield calculated in step four and multiply by 100. Remember, percent yield should always be under 100% and units in the yield must be the same because they have to cancel.
What is the density of nitrogen gas at STP?
1.25 g/L 1. Write out all units given. T= 273.15 P= 1 atm R= 0.08206 L atm V= 22.4 liters 2. Remember the simplest density formula. mass/volume= p 3. Plug in and solve for density. 28.02/22.4= 1.25 g/L
Consider the following reaction: 2Al + 6HCl ⟶2AlCl3 + 3H2 This reaction has a yield of 82.5%. How many moles of HCl are needed to produce 14.0 L of H2 at 351 K and 1.11 atm?
1.31 mol 1. Write out all units given. Actual yield= .85 P=1.11 atm V=14 L n= x R= 0.08206 T= 351 K 2. Determine moles needed. 6 mol HCl yields 3 moles of H2 3. Use PV=nRT to calculate moles of H2. n= PV/RT 1.11 x 14/0.08206 x 351= 0.55 mol H2 4. Calculate the theoretical yield of H2. 0.55/0.825= 0.66 mol H2 5. Use moles of H2 to determine how many moles of HCl are needed. 0.66 mol H2 x 6 mol HCl/3 mol H2= 1.31 mol HCl
If you have 8.3 x 10^28 molecules of CO2, how many moles of CO2 are present?
1.4 x 10^5 moles of CO2 Explanation: 8.3 x 10^28/Avogadro's number= 1.4 x 10^5
A mixture of three gases (A, B, and C) is at a total pressure of 7.20 atm. The partial pressure of gas A is 1.73 atm; that of gas B is 3.66 atm. What is the partial pressure of gas C? Please answer in atm.
1.81 atm 1. Write out pressures given. pA= 1.73 atm pB= 3.66 atm ppC= ? 2. Add pressures given. 1.73 + 3.66 + x= 7.2 atm 3. Solve for x. 7.2- 1.73-3.66= 1.81 atm
Consider the following reaction: CH4(g) + 2O2(g)⟶CO2(g) + 2H2O(l) What is the final volume if 10 L of methane (CH4) reacts completely with 20 L of oxygen?
10 L Avogadro's Principle tells us that if pressure and temperature remain constant, volume is proportional to the number of moles of a gas. So, we can use the chemical equation to relate volumes of gases instead of moles. Make sure you take into account the phases of the reactants and the products.
A 22.4 L vessel contains 0.02 mol H2 gas, 0.02 mol N2 gas, and 0.1 mol NH3 gas. The total pressure is 700 Torr. What is the partial pressure of the H2 gas?
100 torr 1. Write out all units given. v= 22.4 L n= 0.02 mol H2 n= 0.02 mol N2 n= 0.1 mol NH3 P= 700 torr 2. Remember that. ppH2 + ppN2 + ppNH3= 700 torr 3. Calculate the total number of moles. 0.02 mol + 0.02 mol + 0.1 mol= 0.14 mol total 4. Calculate the mole fraction of H2 using moles given/ moles total. 0.02 mol H2/0.14 mol= 0.143 5. Multiply mole fraction by total pressure.' 0.143 x 700= 100.1 torr
What volume of pure oxygen gas (O2) measured at 546 K and 1.00 atm is formed by complete dissociation of 0.5 mol of Ag2O? 2Ag2O(s) ⟶4Ag(s) + O2(g)
11.2 L 1. Balance the equation. 2. Write out units given. V= x T= 546 K P= 1 atm n= 0.5 mol Ag2O R= 0.08206 L atm 3. Determine moles needed. 2 moles Ag2O yields 1 mole of O2 4. Use PV=nRT to determine V of Ag2O V= nRT/P 0.5 x 0.08206 x 546/1=22.4 L Ag2O 5. Calculate volume of O2 needed. 22.4 L Ag2O x 1 mol of O2/2 mol Ag2O= 11.2 L O2
What is the volume of 0.500 mol of O2 (g) at STP?
11.2 L 1. Write out all units given. P= 1 atm V= x n= 0.5 mol R= 0.08206 T= 273.15 degrees Kelvin. 2. Use PV=nRT to solve for V. V= nRT/P 3. Plug in and solve. 0.5 x 0.08206 x 273.15/1= 11.2 L
Consider the following reaction: 2 C6H6 + 15 O2⟶ 12 CO2 + 6 H2O 39.7 grams of C6H6 are allowed to react with 105.7 g of O2. How much CO2 will be produced by this reaction?
116.3 g 1. Balance the equation. 2. Determine moles needed. 2 moles of C6H6 and 15 moles of O2 yields 12 moles of CO2 3. Calculate limiting reactant with grams given. 39.7 g of C6H6 x 1 mol C6H6/78.11 g C6H6 x 12 mol CO2/2 mol C6H6 x 44.01 g CO2 /1 mol CO2 = 134 g of CO2 105.7 g x 1 mol O2/32 g O3 x 12 mol O2/15 mol O2 x 44.01 g CO2/1 mol CO2 = 116.29 g CO2 (limiting reactant)
Consider the UNBALANCED reaction below. Al2(SO4)3 + NaOH⟶Al(OH)3 + Na2SO4 Balance this equation using the smallest possible integers. What is the sum of the coefficients in the balanced equation?
12
When aluminum metal is heated with manganese oxide, the following reaction occurs: Al + MnO2⟶Al2O3 + Mn Balance this equation. What is the sum of the coefficients of ALL species in the balanced chemical equation?
12
What is the mass of oxygen gas in a 16.6 L container at 34.0°C and 6.22 atm?
131 g 1. Write out all units given. V= 16.6 L P= 6.22 atm T= 307.15 K n= x R= 0.08206 L Atm 2. Calculate moles. n= PV/RT 6.22 x 16.6/0.08206 x 307.15= 4.09 mol of O2 3. Calculate mass. 4.09 mol O2 x 32 g O2/1 mol O2= 131 g O2
A sample of nitrogen gas is contained in a piston with a freely moving cylinder. At 0°C, the volume of the gas is 371 mL. To what temperature must the gas be heated to occupy a volume of 557 mL?
137°C 1. Write out all units given. T1= 273.15 V1=371 mL T2= x V2= 557 mL 2. Remember V1/T1= V2/T2 371/273.15= 557/x 3. Cross multiply and divide. 371x= 273.15 x 557 x= 409.8- 273.15= 137 degrees celsius
Calculate the volume of methane (CH4) produced by the bacterial breakdown of 3.87 kg of sugar (C6H12O6) at 258 K and 726 torr. C6H12O6(aq)→3CH4(g) + 3CO2(g)
1430 L 1. Balance the equation. 2. Write out all units given. P= 726 torr=0.955 atm V= x T= 258 k 3. Change kg to g. 3.87 x 1000= 3870 g of C6H12O6 4. Determine moles needed. 1 mol of sugar yields 3 moles of CH4 5. Calculate moles of sugar. 3870 g sugar x 1 mol sugar/180 g sugar=21.5 mol of Sugar 6. Calculate moles of CH4 created. 21.5 mol Sugar x 3 mol CH4/1 mol Sugar= 64.5 mol CH4 7. Use PV=nRT to calculate volume. V= nRT/P 8. Plug in and solve. 0.08206 x 64.5 x 258/0.955 atm= 1429.90 L
Consider the reaction: 4KO2 (s) + 2CO2 (g)→2K2CO3 (s) + 3O2 (g) How much KO2 is needed to react with 75.0 L of carbon dioxide at -25˚C and 215 kPa?
15.6 mol 1. Balance the equation. 2. Determine moles needed. 4 moles of KO2 and 2 moles of CO2 yields 2 mols of K2CO3 and 3 moles of O2. 3. Write out all units given. p= 215 kPa v= 75 L T= 248.15 n= R= 8.314 3. Use PV=nRT to determine how many moles of given gas you have. n= PV/RT 4. Plug in and solve for n. n= 215 x 75/8.314 x 248.15= 7.81 mol of CO2 5. Use balanced equation to turn moles of CO2 into moles of KO2 needed. 7.81 mol x 4 mol KO2/2 mol CO2= 15.62 mol CO2
At 80.0°C and 12.0 torr, the density of camphor vapor is 0.0829 g/L. What is the molar mass of camphor?
152 g/mol 1. Write out all units given. P= 12 torr T= 353.15 p= 0.0829 R= 62.36 torr 2. Remember density formula for molar mass. m= pRT/P 3. Plug and solve for molar mass. 0.0829 x 62.36 x 353.15/12= 151.5 g of Camphor
Consider the following reaction: 2HCl + Na2CO3 ⟶2NaCl + H2O + CO2 For this reaction, 179.2 L of CO2 is collected at STP. How many moles of NaCl are also formed?
16.0 moles 1. Balance the equation. 2. Determine moles needed. 2 mol NaCl and 1 mol CO2 3. Write out all units given. V= 179.2 L P= 1 atm n= x R= 0.08206 T= 273.15 K 4. Calculate moles of CO2. n= PV/RT 1 x 179.2/0.08206 x 273.15= 8 mol of CO2 5. Use moles of CO2 collected to calculates mol of NaCl formed. 8 mol CO2 x 2 mol NaCl/1 mol CO2= 16 mol NaCl
What is the molar mass of a gas if 0.473 g of the gas occupies a volume of 376 mL at 23.0°C and 1.90 atm?
16.1 g/mol 1. Write gown all units given. g= 0.473 g V= .376 L T= 296.15 P= 1.9 atm R= 0.08206 n= x 2. Calculate moles. n= PV/RT 1.9 x .376/0.08206 x 296.15= 0.029 mol 3. Calculate grams by dividing moles given by moles calculated in step 3. 0.473/0.029= 16.31 g/Mol
If 100.0 grams of copper (Cu) completely reacts with 25.0 grams of oxygen, how much copper (II) oxide (CuO) will form from 140.0 grams of copper and excess oxygen? (Note: CuO is the only product of this reaction.)
175 g CuO 1. Add grams of two reactants given to determine final grams of product created. 100 g Cu + 25 g of O = 125 g of CuO 2. Write a ratio using reactants and the product. 100 g of Cu/125 g CuO = 140 Cu/x g CuO 4. Cross multiply and divide. 125 x 140= 100x 5. Solve for x 125 x 140= 17500 17500/100= 175 g CuO
When the equation PbS + O2 ⟶PbO + SO2 is balanced, the coefficients are ________, respectively.
2,3,2,2
What pressure (in atm) is exerted by 0.509 mol of hydrogen gas in a 6.03 L container at 34˚C?
2.13 atm 1. Write out all units. P= x V= 6.03 L n= 0.509 mol R= 0.08206 L atm T= 307.15 degrees Kelvin 2. Use PV=nRT do solve for P. P= nRT/V 3. Plug in values and solve for P. 0.509 x 0.08206 x 307.15/6.03= 2.127 atm
One method of estimating the temperature of the center of the sun is based on the assumption that the center consists of gases that have an average molar mass of 2.00 g/mol. If the density of the center of the sun is 1.40 g/cm3 at a pressure of 1.30 x 109 atm, calculate the temperature.
2.26 x 10^7 °C Pray this isn't on the test.
Consider the following chemical reaction: CH4 + 2O2 → CO2 + 2H2O 4 moles of methane (CH4) are combined with 5 moles of oxygen (O2). How much carbon dioxide (CO2) will form?
2.5 moles 1. Make sure your equation is balanced. 2. Determine initial mole ratios from equation. 1 mole of Methane and 2 moles of oxygen yields 1 mole of CO2. 3. Use moles given in the problem to determine limiting reactant. 4 moles of CH4 x 1 mole of CO2/1 mole of Methane= 4 moles of CO2 5 moles of O2 x 1 mole of CO2/2 moles of O2= 2.5 moles of CO2 - Limiting reactant 4. Use your limiting reactant to determine how many moles of CO2 will form. 2.5 moles of CO2
Consider the following reaction: 2 CO + O2⟶ 2 CO2 How much oxygen is required to convert 35 g of CO into CO2?
20 g 1. Balance the equation. 2. Determine moles needed. 2 moles of CO and 1 mole of O2 yields 2 moles of CO2 3. Calculate moles of grams given. 35 g CO x 1 mol CO /28.01 g CO= 1.24 mol of CO 4. Calculate moles of needed element using the answer from step 3. 1.24 mol of CO x 1 mol of O2/2 mol CO= 0.62 mol of O2 5. Change moles to grams. 0.62 mol O2 x 32 g O2/1 mol O2= 19.84 g of CO2
At STP, 1 mole of gas occupies how much volume (in L)?
22.4 L Remember this.
A gas is enclosed in a 10.0 L tank at 1200 mmHg pressure. Which of the following is a reasonable value for the pressure when the gas is pumped into a 5.00 L vessel?
2400 mmHg 1. Write out all units given. P1= 1200 mmHg V1= 10 L P2= x V2= 5 L 2. Think practically about what is given. Remember relationship between pressure and volume. Volume has doubled so pressure will also double. 1200 x 2= 2400 mmHg
Consider the following chemical reaction: N2 + 3H2 → 2NH3 If 5.00 g of hydrogen gas reacts with excess nitrogen, what mass of ammonia (NH3) will be produced? This reaction is expected to have a 90% yield.
25.5 g 1. Make sure your equation is balance. 2. Determine mole ratio using equation given. 1 mole of N2 and 3 moles of H2 yields 2 moles of NH3 3. Use grams given in equation to determine how many moles of NH3 will be produced. 5 g of H2x 1 mole of H2/2.016 g H2 x 2 mol NH3/3 mole of H2= 1.653 mol of NH3 4. Use moles calculated to determine grams of NH3. 1.653 mols NH3 x 17.031 g of NH3= 28.1 g 5. Multiply grams calculated by percent yield. 28.1 x .90= 25.29 g
Consider the following reaction: N2(g) + 3H2(g) →2NH3(g) If the reaction is carried out at constant temperature and pressure, how much H2 is required to react with 9.8 L of N2?
29.4 L 1. Balance the equation. 2. Determine moles needed. 1 mole of N2 and 3 moles of H2 3. Calculate moles of H2. 9.8 L N2 x 3 liters of H2/1 Liter of N2= 29.4 H2 4. Remember to treat liters as moles in these questions.
Consider the following chemical reaction: 2H2 + O2 → 2H2O How many moles of water are created if 3 moles of hydrogen react completely with excess oxygen?
3 moles 1. Make sure your equation is balanced. 2. Determine initial mole ratio from equation. 2 moles of H2 and 1 mole of O2 yields 2 moles of water 3. Use moles given in the problem to determine how many moles of water will be created. 3 moles H2 x 2 moles H2/2 moles of H20= 3 moles of H20
Consider the following reaction: N2 + 3 H2 ⟶ 2 NH3 How many MOLECULES of NH3 can be produced from the reaction of 74.2 g of N2 and 14.0 moles of H2?
3.19 x 10 ^24 molecules 1. Balance the equation. 2. Determine moles needed. 1 mol of N2 and 3 mol of H2 yield 2 moles of NH3 3. Calculate moles from gram given. 74.2 g N2/28.02 = 2.64 mol of N2 14 mol of H2 4. Use moles calculated from step 3 to determine mol of NH3 and limiting reactant. 2.64 mol of N2 x 2 mol NH3/ 1 mol N2= 5.28 mol NH3 (Limiting reactant) 14 mol of H2 x 2mol NH3/3 mol H2= 21 mol of NH3 5. Multiply moles calculated in step 4 from limiting reactant by Avogadro's number to get amount in molecules. 5.28 mol x 6.022 x 10^23=3.17 x 10^24
What is the ratio of the average velocity of hydrogen molecules to that of neon atoms at the same temperature and pressure?
3.2 1. Set up a ratio of Vrms of H2 to Vrms of Ne. Remember derivation!!! Vrms Ne/Vrms H2 = sqrt Molar mass of Ne/Molar mass of H2 2. Calculate molar mass of both. H2= 2.016 Ne= 20.17 3. Plug into equation and solve. sqrt 20.17/2.016= 3.16
Methane (CH4) is a greenhouse gas. How many hydrogen ATOMS are in 1.4 moles of methane?
3.4 x 10^24 atoms of Hydrogen 1.4 moles x Avogadro's number= 8.4308 x 10^23 4 atoms of hydrogen in methane: 4 X 8.4308 x 10^23= 3.4 x 10^24 atoms of Hydrogen
Consider the following reaction: 4Fe(s) + 3O2(g)⟶2Fe2O3(s) If 12.50 g of iron (III) oxide (rust) are produced from 8.74 g of iron, how much oxygen gas is needed for this reaction?
3.76 g Oxygen 1. Balance the equation. 2. Identify moles needed. 4 moles of Fe and 3 moles of O2 yields 2 moles of Fe2O3 3. Calculate moles present to produce given. 12.5g of Fe2O3/159.7= 0.0782 mol of Fe2O3 4. Use moles calculated in step 3 to determine how much of target element is needed. 0.0782 mol of Fe2O3 x 3 mol of O2 /2 mol Fe2O3= 0.1173 mol of O3 5. Change moles to grams. 0.1173 x 32= 3.75 g of O2
Consider the combustion of methanol at some high temperature in a constant-pressure reaction chamber: 2CH3OH (g) + 3O2 (g)⟶2CO2 (g) + 4H2O (g) If you react 10 L of oxygen gas with 18 L of methanol gas, what will the final volume be?
31.3 L 1. Balance the equation. 2. Write out all units given. VO2= 10 L VCH3OH= 18 L Vf= ? 2. Determine moles needed. 2 moles CH3OH and 3 moles of O2 yields 2 moles CO2 and 4 moles of H20 3. Calculate limiting reactant for both products. 10 L O2 x 2 L CO2/3 L O2= 6.6 L CO2 10 L O2 x 4 L H2O/ 3 L O2= 13.3 L H20 4. Calculate how much methanol is unused. 10 L O2 x 2 L CH3OH/3 L O2= 6.67 CH3OH unused 5. Calculate how much methanol remains. 18-6.67= 11.33 L CH3OH 6. Add up all the volumes to get your final volume. 13.33 + 11.33+ 6.67= 31.33 L
A 125 mL sample of a gas at 27˚C and 900 Torr is heated to 60˚C and the pressure is allowed to drop to 400 Torr. What volume will the gas occupy at these new conditions?
312 mL 1. Write down all units given. V1=125 mL= 0.125 L T1= 300.15 k P1= 900 torr V2= x T2= 333.15 K P2= 400 torr 2. Determine which gas law to use. Combined gas law= P1V1/T1= P2V2/T2 3. Plug in factors. 900 torr x 125 mL/300.15k= 374.81 400x/333.15=374.81 4. Solve for X 400x/333.15=374.81 400x= 124,868.816 124,868.816/400= 312.17 mL
What mass of O2 is required to produce 14.5 g of CO2 if the reaction has a 65.0% yield? CH4(g) + 2O2(g) ⟶ CO2(g) + 2H2O(g)
32.4 g 1. Balance the equation. 2. Calculate the theoretical yield of given. 14.5/.65= 22.3 g CO2 3. Determine moles needed. 2 moles of O2 yields One mole of CO2 4. Use theoretical yield to calculate moles of CO2 created. 22.3 g CO2 x 1 mol CO2/44.01 g CO2= 0.507 mol CO2 5. Use moles calculated in step 3 to determine mass of O2 needed. 0.507 mol CO2 x 2 mol O2/1 mol CO2 x 32 g O2/1 mol O2= 32.4 g
What is the percent yield for the following reaction if 10g H2 reacts to give 19g of NH3? N2 + 3 H2 ---> 2 NH3
34% 1. Make sure your equation is balanced. 2. Determine moles from equation given. 3 moles of H2 to make 2 moles of NH3. 3. Using grams given of reactant, calculate limiting reactant/amount of moles of product created. 10 g H2 x 1 mol H2/2.016 g H2 x 2 mol NH3/3 mol H2= 3.306 mol of NH3 4. Use moles of product calculated to determine theoretical yield. 3.306 mol of NH3 x 17.031 g of NH3/1 mol of NH3= 56.31 g of NH3 5. Use percent yield = actual/theoretical x 100 to determine your answer. 19/56.31 x 100= 33.73%
The density of a gaseous compound of phosphorous is 0.943 g\L-1 at 461 K when its pressure is 708 Torr. What is the molar mass of the compound? Please answer in g/mol.
38.3 g/mol 1. Write out all units given p= 0.943 g/L P= 708 torr T= 461 K M=? R=62.36 torr/L 2. Use density equation to determine the molar mass. M= pRT/P 3. Plug in and solve. 0.943 x 62.36 x 461/708= 38.2 g/mol
The measurement 4.7 x 10^-3 m could also be written as...
4.7 mm
How many atoms of hydrogen are contained in 2 moles of methane (CH4)?
4.82 x 10^24 atoms 1. Determine how many atoms of CH4 there are in total. 2 x 6.02 x 10^23= 1.204 x 10^24 atoms of CH4 2. Multiply number of hydrogen present in the atoms of CH4 4 x 1.204 x 10^24= 4.816 x 10^24
A sample of gas at 200˚C has a volume of 5 L. The temperature changes and the volume expands to 10 L. What is this new temperature?
400˚C 1. Determine which gas law to use. Charles' law= Volume and Temperature are directly proportional. V1T1=V2T2 2. Set up the problem. 5 x 200= 10 * x 3. Cross multiply and divide. 5/200 x 10/x 200x10= 2000 2000/5= 400 degrees celsius
The chemical formula for ethanol is CH3CH2OH. What is the mass for one molecule of ethanol? What about one mole of ethanol?
46 amu, 46 g 1. Calculate molar mass of the entire molecule. 46.07 g/mol 2. Remember that mass of one mole of a substance is equal to that substances molecular weight. 46 g.
If you have 44.8 L of nitrogen gas at standard temperature and pressure, how much will it weigh?
56 g 1. Remember that the volume of an ideal at STP is 22.4 L. 22.4 x 2= 44.8 L meaning you have 2 moles of N2. 2. Convert moles of N2 to mass. 2 mol N2 x 28.02 /1 mol N2= 56 g N2
Consider the following UNBALANCED chemical equation: Ca(OH)2(aq) + H3PO4(aq)⟶Ca3(PO4)2(s) + H2O(l) What is the coefficient for H2O when the reaction is balanced using the smallest possible integers?
6
A 5.00 L sample of a gas exerts a pressure of 1040 torr at 50.0°C. In what volume would the same sample exert a pressure of 1.00 atm at 50.0°C?
6.84 L 1. Write out all units given. V1= 5 L P2=1040 torr/760=1.36 atm T1= 323.15 V2= x P2= 1 atm T2= 323.15 2. Remember P1V1/T1= P2V2/T1 5 x 1.36/323.15 = 1X/323.15 3. Solve for x x= 6.84 L
A sample that contains 0.023 mol of a gas at 80˚C has a pressure of 800 Torr. What is the volume (in mL)?
633 mL 1. Write out all units. P= 800 torr = 1.052 atm V= x n= 0.023 mol R= 0.08206 T= 353.15 K 2. Use PV=nRT to solve for V. V= nRT/P 3. Plug in values to solve for V. 0.023 x 0.08206 x 353.15/1.052= 0.633 L x 1000= 633 mL
A flask containing 163 cm3 of hydrogen was collected under a pressure of 26.7 kPa. What pressure would have been required for the volume of the gas to have been 68 cm3, assuming the temperature is held constant?
64.0 kPa 1. Write out all units given. V1= 163 cm^3 P1= 26.7 kPa P2= x V2= 68 cm^3 2. Remember P1V1=P2V2 163 x 26.7= 68x 3. Plug in and solve for x. x= 64 kPa
A chemist has synthesized a greenish-yellow gaseous compound that contains only chlorine and oxygen and has a density of 7.71 g/L at 36.0°C and 2188.8 mmHg. What is the molar mass of the compound?
67.9 g/mol 1. Write out all units given. p= 7.71 T= 309.15 P= 2188 mmHg R= 62.4 2. Remember the formula for molar mass using density. m= pRT/P 3. Plug in and solve. 7.71 x 62.4 x 309.15/2188= 67.97 g/mol
Calculate the root mean square speed of fluorine gas at 425˚C. Please answer in units of m/s.
677 m/s 1. Remember Vrms= sqrt 3RT/M 2. Calculate molar mass of F2. F2= 37.997 g/mol 3. Convert g calculated to kg/mol. 37.997 x 0.00100= 0.037997 kg/mol F2 4. Remember that this is an energy/speed question so use R that is in Joules! 8.314 J 5. Plug in and solve. sqrt 3 x 8.314x 698.15/0.037997= 676.8
If you have 10 molecules of SO2 gas in 125000 molecules of air, what is the concentration of SO2 in parts per million?
80 ppm Yes. This is a stoichiometry problem. Let's make a ratio between the molecules of SO2 gas and the total molecules in the air, then convert that to molecules of SO2 gas per million molecules of gas: 10 molecules SO2/125000 molecules gas in mixture x 1x10^6= 80 ppm
The height of water in a water barometer is 883 cm at 20˚C. The density of water at 20˚C is 0.998 g•cm-3. What is the pressure?
86.4 Pa 1. Use pressure= density x gravitational constant x height of substance (P= pgh) 0.998x 981 x 883= 8644 kPa 2. Change units to pascals. 8644/100=86.44 pa
A sample of gas in a closed container at a temperature of 76°C and a pressure of 5.0 atm is heated to 399°C. What pressure does the gas exert at the higher temperature?
9.6 atm 1. Write out all units given. T1= 349.15 T2= 672.15 P1= 5 atm P2= x 2. Remember Gay-Lussac's Law. P1/T1= P2/T2 3. Plug in and solve. 5/349.15= 672.15/x x= 9.6 atm
Upon heating, potassium chlorate produces potassium chloride and oxygen. 2KClO3 ⟶2KCl + 3O2 What mass of oxygen would be produced upon thermal decomposition of 25 g of potassium chlorate (KClO3)? The molecular weight (MW) of potassium chlorate is 122.5 g/mol.
9.8 g 1. Balance the equation. 2. Determine moles needed. 2 moles of KClO3 yields 2 moles of KCl and 3 moles of O2 3. Calculate moles of KClO3 from problem. 25 g KClO3 x 1 mol KClO3/122.5 g/mol = 0.204 mol of KClO3 4. Use moles calculated in step 3 to determine moles of needed element. 0.204 mol KClO3 x 3 mol of O2/2 mol of KClO3= 0.306 mol of O2 5. Change moles to grams. 0.306 mol x 32 g O2/1 mol O2= 9.78 g O2
The pressure exerted on a 200. mL sample of hydrogen gas at constant temperature is increased from 4.87 atm to 9.91 atm. What will be the final volume of the sample? Express your answer in mL.
98.3 mL 1. Write down all the units you are given. v1= 200 mL or .200 L p1= 4.87 atm p2= 9.91 atm v2= x 2. Determine which gas law to use. Boyle's law= Volume and pressure are inversely proportional. P1V1=P2V2 3. Set up the problem. 4.87 x .200= 9.91x 0.974= 9.91x 0.974/9.91=0.9828 4. Change answer back to mL. 0.9828 x 1000= 98.28 mL
Which statement is false? A. At a given temperature, according to the kinetic molecular theory, the average velocity of He atoms is the same as the average velocity of F2 molecules. B. The average kinetic energies of molecules of samples of different ideal gases at the same temperature are the same. C. The molecules of a gas move in straight paths until they collide with other molecules or the walls of their containers, according to the kinetic molecular theory. D. The molecules of an ideal gas are relatively very far apart on the average.
A. At a given temperature, according to the kinetic molecular theory, the average velocity of He atoms is the same as the average velocity of F2 molecules. This statement is false. At a given temperature, according to the kinetic molecular theory, the average kinetic energy is the same for all gas molecules, but the average velocity will be greater for gas molecules with smaller mass. Helium atoms are lighter than fluorine gas molecules; therefore, the average velocity of He atoms is greater than the average velocity of F2 molecules.
From the molecular perspective, what is changing as the pressure is increased (at constant) temperature? A. The molecules get closer together. B. The space between the molecules does not change, but the molecules are moving faster. C. The measured pressure increases. D. The measured volume decreases. E. A, C and D are all an acceptable answer for this question.
A. The molecules get closer together.
What are some important things to remember when reading a skeletal drawing?
All the points on the line are carbons. All Carbon wants 4 bonds. If it is bound to an H, do not draw it. If it is bound to any other element, draw it.
Mole
Avogadro's number or 6.022 x 10^23 elementary entities
In an improved version of the gas law, V is replaced by (V - nb). Which of the following would you predict has the largest b? A. C2H4 B. C4H10 C. C2H6 D. C3H8 E. CH4
B. C4H10 Remember: b is the number of particles per volume, so the largest molecule will be your largest b value.
Referring to the data table, under what conditions does the Ideal Gas equation best model the real gas behavior? A. high pressure B. low pressure C. high temperature D. low temperature
B. low pressure
Elementary entities
Building blocks, molecules, ions, atoms, electrons
The analysis of a hydrocarbon revealed that it was 85.6281% C and 14.3719% H by mass. When 3.22 g of the gas was stored in a 1.2 L flask at -190.842°C, it exerted a pressure of 491 torr. What is the molecular formula of the hydrocarbon?
C2H4 1. Write down all units given. 85.6281% C 14.3719% H mw= 3.2 g V= 1.2 L P= 0.646 atm R= 0.08206 T= 82.308 K 2. Calculate moles. n=PV/RT 0.646 atm x 1.2/0.08206 x 82.308= 0.115 mol 3. Find out how much 1 mole of the gas weighs by dividing moles given by moles calculated in step 3. 3.22/0.115= 27.82 g/mol 4. Calculate molar masses of answer choices to see which one gives you the molar mass calculated in step 3. C2H4
Which of the following has the greatest number of ATOMS? A. 3.05 moles of argon B. These all have the same number of atoms. C. 3.05 moles of CH4 D. 3.05 moles of water
CH4 Argon consists of a single atom of argon. So... 3.05 moles of argon = 3.05 moles of atoms Water is a molecule made of 3 atoms. So... 3.05 moles of water = 3 x 3.05 = 9.15 moles of atoms CH4 is a molecule made of 5 atoms. So... 3.05 moles of CH4 = 5 x 3.05 = 15.25 moles of atoms. Therefore, 3.05 moles of CH4 has the most atoms.
Which has the higher mass density (g/L): a sample of O2 with a volume of 10 L, or a sample of Cl2 with a volume of 3 L? Both samples are at the same temperature and pressure.
Cl2 At the same temperature and pressure the two gases have the same number density (moles per volume). That is, both gases have the same number of molecules in one L of the gas. However, as the mass of Cl2 is more than twice as large as O2, the mass density of Cl2 will be about twice as high.
Elemental analysis shows that a compound with the molecular weight of 90.08 g/mol contains 40% Carbon, 6.71% Hydrogen, and 53.3% Oxygen. What are the empirical and molecular formulas?
Empirical= CH2O Molecular= C3H6O3 Start with empirical: 1. Assume a mass that's easy to work with: 100g. 2. Write out your ratio in grams: 40 g C: 6.71 g of H: 53.3 g of O 3. Calculate ratio in moles by dividing percentages given by molar mass of the element. 40g C/12.011= 3.33 6.71 g H/1.008= 6.65 53.3g O/16= 3.33 3.33:6.6:3.33 4. Divide through by the smallest integer. 3.33/3.33= 1 6.66/3.33= 2 Mole ratio: 1:2:1 5. Write out your answer. CH2O For Molecular: 1. Use the molecular weight given to determine what number is needed to multiply by empirical formula to get the molecular formula. 90.08= x(12.011 + 2 x 1.008 + 16) 2. Solve for x. X=3 3. Multiply through by X. 3 x CH2O= C3H6O3
In a mixture containing two components, the number of moles of one of the components divided by the sum of the number of moles of both components is called
Mole fraction The mole fraction of a component in a solution is its mole fraction divided by the sum of the mole fractions of all of the components of the solution.
NO2 is found to have a molar mass of 92 g/mol. What is the molecular formula for the compound?
N2O4 1. Take mass given in problem and divide it by molar mass of compound given. 92/46.0055 (molar mass of NO2) =2 3. Multiply by the quotient determined from step one. 2 x NO2= N2O4
A compound of nitrogen and oxygen is analyzed. It is found to be 30.4% nitrogen and 69.6% oxygen by mass. What is the empirical formula of the compound?
NO2 1. Assume a mass that's easy to work with. 100 g 2. Write out your ratio in grams: 30.4 g N: 69.6 g O 3. Calculate ratio in moles by dividing percentages given by molar mass of the element. N= 30.4/14.01= 2.17 mol N O= 69.6/16= 4.35 mol O 4. Divide through by the smallest integer. 2.17/2.17 4.35/2.17 NO2
How is number density related to molar volume?
Number density and molar volume are inversely proportional.
Empirical formula
a chemical formula that shows the composition of a compound in terms of the relative numbers and kinds of atoms in the simplest ratio or smallest integer representation of the molecular formula, like a "repeat unit" Ex: If C10H12O8 is the molecular formula, the empirical formula would be C5H6O4.
Percent yield formula
actual yield/theoretical yield x 100
The actual volume occupied by most gases at high pressures is larger than that predicted by the ideal gas law. This is because....
as the gas particles get closer together it is no longer possible to ignore the volume occupied by the particles themselves.
If the volume of a gaseous system is increased by a factor of 3 and the temperature is raised by a factor of 6, then the pressure of the system will __________ by a factor of __________.
increase by 2 Tripling the volume will decrease the pressure by a factor of 3 and sextupling the temperature will increase the pressure by a factor of 6, resulting a double the original pressure. Pressure x 1/3 x 6 = 2 times the original Pressure
For a given gas sample with a constant amount of gas at a constant temperature, the pressure and the volume of the sample are
inversely proportional
The mole concept is important in chemistry because...
it allows us to count atoms and molecules by weighing macroscopic amounts of material.
In an improved version of the gas law, V is replaced by (V - nb). The two quantities n and b in this equation represent, respectively, the...
number of moles of gas molar volume of the particles
It's best to think of the ideal gas law as a ___________ ___________ derived from the ____________ ___________ __________ so that we can understand real world insight from where the model fails.
physical model kinetic molecular theory
Molecular formula
shows the types and numbers of atoms combined in a single molecule of a molecular compound, actual formula mostly for organic compounds Ex: C10H12O8
SATP
standard ambient temperature and pressure; exactly 25 degrees celsius (298.15) and 100 kPa
Number density refers to...
the number of particles (moles) per volume.
You have a sample of H2 gas and Ar gas at the same temperature and pressure, but the H2 gas has twice the volume of the Ar gas. Assuming the gases behave ideally, which gas has the larger NUMBER DENSITY (gas particles per volume)?
they are the same At the same temperature and pressure, the number density must be the same. Even though the H2 gas must have twice the number of moles (since it has twice the volume), the two gases have the same number of molecules per a given volume (such as 1 liter).
STP
Standard Temperature and Pressure. 273 Kelvin (0 Celsius), 1 atmosphere (760 torr, 760 kPA).
You fill two balloons, A and B, with helium gas at the same temperature and pressure. You notice that balloon B has a larger volume than balloon A. What must be true about the masses of the gas inside each of the balloons?
The mass of the gas in balloon A is less than the mass of the gas in balloon B. According to Avogadro's law, the volume and amount/mass of a gas are directly proportional at a given temperature and pressure. Balloon B has a larger volume, therefore it has more moles of gas than balloon A and since we know that both of the balloons are filled with only helium, we can say that balloon B has a greater mass.
What would happen to the moles of a gas present if you decreased the volume by 1/2? Which law supports this?
The moles of the gas present would decrease by 1/2 Avogadro's law= Pressure and Volume are directly proportional
Two, sealed 1 L containers full of gas are at room temperature. Container A has a pressure of 4 atm and Container B has a pressure of 2 atm. What must be true about the number densities of the gases in these containers?
The number density of the gas in container A is twice the number density of the gas in container B. From the ideal gas law we know that pressure is governed by the number density of a gas sample. Container A has twice the pressure which means it must have twice the number density of gas as container B. Furthermore, since the volumes are identical we could say that container A has twice the gas moles as container B.
What would happen to the pressure of a gas if the volume is doubled? What law supports this?
The pressure would decrease by a factor of 2. Boyle's law: Pressure and volume are inversely proportional
What would happen to the volume of a gas if the pressure was doubled? What law supports this?
The volume would decrease by a factor of 2. Boyle's law: Pressure and volume are inversely proportional
What would happen to the volume if you doubled the moles of the gas present? Which law supports this?
The volume would double Avogadro's law= Pressure and volume are directly proportional
What would happen to the volume of the gas if the temperature was doubled? Which law supports this?
The volume would double. Charles' law: Volume and Temperature are directly proportional.
For a given gas sample with a constant amount of gas at a constant pressure, the temperature and the volume of the sample are
directly proportional
For a given gas sample at constant temperature and pressure, the moles of gas and the volume of the sample are
directly proportional.
Equal masses of hydrogen, oxygen, and nitrogen gas are all in the same container. Which of the three gases must have the highest partial pressure?
hydrogen Hydrogen is the lightest gas in the mixture. If equal masses of each gas are placed in the container, that would mean that hydrogen gas is present in the greatest number of moles (oxygen would have the fewest moles). Therefore, when calculating partial pressures, the hydrogen gas would have the largest mole fraction and therefore the largest partial pressure.
Partial pressure is a(n) _________ because pressure results from the collisions of all the gas molecules with the walls of the gas container. However, it is still important because we use partial pressures to describe concentrations of gases since the partial pressure is ________ the number of moles of that particular gas per unit volume.
idea proportional to Partial pressure is a concept and is not directly measurable. The partial pressure of a gas is proportional to the number of moles of that gas per unit volume.
