CHEM 1150 Williford FINAL EXAM PREP
What are the possible values of ml for l = 4?
-4,-3,-2,-1,0,1,2,3,4
The American Dental Association recommends that an adult female should consume 3.0 mg of fluoride (F-) per day to prevent tooth decay. If the fluoride is consumed in the form of sodium fluoride (45.24% F), what amount (g) of sodium fluoride contains the recommended amount of fluoride?
.0066g
What is the concentration (M) of the unknown H3PO4 solution when 30.00-mL sample of H3PO4 solution is titrated with 26.38 mL of 0.100 M NaOH solution. The neutralization reaction is: H3PO4(aq) + 3 NaOH(aq) →3 H2O(l) + Na3PO4(aq)
.0293 M
Calculate the molarity of the H2SO4solution if it takes 40.0 mL of H2SO4to neutralize 0.364 g of NaHCO3based on the following reaction. H2SO4(aq)+2 NaHCO3(s)→ Na2SO4(aq)+2H2O(l)+2CO2(g)
.0543 M
Calculate the molarity of theHClsolutionif it takes22.8mL of 0.150M Ca(OH)2to neutralize 35.0 mL of HCl based on the following reaction?Ca(OH)2(aq)+2 HCl(aq)→CaCl2(aq)+ 2 H2O (l)
.195 M
The smallest atoms can themselves exhibit quantum-mechanical behavior. Calculate the de Broglie wavelength (in nm) of a helium atom traveling 475 m/s. (mass of H atom = 6.649 x 10-24 g)
.210 nm
How many grams of magnesium solid can be produced when 25.0 mL of 0.75 M MgSO4 completely reacts with excess potassium solid? Refer to the equation: 2 K(s) + MgSO4(aq) → K2SO4(aq) + Mg(s)
.46 g
How many grams of MgCl2are in 25.0 mL of a 0.30 M MgCl2solution?
.71 g
What is the molarity of a solution containing 18.00 grams of hydrochloric acid dissolved in 500 mL of solution?
.987 M
What are the possible values of l for n = 5?
0,1,2,3,4
Glycerol is a syrupy liquid used in cosmetics and soaps. A 3.25-L sample of pure glycerol has a mass of 9.04 lb. What is the density of glycerol in g/cm3?
1.26 g/cm3
An automobile gasoline tank holds 21 kg of gasoline. When the gasoline burns, 84 kg of oxygen is consumed, and carbon dioxide and water are produced. What is the total combined mass of carbon dioxide and water that is produced?
105 kg products
How many grams of water freezes when liquid water releases 3.85 kJ of heat? (ΔHfusion = 6.01 kJ/mol)
11.5 g
How many mL of 1.2 M HCl is needed to dissolve 5.8 g of Al(OH)3based on the following reaction? Al(OH)3(aq)+ 3HCl (aq)→ AlCl3(aq)+ 3H2O(l)
190 mL
Calculate the frequency of the electromagnetic radiation in a microwave oven that emits 2.0 x 109 J per photon.
3.0x10^42 Hz
A proton in a linear accelerator has a de Broglie wavelength of 122 pm. What is the speed of the proton? (mass of proton = 1.673 x 10-24 g)
3250 m/s
How many milliliters of 0.139 M H2SO4is needed to completely neutralize 43.2 mL of 0.236 M NaOH based on the following reaction? H2SO4(aq)+ 2NaOH(aq)→ Na2SO4(aq)+ 2H2O(l)
36.7 mL
How many grams of lead (Cs Pb = 0.126 J/g∙°C) will release 4140 J of heat when the temperature changes from 95.7°C to 18.4°C?
425 g
How many milliliters of 0.41 M HCl is needed to completely react with 32.5 mL of 0.27 M Mg(OH)2? The neutralization reaction is: 2 HCl(aq) + Mg(OH)2(aq) → MgCl2(aq) + 2 H2O(l)
43 mL
What is the molarity of a solution containing 2.50 moles of sodium chloride in 50.0 mL of solution?
5 M
How many milliliters of 1.5 M NaOH solution contains 32.5 g NaOH?
540 mL
Calculate the wavelength (nm) of the electromagnetic radiation emitted by neon gas that has an energy of 3.10 x 10-19 J per photon.
641 nm
A new penny has a mass of 2.49 g and a volume of 0.349 cm3. Is the penny made of pure copper (dCu = 8.96 g/ cm3)?
7.13 g/cm3. No
Write the molecular equation, ionic equation, and net ionic equation for each pair of reactants. a. Pb(NO3)2(aq) + K2SO4(aq) b. HBr(aq) + LiHCO3(aq)
A) Pb(NO3)2(aq) + K2SO4(aq) → 2 KNO3(aq) + PbSO4(s) Pb2+(aq) + 2 NO3-(aq) + 2 K+(aq) + SO42-(aq) → 2 K+(aq) + 2 NO3-(aq) + PbSO4(s) Pb2+(aq) + SO42-(aq) → PbSO4(s) B) HBr(aq) + LiHCO3(aq) → LiBr(aq) + H2O(l) + CO2(g) H+(aq) + Br-(aq) + Li+(aq) + HCO3-(aq) → Li+(aq) + Br-(aq) + H2O(l) + CO2(g) H+(aq) + HCO3-(aq) → H2O(l) + CO2(g)
Calculate the effective nuclear charge experienced by the valence electrons of each atom. a. K b. Ca c. O d. C
A) +1 B) +2 C) +6 D) +4
Calculate the energy change when an electron in a hydrogen atom makes each transition: a. n=3→n=1 b. n=2→n=4 c. n=4→n=3
A) -1.94x10^-18 J B) 4.09x10^-19 J C) -1.06x10^-19 J
Calculate the designated quantities using stoichiometry of the following reaction: C3H8(g) + 5 O2(g) → 3 CO2(g) + 4 H2O(g) ΔH°rxn = −2217 kJ a. How much heat (kJ) is produced when 84.7 g of oxygen (O2) burns? b. How much heat (kJ) is released when 135 g of CO2 is produced? c. What mass (grams) of C3H8 must burn to emit 267 kJ of heat? d. How many grams of H2O are produced if 35800 J of heat is released?
A) -1170 kJ B) -2270 kJ C) 5.31 g D) 1.16 g
Use bond energy to calculate enthalpy of reaction (ΔHrxn) for each reaction. Refer to bond energy chart in the book. a. H2C=CH2(g) + H2(g) → H3C-CH3(g) b. CH3CH2OH(g) + 3 O2(g) → 2 CO2(g) + 3 H2O(g) c. 4 NH3(g) + 7 O2(g) → 4 NO2(g) + 6 H2O(g) d. CO(g) + H2O(g) → H2(g) + CO2(g)
A) -126 kJ (exothermic) B) -1025 kJ (exothermic) C) -562 kJ (exothermic) D) 90 kJ (endothermic)
Calculate the energy of an electron in a hydrogen atom on each energy level: a. n=3 b. n=5
A) -2.42x10^-19 J B) -8.72x10^-20 J
Use Hess's Law to calculate enthalpy of reaction (ΔHrxn) for each reaction. a. Reaction: Fe2O3(s) + 3 CO(g) → 2 Fe(s) + 3 CO2(g) ΔH = ? Use the following reactions and given ΔH values: 2 Fe(s) + 32 O2(g) → Fe2O3(s) ΔH = −824.2 kJ CO(g) + 12 O2(g) → CO2(g) ΔH = −282.7 kJ b. Reaction: CaO(s) + CO2(g) → CaCO3(s) ΔH = ? Use the following reactions and given ΔH values: Ca(s) + CO2(g) + 12 O2(g) → CaCO3(s) ΔH = −812.8 kJ 2 Ca(s) + O2(g) → 2 CaO(s) ΔH = −1269.8 kJ c. Reaction: 5 C(s) + 6 H2(g) → C5H12(l) ΔH = ? Use the following reactions and given ΔH values: C5H12(l) + 8 O2(g) → 5 CO2(g) + 6 H2O(g) ΔH = −3244.8 kJ C(s) + O2(g) → CO2(g) ΔH = −393.5 kJ 2 H2(g) + O2(g) → 2 H2O(g) ΔH = −483.5 kJ d. Reaction: CH4(g) + 4 Cl2(g) → CCl4(g) + 4 HCl(g) ΔH = ? Use the following reactions and given ΔH values: C(s) + 2 H2(g) → CH4(g) ΔH = −74.6 kJ C(s) + 2 Cl2(g) → CCl4(g) ΔH = −95.7kJ H2(g) + Cl2(g) → 2 HCl(g) ΔH = −92.3 k
A) -23.9 kJ (exothermic B) -117.9 kJ (exothermic) C) -173.2 kJ (exothermic) D) -205.7 kJ (exothermic)
Calculate the designated quantities related to pressure-volume work. a. How much work (in J) is required to expand the volume of a pump from 0.0 L to 2.5 L against an external pressure of 1.1 atm? b. The average human lung expands by about 0.50 L during each breath. If this expansion occurs against an external pressure of 1.0 atm, how much work (in J) is done during the expansion? c. The air within a piston equipped with a cylinder absorbs 565 J of heat and expands from an initial volume of 0.10 L to a final volume of 0.85 L against an external pressure of 1.0 atm. What is the change in internal energy of the air within the piston? d. A gas is compressed from an initial volume of 5.55 L to a final volume of 1.22 L by an external pressure of 1.00 atm. During the compression the gas releases 124 J of heat. What is the change in internal energy of the gas?
A) -280 J B) -51 J C) -76J D) 439 J
Determine the designated quantities of enthalpy and/or internal energy. a. When 1 mol of a fuel burns at constant pressure, it produces 3452 kJ of heat and does 11 kJ of work. What are ΔE and ΔH for the combustion of the fuel? b. The change in internal energy for the combustion of 1.0 mol of octane at a pressure of 1.0 atm is 5084.3 kJ. If the change in enthalpy is 5074.1 kJ, how much work is done during the combustion?
A) -3463 kJ B) 10 kJ
Calculate the frequency of the light emitted when an electron in a hydrogen atom makes each transition: a. n=4→n=2 b. n=5→n=1 c. n=5→n=4
A) -4.09x10^-19 J B) -2.09x10^-18 J C) -4.91x10^-20 J
Calculate the internal energy (ΔE) in kilojoules for each sample. a. A system releases 622 kJ of heat and does 105 kJ of work on the surroundings. b. A system absorbs 196 kJ of heat, and the surroundings does 117 kJ of work on the system. c. The gas in a piston (defined as the system) warms and absorbs 655 kJ of heat. The expansion performs 344 kJ of work on the surroundings. d. The air in an inflated balloon (defined as the system) warms over a toaster and absorbs 915 J of heat. As it expands, it does 77 kJ of work.
A) -727 kJ B) 313 kJ C) 311 kJ D) -76 kJ
Calculate the molarity of each solution. a. 0.38 mol of LiNO3 in 6.14 L of solution b. 72.8 g C2H6O in 2.34 L of solution c. 12.87 mg KI in 112.4 mL of solution
A) .062 M B) .675 M C) 6.898x10^-4 M
Calculate to the correct number of significant figures. a. 0.004 + 0.09879 b. 1239.3 + 9.73 + 3.42 c. 2.4 − 1.777 d. 532 + 7.3 − 48.523
A) .103 B) 1252.5 C) .6 D) 491
Calculate the designated quantities in the following stoichiometry examples referring to the equation: 6 Li(s) + N2(g) → 2 Li3N(s) a. What mass of lithium (in g) reacts completely with 58.5 mL of N2 gas at STP? b. How many mL of N2 reacted to produce 0.689 g Li3N at STP? c. What mass of lithium (in g) reacts completely with 58.5 mL of N2 at 21.3°C and 0.95 atm? d. How many mL of N2 reacted to produce 0.689 g Li3N at 40.0°C and 763 mmHg?
A) .109 g B) 222 mL C) 2.30x10^-3 mol D) 253 mL
Calculate moles of sodium in 8.5 g of each sodium-containing food additive (use subscript ratio method). a. NaCl (table salt) b. Na3PO4 (sodium phosphate) c. NaC7H5O2 (sodium benzoate) d. Na2C6H6O7 (sodium hydrogen citrate)
A) .15 mol B) .16 mol C) .059 mol D) .072 mol
Calculate the number of moles in each sample. a. 11.8 g Ar b. 55.98 g CF2Cl2
A) .295 mol Ar B) .4630 mol CF2Cl2
Calculate the quantities in the following dilutions. a. If 3.5 L of a 4.8 M SrCl2 solution is diluted to 45 L, what is the molarity of the diluted solution? b. To what volume (mL) should you dilute 50.0 mL of a 12 M stock HNO3 solution to obtain a 0.100 M HNO3 solution?
A) .37 M B) 6000 mL
Determine the number of moles in each sample. a. 55.98 g CF2Cl2 b. 1.95 x 1025 molecules CaO
A) .4630 mol B) 32.4 mol
Determine the oxidation number for each element, ion, or element in a compound. a. Cl2 b. Fe3+ c. CuCl2 d. CH4 e. Cr2O72-
A) 0 B) +3 C) -1 D) +1 E) -2
Assign formal charges to each element in each molecule or ion. a. CCl4 b. NF3 c. SO2 d. NO3-
A) 0 B) 0 C) 0 D) -1
Determine the limiting reactant in each pair based on the chemical reaction: 4 Al(s) + 3 O2(g) → 2 Al2O3(s) a. 1.0 mole Al reacts with 1.0 mole O2 b. 4.2 mole Al reacts with 3.6 mole O2 c. 18.5 g Al reacts with 15.7 g O2 d. 0.485 g Al reacts with 1.06 g O2 Calculate the theoretical yield (g) of Al2O3 for each sample
A) 1.0 mol Al is the limiting reactant B) 3.6 mol O2 is the limiting reactant C) 15.7 g O2 is the limiting reactant D) 0.985 g Al is the limiting reactant Theoretical Yield A) 51g B) 240 g C) 33.3 g D) 1.86 g
Use the prefix multipliers to express each measurement without any exponents. a. 1.2 × 10−9 m b. 2.2 × 10−3 s c. 1.5 × 109 g d. 3.5 × 10-6 L e. 8.3 x 103 m
A) 1.2 nm B) 2.2 ms C) 1.5 Gg D) 3.5 µL E) 8.3 km
Calculate the wavelength of the light emitted when an electron in a hydrogen atom makes each transition: a. n=2→n=1 b. n=4→n=1 c. n=5→n=3
A) 1.21x10^-7 m B) 9.74x10^-8 m C) 1.28x10^-6 m
Calculate the following quantities using molarity as a conversion factor. a. How many moles of KCl are in 0.556 L of 2.3 M KCl solution? b. How many milliliters of 0.200 M ethanol solution contains 0.0450 moles ethanol (C2H5OH)? c. What mass (g) of CaCl2 is needed to make 5.5 L of a 0.300 M CaCl2 solution?
A) 1.3 mol B) 225 ml C) 180 g
Round each number to three significant figures. a. 1.548937 × 107 b. 2.3499999995 c. 0.000045389 d. 79,845.82
A) 1.55× 10^7 B) 2.35 C). 0.0000454 D) 79,800
Calculate the mass (g) of each sample. a. 4.91 x 1021 atoms Pt b. 1.91 x 1017 molecules H2O
A) 1.59 g Pt B) 5.72x10^-6 g H2O
Use simple gas laws to calculate the designated quantities in each sample. a. A sample of gas has an initial volume of 13.9 L at a pressure of 1.22 atm. If the sample is compressed to a volume of 10.3 L, what is its pressure? b. A syringe containing 1.55 mL of oxygen gas is cooled from 95.3 °C to 0.0 °C. What is the final volume of oxygen gas? c. A sample of gas in a cylinder initially at 22.3°C is warmed causing the pressure to increase from 745 mm Hg to 1235 mm Hg. What is its temperature (°C) at the final pressure? d. A cylinder with a moveable piston contains 0.553 mol of gas and has a volume of 253 mL. What is its volume if we add 0.365 mol of gas to the cylinder? e. A weather balloon is inflated to a volume of 28.5 L at a pressure of 748 mmHg and a temperature of 28.0 °C. The balloon rises in the atmosphere to an altitude of approximately 25,000 ft, where the pressure is 385 mmHg and the temperature is −15.0 °C. Assuming the balloon can freely expand, calculate the volume of the balloon at this altitude.
A) 1.65 atm B) 1.15 mL C) 216 C D) 420 mL E) 47.5 L
Convert each temperature. a. 212 °F to °C b. 22 °C to K c. 0.00 K to °F d. 2.735 K to °C
A) 100 C B) 295 K C) -459.67 F D) -270.42 C
Calculate the percentage as indicated in the following examples. a. What is the percent (m/m) of KCl in a solution containing 27.8 g KCl in 250. g of solution? b. What is the percent (v/v) of NH3 in a cleanser containing 135 mL NH3 and 775 mL water? c. What is the percent (m/v) of HCl in a concentrated HCl solution contains 125 g of HCl in 0.550 L of solution?
A) 11.1% (m/m) B) 14.8% (v/v) C) 22.7% (m/v)
Calculate the designated quantities using the heat equation. a. How much heat (kcal) is required to warm 1.50 L of water from 25.0 °C to 100.0 °C? (Assume a density of 1.00 g/mL for the water; Cs water = 1.00 𝑐𝑎𝑙𝑔 ∙ °𝐶 ). b. How much heat (kJ) is released when 0.985 kg of sand cools from 120.0 °C to 18.5 °C? (Cs sand = 0.830 𝐽𝑔 ∙ °𝐶 ) c. What mass (g) of Pyrex glass absorbs 1.95 × 103 J of heat when the temperature changes from 23.0 °C to 55.4°C? (Cs Pyrex glass = 0.750 𝐽𝑔 ∙ °𝐶 ) d. What is the final temperature of a 32.5 g sample of gold initially at 19.7°C upon absorbing 235 J of heat? (Cs gold = 0.750 𝐽𝑔 ∙ °𝐶 )
A) 113 kcal B) -83.0 kJ C) 80.2 g D) +19.7 C
Determine the number of protons and electrons in each ion. a. Al3+ b. Se2- c. Br-
A) 13 P/10 E B) 34 P/ 36 E C) 35 P/ 36 E
Use molar volume of a gas to determine the following quantities. a. Calculate the volume (in L) occupied by 0.785 moles of oxygen at STP. b. Calculate the mass (in g) of 5.25 L of CO2 at STP. c. Calculate the volume (in mL) of 0.336 g of neon at STP
A) 17.6 L B) 10.3 g C) 373 mL
Calculate the designated quantities related to nutritional calories using the following relationships: 4 kcal/g carbohydrate, 4 kcal/ g protein, 9 kcal/g fat. a. How many total Calories (kcal) in a granola bar that contains 5 g fat, 45 g carbohydrates, and 9 g protein? b. How many grams of carbohydrates in a serving of peanut butter containing 16 g fat, 7 g protein, and 190 kcal total?
A) 180 kcal carbs B) 4.5 g
Calculate the molar mass of the following compounds. a. MgBr2 b. Ca(NO3)2
A) 184.11 g B) 164.10 g
Write the full electron configuration for each element. a. C b. P c. Ag d. Rn
A) 1s2 2s2 2p2 B) 1s2 2s2 2p6 3s2 3p3 C) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s14d10 (exception) D) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s24d10 5p6 6s2 4f14 5d10 6p6
How many resonance structures can be drawn for each molecule or ion. a. SO2 b. NO3-
A) 2 B) 3
Write and balance the chemical equation for following reactions that contain an acid or a base. a. HCl(aq) + Mg(s) b. HI(aq) + Ba(OH)2(aq) c. HC2H3O2(aq) + NaHSO3(aq) d. HCl(aq) + K2CO3(aq) e. (NH4)2SO4(aq) + Ca(OH)2(aq)
A) 2 HCl(aq) + Mg(s) → MgCl2(aq) + H2(g) B) 2 HI(aq) + Ba(OH)2(aq) → BaI2(aq) + 2 H2O(l) C) : HC2H3O2(aq) + NaHSO3(aq) → NaC2H3O2(aq) + H2O(l) + SO2( D) 2 HCl(aq) + K2CO3(aq) → 2 KCl(aq) + H2O(l) + CO2(g) E) (NH4)2SO4(aq) + Ca(OH)2(aq) → CaSO4(s) + 2 H2O(l) + 2 NH3(aq)
Write the following chemical reactions based on the type of reaction then balance. a. Combination of sodium solid and chlorine gas b. Single replacement with potassium solid and magnesium sulfate (aq) c. Decomposition of solid aluminum oxide into its elements d. Combustion of liquid ethanol (C2H5OH) e. Metathesis of Lithium sulfide (aq) and ammonium nitrate (aq) Determine the states of the products for each reaction in the question above.
A) 2 Na(s) + Cl2(g) → 2 NaCl(s) B) 2 K(s) + MgSO4(aq) → K2SO4(aq) + Mg(s) C) 2 Al2O3(s) → 4 Al(s) + 3 O2(g) D) C2H5OH(l) + 3 O2(g) → 2 CO2(g) + 3 H2O(l E) Li2S(aq) + 2 NH4NO3(aq) → 2 LiNO3(aq) + (NH4)2S(aq)
Balance the following equations. a. SO2(g) + O2(g) → SO3(g) b. Fe2O3(s) + H2(g) → Fe(s) + H2O(l) c. Na3PO4(aq) + Cu(NO3)2(aq) → NaNO3(aq) + Cu3(PO4)2(s) d. C8H18(l) + O2(g) → CO2(g) + H2O(l) e. KClO3(s) → KCl(s) + O2(g) Identify the type of chemical reaction for each reaction in the question above.
A) 2,1,2/ Combustion B) 1,3,2,3/ Single replacement C) 2,3,6,1/ Double replacement D) 2,25,16,18/ Combustion E) 2,2,3/ Decomposition
Use the Ideal gas law to calculate the designated quantities in each sample. a. A cylinder contains 28.5 L of oxygen gas at a pressure of 1.8 atm and a temperature of 298 K. How much gas (in moles) is in the cylinder? b. What volume does 1.25 g of argon gas at a pressure of 1.05 atm and a temperature of 322 K occupy? c. What is the temperature (°C) of 0.0520 mol of gas at 988 mmHg in a 1180 mL container? d. Pressurized carbon dioxide inflators can be used to inflate a bicycle tire in the event of a flat. These inflators use metal cartridges that contain 16.0 g of carbon dioxide. At 298 K, to what pressure (in mmHg) can the carbon dioxide in the cartridge inflate a 3.45-L mountain bike tire?
A) 2.1 mol B) 787 mL C) 86 C D) 1960 mmHg
Perform each unit conversion. a. 28.9 nm to m b. 1432 cm3 to L c. 121100 µg to mg
A) 2.89×10^-8 m B) 1.432 L C) 121.1 mg
Calculate the mass (g) of each sample. a. 2.3 x 103 mol Sb b. 1.23 mol XeF2
A) 2.8x10^5 g Sb B) 208 g XeF2
Calculate the wavelength of each frequency of electromagnetic radiation: a. 100.2 MHz (typical frequency for FM radio broadcasting) b. 1070 kHz (typical frequency for AM radio broadcasting) (assume four significant figures) c. 835.6 MHz (common frequency used for cell phone communication)
A) 2.994 m B) 280.4 m C) .3590 m
Calculate the designated quantities in the following gas mixture examples. a. A gas mixture with a total pressure of 745 mmHg contains each of these gases at the indicated partial pressures: CO2, 125 mmHg; Ar, 214 mmHg; and O2, 187 mmHg. The mixture also contains helium gas. What is the partial pressure of the helium gas? b. A 275-mL flask contains pure helium at a pressure of 752 torr. A second flask with a volume of 475 mL contains pure argon at a pressure of 722 torr. If the two flasks are connected through a stopcock and the stopcock is opened, what is the partial pressure of each gas and the total pressure? c. Air is approximately 21% oxygen. What volume of air contains 10.0 g of oxygen gas at 273 K and 1.00 atm? d. A heliox deep-sea diving mixture contains 2.0 g of oxygen to every 98.0 g of helium. What is the partial pressure of oxygen when this mixture is delivered at a total pressure of 8.5 atm?
A) 219 mmHg B) 733 torr C) 33 L D) 0.022 atm
Convert 235 mmHg to the following units. a. Torr b. atm c. in Hg
A) 235 Torr B) .309 atm C) 9.25 Hg
Calculate the amount of heat (kJ) absorbed or released for each change. a. Convert 59.2 g of ice at -15.0°C to water at 21.0°C b. Convert 263 g of steam at 145.0°C to water at 67.0°C c. Convert 125 g of steam at 115.0°C to ice at -5.5°C
A) 26.8 kJ B) -652 kJ C) 334 kJ
Determine the atomic orbitals that would overlap to form a covalent bond in each pair of elements. a. C-H b. S-O c. N-Cl d. P-Se
A) 2p of C and 1s of H B) 3p of S and 2p of O C) 2p of N and 3p of Cl D) 3p of P and 4p of Se
Calculate the number of atoms in each sample. a. 5.52 mol of sulfur b. 1.84 x 10-2 mol of silver
A) 3.32x10^24 atoms S B) 1.11x10^22 atoms Ag
Calculate the following quantities based on the chemical reaction: 3 N2H4(g) → 4 NH3(g) + N2(g) a. Moles of NH3 produced by 2.6 moles of N2H4 b. Moles of N2H4 that reacts to form 6.8 moles of N2 c. Moles of NH3 produced from 65.3 g N2H4 d. Grams of N2 produced when 0.874 moles of NH3 are produced e. Grams of N2H4 that reacts to form 4.88 g of NH3 f. Grams of NH3 produced when 27.9 g N2 are produced
A) 3.5 mol B) 20 mol C) 2.72 mol D) 6.02 g E) 6.89 g F) 67.8 g
Calculate the amount of heat (kJ) absorbed or released for each change of state. a. Vaporization of 150. g H2O (ΔHvap = 40.7 kJ/mol): b. Freezing of 8.5 g benzene (C6H6) (ΔHfusion = 9.90 kJ/mol): c. Sublimation of 50.0 g CO2 (ΔHsublimation = 32.3 kJ/mol)
A) 339 kJ B) -1.1 kJ C) 36.7 kJ
Perform each unit conversion. a. 1.4 in to mm b. 116 ft to cm c. 1845 kg to lb d. 815 yd to km
A) 36 mm B) 3540 cm C) 4067 lbs D) .745 km
How many significant figures are in each number? a. 0.1111 s b. 0.007 m c. 108,700 km d. 1.563300 × 1011 m e. 30,800 µL
A) 4 B) 1 C) 4 D) 7 E) 3
Calculate the energy of a photon of electromagnetic radiation at each wavelength: a. 475 nm b. 32.7 cm c. 0.684 km
A) 4.18x10^-19 J B) 6.08x10^-25 J C) 2.91x10^-28 J
Calculate the frequency of each wavelength of electromagnetic radiation: a. 632.8 nm (wavelength of red light from helium-neon laser) b. 0.503 mm (wavelength of maximum solar radiation) c. 0.052 nm (wavelength contained in medical X-rays
A) 4.741x10^14 Hz B) 5.96x10^14 Hz C) 5.8x10^18 Hz
Calculate the atoms (or molecules) in each sample. a. 3.78 g He b. 85.26 g CCl4
A) 5.69x10^23 atoms He B) 3.338x10^23 molecules CCl4
Calculate the number of moles in each sample. a. 3.7 x 1024 atoms of aluminum b. 7.9 x 1031 molecules of SO3
A) 6.1 mol Al B) 1.3x10^8 mol SO3
Calculate the energy of a photon of electromagnetic radiation at each frequency: a. 0.935 GHz b. 4860 kHz c. 7530 MHz
A) 6.20x10^-25 J B) 3.22x10^-27 J C) 4.99x10^-24 J
Calculate to the correct number of significant figures. a. 89.3 × 77.0 × 0.08 b. (5.01 × 105) ÷ (7.8 × 102) c. 4.005 × 74 × 0.007 d. 453 ÷ 2.031
A) 600 B) 640 C) 2 D) 223
Calculate the mass percent composition of nitrogen in each nitrogen-containing compound. a. N2O b. NO c. NO2 d. Al(NO3)3
A) 63.65% B) 46.68% C) 30.45% D) 19.73%
Calculate the pressure of gas in each sample in an open-ended manometer. a. Atmospheric pressure is 751.5 mmHg and the height difference is 2.00 cm being lower on the open side of the manometer. b. Atmospheric pressure is 761.2 mmHg and the height difference is 4.53 cm being higher on the open side of the manometer. c. Atmospheric pressure is 757.4 mmHg and the height difference is 6.21 cm being lower on the open side of the manometer.
A) 731.5 mmHg B) 806.5 mmHg C) 695.3 mmHg
Calculate the designated quantities in the following calorimetry examples. (Cs water = 4.18 𝐽𝑔 ∙ °𝐶 ) a. We submerge a silver block (Cs Ag = 0.230 𝐽𝑔 ∙ °𝐶 ), initially at 58.5 °C, into 100.0 g of water at 24.8 °C, in an insulated container. The final temperature of the mixture upon reaching thermal equilibrium is 26.2 °C. What is the mass of the silver block? b. A 2.74-g sample of a substance suspected of being pure gold is warmed to 72.1 °C and submerged into 15.2 g of water initially at 24.7 °C. The final temperature of the mixture is 26.3 °C. What is the heat capacity of the unknown substance? Could the substance be pure gold if the specific heat of pure gold is 0.129 𝐽𝑔 ∙ °𝐶 ? c. We submerge a 2.85-g lead weight (Cs Pb = 0.130 𝐽𝑔 ∙ °𝐶 ), initially at 10.3 °C, in 7.55 g of water at 52.3 °C in an insulated container. What is the final temperature of both substances at thermal equilibrium? d. When 1.00 g of sodium bicarbonate (s) is combined with enough HCl to make 50.0 mL of solution in a coffee-cup calorimeter, all of the sodium bicarbonate reacts, lowering the temperature of the solutions from 22.5°C to 18.7°C. Assume the density of the solution is 1.00 g/mL and specific heat of the solution is 4.184 𝐽𝑔 ∙ °𝐶 What is the heat of the reaction? Is it an endothermic or exothermic reaction? e. When 100 mL of 0.200 M NaCl(aq) and 100 mL of 0.200 M AgNO3(aq), both at 21.9 °C, are mixed in a coffee cup calorimeter, the temperature increases to 23.5 °C as solid AgCl forms. Assume the density of the solution is 1.00 g/mL and specific heat of the solution is 4.184 𝐽𝑔 ∙ °𝐶. How much heat is produced by this precipitation reaction? Is this an endothermic or exothermic reaction?
A) 78.7 g B) Cs Au= .810 Cannot be pure gold since the specific heat is different. C) 51.8 C D) 795 J (endothermic) E) -1340 J (exothermic)
Calculate the designated quantities in the following examples. a. What is the density (in g/L) of hydrogen gas at 20.0 °C and a pressure of 113 atm? b. A sample of N2O gas has a density of 2.85 g/L at 298 K. What is the pressure of the gas (in mmHg)?
A) 9.44 g/l B) 1.20x10^3 mmHg
Calculate the following concentrations. a. What is the concentration (in ppm) of mercury in a 50.0-g sample of industrial wastewater determined to contain 0.48 mg of mercury? b. What is the concentration (in ppb) of lead in a 1.50 kg bottle of drinking water containing 12 µg of lead?
A) 9.6 ppm B) 8.0 ppb
Determine the number of moles of oxygen atoms in each sample. a. 4.88 mol H2O2 b. 2.15 mol N2O c. 0.0237 mol H2CO3 d. 24.1 mol CO2
A) 9.76 mol B) 2.15 mol C) .0711 mol D) 48.2 mol
Calculate the following quantities using percent as a conversion factor. a. How many milliliters of ethanol are present in a 750-mL bottle of wine that is 12% (v/v) ethanol? b. How many grams of a 17.5% (m/m) I2 solution contains 0.375 g of I2? c. How many grams of NaCl are in 500. mL of saline containing 0.90% (m/v) NaCl?
A) 90 mL B) 2.14 g C) 4.5 g
Calculate the amount of heat absorbed or released for each change of temperature. a. Warming 35.0 g of water from 20.0°C to 85.0°C (Cs water = 4.184 J/g∙°C) b. Cooling 158 g of aluminum from 100.0°C to 15.7°C (Cs Al = 0.215 cal/g∙°C) c. Warming 17.4 g of ice from -12.5°C to 0.0°C (Cs ice = 2.09 J/g∙°C
A) 9510 J B) -2860 cal C) 455 J
Calculate the mass (g) of each sample. a. 15.7 mol HNO3 b. 6.27 × 1020 molecules H2O2
A) 989 g B) .0354 g
List the common name of the group of each element. a. Ca b. Br c. Li d. Ne
A) Alkaline Earth Metal B) Halogen C) Alkali Metal D) Noble Gas
Which statements are inconsistent with Dalton's atomic theory as originally stated? Why? a. All carbon atoms are identical. b. An oxygen atom combines with 1.5 hydrogen atoms to form a water molecule. c. Two oxygen atoms combine with a carbon atom to form a carbon dioxide molecule. d. The formation of a compound often involves the destruction of one or more atoms.
A) Atoms of the same element can have different numbers of neutrons B) Atoms combine in whole number ratios to form compounds C) True D) Atoms cannot be created or destroyed
Write a formula for each molecular compound. a. boron tribromide b. dichlorine monoxide c. xenon tetrafluoride d. phosphorus pentachloride e. carbon tetrabromide f. dinitrogen pentoxide
A) BBr3 B) Cl2O C) XeF4 D) PCl5 E) CBr4 F) N2O5
Calculate the empirical formula for each natural flavor based on its elemental mass percent composition. a. methyl butyrate (component of apple taste and smell): 58.80% C, 9.87% H, and 31.33% O b. vanillin (responsible for the taste and smell of vanilla): 63.15% C, 5.30% H, and 31.55% O c. ascorbic acid (vitamin C): 40.92% C, 4.58% H, and 54.50% O
A) C5H10O2 B) C8H8O3 C) C3H4O3
The empirical formula and molecular formula mass of several compounds are listed. Find the molecular formula of each compound. a. C4H9,114.22 g/mol b. CCl, 284.77 g/mol c. C3H2N, 312.29 g/mol
A) C8H18 B) C6CI6 C) C18H12N6
Determine the empirical formula for the compound represented by each molecular formula. a. C5H10 b. C6H12O6 c. NH3 d. N2O4
A) CH2 B) CH2O C) NH3 D) NO2
Write a formula for the ionic compound that forms between each pair of elements. a. Ca and N b. Mg and I c. Cs and SO42- d. aluminum and nitrate e. silver and bromine f. potassium and phosphate
A) Ca3N2 B) MgI2 C) Cs2SO4 D) Al(NO3)3 E) AgBr F) K3PO4
Which set of quantum numbers cannot occur together to specify an orbital? a. n=2, l=1, ml=−1 b. n=3, l=2, ml=0 c. n=3, l=3, ml=2 d. n=4, l=3, ml=-4
A) Can B) Can C) Cannot- I value D) Cannot- MI value
Classify each change as physical or chemical. a. sugar burns when heated on a skillet b. sugar dissolves in water c. a platinum ring becomes dull due to scratches on the surface d. a silver surface becomes tarnished after exposure to air for a long period of time
A) Chemical B) Physical C) Physical D) Chemical
Choose the element with the higher first ionization energy in each pair. a. Si or Cl b. P or Sb c. Ga or Ge d. P or I
A) Cl B) P C) Ge D) Not enough info
Which compounds contain ionic bonds and which contain covalent bonds? a. CO2 b. CuCl2 c. NaI d. PCl
A) Covalent Bonds B) Ionic Bonds C) Ionic Bonds D) Covalent Bonds
Determine whether each transition in the hydrogen atom corresponds to absorption or emission of energy. a. n=3→n=1 b. n=2→n=4 c. n=4→n=3
A) Emission B) Absorption C) Emission
A flask at room temperature contains exactly equal amounts (in moles) of nitrogen and xenon. a. Which of the two gases exerts the greater partial pressure? b. The molecules or atoms of which gas will have the greater average velocity? c. The molecules of which gas will have the greater average kinetic energy? d. If a small hole were opened in the flask, which gas would effuse more quickly?
A) Equal pressure bc equal moles B) N2 bc smaller mass C) Equal bc at the same temp D) N2 bc smaller mass
Classify the following numbers as measured or exact. a. 12 hats b. 39.8 g c. 350°F d. 1 m = 100 cm
A) Exact (counted) B) Measured C) Measured D) Exact (definition)
Determine if the following changes or reactions are endothermic or exothermic processes. Describe the heat flow in terms of system and surroundings. a. Dissolving of a calcium chloride in water in a hot pack. b. Dissolving of ammonium nitrate in water in a cold pack c. Combustion of acetylene in oxygen in a welding torch. d. Evaporation of sweat from your skin.
A) Exothermic; heat is released from system (reaction) to surroundings (system → surroundings) B) Endothermic; heat is absorbed by the system (reaction) from the surroundings (system ← surroundings) C) Exothermic; heat is released from system (reaction) to surroundings (system → surroundings) D) Endothermic; heat is absorbed by the system (sweat) from the surroundings (skin) (system ← surroundings)
Which solid in each pair has the higher melt point? a. Fe(s) or CCl4(s) b. HCl(s) or KCl(s) c. Ne(s) or Ti(s) d. H2O(s) or H2S(s)
A) Fe(s) > CCl4(s), Fe(s): metallic solid, CCl4(s): molecular solid B) HCl(s) < KCl(s), HCl(s): molecular solid, KCl(s): ionic solid C) Ne(s) < Ti(s), Ne(s): nonbonding atomic solid, Ti(s): metallic solid D) H2O(s) > H2S(s), H2O(s): molecular solid, hydrogen bonding, H2S(s): molecular solid, dipole-dipole
Write the chemical symbol or name of the following transition metal ions. a. Iron (III) ion b. Cobalt (II) ion c. Cu+ d. Zn2+
A) Fe3+ B) Co2+ C) Copper (I) ion D) Zinc ion
Write the symbol for each element and classify it as a metal, nonmetal, or metalloid. a. Germanium b. Fluorine c. Sodium d. Manganese
A) Ge- Metalloid B) F- Nonmetal C) Na- Metal D) Mn- Metal
Predict the number of bonds each element would make. a. C b. Br c. S d. N
A) Group 4A so needs 4 bond B) Group 7A so needs 1 bond C) Group 6A so needs 2 bonds D)Group 5A so needs 3 bonds
Provide the formula for each acid. a. hydrofluoric acid b. phosphoric acid c. hydrobromic acid d. sulfurous acid e. acetic acid f. choric acid
A) HF (aq) B) H3PO4 (aq) C) HBr (aq) D) H2S03 (aq) E) HC2H3O2 (aq) F) HCLO3 (aq)
Write the full orbital diagram for each element. a. He b. S c. Ca d. N
A) He: [↑↓] B) S: [↑↓] [↑↓] [↑↓][↑↓][↑↓] [↑↓] [↑↓][↑ ][↑ C) Ca: [↑↓] [↑↓] [↑↓][↑↓][↑↓] [↑↓] [↑↓][↑↓][↑↓] [↑↓ D)N: [↑↓] [↑↓] [↑ ][↑ ][↑ ]
Name each acid. a. HI(aq) b. HNO3(aq) c. H2CO3(aq) d. HCl(aq) e. HClO2(aq) f. H2SO4(aq)
A) Hydroiodic acid B) Nitric Acid C) Carbonic Acid D) Hydrochloric Acid E) Chlorous Acid F) Sulfuric Acid
Classify the following properties as intensive or extensive. a. Density of iron is 7.86 g/mL b. A sample of ethanol is 15.0 mL c. The boiling point of water is 100°C d. A sample of sodium chloride is 36.45 g
A) Intensive B) Extensive C) Intensive D) Extensive
Classify each solid. a. CaCl2(s) b. CO2(s) c. Ni(s) d. I2(s) e. SiO2(s) f. Ar(s
A) Ionic Solid B) Molecular Solid C) Metallic Solid D) Molecular Solid E) Network Covalent Solid F) Molecular Solid
Use the periodic table to determine the element corresponding to each electron configuration. a. [Ar] 4s2 3d10 4p6 b. [Kr]5s2 4d10 5p2 c. [Xe]6s2 d. [Ar] 4s1 3d5
A) Kr B) Sn C) Ba D) Cr
Name each ionic compound. a. Mg3N2 b. PbI2 c. Ba(NO3)2 d. KBr e. SnCl4 f. Na2O g. CuNO2 h. Ba(OH)2 i. Li2S j. FeCO3 k. Pb(C2H3O2)2 l. NH4I
A) Magnesium Nitride B) Lead (II) iodide C) Barium nitrate D) Potassium bromide E) Tin (IV) chloride F) Sodium oxide G) Copper (I) nitrite H) Barium hydroxide I) Lithium sulfide J) Iron (II) carbonate K) Lead (II) acetate L)Ammonium iodide
Classify each substance as a pure substance or a mixture. If it is a pure substance, classify it as an element or a compound. If it is a mixture, classify it as homogeneous or heterogeneous. a. wine b. beef stew c. iron d. carbon monoxide
A) Mixture/Homogeneous B) Mixture/Heterogeneous C) Pure Substance/Element D)Pure Substance/ Compound
Classify each compound as ionic or molecular. a. CF2Cl2 b. NiBr2 c. CCl4 d. NaO
A) Molecular B) Ionic C) Molecular D) Ionic
Write the chemical symbol or name of the following polyatomic ions. a. Nitrate b. Sulfite c. Ammonium d. CO32- e. PO43- f. NO2-
A) NO3- B) SO32- C) NH4+ D) Carbonate E) Phosphate F) Nitrite
Write the formula for each ionic compound. a. sodium sulfide b. calcium carbonate c. silver nitride d. potassium sulfate e. manganese (II) phophide f. barium sulfite g. copper(I) chloride h. ammonium phosphate i. potassium fluoride j. calcium hydroxide k. iron(II) phosphide l. zinc acetat
A) NaS B) CaCO3 C) Ag3N D) K2SO4 E) Mn3P2 F) BaSO3 G)CuCl H) (NH4)3PO4 I) KF J) Ca(OH)2 K) Fe3P2 L) Zn(C2H3O2)2
Arrange the following solids in order of increasing melting points. a. C(s, diamond) b. Kr(s) c. NaCl(s) d. H2O(s
A) Network Covalent Solid B) Nonbonding Atomic Solid C) Ionic Solid D) Molecular Solid
A chemist decomposes samples of several compounds; the masses of their constituent elements are shown. Calculate the empirical formula for each compound. a. 1.245 g Ni, 5.381 g I b. 2.128 g Be, 7.557 g S, 15.107 g O c. 2.677 g N, 7.643 g O
A) NiI2 B) BeSO4 C) N2O5
Which statements are inconsistent with Rutherford's nuclear theory as originally stated? Why? a. Since electrons are smaller than protons, and since a hydrogen atom contains only one proton and one electron, it must follow that the volume of a hydrogen atom is mostly due to the proton. b. A nitrogen atom has 7 protons in its nucleus and 7 electrons outside of its nucleus. c. A phosphorus atom has 15 protons in its nucleus and 150 electrons outside of its nucleus. d. The majority of the mass of a fluorine atom is due to its 9 electrons.
A) Not volume but mass is due to the proton. The volume is mostly empty space. B) True C) The number of electrons equals the number of protons to make a neutral atom. D) The mass of electrons is very small (negligible). The mass is due to the particles in the nucleus.
Classify each statement as an observation, a law, or a theory. a. Chlorine is a highly reactive gas. b. If elements are listed in order of increasing mass of their atoms, their chemical reactivity follows a repeating pattern. c. Neon is an inert (or nonreactive) gas. d. The reactivity of elements depends on the arrangement of their electrons
A) Observation B)Law C)Observation D)Theory
Choose the more metallic element in each pair. a. Sb or Pb b. K or Ge c. Ge or Sb d. As or Sn
A) Pb B) K C) Not enough info D) Sn
Determine which elements are oxidized and reduced in the following reactions. a. Pb(s) + Br2(l) → PbBr2(s) b. 2 K(s) + MgSO4(aq) → K2SO4(aq) + Mg(s) c. Fe2O3(s) + 3 H2(g) → 2 Fe(s) + 3 H2O(l)
A) Pb is oxidized, Br is reduced B) K is oxidized, Mg is reduced, SO4 did not chang C) Fe is reduced, H is oxidized, O did not change
Classify each property as physical or chemical. a. the boiling point of ethyl alcohol b. Argon is an inert gas c. the tendency of iron to rust d. Gold is yellow in color
A) Physical B) Chemical C) Chemical D) Physical
Determine if a bond between each pair of atoms would be pure covalent, polar covalent, or ionic. a. C and N b. N and S c. K and F d. N and N
A) Polar Covalent B) Polar Covalent C) Ionic D) Nonpolar Covalent
Determine the formula of each compound based on their unit cell. a. Re ions on each corner and oxygen ions on each edge b. Li ions with 8 in the middle and oxygen ions on each corner and each face c. Ag ions with 4 in the middle and iodide ions on each corner and each face
A) Re: 8 corners x 1/8 = 1 Re ions, O: 12 edges x ¼ = 3 O ions, ReO3 B) Li: 8 middles = 8 Li ions, O: 8 corners x 1/8 + 6 faces x ½ = 4 O ions, Li2O C) Ag: 4 middles = 4 Ag ions, I: 8 corners x 1/8 + 6 faces x ½ = 4 I ions, AgI
Which combinations of n and l represent real orbitals, and which do not exist? a. 1s b. 2p c. 4s d. 2d
A) Real B) Real C) Real D) Not Real
Choose the element with the more negative (more exothermic) electron affinity in each pair. a. Mg or S b. K or Cs c. Si or P d. Ga or Br
A) S B) K C) Si D) Ga
Write the chemical symbol for the ion formed by each atom. a. S b. Li c. Br d. P e. Al f. Ca
A) S2- B) Li+ C) Br- D) P3- E) Al3+ F) Ca2+
Choose the larger atom in each pair. a. Sn or Si b. Br or Ga c. Sn or Bi d. Se or Sn
A) Sn B) Ga C) Not enough info D) Sn
Which is the larger species in each pair? a. Sr or Sr2+ b. N or N3- c. Ni or Ni2+ d. S2- or Ca2+
A) Sr B) N3- C) Ni D) S2-
Name each molecular compound. a. SO3 b. S2F4 c. CO2 d. BrF5 e. NO f. N4Se4
A) Sulfur trioxide B) Disulfur tetrafluoride C) Carbon dioxide D) Bromine pentafluoride E) Nitrogen monoxide F) Tetranitrogen tetraselenide
Which statements about subatomic particles are false? a. Proton and electrons have charges of the same magnitude but opposite sign. b. Protons have about the same mass as neutrons. c. Some atoms don't have any protons. d. Protons and neutrons have charges of the same magnitude but opposite signs.
A) True B) True C)False, all atoms have one or more protons. D) False, neutrons are neutral.
Answer the following questions regarding real gases include explanations. a. Which postulate of the kinetic molecular theory breaks down under conditions of high pressure? b. Which postulate of the kinetic molecular theory breaks down under conditions of low temperature?
A) Under high pressure, the volume of the atoms or molecules of gas is not negligible b/c the molecules are squeezed close together with very little empty space in between. B) At low temperature, the molecules of gas are moving slowly with less energy so they tend to "stick" together when they collide due to intermolecular forces (IMF).
Determine if the following pairs of solute and solvent will form a solution (homogeneous mixture). a. NH3 and water b. NaCl and C6H14 c. I2 and CCl4 d. SO3 and CH3OH
A) Yes, both are polar B) No, NaCl is ionic/polar and C6H14 is nonpolar C) Yes, both are nonpolar D) No, SO3 is nonpolar and CH3OH is polar
Write the abbrievated electron configuration for each element. a. O b. Si c. Cd d. Po
A) [He] 2s2 2p4 B) [Ne] 3s2 3p2 C) [Kr] 5s2 4s10 D) [Xe] 6s2 4f14 5d10 6p4
Write the electron configuration for each ion. a. Cl- b. K+ c. Se2- d. Mo3+
A) [Ne] 3s2 3p6 B) [Ne] 3s2 3p6 C) [Ar] 4s2 3d10 4p6 D) [Kr] 5s0 4d3
Name the following changes of state and determine if an endothermic or exothermic process. a. Gas → liquid b. Solid → gas c. Liquid → solid d. Gas → solid e. Liquid → gas f. Solid → liquid
A) condensation, exothermic B) sublimation, endothermic C) freezing, exothermic D) deposition, exothermic E) vaporization, endothermic F) melting, endothermic
Is each process exothermic or endothermic? Indicate the sign of ΔH. a. Dry ice evaporating b. A sparkler burning c. Reaction in a chemical cold pack
A) endothermic, +ΔH B) exothermic, - ΔH C) endothermic, + ΔH
Determine if the following compounds are soluble or insoluble in water. a. FeCO3 b. K3PO4 c. AgI d. Cu2SO4
A) insoluble, (s) B) soluble, (aq) C) insoluble, (s) D) soluble, (aq)
Identify the type of element (metal, nonmetal, metalloid), amount of band gap expected (large, small, no gap), and type of conductor (good, semi, poor). a. Zn b. As c. Bi d. P
A) metal, no gap, good conductor B) metalloid, small gap, semiconductor C) metal, no gap, good conductor D) nonmetal, large gap, poor conductor
Determine the type of IMF that are present in each element, compound, or mixture. a. Kr b. NCl3 c. SiH4 d. HF e. NaCl f. CH2F2 g. Li2S - water
A) nonpolar, dispersion B) polar, dispersion and dipole-dipole C) nonpolar, dispersion D) polar, dispersion and hydrogen bonding E) ionic, ionic bonding F) polar, dispersion and dipole-dipole G) ionic compound with polar compound, ion-dipole
Determine i) the type of unit cell for each sample, ii) the number of atoms in the unit cell, and iii) the coordination number. a. Po has an atom at each corner: b. W has an atom at each corner and one in the center: c. Ni has an atom at each corner and each face
A) simple cube, 8 corners x 1/8 = 1 atom, 6 B) body centered cube, 8 corners x 1/8 + 1 center = 2 atoms, 8 C) face centered cube, 8 corners x 1/8 + 6 faces x ½ = 4 atoms, 12
Use Standard Enthalpy of Formation (ΔH°f) to calculate enthalpy of reaction (ΔHrxn) for each reaction. a. 2 H2S(g) + 3 O2(g) → 2 H2O(l) + 2 SO2(g) *refer to appendix in the back of the book for ΔH°f values b. 2 SO2(g) + O2(g) → 2 SO3(g) c. 6 CO2(g) + 6 H2O(g) → C6H12O6(s) + 6 O2(g) d. N2O4(g) + 4 H2(g) → N2(g) + 4 H2O(g
A) ΔHrxn = -1124 kJ (exothermic) B) ΔHrxn = -197.8 kJ (exothermic) C) ΔHrxn = 2538.8 kJ (endothermic) D) ΔHrxn = -936.36 kJ (exothermic)
Arrange these compounds in order of increasing strength of IMF. a. CH3CH2Cl b. CH4 c. CH3CH2CH2OH d. CH3CH3 e. CH3CH2OH
CH4 < CH3CH3 < CH3CH2Cl < CH3CH2OH < CH3CH2CH2O A) polar, dipole-dipole B) nonpolar, dispersion, 16.04 g/mol C) polar hydrogen bonding, 60.09 g/mol D) nonpolar, dispersion, 30.07 g/mol E) polar, hydrogen bonding, 46.07 g/mol
Arrange these compounds in order of increasing viscosity: a. PH3 b. HBr c. CH3OH d. Cl2
Cl2 < PH3 < HBr < CH3OH A) polar, dipole-dipole, 33.99 g/mol B) polar, dipole-dipole, 80.91 g/mol C) polar, hydrogen bonding D) nonpolar, dispersion
A sample of N2O effuses from a container in 42 seconds. How long will it take the same amount of gaseous I2 to effuse from the same container under identical conditions?
Graham's Law: 𝒓𝒂𝒕𝒆 𝑨𝒓𝒂𝒕𝒆 𝑩=√𝑴𝑩𝑴𝑨 Rate ∝ 𝟏𝒕𝒊𝒎𝒆 so 𝒕𝒊𝒎𝒆 𝑨𝒕𝒊𝒎𝒆 𝑩=√𝑴𝑨𝑴𝑩 𝑨=𝑰𝟐; 𝑩=𝑵𝟐𝑶 𝒕𝒊𝒎𝒆 𝑰𝟐𝟒𝟐𝒔=√𝟐𝟓𝟑.𝟖𝒈/𝒎𝒐𝒍𝟒𝟒.𝟎𝟐𝒈/𝒎𝒐𝒍=𝟐.𝟒𝟎𝟏𝟏𝟓𝟖 𝒕𝒊𝒎𝒆𝑰𝟐=𝟐.𝟒𝟎𝟏𝟏𝟓𝟖×𝟒𝟐𝒔=𝟏𝟎𝟎.𝟖 𝒔=𝟏.𝟎×𝟏𝟎𝟐 𝒔 Since 𝑰𝟐 is larger it should take longer to effuse.
Arrange these compounds in order of increasing boiling point: a. H2S b. H2Se c. H2O
H2S < H2Se < H2O because strongest IMF has the highest BP A) H2S: polar, dipole-dipole, 34.09 g/mol B) H2Se: polar, dipole-dipole, 80.98 g/mol C) H2O: polar, hydrogen bonding
Put the following elements in order of increasing electronegativity: Mg, F, S, K
K<Mg<S<F
Calculate the root mean square velocity and kinetic energy of CO, CO2, and SO3 at 298 K. Which gas has the greatest velocity? The greatest kinetic energy? The greatest effusion rate?
KE is the same for all of them because it only depends on temperature: KE = 3/2 RT = (3/2) (8.314J/mol*K) (298K) = 3.72 x 103 J Vrms is different for each gas because it depends on both temperature and molar mass: *be sure to use molar mass in kilograms in this equation. 𝑪𝑶: 𝑽𝒓𝒎𝒔=√𝟑𝑹𝑻𝑴=√𝟑(𝟖.𝟑𝟏𝟒𝑱/𝒎𝒐𝒍∙𝑲)(𝟐𝟗𝟖𝑲)(𝟐𝟖.𝟎𝟏𝒈/𝒎𝒐𝒍)(𝟏 𝒌𝒈𝟏𝟎𝟎𝟎𝒈)=𝟓𝟏𝟓 𝒎/𝒔 𝑪𝑶𝟐: 𝑽𝒓𝒎𝒔=√𝟑𝑹𝑻𝑴=√𝟑(𝟖.𝟑𝟏𝟒𝑱/𝒎𝒐𝒍∙𝑲)(𝟐𝟗𝟖𝑲)(𝟒𝟒.𝟎𝟏𝒈/𝒎𝒐𝒍)(𝟏 𝒌𝒈𝟏𝟎𝟎𝟎𝒈)=𝟒𝟏𝟏 𝒎/𝒔 𝑺𝑶𝟑: 𝑽𝒓𝒎𝒔=√𝟑𝑹𝑻𝑴=√𝟑(𝟖.𝟑𝟏𝟒𝑱/𝒎𝒐𝒍∙𝑲)(𝟐𝟗𝟖𝑲)(𝟖𝟎.𝟎𝟕𝒈/𝒎𝒐𝒍)(𝟏 𝒌𝒈𝟏𝟎𝟎𝟎𝒈)=𝟑𝟎𝟓 𝒎/𝒔
Order the following states in order of i) increasing kinetic energy and ii) increasing strength of IMF. a. Solid b. Liquid c. Gas
Kinetic Energy) solid < liquid < gas IMF) gas < liquid < solid
Two samples of sodium chloride are decomposed into their elements. Once sample produces 6.98 g of sodium and 10.7 g of chlorine, and the other sample produces 11.2 g sodium and 17.3 g of chlorine. Are these results consistent with the law of definite proportions or the law of multiple proportions?
Law of Definite Proportions
Two samples of a compound containing sulfur and fluorine are decomposed into their elements. Once sample produces 4.45 g of fluorine and 1.25 g of sulfur, and the other sample produces 4.43 g fluorine and 1.87 g of sulfur. Are these results consistent with the law of definite proportions or the law of multiple proportions?
Law of Multiple Proportions
Determine the limiting reactant, theoretical yield of product (in grams) and percent yield when 0.233 g Ti reacts with 0.288 g F2 producing 0.376 g TiF4 according to the equation: Ti(s) + 2 F2(g) → TiF4(s)
Limiting Reactant) .00487 mol Theoretical Yield) .469 g Percent Yield) 80.2% TiF4
Determine the limiting reactant, theoretical yield, and percent yield for the reaction when 10.1 g of Mg reacts with 10.5 g O2 forming 11.9 g MgO. The balanced equation for the reaction is: 2 Mg(s) + O2(g) → 2 MgO(s
Limiting Reactant) .415 mol Theoretical Yield) 16.7 g Percent Yield) 71.3% MgO
Order the following ionic compounds by increasing lattice energy: KCl, SrO, RbBr, CaO.
RbBr < KCl < SrO < CaO *Lattice energy increases with greatest charges and smallest size
Arrange these compounds in order of increasing surface tension: a. CH3CH2CH2CH3 b. CH3CH2CH2CH2CH3 c. SiH4 d. CO2
SiH4 < CO2 < CH3CH2CH2CH3 < CH3CH2CH2CH2CH A) nonpolar, dispersion, 58.12 g/mol B) nonpolar, dispersion, 72.15 g/mol C) nonpolar, dispersion, 32.03 g/mol D)nonpolar, dispersion, 44.01 g/mol
Which element would have characteristics most similar to Mg? a. Na b. Al c. Sr d. Ar
Sr
Arrange this isoelectronic series in order of increasing atomic radius: Se2−,Sr2+,Rb+,Br−.
Sr2+ < Rb+ < Br- < Se2-
Pick the compound with the highest boiling point in each pair. a. NH3 or CH4 b. HCl or NaCl c. CS2 or CO2 d. CO2 or NO2
Strongest IMF has highest BP A) NH3 > CH4, NH3: polar, hydrogen bonding, CH4: nonpolar, dispersio B) HCl < NaCl, HCl: polar, dipole-dipole, NaCl: ionic, ionic bonding C) CS2 > CO2, CS2: nonpolar, dispersion, 76.15 g/mol, CO2: nonpolar, dispersion, 44.01 g/mol D) CO2 < NO2: CO2: nonpolar, dispersion, NO2: polar, dipole-dipole
Pick the compound with the highest vapor pressure in each pair. a. Br2 or I2 b. H2S or H2O c. PF3 or PH3
Weakest IMF has highest vapor pressure A) Br2 > I2, Br2: nonpolar, dispersion, 159.80 g/mol, I2: nonpolar, dispersion, 253.80 g/mol B) H2S > H2O, H2S: polar, dipole-dipole, H2O: polar, hydrogen bonding C) PF3 < PH3, PF3: polar, dipole-dipole, 87.97 g/mol, PH3: polar, dipole-dipole, 33.99 g/mo
Arrange these substances in order of increasing i) surface tension, ii) viscosity, and iii) vapor pressure. a. CH3CH2OH at 20°C b. CH3CH2OH at 40°C c. CH3CH2OH at 60°C
i) 60°C < 40°C < 20°C because ST increases with decreasing temp ii) 60°C < 40°C < 20°C because viscosity increases with decreasing temp iii) 20°C < 40°C < 20°C because VP increases with increasing tem
List these types of electromagnetic radiation from smallest to greatest (i) wavelength, (ii) frequency, and (iii) energy per photon: a. gamma rays b. radio waves c. microwaves d. visible light
i) wavelength: gamma < visible < microwaves < radio ii) frequency: radio < microwaves < visible < gamma iii) energy: radio < microwaves < visible < gamma