CHM2045C Chemistry I Final Exam
Identify the color of a flame test for potassium.
violet
How many of the following compounds are soluble in water? Cu(OH)2 NaNO3 NH4Cl Li2S
3
How many of the following molecules are polar? PCl₅ COS XeO₃ SeBr₂
3
How many atoms of oxygen are in 2.50 moles of CO2?
3.01 X 10²⁴ atoms
Which of the following (with specific heat capacity provided) would show the smallest temperature change upon gaining 200.0 J of heat?
50.0 g Al, CAl = 0.903 J/g°C
Automotive air bags inflate when sodium azide decomposes explosively to its constituent elements: 2NaN3(s) → 2Na(s) + 3N2(g) How many grams of sodium azide are required to produce 33.0 g of nitrogen?
51.1
If the temperature is 128°F, what is the temperature in degrees celsius?
53.3°C (°C x 1.8) + 32 = °F °C + 273.15 = °F
All of the following statements about the quantum numbers are true EXCEPT A) ml may take integral values of +l to -l, including zero B) n may take integral values from 1 to infinity C) l may take integral values from 1 to n-1 D) ms may take only the values of +¹/₂ and -¹/₂. E) ml has -l to +l values or 2l + 1 possible values (HINT: verify this by plugging in l values to see if it is true)
C) l may take integral values from 1 to n-1
Which of the following sets of quantum numbers (n, l, ml, ms) refers to a 3d orbital? A) n = 4, l = 2, ml = 1, ms = -¹/₂ B) n = 2, l = 0, ml = 0, ms = -¹/₂ C) n = 5, l = 4, ml = 3, ms = +¹/₂ D) n = 3, l = 2, ml = 2, ms = +¹/₂ E) n = 4, l = 3, ml = 2, ms = +¹/₂
D) n = 3, l = 2, ml = 2, ms = +¹/₂
Give the electron geometry, molecular geometry, and hybridization for CH₃⁻.
eg = tetrahedral; mg = trigonal pyramidal; sp³
Determine the electron geometry (eg) and molecular geometry (mg) of NCl₃.
eg=tetrahedral, mg=trigonal pyramidal
spin quantum number (ms)
electron spin
On the electromagnetic spectrum, visible light is immediately between two other wavelengths. Name them.
infrared and ultraviolet
The Scientific Method
is based on continued observation and experiment.
Write the name for PbS.
lead (II) sulfide
For a particular process that is carried out at constant pressure, q = 145 kJ and w = -35 kJ. Therefore,
∆E = 110 kJ and ΔH = 145 kJ.
Determine the density of CO2 gas at STP.
1.96 g/L
Identify a substance that is not in its standard state. A) H2 B) Ba C) Ne D) O2 E) CO
E) CO
Determine the oxidation state of P in PO₃³⁻.
+3
For ΔEsys to always be -, what must be true?
-w > +q
Which element has the ground-state electron coniguration [ Xe] 6s2 4f7
Eu
Determine the number of protons, neutrons and electrons in the following: 65 X 29
p+ = 29 n° = 36 e- = 29
Define vapor pressure.
partial pressure of water in a gaseous mixture
Identify the charges of the protons, neutrons, and electrons.
protons +1, neutrons 0, electrons -1
Which of the following signs on q and w represent a system that is doing work on the surroundings, as well as gaining heat from the surroundings?
q = +, w = -
Which of the following visible colors of light has the longest wavelength?
red
principal quantum number (n) **schodinger**
represents the relative overall energy of each orbital
What are the possible orbitals for n = 3?
s, p, d
Draw the valence orbital diagram that represents the ground state of Se²⁻.
↑↓ ↑↓ ↑↓ ↑↓ ⁻⁻⁻⁻ ⁻⁻⁻⁻ ⁻⁻⁻⁻ ⁻⁻⁻⁻ 4s 4p
A 35.6 g sample of ethanol (C2H5OH) is burned in a bomb calorimeter, according to the following reaction. If the temperature rose from 35.0 to 76.0°C and the heat capacity of the calorimeter is 23.3 kJ/°C, what is the value of H°rxn? The molar mass of ethanol is 46.07 g/mol. C2H5OH(l) + 3 O2(g) → 2 CO2(g) + 3 H2O(g) ΔH°rxn = ?
-1.24 × 103 kJ/mol
Use the standard reaction enthalpies given below to determine ΔH°rxn for the following reaction: P4(g) + 10 Cl2(g) → 4PCl5(s) ΔH°rxn = ? Given: PCl5(s) → PCl3(g) + Cl2(g) ΔH°rxn = +157 kJ P4(g) + 6 Cl2(g) → 4 PCl3(g) ΔH°rxn = -1207 kJ
-1835 kJ
Calculate the energy change associated with the transition from n=4 to n=1 in the hydrogen atom.
-2.04 × 10-18 J
Calculate the change in internal energy (ΔE) for a system that is giving off 25.0 kJ of heat and is changing from 12.00 L to 6.00 L in volume at 1.50 atm pressure. (Remember that 101.3 J = 1 L∙atm)
-24.1 kJ
Use the bond energies provided to estimate ΔH° rxn for the reaction below. XeF₂ + 2 F₂ → XeF₆ ΔH°rxn = ? Bond Bond Energy (kJ/mol) Xe-F 147 F-F 159
-270 kJ
Use the bond energies provided to estimate ∆H°rxn for the reaction below XeF₂ + 2F₂ → XeF₆ ∆H°rxn=? Xe-F 147 kJ/mol F-F 159 kJ/mol
-270 kJ
Use the ΔH°f and ΔH°rxn information provided to calculate ΔH°f for SO3(g): ΔH°f (kJ/mol) 2 SO2(g) + O2(g) → 2 SO3(g) ΔH°rxn = -198 kJ SO2(g) -297
-396 kJ/mol
Use the information provided to determine ΔH°rxn for the following reaction: ΔH°f (kJ/mol) 3 Fe2O3(s) + CO(g) → 2 Fe3O4(s) + CO2(g) ΔH°rxn = ? Fe2O3(s) -824 Fe3O4(s) -1118 CO(g) -111 CO2(g) -394
-47 kJ
Use the ΔH°f information provided to calculate ΔH°rxn for the following: ΔH°f (kJ/mol) SO2Cl2(g) + 2 H2O(l) → 2 HCl(g) + H2SO4(l) ΔH°rxn = ? SO2Cl2(g) -364 H2O(l) -286 HCl(g) -92 H2SO4(l) -814
-62 kJ
Calculate the change internal energy (ΔE) for a system that is giving off 65.0 kJ of heat and is performing 855 J of work on the surroundings.
-65.9 kJ
Trigonal Bipyramidal (electron geometry)
-Five electron groups surround the central atom -molecular geometry **(no lone pairs) equatorial positions is 120° equatorial positions is 120° Trigonal Bipyramidal -Hybridization: sp³d
tetrahedral (electron geometry)
-Four electron groups surround the central atom -molecular geometry **(no lone pairs) 109.5° tetrahedral **(1 lone pair) 107° trigonal pyramid **(2 lone pairs) 104.5° Bent -Hybridization: sp³
Octahedral (electron geometry)
-Six electron groups surround the central atom -molecular geometry **(no lone pairs) 90° Octahedral -Hybridization: sp³d²
Trigonal Planar (electron geometry)
-Three electron groups surround the central atom -molecular geometry **(no lone pairs) 120° Trigonal Planar **(1 lone pair) <120° Bent -Hybridization: sp²
Linear (electron geometry)
-Two electron groups surround the central atom -molecular geometry **(no lone pairs) 180° linear -Hybridization: sp
Draw the best Lewis structure for BrO4⁻ and determine the formal charge on bromine.
0
In the best Lewis structure for NO+, what is the formal charge of the N atom?
0
In the best Lewis structure for NO+, what is the formal charge on the N atom? A) -1 B) +2 C) 0 D) +1
0
For n = 3, what are the possible sublevels?
0, 1, 2
A gas bottle contains 0. 250 mol of gas at 730 mm Hg pressure. If the final pressure is 1.15 atm, how many moles of gas were added to the bottle?
0. 0493 mol
How many grams of KCl(s) are produced from the thermal decomposition of KClO3(s) which produces 50.0 mL of O2(g) at 25°C and 1.00 atm pressure according to the chemical equation shown below? 2 KClO3(s) → 2 KCl(s) + 3 O2(g)?
0. 102 g
A sample of oxygen is collected over water at a total pressure of 644.3 mmHg at 10°C. The vapor pressure of water at 10°C is 9.2 mmHg. the partial pressure of the O2 is
0.8357 atm
Give the temperature and pressure at STP.
0°C and 1.00 atm
Draw the best Lewis structure for CH3 +1 . What is the formal charge on the C?
1
How many of the following molecules have sp² hybridization on the central atom? HCN SO₂ OCl₂ XeCl₂
1
What answer should be reported, with the correct number of significant figures, for the following calculation? (433.621 - 333.9) × 11.900
1.19 × 10³
How many molecules of sucrose (C12H22O11, molar mass = 342.30 g/mol) are contained in 14.3 mL of 0.140 M sucrose solution?
1.21 × 1021 molecules C12H22O11
How much energy is required to decompose 612 g of PCl3, according to the reaction below? The molar mass of PCl3 is 137.32 g/mol and may be useful. 4 PCl3(g) → P4(s) + 6 Cl2(g) ΔH°rxn = +1207 kJ
1.34 × 103 kJ
The density of nitric oxide (NO) gas at 1.21 atm and 54.1°C is ________ g/L.
1.35
Calculate the atomic mass of silver if silver has 2 naturally occurring isotopes with the following masses and natural abundances: Ag-107 106.90509 amu 51.84% Ag-109 108.90476 amu 48.46%
108.19 amu
Calculate the atomic mass of silver if silver has 2 naturally occurring isotopes with the following masses and natural abundances: Ag-107 106.90509 amu 51.84% Ag-109 108.90476 amu 48.46%
108.19 amu
Consider the following compound. How many sigma and pi bonds does it contain? CH3CHCHCO2H
11 sigma, 2 pi
Convert 1.50 atm to mm Hg.
1140 mm Hg
Sodium metal and water react to form hydrogen and sodium hydroxide. If 5.98 g of sodium react with water to form 0.26 g of hydrogen and 10.40 g of sodium hydroxide, what mass of water was involved in the reaction?
4.68 g
When 14.0 g of zinc metal reacts with excess HCl, how many liters of H2 gas are produced at STP?
4.80 L
Determine the velocity of a medicine ball (m = 10.0 kg) with a wavelength of 1.33 × 10-35 m.
4.98 m/s
How many photons are contained in a flash of green light (525 nm) that contains 189 kJ of energy?
4.99 × 10²³ photons
Determine the oxidizing agent in the following reaction. Ni(s) + 2 AgClO4(aq) → Ni(ClO4)2(aq) + 2 Ag(s)
Ag
Which of the following contains BOTH ionic and covalent bonds?
BaSO₄
Which of the following does NOT describe a metal? A) found on the left side of the periodic table. B) good conductor of heat C) tends to gain electrons D) good conductor of electricity E) forms ionic compounds with nonmetals
C
Consider the molecule below. Determine the hybridization at each of the 2 labeled carbons. H Cl H ↑ ↑ ↑ C⇌C(¹) H←C(²)→H ↓ ↓ ↑ H C ⇌ C ↓ ↓ H F
C(¹)= sp², C(²) = sp³
Which of the following statements is TRUE? A) It is possible for two electrons in the same atom to have identical values for all four quantum numbers. B) An orbital that penetrates into the region occupied by core electrons is more shielded from nuclear charge than an orbital that does not penetrate and therefore has a lower energy. C) An orbital that penetrates into the region occupied by core electrons is less shielded from nuclear charge than an orbital that does not penetrate and therefore has a lower energy. D) Two electrons in the same orbital can have the same spin. E) None of the above are true.
C) An orbital that penetrates into the region occupied by core electrons is less shielded from nuclear charge than an orbital that does not penetrate and therefore has a lower energy.
Of the following, which element has the highest irst ionization energy? A) Sr B) K C) Br D) Te
C) Br
In which of the following sets do all species have the same number of electrons? A) Br, Br-, Br+ B) K+, Rb+, Cs+ C) F-, Ne, Mg2+ D) Ge, Se2-, Br- E) none of the above
C) F-, Ne, Mg2+
Which ion does not have a noble gas coniguration in its ground state? A) As3- B) Sc3+ C) Ga3+ D) Al3+
C) Ga3+
Which of the following is a gas-evolution reaction? A) 2 H2(g) + O2(g) → 2 H2O(g) B) LiCl(aq) + NaNO3(aq) → LiNO3(aq) + NaCl(g) C) NH4Cl(aq) + KOH(aq) → KCl(aq) + NH3(g) + H2O(l) D) 2 C2H6(l) + 7 O2(g) → 4 CO2(g) + 6 H2O(g) E) None of the above are gas-evolution reactions.
C) NH4Cl(aq) + KOH(aq) → KCl(aq) + NH3(g) + H2O(l)
Choose the paramagnetic species from below A) Cd B) Ca C) Nb³⁺ D) O²⁻ E) Zn²⁺
C) Nb³⁺
All of the following are postulates of the kinetic molecular theory of gases EXCEPT A) The volumes of the molecules are negligible compared with the volume of the container. B) The collisions between molecules are elastic. C) The gas molecules are in rapid motion D) At a constant temperature, each molecule has the same kinetic energy. E) The gas molecules are in constant motion.
C) The gas molecules are in rapid motion
Which of the following statements is FALSE? A) Nonmetals tend to gain electrons. B) Anions are usually larger than their corresponding atom. C) The halogens tend to form 1+ ions. D) Metals tend to form cations. E) Atoms are usually larger than their corresponding cation.
C) The halogens tend to form 1+ ions.
What is the compound ethanol?
C2H5OH
Place the following elements in order of increasing atomic radius P Ba Cl
Cl < P < Ba
Place the following elements in order of increasing atomic radius. P Ba Cl
Cl < P < Ba
A 0.334 g sample of an unknown halogen occupies 109 mL at 398 K and 1.41 atm. What is the identity of the halogen?
Cl₂
The element that corresponds to the electron configuration 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d⁵ is
Cr
Combustion analysis of an unknown compound containing only carbon and hydrogen produced 0.2845 g of CO2 and 0.1451g of H2O. What is the empirical formula of the compound?
C₂H₅
Combustion analysis of 1.200 g of an unknown compound containing carbon, hydrogen, and oxygen produced 2.086 g of CO2 and 1.134 g of H2O. What is the empirical formula of the compound?
C₃H⁸O₂
What is the empirical formula of a compound that is 62.0% C, 10.4% H, and 27.5% O by mass?
C₃H₆O
Determine the molecular formula of a compound that has a molar mass of 183.2 g/mol and an empirical formula of C₂H₅O₂.
C₆H₁₅O₆
Which substance has a standard enthalpy of formation equal to zero at 25°C? A) Br2 (g) B) C2H6 (g) C) C2H6 (l) D) Br2 (l) E) Br2 (s)
D
Kinetic Energy
Due to motion
Which molecule or compound below contains an ionic bond? A) CO₂ B) SiF₄ C) OCl₂ D) C₂Br₄ E) NH₄NO₃
E
Which statement is FALSE? A) Enthalpy is the sum of a system's internal energy and the product of pressure and volume. B) An exothermic reaction gives heat off heat to the surroundings. C) Endothermic has a positive ΔH. D) ΔHrxn is the heat of reaction. E) ΔErxn is a measure of heat
E
Choose the thermochemical equation that illustrates ΔH°f for Li2SO4. A) Li2SO4(aq) → 2 Li+(aq) + SO42-(aq) B) 2 Li+(aq) + SO42-(aq) → Li2SO4(aq) C) 8 Li2SO4(s) → 16 Li(s) + S8(s, rhombic) + 16 O2(g) D) 16 Li(s) + S8(s, rhombic) + 16 O2(g) → 8 Li2SO4(s) E) 2 Li(s) + 1/8 S8(s, rhombic) + 2 O2(g) → Li2SO4(s)
E) 2 Li(s) + 1/8 S8(s, rhombic) + 2 O2(g) → Li2SO4(s)
Which of the following statements is TRUE? A) It is not possible for two atoms to share more than two electrons. B) Single bonds are shorter than double bonds. C) A pair of electrons involved in a covalent bond are sometimes referred to as "lone pairs." D) A covalent bond is formed through the transfer of electrons from one atom to another.
E) A covalent bond has a lower potential energy than the two separate atoms.
Which of the following statements is TRUE? A) Cars that run on hydrogen fuel cells are environmentally friendly. B) The more energy produced per kg of CO2 produced, the better the fuel. C) Acid rain is one of the problems associated with the combustion of fossil fuels. D) The burning of fossil fuels contributes to global warming. E) All of the above are true.
E) All of the above are true.
Which of the following statements is TRUE? A) ΔErxn can be determined using constant volume calorimetry. B) Energy is neither created nor destroyed, excluding nuclear reactions. C) ΔHrxn can be determined using constant pressure calorimetry. D) State functions do not depend on the path taken to arrive at a particular state. E) All of the above are true.
E) All of the above are true.
The chemical formula for lithium peroxide is
Li2O2.
Aufbau principle
Only two electrons, with opposing spins, are allowed in each orbital
van der waals equation
The constant 'a' provides a correction for the intermolecular forces. Constant b is a correction for finite molecular size and its value is the volume of one mole of the atoms or molecules.
Choose the best Lewis structure for ICl₅
all single bonds and 1 unpaired e⁻ group
ate → ? ite → ?
ate → ic ite → ous
isotope
atoms that have the same number of protons but have different numbers of neutrons
Using the VSEPR model, the molecular geometry of the central atom in SO₂ is ________.
bent
Determine the name for H2CO3.
carbonic acid
angular momentum quantum number (l)
describes the shape of an orbital
magnetic quantum number (ml)
distinguishes the orbitals available within a subshell
What type of chemical reaction is listed below? Al(C₂H₃O₂)₃(aq) + (NH₄)₃PO₄(aq) → AlPO₄(s) + 3 NH₄C₂H₃O₂(aq)
double replacement
Potential Energy
due to position or composition
Describe the shape of a p orbital.
dumbbell shaped
Determine the electron geometry (eg) and molecular geometry (mg) of NCl₃⁻.
eg=tetrahedral, mg=trigonal pyramidal
Determine the electron geometry (eg), molecular geometry (mg), and polarity of SO2
eg=trigonal planar, mg=bent, polar
Systematic error is defined as
error that tends to be too high or too low.
Hund's Rule
fill the electons singly first then fill with parallel spins is known ↑↓
Place the following types of electromagnetic radiation in order of increasing wavelength. visible light gamma rays microwaves
gamma rays < visible light < microwaves
Give the hybridization for the C in HCCH.
sp
Gas is sold for $1.399 per liter in Toronto, Canada. Your car needs 12.00 gallons. How much will your credit card be charged in Canadian dollars?
$63.54
values of internal energy (∆E)
(+) energy is absorbed by system (-) energy is released by system
values of heat (q)
(+) heat is absorbed by system (-) heat is released by system
values of work (w)
(+) work is done unto the system (-) work is done by the system
Dalton's atomic theory consisted of all the following postulates
*Atoms of different elements have different properties *Elements are composed of indivisible particle called atoms *Atom combine in fixed ratios of whole numbers *in chemical changes, atoms are not destroyed, created, or changed.
Draw the best Lewis structure for the free radical, NO2. What is the formal charge on the N?
+1
Determine the oxidation state of nitrogen in NO2.
+4
What is the oxidation number of the chromium atom in K2CrO4 ?
+6
Calculate ∆E of a gas for a process in which the gas absorbs 18 J of heat and does 19 J of work by expanding
-1 J
Give the numbers for ml for a d orbital.
-2, -1, 0, 1, 2
Two solutions, initially at 24.60°C, are mixed in a coffee cup calorimeter (Ccal = 15.5 J/°C). When a 100.0 mL volume of 0.100 M AgNO3 solution is mixed with a 100.0 mL sample of 0.200 M NaCl solution, the temperature in the calorimeter rises to 25.30°C. Determine the H°rxn for the reaction as written below. Assume that the density and heat capacity of the solutions is the same as that of water. NaCl(aq) + AgNO3(aq) → AgCl(s) + NaNO3(aq) H°rxn = ?
-59.7 kJ
What is the oxidation number change for the bromine atom in the following unbalanced reduction half reaction: BrO₃₋(aq) + H⁺(aq) → Br⁻(aq)+ H20 (l)
-6
A student prepared a stock solution by dissolving 10.0 g of KOH in enough water to make 150. mL of solution. She then took 15.0 mL of the stock solution and diluted it with enough water to make water to make 65.0 mL of a final solution. What is the concentration of KOH for the final solution?
0. 274 M
Determine the concentration of a solution prepared by diluting 20.0 mL of a 0.200 M CsCl to 250.0 mL.
0.0160 M
Determine the concentration of a solution prepared by diluting 25.0 mL of a stock 0.188 M Ca(NO3)2 solution to 150.0 mL.
0.0313 M
Consider the following balanced reaction. What mass (in g) of CO2 can be formed from 288 mg of O2? Assume that there is excess C3H7SH present. C3H7SH(l) + 6 O2(g) → 3 CO2(g) + SO2(g) + 4 H2O(g)
0.198 g CO2
In a container containing CO, H2, and O2, what is the mole fraction of CO if the H2 mole fraction is 0.22 and the O2 mole fraction is 0.58?
0.20
The titration of 25.0 mL of an unknown concentration H2SO4 solution requires 83.6 mL of 0.12 M LiOH solution. What is the concentration of the H2SO4 solution (in M)?
0.20 M
The partial pressure of CH4, N2, and O2 in a sample of gas were found to be 155 mmHg 434 mmHg, and 685 mmHg, respectively. What is the mole fraction of nitrogen?
0.341
A stock solution of HNO3 is prepared and found to contain 13.5 M of HNO3. If 25.0 mL of the stock solution is diluted to a final volume of .500L, the concentration of the diluted solution is ________M.
0.675 M V₁M₁ = V₂M₂
How many moles of oxygen are formed when 58.6 g of KNO3 decomposes according to the following reaction? The molar mass of KNO3 is 101.11 g/mol. 4 KNO3(s) → 2 K2O(s) + 2 N2(g) + 5 O2(g)
0.724 mol O2
What is the volume of 5.60 g of O2 at 7.78 atm and 415K?
0.766 L
What is the molarity of a NaOH solution if 28.2 mL of a 0.355 M H2SO4 solution is required to neutralize a 25.0-mL sample of the NaOH solution?
0.801
How many of the following species are paramagnetic? Sc3 ⁺ Cl⁻ Ba2⁺ Se
1
How many of the following species are paramagnetic? Sc³⁺ Cl⁻ Ba²⁺ Se
1
How many valence electrons do the alkali metals possess
1
A piece of metal ore weighs 8.25 g. When a student places it into a graduated cylinder containing water, the liquid level rises from 21.25 mL to 26.47 mL. What is the density of the ore?
1. 58 g/mL
A car averages 19.2 miles per gallon of gasoline. How many liters of gasoline will be needed for a trip of 855 km? 1 qt = 0.946 L 1 mile = 1.609 km 4 qt = 1 gal 1 ft = 12 in
1.05 x 10² L
What is the quantity of heat envolved at constant pressure when 67.7 g H2O(l) is formed from the combustion of H2(g) and O2(g)? H₂(g) + ¹/₂O₂(g) → H₂O(l) ∆H°= -285.8 kJ
1.07 x 10³ kJ
A mixture of 10.0 g of Ne and 10.0 g Ar have a total pressure of 1.6 atm. What is the partial pressure of Ne?
1.1 atm
A mixture of 10.0 g of Ne and 10.0 g Ar have a total pressure of 1.6 atm. What is the partial pressure of Ne?
1.1 atm M(i)/M(tot)=p(i)/p(tot)=n(i)/n(tot)=V(i)/V(tot)
Calculate the kinetic energy of a 150 g baseball moving at a speed of 39. m/s ( 87 mph).
1.1 × 10² J
How many atoms of oxygen are contained in 47.6 g of Al2(CO3)3? The molar mass of Al2(CO3)3 is 233.99 g/mol.
1.10 × 1024 O atoms
Based on the balanced chemical equation shown below, determine the molarity of a solution containing Fe2+(aq), if 40.00 mL of the Fe2+(aq) solution is required to completely react with 30.00 mL of a 0. 250 M potassium bromate, KBrO3(aq), solution. The chemical equation for the reaction is 6 Fe2+(aq) + BrO3-(aq) + 6 H+(aq) → 6 Fe3+(aq) + Br-(aq) + 3 H2O(l).
1.12 M
How many oxygen atoms are there in 7.00g of sodium dichromate, Na2Cr207?
1.13 x 10²³ oxygen atoms
The density of nitric oxide (NO) gas at 1.21 atm and 54.1°C is _________g/L
1.35 g/L
In the laboratory, hydrogen gas is usually made by the following reaction: Zn(s) + 2 HCl(aq) → H2(g) + ZnCl2(aq) How many liters of H2 gas, collected over water at an atmospheric pressure of 752 mm Hg and a temperature of 21.0°C, can be made from 3.566 g of Zn and excess HCl? The partial pressure of water vapor is 18.65 mm Hg at 21.0°C.
1.36 L
What is the pressure of a 92.1L gas sample containing 6.96 mol of gas at 16.9°C? r=0.0821 L*atm/K*mol 1atm = 760 torr
1.37 x 10³ mmHg
How many xenon atoms are contained in 2.36 moles of xenon?
1.42 × 10²⁴ xenon atoms Avogadros number: 6.02214 x 10²³
Determine the molarity of a solution formed by dissolving 97.7 g LiBr in enough water to yield 750.0 mL of solution.
1.50 M
Determine the molarity of a solution formed by dissolving 97.7 g LiBr in enough water to yield 750.0 mL of solution.
1.50 M M=mol of solute/liters of solution
A sample of copper absorbs 43.6 kJ of heat, resulting in a temperature rise of 75.0°C, determine the mass (in kg) of the copper sample if the specific heat capacity of copper is 0.385 J/g°C.
1.51 kg
a piece of metal ore weighs 8.25g. when a student places it into a graduated cylinder containing water, the liquid level rises from 21.25 mL to 26.47 mL. What is the density of the ore?
1.58 g/mL
How many anions are there in 2.50 g of MgBr2?
1.64 × 10²² anions
Determine the density of CO2 gas at STP.
1.96 g/L STP L = 22.4L
Calculate the wavelength of an electron (m = 9.11 × 10-28 g) moving at 3.66 × 106 m/s.
1.99 × 10⁻¹⁰ m
Using the following equation for the combustion of octane, calculate the amount of moles of oxygen that reacts with 100.0 g of octane. The molar mass of octane is 114.33 g/mole. The molar mass of carbon dioxide is 44.0095 g/mole. 2 C8H18 + 25 O2 → 16 CO2 + 18 H2O ΔH°rxn = -11018 kJ
10.93 moles
How many milliliters of 0.200 M FeCl3 are needed to react with an excess of Na2S to produce 1.38 g of Fe2S3 if the percent yield for the reaction is 65.0%? 3 Na2S(aq) + 2 FeCl3(aq) → Fe2S3(s) + 6 NaCl(aq)
102 mL
Calculate the wavelength of light associated with the transition from n=1 to n=3 in the hydrogen atom.
103 nm
Silver has an atomic mass of 107.868 amu. The Ag-109 isotope (108.905 amu) is 48.16%. What is the amu of the other isotope?
106.905 amu
Determine the volume of hexane that contains 5.33 × 1022 molecules of hexane. The density of hexane is 0.6548 g/mL and its molar mass is 86.17 g/mol.
11.6 mL
A flexible vessel contains 47L of gas where the pressure is 1.6 atm. what will the volume be when the pressure is 0.67 atm, the temperature remaining constant?
112L
Give the number of pairs of valence electrons for BF₃?
12
Give the approximate bond angle for a molecule with a trigonal planar shape.
120°
What is the O-B-O bond angle in BO₃³⁻?
120°
If the percent yield for the following reaction is 65.0%, how many grams of KClO3 are needed to produce 32.0 g of O2? 2 KClO3(s) → 2 KCl(s) + 3 O2(g)
126 g
the outside temperature is 55°C what is the temp in °F?
131°F
Give the number of neutrons in Al+3.
14
A balloon contains 0.76 mol N2, 0.18 mol O2, 0.031 mol He and 0.026 mol H2 at 7 49 mm Hg. What is the partial pressure of O2?
140 mm Hg
Balance the chemical equation given below, and determine the number of milliliters of 0.00300 M phosphoric acid required to neutralize 45.00 mL of 0.00150 M calcium hydroxide. _____ Ca(OH)2(aq) + _____ H3PO4(aq) → _____ Ca3(PO4)2(aq) + _____ H2O(l)
15.0 mL
How many moles of C8H18 contain 9.25 × 10²⁴ molecules of C8H18 ?
15.4 moles
A 21.8 g sample of ethanol (C2H5OH) is burned in a bomb calorimeter, according to the following reaction. If the temperature rises from 25.0 to 62.3°C, determine the heat capacity of the calorimeter. The molar mass of ethanol is 46.07 g/mol. C2H5OH(l) + 3 O2(g) → 2 CO2(g) + 3 H2O(g) ΔH°rxn = -1235 kJ
15.7 kJ/°C
Use the standard reaction enthalpies given below to determine ΔH°rxn for the following reaction: Given: SO2(g) → S(s) + O2(g) ΔH°rxn = +296.8 kJ 2 SO2(g) + O2(g) → 2 SO3(g) ΔH°rxn = -197.8 kJ 4 SO3(g) → 4 S(s) + 6 O2(g) ΔH°rxn = ?
1583 kJ
Use the standard reaction enthalpies given below to determine ΔH°rxn for the following reaction: 4 SO3(g) → 4 S(s) + 6 O2(g) ΔH°rxn = ? Given: SO2(g) → S(s) + O2(g) ΔH°rxn = +296.8 kJ 2 SO2(g) + O2(g) → 2 SO3(g) ΔH°rxn = -197.8 kJ
1583 kJ
What is the volume (in cm3) of a 43.6 g piece of metal with a density of 2.71 g/cm3?
16.1 cm³ D=m/v 1 mL = 1 cm³
a mass of mercury occupies 0.950 L. What volume would an equal mass of ethanol occupy? the density of mercury is 13.546 g/mL and density of ethanol is 0.789 g/mL.
16.3 L
A 14.01 g sample of N2 reacts with 3.02 g of H2 to form ammonia (NH3). If ammonia is the only product, what mass of ammonia is formed?
17.01 g
Calculate the mass percent composition of lithium in Li3PO4.
17.98 %
Give the complete electronic coniguration for S²⁻.
1s² 2s² 2p⁶ 3s² 3p⁶
The complete electron coniguration of gallium, element 31, is ________
1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p¹
Identify the number of electron groups around a molecule with sp hybridization.
2
Write a balanced equation to show the reaction of gaseous ethane with gaseous oxygen to form carbon monoxide gas and water vapor.
2 C₂H₆(g) + 5 0₂(g) → 4 CO(g) + 6 H₂O(g)
Give the complete ionic equation for the reaction (if any) that occurs when aqueous solutions of lithium sulfide and copper (II) nitrate are mixed.
2 Li+(aq) + S2-(aq) + Cu2+(aq) + 2 NO3-(aq) → CuS(s) + 2 Li+(aq) + 2 NO3-(aq)
Identify the number of bonding pairs and lone pairs of electrons in water.
2 bonding pairs and 2 lone pairs
What are the coefficients in front of NO3-(aq) and Cu(s) when the following redox equation is balanced in an acidic solution: ___ NO3-(aq) + ___ Cu(s) → ___ NO(g) + ___ Cu2+(aq)?
2, 3
Predict the charge that an ion formed from sulfur would have.
2-
For a process at constant pressure, 49,600 calories of heat are released. This quantity of heat is equivalent to
2.08 × 105 J.
Balance the chemical equation given below, and determine the number of moles of iodine that reacts with 30.0 g of aluminum. _____ Al(s) + _____ I2(s) → _____ Al2I6(s)
2.22 mol
Round the following number to four significant figures and express the result in standard exponential notation: 0.00222755
2.228 × 10-3
What is the concentration (M) of CH₃OH in a solution prepared by dissolving 16.8 g of CH₃OH in sufficient water to give exactly 230 mL of solution?
2.28
How much heat is gained by iron when 98.6g of iron is warmed from 25.1°C to 77.8°C? The specific heat of iron is 0.449J/(g*°C).
2.33x10³J
Give the theoretical yield, in moles, of CO2 from the reaction of 4.00 moles of C8H18 with 4.00 moles of O2. 2 C8H18 + 25 O2 → 16 CO2 + 18 H2O
2.56 moles
How many grams of NaOH (MW = 40.0) are there in 250.0 mL of a 0.275 M NaOH solution?
2.75
According to the following balanced reaction, how many moles of NO are formed from 8.44 moles of NO2 if there is plenty of water present? 3 NO2(g) + H2O(l) → 2 HNO3(aq) + NO(g)
2.82 moles NO
What is the concentration of an AlCl3 solution if 150. mL of the solution contains 450. mg of Cl- ion?
2.82 × 10⁻² M
A 26.8L sample of nitrogen at 6.10 atm and 23°C is simultaneously expanded to 59.4L and heated to 34°C. What is the new pressure of the gas?
2.85 atm
How many grams of Li3N can be formed from 1.75 moles of Li? Assume an excess of nitrogen. 6 Li(s) + N2(g) → 2 Li3N(s)
20.3 g Li₃N
A sample of pure calcium fluoride with a mass of 15.0 g contains 7.70 g of calcium. How much calcium is contained in 40.0 g of calcium fluoride?
20.5 g
Calculate the molar mass of Al(C2H3O2)3.
204.13 g/mol
What is the Celsius temperature of 100.0 g of chlorine gas in a 55.0-L container at 800 mm Hg?
228°C
If a room requires 25.4 square yards of carpeting, what is the area of the floor in units of ft sq. (3ft = 1yd)
229 ft²
When 50.0 mL of 0.400 M Ca(NO3)2 is added to 50.0 mL of 0.800 M NaF, CaF2 precipitates, as shown in the net ionic equation below. The initial temperature of both solutions is 23.0°C. Assuming that the reaction goes to completion, and that the resulting solution has a mass of 100.00 g and a specific heat of 4.18 J/(g ∙ °C), calculate the final temperature of the solution. Ca2+(aq) + 2 F-(aq) → CaF2(s) ΔH° = -11.5 kJ
23.55°C
When 11.0g of calcium metal is reacted with water, 5.00g of calcium hydroxide is produced. Using the following balanced equation, calculate the percent yield for the reaction: Ca(s) + 2H₂O(l) →Ca(OH)₂(aq) + H₂(g)
24.6%
Give the percent yield when 28.16 g of CO2 are formed from the reaction of 4.000 moles of C8H18 with 4.000 moles of O2. 2 C8H18 + 25 O2 → 16 CO2 + 18 H2O
25.00%
Give the percent yield when 28.16 g of CO2 are formed from the reaction of 4.000 moles of C8H18 with 4.000 moles of O2. 2 C8H18 + 25 O2 → 16 CO2 + 18 H2O
25.00%
Determine the theoretical yield and the percent yield if 21.8 g of K2CO3 is produced from reacting 27.9 g KO2 with 29.0 L of CO2 (at STP). The molar mass of KO2 = 71.10 g/mol and K2CO3 = 138.21 g/mol. 4 KO2(s) + 2 CO2(g) → 2 K2CO3(s) + 3 O2(g)
27.1 g, 80.4 % yield
What volume of 0.305 M AgNO3 is required to react exactly with 155.0 mL of 0.274 M Na2SO4 solution? Hint: You will want to write a balanced reaction.
278 mL
A 1.00L sample of a gas a STP has a mass of 1.25g. The molar mass of the gas is
28.0 g/mol
Naturally occurring element X exists in three isotopic forms: X-28 (27.976 amu, 92.21%), X-29 (28.976 amu, 4.70%), X-30 (29.974 amu, 3.09% abundance). Calculate the atomic weight of X.
28.1 amu
Calculate the mass percent composition of sulfur in Al2(SO4)3.
28.12 %
Calculate the mass percent composition of sulfur in Al₂(SO₄)₃.
28.12%
Determine the final temperature of a gold nugget (mass = 376 g) that starts at 398 K and loses 4.85 kJ of heat to a snowbank when it is lost. The specific heat capacity of gold is 0.128 J/g°C.
297 K
How many lone pairs are on the Br atom in BrCl⁻?
3
How many lone pairs are on the Br atom in BrCl₂-?
3
How many of the following elements have 1 unpaired electron in the ground state? B Al O F
3
How many of the following molecules have sp³d² hybridization on the central atom? SeF₆ XeF₄ IF₅ AsCl₅
3
How many unpaired electrons are present in the ground state As atom?
3
What value of l is represented by a f orbital?
3
ow many of the following molecules contain at least one pi bond? C₂H₆ Cl₂CO C₂Cl₄ HCN
3
Identify the number of bonding pairs and lone pairs of electrons in N2.
3 bonding pairs and 2 lone pairs
According to the following reaction, what mass of PbCl2 can form from 235 mL of 0.110 M KCl solution? Assume that there is excess Pb(NO3)2. 2 KCl(aq) + Pb(NO3)2(aq) → PbCl2(s) + 2 KNO3(aq)
3.59 g
Calculate the energy of the green light emitted, per photon, by a mercury lamp with a frequency of 5.49 × 10¹⁴ Hz.
3.64 × 10⁻¹⁹ J
Determine the specific heat capacity of an alloy that requires 59.3 kJ to raise the temperature of 150.0 g alloy from 298 K to 398 K.
3.95 J/g°C
If the percent yield for the following reaction is 75.0%, and 45.0 g of NO2 are consumed in the reaction, how many grams of nitric acid, HNO3(aq) are produced? 3 NO2(g) + H2O(l) → 2 HNO3(aq) + NO(g)
30.8 g
What is the wavelength of a photon having a frequency of 9.97 x 10¹⁴ Hz? c= 3.00 x 10⁸ m/s h= 6.63 x 10⁻³⁴ J*s
301 nm
A piece of iron (mass = 25.0 g) at 398 K is placed in a styrofoam coffee cup containing 25.0 mL of water at 298 K. Assuming that no heat is lost to the cup or the surroundings, what will the final temperature of the water be? The specific heat capacity of iron = 0.449 J/g°C and water = 4.18 J/g°C.
308 K
The outside temperature is 35°C, what is the temperature in K?
308 K
Based on the balanced chemical equation shown below, determine the mass percent of Fe3+ in a 0.7450g sample of iron ore, if 22.40 mL of a 0.1000 M stannous chloride, SnCl2(aq), solution is required to completely react with the Fe3+ present in the ore sample. The chemical equation for the reaction is 2 Fe3+(aq) + Sn2+(aq) → 2 Fe2+(aq) + Sn4+(aq).
33.58%
Calculate the change in internal energy (∆E) for a system that is absorbing 35.8kJ of heat and is expanding from 8.00 to 24.0L in volume at 1.00 atm. (remember that 101.3 J = 1 L * atm
34.2 kJ
A sample of N2 effuses in 220 s. How long will the same size sample of Cl2 take to effuse?
350 s
What volume will a balloon occupy at 1.0 atm, if the balloon has a volume of 8.8 L at 4.4 atm?
39 L
4.02 seconds contain this many nanoseconds.
4.02 X 10⁻⁹ ns
What pressure (in atm) will 0.44 moles of CO2 exert in a 2.6 L container at 25°C?
4.1 atm
What pressure will 14.0 g of CO exert in a 3.5 L container at 75°C?
4.1 atm
How many moles of Cs are contained in 595 kg of Cs?
4.48 x 10³ moles Cs
What is the wavelength of light emitted when the electron in a hydrogen atom undergoes a transition form level n=7 to level n=5? c= 3.00 x 10⁸ m/s h= 6.63 x 10⁻³⁴ J*s Rh= 2.179 x 10⁻¹⁸ J
4.66 x 10⁻⁶m
What volume of benzene (C6H6, d= 0.88 g/mL, molar mass = 78.11 g/mol) is required to produce 1.5 x 103 kJ of heat according to the following reaction? 2 C6H6(l) + 15 O2(g) → 12 CO2(g) + 6 H2O(g) ΔH°rxn = -6278 kJ
42 mL
How many grams of NaCl are required to make 500.0 mL of a 1.500 M solution?
43.83 g
Calculate the wavelength (in nm) of the blue light emitted by a mercury lamp with a frequency of 6.88 × 1014 Hz
436 nm
Calculate the wavelength (in nm) of the blue light emitted by a mercury lamp with a frequency of 6.88 × 1014 Hz
436 nm c= 3.00 x 10⁸ m/s
What is ∆H° of the following reaction? CO₂(g) + 2CH₄(g)→C₃H₈(g) + O₂(g) CO₂(g) -393.5 CH₄(g) -74.9 C₃H₈(g) -104.7
438.6 kJ
A large balloon is initially filled to a volume of 25.0 L at 353 K and a pressure of 2575 mm Hg. What volume of gas will the balloon contain at 1.35 atm and 253 K?
45.0 L
What pressure would a gas mixture in a 10.0 L tank exert if it were composed of 48.5 g He and 94.6 g CO2 at 398 K?
46.6 atm
How many grams of oxygen are formed when 6.21 moles of KOH are formed? 4 KO(s) + 2 H2O(l) → 4 KOH(s) + O2(g)
49.7 g O2
An alligator is 152.4 cm long. How long is he in feet?
5.00 ft 2.54 cm = 1 in
What is the mass (in kg) of 6.89 × 1025 molecules of CO2? The molar mass of CO2 is 44.01 g/mol.
5.04 kg
To what volume will a sample of gas expand if it is heated from 50.0∘C and 2.33 L to 500.0°C?
5.58 L
Electromagnetic radiation with a wavelength of 525 nm appears as green light to the human eye. The frequency of this light is ________s⁻¹ .
5.71 x 10¹⁴ s⁻¹
How many grams of AgNO3 are needed to make 250 mL of a solution that is 0.135 M?
5.73 g
An atom of 131 Xe contains ________ electrons.
54
How many protons (p) and neutrons (n) are in an atom of Cesium-133?
55 p, 78 n
How many protons (p) and neutrons (n) are in an atom of barium-130?
56 protons 74 neutrons
Determine the theoretical yield of HCl if 60.0 g of BCl3 and 37.5 g of H2O are reacted according to the following balanced reaction. A possibly useful molar mass is BCl3 = 117.16 g/mol. BCl3(g) + 3 H2O(l) → H3BO3(s) + 3 HCl(g)
56.0 g HCl
What is the molar mass of 1-butene if 5.38 × 1016 molecules of 1-butene weigh 5.00 μg?
56.0 g/mol
According to the following reaction, how much energy is evolved during the reaction of 2.50 L B2H6 and 5.65 L Cl2 (Both gases are initially at STP)? The molar mass of B2H6 is 27.67 g/mol. B2H6(g) + 6 Cl2(g) → 2 BCl3(g) + 6 HCl(g) ΔH°rxn = -1396 kJ
58.7 kJ
Calcium acetate reacts with phosphoric acid to form calcium phosphate and acetic acid. What is the coefficient for acetic acid when the equation is balanced using the lowest, whole numbered coefficients.
6
How many significant figures are in 0.00523980 mL?
6
How many valence electrons does a neutral tellurium atom have?
6
Calculate the frequency of the green light emitted by a hydrogen atom with a wavelength of 486.1 nm.
6.17 × 10¹⁴ s⁻¹
Given: Fe₂O₃(s) + 3CO(g) → 2 Fe(s) + 3CO₂(g) ∆H°= -26.8 kJ FeO(s) + CO(g) → Fe(s) + CO₂(g) ∆H°= -16.5kJ Fe₂O₃(s) + CO(g) → 2FeO(s) + CO₂(g) ∆H°=?
6.2 kJ
What is the mass of a single fluorine molecule, F2?
6.310 × 10⁻²³ g
Magnesium burns in air with a dazzling brilliance to produce magnesium oxide: 2Mg(s) + O2(g) → 2MgO(s) When 4.00 g of magnesium burns, the theoretical yield of magnesium oxide is ________ g.
6.63
What is the average speed (actually the root-mean-square speed) of a neon atom at 27°C?
609 m/s
What is the pressure in a gas container that is connected to an open-end U-tube manometer if the pressure of the atmosphere is 742 torr and the level of mercury in the arm connected to the container is 8.60 cm higher than the level of mercury open to the atmosphere?
656 mm Hg
What is the maximum number of f orbitals that are possible?
7
If the walls in a room are 955 square feet in area, and a gallon of paint covers 15 square yards, how many gallons of paint are needed for the room? (3 ft = 1 yd)
7.1 gallons
Balance the chemical equation given below, and calculate the volume of nitrogen monoxide gas produced when 8.00 g of ammonia is reacted with 12.0 g of oxygen at 25°C? The density of nitrogen monoxide at 25°C is 1.23 g/L. _____ NH3(g) + _____ O2(g) →______ NO(g) + _____ H2O(l)
7.32 L
How many moles of Li F are contained in 258.6 mL of 0.0296 M Li F solution?
7.65 × 10-3 mol
How many moles of potassium are contained in 300 g of potassium?
7.67 moles
The action of some commercial drain cleaners is based on the following reaction: 2 NaOH(s) + 2 Al(s) + 6 H2O(l) → 2 NaAl(OH)4(s) + 3 H2(g) What is the volume of H2 gas formed at STP when 6.32 g of Al reacts with excess NaOH?
7.87 L
What is the wavelength of an electron traveling at 3.04% of the speed of light? (m e=9.109 x 10⁻³¹ kg, c = 3.00 x 10⁸ m/s, h = 6.63 x 10⁻³⁴ J*s)
7.98 x 10⁻¹¹ m
A 0. 286-g sample of gas occupies 125 mL at 60. cm of Hg and 25°C. What is the molar mass of the gas?
71 g/mol
What is the mass of 3.91 x 10^24 molecules of SeO2? The molar mass is 110.96 g/mol.
721 g
what is the mass of 3.91 x 10^24 molecules of SeO2? the molar mass of SeO2 is 110.96 g/mol.
721g SeO₂
How much energy is evolved during the formation of 98.7 g of Fe, according to the reaction below? Fe2O3(s) + 2 Al(s) → Al2O3(s) + 2 Fe(s) ΔH°rxn = -852 kJ
753 kJ
Two samples of potassium iodide are decomposed into their constituent elements. The first sample produced 13.0 g of potassium and 42.3 g of iodine. If the second sample produced 24.4 kg of potassium, how many kg of iodine were produced?
79.4 kg
If an object has a density of 8.65 g/cm³, what is its density in units of kg/m³?
8.65 × 103 kg/m³
A 55.0-L steel tank at 20.0°C contains acetylene gas, C2H2, at a pressure of 1.39 atm. Assuming ideal behavior, how many grams of acetylene are in the tank
82.9 g
What mass of NO2 is contained in a 13.0 L tank at 4.58 atm and 385 K?
86.7 g
Determine the mass of an object that has a volume of 88.6 mL and a density of 9.77 g/mL.
866 g
To what temperature must a balloon, initially at 25°C and 2.00 L, be heated in order to have a volume of 6.00 L?
894 K °C + 273.15 = °K
How many molecules of H2S are required to form 79.0 g of sulfur according to the following reaction? Assume excess SO2. 2 H2S(g) + SO2(g) → 3 S(s) + 2H2O(l)
9.89 × 1023 molecules H2S
If the density of ethanol, C2H5OH, is 0.789 g/mL. How many milliliters of ethanol are needed to produce 15.0 g of CO2 according to the following chemical equation? C2H5OH(l) + 3 O2(g) → 2 CO2(g) + 3 H2O(l)
9.95 mL
Give the approximate bond angle for a molecule with an octahedral shape.
90°
A student is preparing to perform a series of calorimetry experiments. She first wishes to determine the calorimeter constant (Ccal) for her coffee cup calorimeter. She pours a 50.0 mL sample of water at 345 K into the calorimeter containing a 50.0 mL sample of water at 298 K. She carefully records the final temperature of the water as 317 K. What is the value of Ccal for the calorimeter?
99 J/K
percent ionic character
= 0, the bond is pure covalent. = 0.1 to 0.4, the bond is nonpolar covalent. = 0.5 to 1.9, the bond is polar covalent. = ≥ 2.0, the bond is ionic.
Choose the compound below that should have the highest melting point according to the ionic bonding model. A) MgF² B) SrI² C) SrF² D) CaCl² E) SrBr²
A
Choose the reaction that illustrates ∆H°f for Mg(NO3)2(s) A)Mg(s) + N₂(g) +3O₂(g) → Mg(NO₃)₂(s) B)Mg²⁺(aq) +2NO₃⁻(aq) → Mg(NO₃)₂(aq) C) Mg(NO₃)₂(s) → Mg(s) + N₂(g) +3O₂(g) D)Mg(s) + 2N(g) + 6O(g) →Mg(NO₃)₂(s) E) Mg(NO₃)₂(aq) → Mg²⁺(aq) + 2NO₃⁻(aq)
A
Which of the following elements is a gas at room temperature? A) helium B) bromine C) sodium D) carbon
A
Which of the following statements is FALSE? A) The alkali metals are fairly unreactive. B) Halogens are very reactive elements. C) Sulfur is a main group element. D) Noble gases do not usually form ions. E) Zn is a transition metal.
A
What is the maximum # of electrons and the shape of A) s orbital B) p orbital C) d orbital D) f orbital
A) 2 e⁻ B) 6 e⁻ C) 10 e⁻ D) 14 e⁻
7.0 g of nitrogen is reacted with 5.0 g of hydrogen to produce ammonia according to the chemical equation shown below. complete the following problems: N₂(g) + 3H₂(g) → 2NH₃(g) A) The theoretical yield of ammonia is? B) The limiting reactant is C) the excess reactant is D) The amount of excess reactant used is E) The amount of excess reactant left over is
A) 8.5g B) N C) H D) 1.5g of H₂ E) 3.5g of H₂
Identify the compound with atoms that have an incomplete octet. A) BF₃ B) CO C) Cl₂ D) ICl₅ E) CO₂
A) BF₃
Which of the following compounds is not an Arrhenius acid? A) C H3CH2NH2 B) H2SO4 C) CH3CO2H D) HNO2
A) C H3CH2NH2
Which of the following processes are endothermic? A) Ca(s) → Ca(g) B) 2Na(s) + ¹/₂ O₂(g) → Na₂O(s) C) 2Br(g) → Br₂(g) D) K⁺(g) + I⁻(g) → KI(s) E) None of the above are endothermic.
A) Ca(s) → Ca(g)
Choose the reaction that illustrates ΔH°f for CsHCO3. A) Cs(s) + 1/2 H2(g) + C(s) + 3/2 O2(g) → CsHCO3(s) B) Cs+(aq) + HCO3 -1(aq) → CsHCO3(s) C) Cs(s) + 2 H(g) + C(s) + 3 O(g) → CsHCO3(s) D) Cs+(aq) + H2O(l) + CO2(g) → CsHCO3(s) E) Cs(s) + H2(g) + C(s) + O2(g) → CsHCO3(s)
A) Cs(s) + 1/2 H2(g) + C(s) + 3/2 O2(g) → CsHCO3(s)
Choose the bond below that is most polar. A) H-F B) C-H C) H-Br D) H-I E) H-Cl
A) H-F
What is the chemical formula os the following compounds and are they ionic/covalent/or both? A) Nitrous Acid B) Tetraphosphorus Decaoxide C) Tin (IV) Sulfate D) Carbonic Acid E)Cobalt (II) Chloride hexahydrate F) Iron (II) Nitride
A) HNO₂, covalent B) P₄O₁₀, covalent C) Sn(SO₄)₂, Both D) H₂CO₃(aq), covalent E) CoCl₂* 6H₂O, both F) Fe₃N₂, Ionic
Which of the following is an oxidation-reduction reaction? A) Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g) B) H I(aq) + LiOH(aq) → Li I(aq) + H2O(l) C) Na Cl(aq) + AgNO3(aq) → Ag Cl(s) + NaNO3(aq) D) Pb(C2H3O2)2(aq) + 2 NaCl(aq) → PbCl2(s) + 2 NaC2H3O2(aq) E) All of the above are oxidation-reduction reactions.
A) Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g)
Which of the following pairs of aqueous solutions will form a precipitate when mixed? A) MgCl2 + KOH B) Li2S + H F C) H I+ LiOH D) K2CO3 + HNO3 E) All of these solution pairs will produce a precipitate.
A) MgCl2 + KOH
What are characteristics of energy?
A) Systems tend to change in order to lower their potential energy. B) The total energy of a system remains constant. C) Energy can be converted from one type to another. D) Energy is the capacity to do work.
Draw the Lewis structure for CO₃²⁻ including any valid resonance structures. Which of the following statements is TRUE A) The CO₃²⁻ ion contains two C-O single bonds and one C-O double bond. B) The CO₃²⁻ ion contains three C-O double bonds. C) The CO₃²⁻ ion contains two C-O single bonds and one C-O triple bond. D) The CO₃²⁻ ion contains one C-O single bond and two C-O double bonds. E) None of the above are true.
A) The CO₃²⁻ ion contains two C-O single bonds and one C-O double bond.
Which of the following is TRUE if ΔEsys = - 115 J? A) The system is losing 115 J, while the surroundings are gaining 115 J. B) Both the system and the surroundings are losing 115 J. C) Both the system and the surroundings are gaining 115 J. D) The system is gaining 115 J, while the surroundings are losing 115 J. E) None of the above are true.
A) The system is losing 115 J, while the surroundings are gaining 115 J.
Which of the following is an example of the law of multiple proportions? A) Two different compounds formed from carbon and oxygen have the following mass ratios: 1.33 g O: 1 g C and 2.66 g O: 1 g C. B) A sample of chlorine is found to contain three times as much Cl-35 as Cl-37. C) Two different samples of table salt are found to have the same ratio of sodium to chlorine. D) The atomic mass of bromine is found to be 79.90 amu. E) Nitrogen dioxide always has a mass ratio of 2.28 g O: 1 g N.
A) Two different compounds formed from carbon and oxygen have the following mass ratios: 1.33 g O: 1 g C and 2.66 g O: 1 g C.
What are the principles of Dalton's Atomic Theory?
A) all atoms of the same element have identical properties that distiguish them from other elements. B) atoms combine in simple whole number ratios to form compounds C) atoms of a specific element cannot change into another element even during chemical reactions. D) One carbon atom will combine with one oxygen atom to form a molecule of carbon monoxide.
Choose the pure substance from the list below A) carbon monoxide B) pomegranate juice C) a casserole D) tea E) sugar water
A) carbon monoxide
Which of the following combinations of quantum numbers is permissible? A) n = 3, l = 2, ml = 1, ms = +¹/₂ B) n = 1, l = 2, ml = 0, ms = -¹/₂ C) n = 4, l = 3, ml = 4, ms = -¹/₂ D) n = 2, l = 1, ml = -1, ms = 0 E) n = 3, l = 3, ml = 1, ms = -¹/₂
A) n = 3, l = 2, ml = 1, ms = +¹/₂
There are 4 quantum numbers. what is there description and permissibility of A) n B) l C) ml D) ms
A) principle quantum # → energy level (1 , 2, 3, 4, depending on row on periodic table) B) angle quantum # → shape of orbit (0, 1, 2, 3, or n-1) C) orientation quantum # → where electrons occur in orbit (-l to +l) D) magnetic quantum # → direction of e⁻ spin (+¹/₂=↑ or -¹/₂=↓)
What element is undergoing reduction (if any) in the following reaction? Zn(s) + 2 AgNO3(aq) → Zn(NO3)2(aq) + 2 Ag(s)
Ag
What element is undergoing reduction (if any) in the following reaction? Zn(s) + 2 AgNO3(aq) → Zn(NO3)2(aq) + 2 Ag(s)
Ag LEO says GER loss electrons oxidation gain electrons redox
How many protons neutrons and electrons do the following elements have? Al, Br-, Ca2+
Al = 13 p, 14 n, 13 e Br- = 35 p, 45 n, 36 e Ca2+ = 20 p, 20 n, 18 e
Write a balanced equation to show the reaction of aqueous aluminum acetate with aqueous ammonium phosphate to form solid aluminum phosphate and aqueous ammonium acetate.
Al(C2H3O2)3(aq) + (NH4)3PO4(aq) → AlPO4(s) + 3 NH4C2H3O2
What are Strong electrolytes?
Ammonium sulfate, calcium chloride, hydrochloric acid, sodium hydroxide, sodium phosphate, zinc nitrate, and tend to be inorganic compounds meaning that they lack carbon atoms.
cation
An atom that has lost an electron.
Place the following in order of decreasing IE₁ (first ionization energy) Cs Mg Ar
Ar > Mg > Cs
Chemical energy
Associated with positions of electrons and nuclei
Thermal Energy
Associated with temperature
Which atom in each group (I and II) has the smallest atomic radius? (I) Ba, Hf, At (II) As, Sb, Bi
At; As
Identify the polyprotic acid. A) HCl B) H2SO4 C) Li Cl D) Ca(OH)2 E) LiOH
B
Identify the states of the halogens at room temperature. A) Fluorine, chlorine, and bromine are solids, and iodine is a solid. B) Fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid. C) Fluorine and iodine are solids, chlorine and bromine are gases. D) Fluorine, chlorine, bromine, and iodine are solids. E) Fluorine is a gas, chlorine and bromine are liquids, and iodine is a solid.
B
Which of the following elements has chemical properties similar to tellurium? A) hydrogen B) sulfur C) nitrogen D) fluorine
B
Which of the following is not a standard state? A) for a liquid, it is 1 atm B) for a liquid, it is 25°F C) for a solid, it is 1 atm D) for a solid, it is 25°C E) for a solution, it is 1 M
B
Identify a characteristic of halogens. A) absorbs water in reactions B) powerful oxidizing agents C) inert D) forms water in reactions E) powerful reducing agents
B)
Which of the following solutions will have the highest concentration of chloride ions? A) 0.10 M LiCl B) 0.10 M AlCl3 C) 0.05 M CaCl2 D) 0.10 M MgCl2 E) All of these solutions have the same concentration of chloride ions
B) 0.10 M AlCl3
Which of the following gas samples would be most likely to behave ideally under the stated conditions? A) SO2 at 2 atm and 0 K B) Ar at STP C) N2 at 1 atm and -70°C D) H2 at 400 atm and 25°C E) CO at 200 atm and 25°C
B) Ar at STP
Which of the following elements is a metalloid? A) Br B) As C) Kr D) Fe E) S
B) As
Which of the following statements is TRUE? A) The smaller a gas particle, the slower it will effuse B) At low temperatures, intermolecular forces become important and the pressure of a gas will be lower than predicted by the ideal gas law. C) The higher the temperature, the lower the average kinetic energy of the sample. D) At a given temperature, lighter gas particles travel more slowly than heavier gas particles. E) None of the above statements are true.
B) At low temperatures, intermolecular forces become important and the pressure of a gas will be lower than predicted by the ideal gas law.
Which of the following contains BOTH ionic and covalent bonds? A) COS B) BaSO4 C) CaCl2 D) SF6 E) None of the above contain both ionic and covalent bonds.
B) BaSO₄
In which of the following sets do all species have the same number of protons? A) F-, Ne, Mg2+ B) Br, Br-, Br+ C) K+, Rb+, Cs+ D) Ge, Se2-, Br-
B) Br, Br-, Br+
Use the molecular orbital diagram shown to determine which of the following are paramagnetic A) C₂²⁻ B) B₂ C) B₂²⁺ D) N₂²⁺ E) B₂²⁻
B) B₂
Which of the following pairs of aqueous solutions will form a precipitate when mixed? A) K₂CO₃ + NaBr B) CaS + Na₂SO₄ C) Na₂SO₄ + KOH D) None of these solution pairs will produce a precipitate. E) All of these solution pairs will produce a precipitate.
B) CaS + Na₂SO₄
In which of the following sets do all species have the same number of electrons? A) K+, Rb+, Cs+ B) F-, Ne, Mg2+ C) Ge, Se2-, Br- D) Br, Br-, Br+
B) F-, Ne, Mg2+
Which compound has the highest carbon-carbon bond strength? A) CH₂CH₂ B) HCCH C) CH₃CH₃ D) all bond strengths are the same
B) HCCH
Which of the following species will have the highest ionization energy? A) S2- B) K+ C) Cl- D) Ar
B) K+
Which element has the highest first electron affinity? A) Ne B) O C) Na D) Mg
B) O
Which of the following processes is endothermic? A) A hot cup of coffee (system) cools on a countertop B) the vaporization of rubbing alcohol C) the combustion of butane D) the freezing of water E) the chemical reaction in a "hot pack" often used to treat sore muscles
B) the vaporization of rubbing alcohol
Place the following in order of increasing radius. Ba2⁺ Te2⁻ I⁻
Ba²⁺ < I⁻ < Te²⁻
Using the VSEPR model, the molecular geometry of the central atom in SO₂ is?
Bent
What is the element symbol for an atom that has 5 protons and 6 neutrons?
Boron-II Isotope
Of the following, which element has the highest first ionization energy? A) Sr B) K C) Br D) Te
Br
Identify the compound with atoms that have an incomplete octet. A) ICl5 B) CO2 C) BF3 D) CO E) Cl2
C
Which ion does NOT have a noble gas configuration in its ground state? A) As³⁻ B) Sc³⁺ C) Ga³⁺ D) Al³⁺
C
Determine the empirical formula for a compound that contains C, H and O. It contains 52.14% C and 34.73% O by mass.
C2H6O
What is the empirical formula of a compound that is 62.0% C, 10.4% H, and 27.5% O by mass?
C3H6O
Which of the compounds C4H10, BaCl2, Ni(NO3)2, SF6 are expected to exist as molecules?
C4H10 and SF6
Determine the molecular formula of a compound that has a molar mass of 183.2 g/mol and an empirical formula of C2H5O2.
C6H15O6
What is the empirical formula of a substance that contains 2.64 g of C, 0.444 g of H, and 3.52 g of O?
CH₂O
Choose the statement that is TRUE. A) Core electrons are the easiest of all electrons to remove. B) Outer electrons eiciently shield one another from nuclear charge. C) Valence electrons are most diicult of all electrons to remove. D) Core electrons efectively shield outer electrons from nuclear charge. E) All of the above are true.
D
Identify acetic acid. A) weak electrolyte, strong acid B) strong electrolyte, weak acid C) nonelectrolyte D) weak electrolyte, weak acid E) strong electrolyte, strong acid
D
Which element has the highest first electron affinity? A) Ne B) Mg C) Na D) O
D
Which of the following is the smallest volume? A) 5.0 × 10⁷ nL B) 0.50 dL C) 2.8 × 10³ mL D) 22 cm³
D
Which of the following sets of units is NOT in the order of increasing size? A) ug < mg < cg B) uL < dL < L C) cPa < dPa < kPa D) pm < mm < nm E) ns < ms < s
D
Which of the following processes is exothermic? A) the ionization of a potassium atom B) the sublimation of dry ice (CO2(s)) C) the breaking of a Br-Br bond D) the reaction associated with ∆H°f for an ionic compound E) All of the above processes are exothermic.
D)
Which of the following solutions will have the highest electrical conductivity? A) 0.10 M NaBr B) 0.10 M NaI C) 0.10 M NaF D) 0.045 M Al2(SO4)3 E) 0.050 M (NH4)2CO3
D) 0.045 M Al2(SO4)3
Which of the following is a precipitation reaction? A) Zn(s) + 2 AgNO3(aq) → 2 Ag(s) + Zn(NO3)2(aq) B) NaCl(aq) + LiI(aq) → NaI(aq) + LiCl(aq) C) H I(aq) + NaOH(aq) → Na I(aq) + H2O(l) D) 2 KI(aq) + Hg2(NO3)2(aq) → Hg2I2(s) + 2 KNO3(aq) E) None of the above are precipitation reactions.
D) 2 KI(aq) + Hg2(NO3)2(aq) → Hg2I2(s) + 2 KNO3(aq)
All of the following elements are nonmetals EXCEPT: A) helium B) nitrogen C) oxygen D) arsenic E) none of the above
D) Arsenic
which of the following are examples of a chemical change A) a candle burns B) copper building materials develop a green patina over time C) rubbing alcohol evaporates D) Both A and B are examples of chemical change E) All of the above are examples of chemical change
D) Both A and B are examples of chemical change
Choose the bond below that is most polar. A) H-I B) H-Br C) H-Cl D) C-H E) H-F
D) C-H
Choose the reaction that represents the combustion of C6H12O2. A) 6 C(s) + 6 H2(g) + O2(g) → C6H12O2(l) B) Mg(s) + C6H12O2(l) → MgC6H12O2(aq) C) C6H12O2(l) → 6 C(s) + 6 H2(g) + O2(g) D) C6H12O2(l) + 8 O2(g) → 6 CO2(g) + 6 H2O(g) E) None of the above represent the combustion of C6H12O2.
D) C6H12O2(l) + 8 O2(g) → 6 CO2(g) + 6 H2O(g)
Choose the bond below that is the strongest. A) C-I B) I-I C) C=O D) C≡N E) C-F
D) C≡N
Which of the following is a correct statement of Charles's law, V/T=k A)The volume of gas varies proportionally with the pressure B) The pressure of a gas sample varies inversely with the volume. C) all gas samples of the same volume at STP contain the same number of atoms. D) The volume of a gas sample varies directly with the absolute temperature E) All gas samples of the same volume at STP contain the same number of molecules.
D) The volume of a gas sample varies directly with the absolute temperature
Which one of the following elements is a poor conductor of heat and electricity? A) lead B) copper C) iron D) fluorine
D) fluorine
Which of the following occur as the energy of a photon increases? A) the speed increases. B) the frequency decreases. C) the wavelength increases D) the wavelength gets shorter. E) None of the above occur as the energy of a photon increases.
D) the wavelength gets shorter.
Which of the following statements is TRUE? A) The emission spectrum of a particular element is always the same and can be used to identify the element. B) The uncertainty principle states that we can never know both the exact location and speed of an electron. C) An orbital is the volume in which we are most likely to find an electron. D) Part of the Bohr model proposed that electrons in the hydrogen atom are located in "stationary states" or particular orbits around the nucleus. E) All of the above are true.
E) All of the above are true.
Which of the following are examples of physical change? A) sugar is dissolved in water B) dry ice sublimes C) tea is brewed D) water freezes E) All of these are examples of physical change
E) All of these are examples of physical change
Which of the following are examples of physical change? A) water freezes B) tea is brewed C) dry ice sublimes D) sugar is dissolved in water E) All of these are examples of physical change.
E) All of these are examples of physical change.
Choose the ground state electron coniguration for Hf2⁺. A) [ Xe] B) [ Xe] 6s2 5d4 C) [ Xe] 6s2 D) [ Xe] 6s2 5d2 E) [ Xe] 5d2
E) [ Xe] 5d2
Which of the following is a molecular element? A) xenon B) argon C) barium D) potassium E) iodine
E) iodine
Give the set of four quantum numbers that could represent the electron lost to form the Rb ION from the Rb atom. A) n = 4, l = 1, ml = 1, ms = +¹/₂ B) n = 6, l = 0, ml = -1, ms = -¹/₂ C) n = 4, l = 1, ml = 0, ms = -¹/₂ D) n = 5, l = 0, ml = 0, ms = -¹/₂ E) n = 5, l = 0, ml = 0, ms = +¹/₂
E) n = 5, l = 0, ml = 0, ms = +¹/₂
Which element has the ground-state electron configuration [Xe] 6s² 4f⁷?
Eu
Abnormally charged elements
Fe +² +³ Cu +¹+² Zn +² Sn +⁴+² Pb +⁴+²
Give the net ionic equation for the reaction (if any) that occurs when aqueous solutions of K2S and Fe(NO3)2 are mixed.
Fe2+(aq) + S2-(aq) → FeS(s)
Give the net ionic equation for the reaction (if any) that occurs when aqueous solutions of H2SO4 and KOH are mixed.
H+(aq) + OH-(aq) → H2O(l)
Identify the gas particle that travels the fastest.
H2
Which of the compounds of H2C2O4, Ca(OH)2, KOH, and H I, behave as acids when they are dissolved in water?
H2C2O4 and H I
A student performs an experiment to determine the density of a sugar solution. She obtains the following results: 1.71 g/mL, 1.73 g/mL, 1.67 g/mL, 1.69 g/mL. If the actual value for the density of the sugar solution is 1.40 g/mL, which statement below best describes her results?
Her results are precise, but not accurate
Name the 7 diatomic elements.
H₂, N₂, O₂, F₂, Cl₂, Br₂, and I₂
Write a balanced equation to show the reaction of sulfurous acid with lithium hydroxide to form water and lithium sulfite.
H₂SO₃(aq) + 2 LiOH(aq) → 2 H₂O(l) + Li₂SO₃(aq)
when a hydrogen electron makes a transition from n = 3 to n = 1, which of the following statements are true? I) energy is emitted II) energy is absorbed III) the electron loses energy IV) the electron gain energy V) the electron cannot make this transition
I and III
Determine the limiting reactant (LR) and the mass (in g) of nitrogen that can be formed from 50.0 g N2O4 and 45.0 g N2H4. Some possibly useful molar masses are as follows: N2O4 = 92.02 g/mol, N2H4 = 32.05 g/mol. N2O4(l) + 2 N2H4(l) → 3 N2(g) + 4 H2O(g)
LR = N2O4, 45.7 g N2 formed
Determine the limiting reactant (LR) and the mass (in g) of nitrogen that can be formed from 50.0 g N2O4 and 45.0 g N2H4. Some possibly useful molar masses are as follows: N2O4 = 92.02 g/mol, N2H4 = 32.05 g/mol. N2O4(l) + 2 N2H4(l) → 3 N2(g) + 4 H2O(g)
LR = N₂O₄, 45.7 g N₂ formed
All samples of a given compound, regardless of their source or how they were prepared, have the same proportions of their constituent elements. Which law does this refer to?
Law of Definite Proportions
Identify the spectator ions in the following molecular equation. Li Cl(aq) + AgNO3(aq) → Ag Cl(s) + LiNO3(aq)
Li+ and NO3-
What is the identity of element Q if the ion Q2+ contains 10 electrons?
Mg
What species is represented by the following information? p+ = 12 n° = 14 e- = 10
Mg2+
Place the following gases in order of increasing density at STP. F2 NH3 N2O4 Ar
NH3 < F2 < Ar < N2O4
Place the following in order of increasing metallic character. Rb Cs K Na
Na < K < Rb < Cs
Pauli exclusion principle
No two electrons can have the same four quantum numbers
A compound is found to be 30.45% N and 69.55 % O by mass. If 1.63 g of this compound occupy 389 mL at 0.00°C and 775 mm Hg, what is the molecular formula of the compound?
N₂O₄
Using Lewis structures and formal charge, which of the following ions is most stable? OCN⁻ ONC⁻ NOC⁻
OCN⁻
ideal gas law
PV = nRT r = 0.08206 L*atm/mol*K
Electromagnetic spectrum in order of decreasing wavelength.
Radio Microwave Infrared Visible Light Ultraviolet Xrays Gamma
Using periodic trends, place the following bonds in order of decreasing ionic character. Sb-F P-F As-F
Sb-F > As-F > P-F
Identify the chemical symbol of element Q atomic number = 34 and atomic mass = 80
Se
STP
T = 273.15 °K or 0°C P = 1 atm V = 22.4 L
Theoretical Yield
The amount of product that can be made from the limiting reactant
Ionic Bonds
The reaction between metals and nonmetals and involve the transfer of electrons from one atom to another
Covalent Bonds
The reaction between two or more nonmetals and involve the sharing of electrons between two atoms
What are the standard states?
The standard state of a chemical substance is its phase (solid, liquid, gas) at 25.0 °C and one atmosphere pressure. This temperature/pressure combo is often called "room conditions."
Stoichiometry
The study of the numerical relationship between chemical quantities in a chemical reaction
Boyle's Law
The volume of a gas is inversely proportional to the pressure of a gas
Molecular Formula
This gives the actual number of atoms of each element in a molecule of a compound
Using the VSEPR model, the electron-domain geometry of the central atom in SF₄ is?
Trigonal BiPyramidal
using the Van Der Waals equation and ideal gas law, determine and compare the pressure of 321.0g of SO2(g) in a 8.50L vessel at 769K. For SO2(s), a = 6.865 L²*atm/mol² b = 0.05679 L/mol. r = 0.0821 L*atm/k*mol
Van Der Waals = 36.8 atm Ideal gas law = 37.2 atm
Hund's rule
When filling degenerate orbitals with e⁻ share one with each position then go back and pair to shell is full
What did Mendeleev contribute to Chemistry?
When the elements are arranged in order of increasing mass, certain sets of properties recur periodically. --periodic table
destructive interference
When waves of equal amplitude from two sources are out of phase interact
Choose the best Lewis structure for PO₄³⁻.
[3 single bonds 1 double bond]³⁻
Give the ground state electron configuration for Rb⁺.
[Ar]4s²3d¹⁰4p⁶
Which of the following represent the Lewis structure for Ca²⁺?
[Ca]²⁺
Give the ground state electron coniguration for the ion of Ba
[Kr] 5s² 4d¹⁰ 5p⁶
In the Van Der Waals Equation the effect of the intermolecular forces is accounted for by
a
A substance composed of two or more elements in a fixed, definite proportion is
a compound.
Ammonia is an example of
a compound.
Two or more substances in variable proportions, where the composition is constant throughout are
a homogeneous mixture.
Distillation is
a process in which the more volatile liquid is boiled off
The Heisenberg Uncertainty principle states that
both the position of an electron and its momentum cannot be known simultaneously very accurately.
Determine the name for CoCl2∙6H2O. Remember that Co forms several ions.
cobalt (II) chloride hexahydrate
The square of the wave function of an electron in an atom
gives the probability of finding the electron in a region of space
A laser emits photons having energy of 3.74 x 10⁻¹⁹ J. What color would be expected for the light emitted by this laser? c= 3.00 x 10⁸ m/s h= 6.63 x 10⁻³⁴ J*s
green
The electronegativity is 2.1 for H and 1.9 for Pb. Based on these electronegativities PbH4 would be expected to
have polar covalent bonds with a partial negative charges on the H atom
What did Thomson contribute to Chemistry?
he discovered that particles have a negative electrical charge (e-)
what are the exceptions to the Ideal Gas Law?
high pressure or low temperature
An aqueous solution of H2S is named
hydrosulfuric acid
Using the VSEPR model, the molecular geometry of the central atom in XeF₂ is ________.
linear
Which of the following quantum numbers describes the orientation of an orbital?
magnetic quantum number
mass density
molar mass/molar volume
mole fraction
mol₁/total mol
Give the set of four quantum numbers that could represent the last electron added (using the Aufbau principle) to the Cl atom.
n = 3, l = 1, ml = 1, ms = +¹/₂
Give the set of four quantum numbers that could represent the last electron added (using the Aufbau principle) to the Cl atom.
n = 3, l = 1, ml =1, ms = +1/2
Give the set of four quantum numbers that could represent the electron lost to form the Rb ION from the Rb atom.
n = 5, l = 0, ml = 0, ms = ⁺¹/₂
Isotopes differ in the number of
neutrons
Which of the following signs on q and w represent a system that is doing work on the surroundings, as well as losing heat to the surroundings?
q = -, w = -
Describe a pi bond.
side by side overlap of p orbitals
Choose the best Lewis structure for XeI₂.
single bonds and 3 unpaired e⁻ group
Give the hybridization for the S in SO3.
sp²
Give the hybridization for the Br in BrF₅.
sp³d²
Using the VSEPR model, the molecular geometry of the central atom in CF4 is ________.
tetrahedral
The law of the conservation of energy states
that energy can neither be created or destroyed.
Dalton's Atomic Theory states
that matter is composed of small indestructible particles.
Charles's Law
the change in the volume of gas is directly dependent on the the temperature of gas.
metallic bonding
the electrostatic attraction between the positively charged atomic nuclei of metal atoms
heat capacity.
the quantity of heat required to change a system's temperature by 1°C
specific heat capacity
the quantity of heat required to raise the temperature of 1 gram of a substance by 1°C
Dalton's Law
the total pressure of a mixture of gases in a container equals the sum of the pressures of each individual gas.
How many of the following numbers contain 3 significant figures? *0.408 *9.040 *0.0400 *9.05 x 10^24
three
Give the name for SnO.
tin (II) oxide
Write the name for Sn(SO4)2. Remember that Sn forms several ions.
tin (IV) sulfate
Using the VSEPR model, the electron-domain geometry of the central atom in SF₄ is ________.
trigonal bipyramidal
how many significant figures are there in the answer for the following problem? [(143.7-121) x 2.06]/.660 = ?
two
Molecules can be described as
two or more atoms chemically joined together.
A sunburn is caused by overexposure to ________ radiation.
ultraviolet
This equation is used to calculate the properties of a gas under nonideal conditions.
van der Waals equation
The distance between adjacent crests is called
wavelength
What did Rutherford contribute to chemistry?
working under Thomson, Rutherford used the gold leaf experiment to discover the nuclear theory of the atom (subatomic particles ie. protons, neutrons, and electrons).
The orbital diagram that represents the ground state of N is
↑↓ ↑↓ ↑ ↑ ↑ ⁻⁻⁻⁻ ⁻⁻⁻⁻ ⁻⁻⁻ ⁻⁻⁻ ⁻⁻⁻ 1s 2s 2p
