Chemistry
Answers to Chapter 2 Study Questions
1. a) An element is a fundamental substance that cannot be broken down into simpler substances by chemical means. It is composed of atoms.
b) A compound is a pure substance made of at least two elements in a fixed ratio.
c) A pure substance can be an element or a compound. It has a constant composition.
2. a) air = mixture, solutionb) titanium = pure, element
c) oak = mixtured) baking soda = pure, compound
e) oxygen = pure, elementf) 7-Up = mixture, solution
g) wine = mixture, solutionh) carbon monoxide = pure, compound
3. a) elementb) compoundc) elementd) mixture
a) and b) contain molecules.
4. solid: particles not moving and close; liquid: particles moving and close; gas: particles moving and far apart; aqueous (dissolved in water): particles moving and far.
5. chemical: Chlorine reacts with sodium to form NaCl. You could also list the formula of any ionic compound chlorine forms, such as MgCl2, CaCl2, etc.
physical: Chlorine is a pale green gas at room temperature. It’s a nonmetal made up of diatomic molecules.
6. Chemical reactions are frequently accompanied by:
a) mass changes and/or bubbles which show that a gas is involved in the reaction.
b) heat changes; heat is evolved in exothermic reactions; heat is used up in endothermic reactions. Exothermic reactions also often result in the production of light and sound.
c) color changes which often signify a change in chemical composition.
d) the formation of a precipitate (formation of a solid from mixing solutions) which represents the formation of an insoluble substance from soluble substances.
7. a) chemicalb) physicalc) physicald) chemical
e) chemicalf) chemicalg) physical
8. a) physical b)chemical c) physical d)physical e) physical f)chemical g)chemical
8. a) Qualitative: This page is colorful. This page contains questions.
b) Quantitative: This page is 8.5 in x 11 in. The page contains 9 questions.
Theory: The questions on this page will be useful in studying for the test.
9. a) beakerb) Erlenmeyer flaskc) graduated cylinderd) pipet
Chemistry
Answers to Chapter 5 Study Questions
1. a) 6.5 x 102b) 5 x 104 c) 2.07 x 105d) 1.0 x 106e) 5.0 x 104
2. a) liters (L) or cm3, graduated cylinder, buret, or volumetric flask
b) grams (g), balancec) meters (m), ruler or meterstick
3. a) 4b) 5c) 2d) 3 e) 4
4. How many significant figures are there in the following numbers or answers?
a) 3 b) 2 c) 2 d) 2 (answer 7.6) e) 2
5. a) 1.24 x 8.2 = 10. b) 6.78 - 3.3 = 3.5
c) 9.999 + 0.22 =10.22d) (5.67 x 103) x (2.1 x 102) = 1.2 x 102, or 120
6. a)
b) % accuracy error = 13%
7. 275 grams x 1 kg/1000g = 0.275 kg
8. 0.286m x 100cm/1m = 28.6 cm
9. 11.8 g x 1cm3/7.87g = 1.59 ml (note cm3 = ml)
10. a) Container A is the most precise because it has the most number of digits and taking into account the fact that it is off by 2 mL, it never varies by more than 0.08 mL. You might also say that Container C is the most precise because its volumes are reproducible, but you don’t know whether the actual differences are more or less than Container A.
b) Container B is the most accurate, since it is consistently closest to the actual volume.
Chemistry
Answers to Chapters 3 & 4 Study Questions
1. Nuclear Atomic Mass Number of Number of Number of Charge
Symbol Number Number Protons Electrons Neutrons
__18__ __40__ ___18______18______22______0__
__19__ 39 19 18 ____20___ __+1__
16 __36__ ___16______18____ 20 -2
2. Rutherford’s experiment supported the ideas that atoms contain a small dense center (nucleus) and are mostly empty space.
3. and
4. a) MG, Group 2, metal, Period 5b) MG, Group 17, nonmetal, Period 4
c) TM, Period 5d) MG, Group 15, nonmetal, Period 3
e) MG, Group 13, metalloid, Period 2f) ITM, Period 7
g) MG, Group 14, metal, Period 5h) TM, Period 6
5. Group 1 = alkali metals; Group 2 = alkaline earth metals; Group 17 = halogens; Group 18 = noble gases.
6. a) Elements made of molecules: O2, N2, Cl2, or any other diatomic element.
b) Compounds made of molecules: CO2, H2O, NH3, or any other covalent compound.
c) Compounds made of ions: NaCl, MgSO4, or any other ionic compound.
7. a) positive, +1b) positive, +2 c) negative, 1 d) negative, 2 e) positive, +1
8. a) covalent, dinitrogen oxide or dinitrogen monoxideb) ionic, potassium oxide
c) covalent, phosphorus trichlorided) ionic, aluminum phosphate
e) covalent, hydrochloric acidf) ionic, ammonium fluoride
g) ionic, lead(II) nitriteh) covalent, sulfurous acid
9. a) calcium carbonateb) zinc sulfidec) copper(I) hydroxide
d) magnesium perchlorate
10. a) K3PO4b) (NH4)2SO4c) Co(OH)2d) FeN
11. a) PI3b) N2O5c) HClO3
12. a)covalent P2O5 b) ionic Mg(NO3)2c) ionic AgO2 d) ionic KOH
13. a)lead II chlorideb)copper I sulfatec) carbón disulfided)hydrofluiric acide)sodium chlorate
Answers to Chapters 6 & 7 Study Questions
1. a) 3(12.0) + 8(1.01) + 3(16.0) = 92.1 g/mole b) 92.1 g c) 6.02 x 1023 molecules
d) = 20.0 g e) = 8.50 moles
2. a) 4 atoms (one N + 3 H)
b) = 2.41 x 1024 atoms
c) = 4.82 x 1023 atoms
3. Molar mass of NaNO2 = 23.0 + 14.0 + 2(16.0) = 69.0 g/mole
% Na = 23.0/69.0 = 33.3% Na; % N = 14.0/69.0 = 20.3% N; %O = 2(16.0)/69.0 = 46.4% O
33.3% Na, 20.3% N and 46.4% O.
(note, questions 4-5 where intentionally skipped)
6. Chemical reactions are frequently accompanied by:
a) bubbles which show that a gas is one of the products of the reaction.
b) heat changes; heat is evolved in exothermic reactions; heat is used up in endothermic reactions. Exothermic reactions also often result in the production of light and sound.
c) color changes which often signify a change in chemical composition.
d) the formation of a precipitate which represents the formation of an insoluble ionic compound from soluble ionic compounds.
7. a) 4Fe(s) + 3 O2(g) 2 Fe2O3(s)
b) B2H6(l) + 3 O2(g) B2O3(s) + 3 H2O(l)
c) 4PH3(g) + 8 O2(g) 6 H2O(l) + P4O10(s)
8. a) PH3(g) and O2(g)b) H2O(l) and P4O10(s)c) 6
9.a) 6 Li(s) + N2(g) 2 Li3N(s)
b) C3H8(g) + 5 O2(g) 3 CO2(g) + 4 H2O(l)
10.
Chemistry
Answers to Chapters 8 & 9 Study Questions
1.
2.
4. a) 3 Na2CO3(aq) + 2 FeCl3(aq) 6 NaCl(aq) + Fe2(CO3)3(s)
5. a) Fe3+(aq) + 3 OH (aq) Fe(OH)3(s)
b) No Reaction ((NH4)2CO3 and LiCl are both soluble)
c) Ni2+(aq) + S2-(aq) NiS(s)
6. a) S, OR; 2 Li(s) + Cl2(g) 2 LiCl(s)
b) DD, P; Sr(NO3)2(aq) + K2SO4(aq) 2 KNO3(aq) + SrSO4(s)
c) C, OR; 2 C3H6(g) + 9 O2(g) 6 CO2(g) + 6 H2O(l)
d) DD; CaCl2(aq) + 2 NaNO3(aq) No reaction (all products are soluble)
e) SR, OR; Fe(s) + MgSO4(aq) No reaction (Mg is more active than Fe)
f) D, OR; 2 KI(l) 2 K(s) + I2(s)
g) SR, OR; 2 Al(s) + 6 HCl(aq) 3 H2(g) + 2 AlCl3(aq)
h) DD, AB; HNO3(aq) + KOH(aq) H2O(l) + KNO3(aq)
7. a) 1.60 mol CrCl3 x = 4.80 moles HCl
b) 0.450 mol HCl x = 7.80 g Cr
c) 12 mol H2 x = 4.8 x 1024 atoms Cr
d) 3.20 g H2 x = 55.0 g Cr
e) 8.30 g HCl x = 12.0 g CrCl3
theoretical yield = 12.0 g CrCl3
% Yield = x 100% = x 100% = 85.0%
f) 6.0 moles Cr x = 6.0 moles CrCl3
12.0 moles HCl x = 4.0 moles CrCl3; therefore, HCl is limiting
4.0 moles CrCl3 x = 4.0 moles Cr used up.
6.0 - 4.0 = 2.0 moles Cr left over.
g) 13.0 g Cr x = 39.5 g CrCl3
43.8 g HCl x = 63.2 g CrCl3
since 39.5 g < 63.2 g, 39.5 g CrCl3 is produced.
8. a)CH4 + 2 O2 -> CO2 + 2 H20
b) 2AL + 3Cl2 -> 2 ALCL3, synthesis
c)2NH4OH + Cu(NO3)2 -> 2 NH4NO3 + Cu(OH)2 , double replacement
d) 2 Fe + 6 HNO3 -> 2Fe(NO3)3 + 3 H2 single replacement
e) 2 H2O -> 2 H2 + O2 decomposition
9.
10.
11.
Chemistry
Answers to Chapter 11 Study Questions
1. Wavelength is inversely proportional to frequency. Energy is proportional to frequency.
2. The new idea in Bohr's model was that electrons can only exist in specific energy states. Bohr's model included an electron orbiting the nucleus as a planet does the sun; according to the quantum mechanical model, we can only define the probability of finding an electron at a given location. When electrons drop from higher energy levels to lower ones, they give off energy in the form of light. The color of light emitted depends on the energy difference between the levels. The greater the energy difference, the shorter the wavelength of light, the more violet the color.
3. The electron configurations of all Group 1 metals end with a single s electron. When these metals lose this s electron, they acquire noble gas electron configurations which end in completed energy levels. They have a strong tendency, therefore, to lose their final single s electrons. This makes them extremely reactive and the metals with the greatest tendency to lose electrons. Group 17 elements need only 1 p electron to complete their outermost energy levels. They have a strong tendency to gain an electron and thus are the most reactive nonmetals. The energy levels of noble gases are all full so these elements have no need to gain or lose electrons and therefore don't react with anything.
4. a) 2n2, where n = Principal Energy level
b) s = 2 e-, p = 6 e-, d = 10 e-, f = 14 e-
c) 2 e-
5. Does not exist: c) 2d; Increasing energy: 1s < 2s < 3d < 4p < 4f
6. a) s b) d c) p d) f
7. a) ns1b) ns2np5c) ns2np6
8. a) ground stateb) impossible c) excited d) excited
9. a) 1s22s22p63s23p4b) 1s22s22p63s23p64s23d9
10. a) [36Kr] 5s2b) [54Xe] 6s24f145d106p2
11. a) Ba (Period 6, Group 2)
b) He (1s2)
c) As (4s24p3 = Period 4, Group 15)
d) C (2s22p2 = Period 2, Group 14)
e) Cl (Period 3, Group 17)
f) Mn (3d5 = 5th transition metal where 3d is being filled)and also Cr (which is 4s13d5)
12.s: p:
As n increases, the size of an orbital increases since the probability of finding an electron farther from the nucleus increases.
13. a) O: () () ()( )( )
1s 2s 2p
b) Ti: () () ()()() () ()()() () ( )( )( )( )( )
1s 2s 2p 3s 3p 4s 3d
14. Chemical properties are shared within a Group but not within a Period. Group number is a good predictor of chemical properties; Period number is not.
15. a) Ar; Na b) Ca; S
16.
17.
Chemistry
Answers to Chapter 12 Study Questions
1. A chemical bond is the force holding atoms together. Atoms form chemical bonds to acquire a stable number of electrons, which is often the same number of electrons as the nearest noble gas. Covalent bonds involve the sharing of electrons between atoms, so that both atoms act as if they've acquired the shared electrons. In ionic bonds, electrons are NOT shared; a metal atom gives up one or more electrons to form a positive ion, a nonmetal gains one or more electrons to form a negative ion, and then the two ions are attracted to each other due to their opposite charges. Both covalent bonding and ionic bonding are strategies by which an atom attains a stable number of electrons.
2. The number of bonds a nonmetal forms usually = 8 – valence.
3. Covalent bonds are found in molecules (covalent compounds and some nonmetal elements) and in polyatomic ions.
4. a) As < P < Nb) Li < C < Oc) K < Mg < B
5.
6. a) Li+b) Na+c) F
7. a) 1s22s22p6; N3, F, Na+, Mg2+, Al3+; Neb) 1s22s22p63s23p6; S2-, Cl, K+, Ca2+; Ar
8. Ca(s) + I2(s) CaI2(s)
9. a) Beb) Cc) F
10. a) b)c) d)
e)f)g)
h) i)
13.
Chemistry
Answers to Chapter 13 Study Questions
1. P1 = 722 mmHg = 0.950 atm; T1 = 22 + 273 = 295 K; P2 = 1.07 atm; T2 = ?
; T2 = T1 = 295 K = 332 K or 59°C
2. a) PT = PH2 + PH2O; Find PH2O in Table from lab report; at 19°C, PH2O = 16 mmHg
PH2 = PT - PH2O = 756 - 16 = 740. mmHg
b) 740 mmHg = 0.974 atm
3. V1 = 600. cm3; T1 = 25°C = 298 K; P1 = 750. mmHg
V2 = 480. cm3; T2 = 41°C = 314 K; P2 = ?
= 988 mmHg
4 methane = CH4. To find grams, use PV = nRT to calculate n. molar mass (CH4) = 16.0 g/mol
n = ?; V = 28.0 L; T = 68 + 273 = 341 K; P = 2.00 atm.
= 2.00 mol; 2.00 mol = 32.0 g
5. 15.0 g CO2= 0.341 mol CO2; 12.0 g CH4 = 0.750 mole CH4;
At constant T and P, ; = 15.7 L
6.
7.
Chemistry
Answers to Chapter 14 Study Questions
1. In gases, the particles are far apart, move independently in the ideal case, take the shape of their container and can have a variety of volumes. In ideal gases, the particles are not attracted to each other at all. In liquids, the particles are close together giving them a constant volume although liquids take the shape of their container. The particles are mobile and attracted to each other. In solids, the particles are close together, fixed in position and strongly attracted to each other. Solids hold their own shape.
2. The heat of vaporization is endothermic because the bonds between particles in a liquid must be broken in order to separate the particles from each other. The breaking of bonds always requires energy.
3. The boiling point of a liquid is the temperature at which both liquid and gas exist in equilibrium; it is the temperature at which the equilibrium vapor pressure of the liquid is equal to external pressure.
Chemistry
Answers to Chapter 15 Study Questions
1. mass percent =
mass KMnO4 = 1.00 mole = 158 g; mass solution = 158 g KMnO4 + 158 g H2O = 316 g
mass percent = = 50.0 %
2. 335 g solution = 0.466 moles
3. 275 mL solution x = 0.138 moles
4. 250 mL solution x = 29.2 g
5. V1 x M1 = V2 x M2; V1 x 2.00 M = 125 mL x 0.350 M
V1 = 125 mL x 0.350 M/2.00 M = 21.9 mL
6. molarity = ; = 2.50 M