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1-1 The Foundations
of Chemistry (a) Biochemistry is the study of the chemistry of living things.
(b) Analytical chemistry studies the quantitative and qualitative composition analysis of substances.
(c) Geochemistry is the study of the properties and reactions of the substances that compose earth’s
(d) Nuclear chemistry is the study of the properties and reactions of atomic nuclei.
(e) Inorganic chemistry is the study of compounds of elements other than carbon; however, simple
carbon compounds are also included, such as CO, CO2, carbonates, and bicarbonates. 1-3 (a) Matter is anything that has mass and occupies space. An example of matter is your textbook.
(b) Kinetic energy is the energy of a moving object or the energy of an object due to its motion. A
bowling ball has kinetic energy as it is rolling down the lane.
(c) Mass is a measure of the amount of matter in an object. The mass of a penny (a copper coin) is
about 1 gram.
(d) An exothermic process is a process that releases heat energy. The combustion of gasoline is an
exothermic process that is used in automobile engines.
(e) An intensive property is a property that is independent of the amount of material present. Density
is an intensive property. 1-5 Law of Conservation of Matter and Energy: The combined amount of matter and energy available in
the universe is fixed. This law recognizes that the energy released in a nuclear reaction comes from
the conversion of matter into energy. The Law of Conservation of Matter and Law of Conservation of
Energy refer to chemical (not nuclear) reactions and physical changes. In chemical reactions and
physical changes, the quantity of mater has no detectable change and energy is neither created nor
destroyed; energy is only converted from one form to another. 1-7 (a) Since energy can be converted from one type to another, a broad definition of exothermic is that
the reaction releases energy. Since light is a form of energy, the production of light from a
fluorescent light is a release of energy.
(b) In a similar manner, the production of light by a glow-in-the-dark object also releases light a form
of energy. 1-9 (a) Exothermic. The gasoline gives off heat and light during combustion or burning.
(b) Exothermic. The ice cream is changing from a liquid to a solid. Heat must be lost for the
particles to slow down and to freeze. This is the opposite of melting.
(c) Endothermic. The chocolate absorbs heat as it melts or changes from a solid to a liquid.
(d) Exothermic. As the temperature of the water drops, the heat energy is leaving the water and
moving into the surroundings.
(e) Exothermic. Water vapor gives off heat as it condenses. The particles must cool to change from a
gas to a liquid
1-1 (f) Exothermic. The match gives off heat as it burns. This heat can be used to light the wick of a
1-11 (a) Law of Conservation of Matter: There is no detectable change in the quantity of matter during an
ordinary chemical reaction or during a physical change. Examples—(i) when magnesium metal
burns in oxygen, the mass of the product (magnesium oxide) is equal to the sum of the masses of
the magnesium and oxygen that combine; (ii) when ice melts, its mass does not change.
(b) Law of Conservation of Energy: Energy cannot be created or destroyed in a chemical reaction or
in a physical change; it can only be converted from one form to another. Example—in a
hydroelectric plant, the mechanical (kinetic) energy of the falling water is converted into
electrical energy; some of the energy is converted into heat.
(c) Law of Conservation of Matter and Energy: The combined amount of matter and energy
available in the universe is fixed. Example—the energy released in a nuclear reaction comes from
the conversion of matter into energy. 1-13 An incandescent light bulb converts electrical energy into light energy. A considerable portion of the
electrical energy used is converted into heat energy. The Law of Conservation of Energy is observed
since the sum of the heat energy and light energy produced is equal to the electrical energy consumed. 1-15 A homogeneous mixture has uniform composition and properties throughout. Among the examples
given in this exercise, carbon dioxide (f) is the only pure substance. All samples of carbon dioxide
would always contain the same ratio of carbon and oxygen. Examples (a), and (e) are homogeneous
mixtures; examples (b), (c), (d), and (g) are heterogeneous mixtures. The heterogeneous mixtures have
large particles that are suspended (mud, noodles, onion), floating (ice), or that are at the bottom of the
container (chocolate chips, chunks of chicken); therefore, they are not homogeneous mixtures. 1-17 (a) A gaseous element is shown in box (i). The substance contains only one element because only
blue spheres are shown, even though the element is diatomic. The substance is a gas because the
particles have the maximum separation.
(b) A gaseous compound is shown in box (v). The substance is a compound because each particle
contains two elements (two blue atoms and one red atom bonded together). The substance is a gas
because the particles have the maximum separation.
(c) A homogeneous gaseous mixture is shown in box (iv). A mixture is shown because there are two
different types of particles (diatomic blue and a compound made of two blue and one red atom).
The substance is a gas because the particles have the maximum separation.
(d) A liquid solution is shown in box (vi). A solution is a homogeneous liquid mixture. A mixture is
shown because there are two different types of particles (a compound made of one red and two
white atoms, with a second compound made of one red, one blue, and four white atoms). The
substance is a liquid because the particles are much closer than in a gas, but the particles are not
as close as a solid or in a regular repeating pattern as a solid.
(e) A solid is shown in box (ii). A solid is shown because the particles are shown very close together
and are in a regular repeating pattern. A crystalline solid is depicted.
(f) A pure liquid is shown in box (iii). The substance is a liquid because the particles are all the same
(maroon), are much closer than in a gas, but the particles are not as close as a solid or in a regular
repeating pattern as a solid. The liquid happens to be diatomic. The liquid is pure because there is
only one type of particle.
1-2 1-19 (a) Salt and water will form a homogeneous mixture, so to separate the salt from the water, you
would need to evaporate or boil away the water to leave the salt behind.
(b) Iron filings and lead can be separated be using a magnet. Iron is attracted to a magnet, while lead
(c) Elemental sulfur can be separated from sugar by using solubility properties. Sugar is soluble in
water, while sulfur is not. Adding water to the mixture and pouring off the solution, sulfur will be
left. 1-21 (a) Chemical properties are exhibited as matter undergoes changes in composition, whereas physical
properties can be observed in the absence of any such change in composition.
Examples of chemical properties—(i) magnesium can combine with oxygen; (ii) gasoline is
Examples of physical properties—(i) water is a colorless liquid at room temperature; (ii) oxygen
is a gas at room temperature and ordinary pressures; (iii) the melting point of bromine is –7.1˚C.
(b) Intensive properties are those properties that are independent of the amount of material examined,
while extensive properties depend on the amount of material examined.
Examples of intensive properties—(i) magnesium can combine with oxygen; (ii) the melting
point of bromine is –7.1˚C.
Examples of extensive properties—(i) the mass of a sample; (ii) the volume of a sample at
(c) Chemical changes occur when one or more substances react resulting in the formation of one or
more new substances. Physical changes most often involve changes in physical state brought
about by the absorption or release of energy
Example of chemical change—(i) alcohol reacting (burning) in oxygen to form carbon dioxide
Examples of physical change—(i) ice melting to water with the absorption of heat; (ii) steam
condensing to liquid water with the release of heat.
(d) Mass is a measure of the amount of matter in an object, while weight is a measure of gravitational
attraction of the earth for an object.
An object having a mass of 454 g has a weight of one pound on Earth and the same object having
a mass of 454 g would have zero weight in a zero gravitational field. 1-23 (a) Chemical process. Iron is combining with oxygen in the presence of water to form a new
(b) Physical process. Water as a solid (ice) is changing to liquid water. Melting does not change the
(c) Chemical process. The wood is changed by the combustion or burning into ash, which is a new
substance with none of the properties of the wood.
(d) Chemical process. The components of the potato are broken down into substances that can be
absorbed by the digestive tract.
(e) Physical process. Dissolving sugar in water does not change the composition. If the water in the
solution were allowed to evaporate, the sugar would be left behind. 1-25 (a) Kinetic energy
(b) Potential energy
(c) Potential energy (d) Kinetic energy
(e) Kinetic energy
(f) Potential energy
1-3 1-27 Both physical and chemical changes have taken place. The outer edge of the sugar cube melted (a
physical change), then the sugar began to burn or oxidize (a chemical change). The heated portion has
a different color and odor. The brown portion contains carbon left as the sugar decomposes. 1-29 (a) 6.50 x 102
(d) 8.600 x 103 1-31 (a)
(f) 1-33 Circumference = πd = (3.141593)(7.41 cm) = 23.3 cm 1-35 (a) 106 1-37 5.31 cm = 5.31 x 10-2 m, 53.1 mm, 5.31 x 10-5 km, and 5.31 x 104 micrometers 1-39 4 qt
? $ = 14 gal x 1 gal x 1.056 qt x
1L 1-41 ? cm = 8.25 in x (b) 6.30 x 10–2
(e) 1.6 x 104 (c) 8.60 x 103
(f) 1.0010 x 10–1 Exact (the result of counting)
Exact (the result of counting)
Exact (counted to the nearest penny)
Not exact (obtained by measurement)
Not exact (obtained by measurement)
Exact (the result of counting) (b) 10-3 (c) 10-2 (d) 10-1 2.54 cm
1 in = 20.955 cm (e) 103 ? cm = 6.25 in x (f) 10-9 2.54 cm
1 in = 15.875 cm 21.0 cm x 15.9 cm = screen size
1-43 ? g = 10.25g + 5.5654g x 105.4g = 121.2 g 1-45 ?g = 1-47 ? g = 3.00 L x 1-49 (a) mass of water = 92.44 g – 78.91 g = 13.53 g water
volume of water = 13.53 g x 1.0000 g = 13.53 cm 8.92 g
x 24.4 cm x 11.4 cm x 7.9 cm = 19601g = 2.0 x 10 g
1000 cm3 1.0056 g
1 cm3 3
= 3.02 x 10 g (if three L has 3 sig. figs.) (b) mass of unknown liquid = 88.42 g – 78.91 g = 9.51 g
density of unknown liquid = V =
= 0.703 g/cm
1.049 g soln. 40.0 g acetic acid
= 104.9 ⇒ 105 g acetic acid
100 g soln. 1-51 ? g = 250.0 mL x 1-53 (a) ? K = 245° C + 273.15° = 518 K
€ 1-4 (b) ? ° C = 25.2 K – 273.15° C = -247.95° C = −248.0˚C !
(c) ? ° F = # –42°C x
& + 32˚F = -43.6 = –44°F
(d) First convert °F to °C, then ° C to K.
? °C =
x (110.0°F – 32˚F) = 43.3°C with only 2 sig figs
? K = (43°C + 273.15°) = 316K
(a) ? ° F = $20˚C x
' + 32°F = 68°F
1.0˚C & so 20˚C or 68°F is higher than 20°F "
(b) ? ° F = $100˚C x
' + 32°F = 212°F so 100˚C or 212°F is higher than 180°F
(c) ? ° C =
x (100°F – 32°F) = 23.6° C so 60°C is higher than 100°F or 23.6° C
(d) ? ° F = %−12˚C x
( + 32°F = 10.4°F so 20˚F is higher than –12° C or 10.4°F
€ 1-57 He: ? ° C = 4.2 K – 273.15°C = −269.0˚C
€? ° F = %
x (− 269.0˚C)( + 32°F = 452.2°F
N2: ? ° C = 77.4 K – 273.15°C = −195.8°C
x (−195.8˚C)⎞⎠ + 32°F = −320.4°F
€? ° F = ⎝
1.0°C 1-59 If °F = 2x and ° C = x, Then 2x = 1.8 x + 32 0.2x = 32 x = 160 to check to see what °F are if ° C = 160
? °F = $
x 160˚C ' + 32°F = 320°F
& 1-61 Temperature change = 32.0°C – 10.0°C = 22.0°C
€ ? J = mass of substance x specific heat x temperature change
= 78.2 g x
x 22.0°C = 7.20 x 10 J
g • °C 1-63 ? J = mass of substance x specific heat x temperature change
=€ 15.5 g x 4.184J
x (38.2°C – 90.0°C) = −3.36 x 10 J or 3.36 x 10 J must be removed
g • °C
1-5 € 1-65 0.997 g
? g H2O = 245 mL x 1 mL = 244 g H2O
? J = 244 g x
? 1-67 4.184J
x (85.°C – 25.°C) = 6.1 x 104 J
g•°C 1 kJ
6.1 x 10 4 J
x 1000 J
2.00 min = 30.5 = 31 kJ/min (a) ? g calcium carbonate = 75.45 g sample x
(b) ? g sample = 18.8 g calcium carbonate x 25.8g calcium carbonate
= 19.5 g calcium carbonate
100g sample 100g sample
= 72.9 g sample
25.8g calcium carbonate € 1-69 103 m
Radius of earth’s orbit (m) = 1.5 x 108 km x 1 km = 1.5 x 1011 m
Radius of hydrogen atom (m) = 0.37 Å x 1 Å
= 3.7 x 10-11 m
Ratio = 1-71 1.5 x 1011 m
= 4.1 x 10
3.7 x 10 m ? km
1.609 km "104.6 km %
' = 1.0 x 10 km/h
€ 1.00 mi
60 min 1.609 km
€= 19.7 km/h
1 min 1000 m 100 cm
= 547 cm/s
(c) 1500. m x
= 274 s ⇒ 4 min 34s
19.7 km 1000 m
1-75 If you wanted the pot or pan to heat up quickly, you would select material that has a small specific
heat value. If you wanted the pot or pan to retain its temperature
once it is hot, then you would select
€ material that would have a higher specific heat. Most individuals desire some of both of these traits but
feel that the first is the more important. 1-73
€ 1-77€ The density of newly-minted penny, g/cm3 :
= 0.027)(8.72 g/cm3) + (1.000 – 0.027)(7.14 g/cm3)
= 0.24 g/cm3 + 6.95 g/cm3 = 7.18 g/cm3 1-79 We know that water must be more dense, because ice floats in water. 1-81 The correct answer is (a). The particles would be the same size but closer together at the lower
1-6 1-83 (a) Let x = the reading on the Celsius thermometer = the reading on the Fahrenheit thermometer
x ° F = ⎝ x°C x
+ 32°F or, without units,
1.0°C ⎠ 1.8
x = 1.0 x + 32 ; 0.8x = -32 ; x = −40°C , x = −40°F
(b) 2x = 1.0 x + 32 ; 2x = 1.8x + 32 ; 0.2x = 32; x = 160°C , 2x = 320°F
–x = 1.0 x + 32 ; 2.8x = –32; x = −11.4°C , -x = +11.4°F 1-85 The balloons filled with substances that are lighter than air will float. Assuming that the balloons are
all the same volume, the He and Ne balloons should float, while the Ar and Kr balloons will sink. 1-87 Students know many chemical terms before they begin to read this textbook. A few of the terms that
they are likely to know are: compound, distillation, and chemical reaction. 1-89 (a) A gas is shown in boxes (iii), (iv), (vii), and (ix). The particles are in the gas phase because the
particles have the maximum separation and are in a random arrangement.
(b) A liquid is shown in boxes (v -the blue particles in the top right of the diagram) and (viii). These
are liquids because the particles are much closer than in a gas, but the particles are not as close as
a solid or in a regular repeating pattern as a solid.
(c) A solid is shown in boxes (i), (ii), (v- the brown particles in the bottom left of the diagram), and
(vi). A solid is shown because the particles shown very close together and are in a regular
(d) An element is shown in boxes (i), (iv), and (vi). The particles are all the same color (blue), even
though the blue atoms in box (vi) are shown as diatomic particles.
(e) A compound is shown in box (iii). The compound depicted here is composed of one blue atom
and one brown atom, since one blue is attached to one brown throughout. The arrangement of the
compound shows that it happens to be in the gaseous state.
(f) A mixture is shown in boxes (ii), (vii), and (ix). Mixtures contain two or more different types of
particles. Boxes (ii) and (vii) contain both blue and brown particles. Box (ix) contains diatomic
blue particles and single brown particles. Boxes (v) and (viii) show two types of particles, but
these particles are not yet mixed.
(g) A pure substance is shown in boxes (i), (iii), (iv), and (vi). A pure substance contains particles
that are identical. 1-91 Chlorine is an element. The atom of chlorine could be represented by a single sphere or 1 Cl. The
diatomic molecule of chlorine would be represented by two adjoining spheres, to depict 2 atoms of
chlorine in the molecule. These two differ in that the molecule has 2 atoms.
It is impossible to draw an atom of methane. Methane is a compound. A single methane molecule is
made up of 1 atom of carbon and 4 atoms of hydrogen. One could draw a carbon or hydrogen atom
from the molecule, but there is no atom of methane; the molecule is the smallest unit.
Methane is a compound, while chlorine is an element.
1-7 1-93 The change from solid, black carbon and colorless oxygen gas to colorless carbon dioxide gas is a
chemical change because a new substance is formed (carbon dioxide). The colors and states of matter
of these substances are physical properties. 1-95 To say the length of an animal is 51 doesn't give any units, so the understanding is very limited. An
animal of 51 meters is very different from an animal of 51 millimeters. One must use units that others
understand. It is also a problem to give a measurement with a unit that has no understanding. To know
the width of a room is 7.36 bleams is also no understanding, since there is no understanding of the unit
'bleams'. 1-97 ? K = 5500° C + 273.15 = 5773K
? ° F = $5500°C x
' + 32˚F = 9932 °F or 9.9 x 10 °F in 2 significant digits
1.0°C & The astronomer is referring to degrees Fahrenheit, but is confused about the significant numbers if
he/she is referring to 5500° C. The astronomer could be the larger number, which does give 10,000
? ° F = $6000°C x
' + 32˚F = 10,832 °F or 1 x 10 °F in 1 significant digit
1.0°C & € € 1-8 2 Chemical Formulas
Stoichiometry 2-1 (a) Stoichiometry is the description of the quantitative relationships among elements in a compound
and among substances as they undergo chemical change.
(b) Composition stoichiometry describes the quantitative relationships among elements in
compounds, e.g., in water, H2O, there are 2 hydrogen atoms for every 1 atom of oxygen. Reaction
stoichiometry describes the quantitative relationships among substances as they undergo chemical
changes. (Reaction stoichiometry will be discussed in Chapter 3.) 2-3 The common ions for each formula unit is listed below:
(a) MgCl2 contains Mg2+ and Cl- ions
(b) (NH4)2CO3 contains NH4+ and CO32- ions
(c) Zn(NO3)2 contains Zn2+ and NO3- ions 2-5 Ethanol -CH3CH2OH Methanol-CH3OH (space-filling; ball-and-stick) (space-filling; ball-and-stick) Both are composed of hydrogen, carbon, and oxygen. Both have an oxygen and hydrogen on the end.
The ethanol molecule has an additional carbon and two hydrogens.
2-7 Organic compounds are those that contain carbon-to-carbon bonds, carbon-to-hydrogen bonds, or
both. Organic formulas given in Table 2-1 include: acetic acid- CH3COOH, methane- CH4, ethaneC2H6, propane- C3H8, butane- C4H10, pentane- C5H12, benzene- C6H6, methanol- CH3OH, ethanolCH3CH2OH, acetone- CH3COCH3, diethyl ether- CH3CH2COCH2CH3. 2-9 Compounds from Table 2-1 that contain only carbon and hydrogen and are not shown in Figure 1-5:
Ball and stick model
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