Chemistry
Final Exam Problems Review Answers
1) For the following
reaction, determine how many moles of oxygen gas (O2), are needed to
react with each of the given moles of glucose (C6H12O6
) listed below:
C6H12O6
(s) +
6 O2(g) à 6 CO2(g) + 6
H20(l)
Ans
- 5
moles C6H12O6 30 mol O2
- 0.05
moles C6H12O6 0.3 mol O2
- 2.5 E
-6 moles C6H12O6 1.5 E -5 mol
O2
2) A mass of 46.0
grams of solid sodium reacts with 38.0 g of fluorine gas to produce 84.0 g of
sodium fluoride [NaF(s)].
a. Write a balanced equation for
this reaction. 2 Na + F2 à 2 NaF
b. Calculate the number of moles
of each substance. 2.00 mol Na, 1.00 mol F2, 2.00 mol NaF
3) If 17.5 g of zinc
metal are reacted with phosphoric acid [H3PO4(aq)], then
zinc (II) phosphate [Zn3(PO4)2 ]
and hydrogen gas are produced. Zn(s) + H3PO4(aq) à Zn3(PO4)2 + H2
a.
How many moles of phosphoric acid are required? 0.178
mol H3PO4
b.
If the phosphoric acid is a 3.00 molar solution, how
many liters are needed? 0.0593 L H3PO4
c.
What mass of zinc phosphate will be produced? 34.4 g Zn3(PO4)2
d.
How many moles of hydrogen gas will be given off? 0.268 mol H2
4) What volume of
0.60 M copper(II) sulfate [CuSO4(aq)] will
react with 45 mL of 1.5 M sodium hydroxide [NaOH(aq)] to form copper(II)
hydroxide [Cu(OH)2(s)] and sodium sulfate [Na2SO4(aq)]?
CuSO4(aq)
+ 2 NaOH(aq) à Cu(OH)2(s) +
Na2SO4(aq)
Ans: 56 mL CuSO4
5) Explain why the
percent yield of a product in a reaction is usually less than 100.
If all chemicals are 100% pure, and all react
with no atoms lost, the yield is 100%. Ususally some atoms are lost or impurities are
present. Therefore the yield is
frequently less than 100%. If the
apparent yield is more than 100%, there must be some impurity present in the
product.
6) If 5.45 g of
potassium chlorate [KClO3(s)] are decomposed to form potassium choride [KCl(s)], then 1.95 g of
oxygen gas (O2) are also given off. 2 KClO3 à 2 KCl +
3 O2
a.
Calculate the theoretical yield of oxygen. 2.13 g O2
b.
Calculate the percent yield of oxygen. % yield = actual
yield = 1.95
x 100 = 91.5%
Theoretical yield 2.13
c.
Explain why the percent yield of oxygen is less than
100.
Some oxygen may have
escaped or dissolved in the water, or the KClO2 may not have reacted
completely.
7) a.
How many grams of FeI2(s) can be formed when 25.7 g of Fe(s)
reacts with 105 g of I2(s)?
b. Which
reactant is the limiter?
c. Which reactant is the
excess? How much excess (in grams)?
8) Using a molecular
model, explain why a gas can be easily compressed, while a liquid and a solid
cannot.
Space between molecules is larger in a gas than
in a solid or liquid, thus the molecules can be pushed closer together.
9) The volume of a
gas is often referred to as one of the properties that can be measured. Is this volume simply the sum of all of the
individual molecular volumes? Explain
your answer.
The volume is not simply the sum of all of the
individual molecular volumes. The
molecules of a gas are much smaller than their entire container. The volume of the gas is the volume a given
amount of gas occupies at a given temperature and pressure.
10) Explain why the air pressure is greater on
A
column of air over a given area at sea level is taller and more massive than a
column of air over an equivalent area of the volcano. Thus air pressure is greater on the beach.
11) Carbon dioxide
does not exist in the liquid state unless the pressure is at least 5.1 atm. Convert this
pressure into units of: a.
torr 3900 torr
b. kPa 520 kPa
12) If air pressure
is reduced from normal sea level values (this happens at higher elevations),
the boiling point of a liquid falls. For instance, water boils at only 95
degrees Celsius if the atmospheric pressure is 634 mm Hg. Explain this in terms of intermolecular
forces and vapor pressure. Also, convert
this pressure into units of:
a.
atm 0.834 atm
b.
13) Hydrogen gas is
collect by bubbling it through water.
Calculate the partial pressure of the hydrogen gas if:
a.
the total pressure is 94,000 Pa, and the partial
pressure of water is 1200
Ans: 94,000 Pa - 1200 Pa
= 92,800 Pa
b.
the total pressure is 100.3 kPa,
and the partial pressure of water is 2600.
Ans: 100,300 Pa - 2600
Pa =
97,700 Pa
c.
Whose gas law helps to solve this problem?
14) In a flask that has a volume of 273 dm3, you
have a sample of two noble gases: neon and xenon. The partial pressure of the neon is 96,950
Pa, and the partial pressure of the xenon is 1.025 atm. What is the total pressure (in kPa) exerted by
these two gases? Ans: 200.8 kPa
15) Explain why you
must convert the temperature to an absolute temperature scale (such as a Kelvin
scale) when you wish to use the direct proportionality of the
temperature-volume law (Charle’s Law).
You
need to convert from Celsius degrees to kelvins to
get a plot of temperature versus volume that intercepts the x-axis at the
origin.
16) Change the
following volumes of gases from the conditions given to the new
conditions. Assume that the pressure and
amount of gas is constant.
a.
85 cm3 at 61 degrees C to 35 degrees C
Ans: 78 cm3
b.
7.3 dm3 form 228 degrees C to –48 degrees
Celsius
Ans: 3.3 dm3
17) An anaesthesiologist is about to administer gas to a
patient. The gas has a temperature of
22.4 degrees C. When the gas enters the
patient’s body, it is warmed to a temperature of 37.2 degrees C. Assuming that the gas does not undergo a
change in pressure, what percentage increase in volume does the gas experience
as it reaches the new temperature?
Ans:
1.05 L
There is a 5.0% increase.
18) A sample of
carbon dioxide has a volume of 2.0 dm3 at a temperature of –10
degrees C. What volume will this sample
have when the temperature is increased to 110 degrees C? Assume that the pressure does not change and
that no carbon dioxide leaks from the sample. Ans: 2.9 dm3
19) Change the following volumes of gases from the
conditions given to the new conditions.
Assume that the temperature and amount of gas is constant.
a.
1.15 E3 cm3 at 75.2 kPa
to 14.0 kPa
Ans: 6.1 E3 cm3
b.
94.7 dm3 at 1.00 kPa
to 100.0 kPa
Ans: 0.947 dm3
20) A 12.7-L sample
of gas is under a pressure of 9.3 kPa. What will be the pressure of the gas when the
volume increases to 20.1 L (assume the temperature is held constant).
Ans: 5.9 kPa
21) A sample of nitrogen gas having a volume of 2.44 dm3 is
collected at a pressure of 67.4 kPa. What volume will the gas occupy when its
pressure is changed to 145.1 kPa if the temperature
is held constant?
Ans:
1.18 dm3
22) What did Avogadro
assume to be true about equal volumes of gases if they are held at the same
temperature and pressure?
Under
the same conditions of temperature and pressure, the volumes of reacting gasses
are in small whole number ratios.
23) Which of the
following sample of gases occupies the largest volume, assuming that each
sample is at the same temperature and pressure—50.0 g of neo, 50.0 g of argon,
or 50.0 g of xenon?
Since
neon has the smallest molar mass, 50.0 g of neon has the largest number of
atoms and therefore the largest volume.
24) What volume of carbon dioxide gas contains the same
number of oxygen atoms as 250.0 cm3 of carbon monoxide gas, if each gas sample
is measured at the same temperature and pressure?
Since
CO2 has twice as many O atoms as CO, 250.0 cm3 of CO2
has twice as many O atoms as 250.0 cm3 of CO. Therefore, 125.0 cm3
of CO2 has the same number of O atoms as 250.0 cm3 of CO.
25) Mathematically,
the ideal gas law may be expressed as PV = nRT. What do each of the letters in this equation
represent? Identify the variables and
constants.
Variables:
P-Pressure T-absolute temperature, V- volume, n-moles
Constant: R-ideal gas law constant
26) What is the volume in dm3 of 12.0 g of
nitrogen gas if the gas is measured at a pressure of 125 kPa
and a temperature of 45 degrees C?
Ans: 9.07 dm3
27) What mass of carbon dioxide will occupy a volume of 5.5
L at a temperature of 5 degrees C and a pressure of 75 kPa?
Ans: 7.9 g CO2
28) What is the temperature of a 0.00893 mol sample of neon
gas that has a volume of 302 mL and a pressure of 715
torr?
Ans:
388K or 115 degrees C
29) A balloon is filled with helium to a volume of 12.5
L. If the temperature of the gas is 25
degrees C and the pressure is 101 kPa, How many moles
of helium are in the balloon? How many
helium atoms are in it?
Ans: Ans: 3.07 E 23 He atoms
30) What is the value
of the standard temperature expressed in degrees C and in Kelvin?
O degrees C or 273 K
31) What is the value
of standard pressure in atm and kPa?
1.00 atm or 101.325 kPa
32) What volume will
25.0 g of ammonia (NH3) occupy at STP?
Ans: 32.9 dm3
33) Cooks sometimes use the fermentation of glucose to
produce the gas required to make bread rise, as shown by the following chemical
equation:
C6H12O6
(s) +
2 O2(g) à 2 CH3COOH(l) + CO2(g) + 2 H20(l)
What mass of glucose is required to produce 150. cm3
of carbon dioxide gas measured at STP?
Ans: 0.60 g C6H12O6
34) When coke (almost
pure carbon) is burned in th presence of air, the
product is carbon dioxide in the following equation:
C(s) + O2(g) à CO2(g)
How many liters of carbon dioxide reproduced from burning
750. g of coke with an excess supply of oxygen?
Assume that the carbon dioxide is measured at STP.
Ans: 1400 L CO2
35) What is meant by the term ideal gas? What is a real gas?
An
ideal gas obeys the five conditions of the kinetic molecular theory; a real
gas does not obey all these conditions.
36) Gases tend to
follow the ideal gas law more closely when their pressure is low than when it
is high. Considering what happens to gas
volume under high pressure, why do you suppose this trend is true?
The
volume of a gas decreases when the pressure increases, so the relative volume
occupied by the gas molecules themselves increases. Under these conditions, a gas violates one of
the basic postulates of the kinetic molecular theory (relating to molecular
volume).
37) According to the kinetic molecular theory, what is
different about a sample of xenon gas at 25 degrees C and another sample at 100
degrees C.
The
atoms are moving faster and have more kinetic energy at 100 degrees C.
38) According to the kinetic molecular theory, what is the
relationship between the average kinetic energy of a sample of gas molecules
and their temperature? Does it matter
what scale you use to express this temperature?
The
average kinetic energy of a sample of gas molecules is directly proportional to
its temperature, provided that the temperature is measured on an absolute
scale.
39) Suppose you have two balloons filled with air. The smaller balloon has a volume of 500. mL, while the larger balloon has a volume of 1500. mL. Both balloons
have the same temperature and pressure.
In which balloon (if either) are the nitrogen molecules moving faster,
the larger balloon or the smaller? How
do you know?
On
average, the nitrogen molecules are moving with the same velocity in both balloons because they are at the same
temperature.
40) In an experiment, it takes an unknown gas 1.5 times
longer to diffuse than the same amount of oxygen gas (O2). Find the molar mass of the unknown gas.
Ans:
x = 72 g/mol
41) Suppose a gas
diffuses 1.41 times as fast as sulfur dioxide at the same temperature and
pressure. What is the molar mass of the
unknown gas?
Ans: 32.3 g/mol
42) Two beams of
ions—germanium-74 and selenium-74—pass between a pair of charged plates. The ions have the same charge, but the
velocity of the germanium is twice as large as the velocity of the selenium
ions. Identify which beam will be
deflected more and explain your reasoning?
The
selenium ions are deflected more because they have a smaller velocity and
therefore have less momentum.
43) Compare and contrast the subatomic particles contained
in the following atoms or ions
(identify number of protons, neutrons, and electrons of
each).
a.
a barium-137 atom and a barium-137 ion with a 2+
charge
barium – 137 atom 56
protons 81 neutrons 56 electrons
barium – 137 ion 56
protons 81 neutrons 54 electrons
b.
a molybdenum-96 atom and a molybdenum-96 ion with a 6+
charge
molybdenum – 96 atom 42
protons 52 neutrons 42 electrons
molybdenum – 96 ion 42 protons
52 neutrons 36 electrons
c.
a phosphorus-31 atom and a phosphorus-31 ion with a 3-
charge
phosphorus – 31 atom 15
protons 16 neutrons 15 electrons
phosphorus
– 31 ion 15 protons
16 neutrons 18 electrons
44) List which of the three types of radiation—alpha, beta,
or gamma—each of the following describes.
a.
is not deflected by a magnet gamma
b.
has a negative charge beta
c.
moves with the greatest speed gamma
d.
consists of ions alpha
e.
is similar to light rays gamma
f.
consists of the same particles as cathode rays beta
45) Describe the problem with
The
electron must be moving around the nucleus.
However, a charge moving in a circle radiates energy. If that is true of the electron, it would
lose energy and spiral into the nucleus, and the atom would collapse.
46) A proton beam, an electron beam, and a neutron beam pass
between two charged plates. Assuming the
velocities to be the same, draw a diagram showing the paths of each type of
subatomic particle.
+
e-
n
p
-
47) How are
chlorine-37 and calcium-40 similar?
Both
contain 20 neutrons in the nucleus.
48) List the number
of protons, neutrons, and electrons in each of the following ions.
a. 7 Li+ p 3
n 4 e 2
b. 24 Mg2+ p 12
n 12 e 10
c. 27 Al3+ p 13
n 14 e 10
d. 31 P 3- p
15 n 16
e 18
e. 127 I- p
53 n 74 e 54
49) A very bright
line in the brightline spectrum of sodium has a
wavelength of 5.90 E –7 m. What is the
frequency of this line?
Ans:
5.08 E 14 Hz
50) A radio station
broadcasts at a frequency of 105.4 MHz.
What is the wavelength of this electromagnetic wave?
Ans: frequency (n) à
3.08 E 15 Hz wavelength (l) à
9.73 E -8 m
Use c = l n and E
= h n
51) What must be done
to a hydrogen atom to change its 2s electron to a 3s electron? What happens when a hydrogen atom with a 3s
electron becomes a hydrogen atom with a 2s electron?
A
specific amount of energy must be added to the atom. The same amount of energy
is given off in the form of a photon.
52) What orbital will the sixth electron of ground-state
nitrogen occupy?
Ans: 2p
53) What are the three quantum numbers that describe an
orbital?
n
(energy level or size)
l (energy sublevel or shape), and
m (orbital or orientation)
54) Which of the following statements about orbitals is false?
a.
Orbitals are
distributed in space round the nucleus.
b.
Orbitals are
regions of space in which electrons are likely to be found.
c.
Orbitals show the
path of the electron.
d.
Orbitals are part
of one model for atomic structure.
Ans: c is false
55) Write the
electron configurations for the following elements: arsenic, krypton, bromine,
and phosphorus.
As 1s22s2p63s23p64s23d104p3
Kr 1s22s2p63s23p64s23d104p6
Br 1s22s2p63s23p64s23d104p5
P 1s22s2p63s23p3
56) Write the electron
configurations for the following ions: K+, O2-, Br 1-
K+ 1s22s2p63s23p6
O2- 1s22s2p6
Br1- 1s22s2p63s23p64s23d104p6
57) Name the elements
that correspond to each of the following electron configurations. (Assume all
are neutral atoms in their ground states.)
a.
1s22s22p1 boron
b.
1s22s2 beryllium
c.
1s22s22p63s23p2 silicon
58) Write the symbol for the elements found in
the following locations.
a.
period 3, Group 3A Ans:
Al
b.
period 1, Group 8A Ans:
He
c.
period 4, Group 2B Ans:
Zn
d.
period 6, Group 5A Ans:
Bi
59) At room
temperature, nitrogen is a nonmetallic gas, and bismuth is a solid metal. Why are they both in Group 5A?
They both have 5
electrons in their outermost s and p orbitals.
60) Name the group of
elements in the periodic table that has the following outer electron
configurations.
a.
s2 alkaline earth metals
b.
s2p5 halogens
c.
electrons filling the d orbital transition
elements (transition metals)
d.
s2p6 nobel gases
61) Use the periodic table to determine the electron
configuration of the valence electrons for the following representative
elements. (The principal quantum number is
the same as the period number, and the total number of s and p
electrons is the same as the Group number.)
a. Ca
b. As
c. Cs
d. Ne
e.
Ans: a. 4s2 b. 4s24p3 c. 6s1 d. 2s22p6 e.
6s26p4
62) Predict the common ions formed when atoms of the
elements listed gain or lose electrons.
Then name the noble gas with which the ion is isoelectronic.
a. Mg b. Cl c. Al
d. K e. S f.
Ba g.
P
Ans: a. Mg 2+ isoelectronic
with neon
b. Cl- isoelectronic
with argon
c.
Al 3+ isoelectronic
with neon
d.
K+ isoelectronic
with argon
e.
S 2‑ isoelectronic
with argon
f.
Ba 2+ isoelectronic
with xenon
g.
P 3- isoelectronic
with argon
63) For which of the following elements do you expect there
to be a very large increase, going from the 2nd to the 3rd
ionization energy (IE): Na, Mg, Al?
Magnesium.
Between the second and third electron taken away, there will be a large
increase in IE because the 3rd electron is a core electron.
64) What trend is
observed for the atomic radius of atoms
a.
going down a group?
Ans:
Increases due to adding additional energy levels and the shielding effect of
the core electrons.
b.
Going left to right across a period?
Ans: Decreases.
Additional electrons are added to the SAME energy level. Increased nuclear charge with the same
distance and no additional shielding will draw the outer level electrons in
more tightly.
65) How does core electron shielding affect the attraction
of the nucleus for the valence electrons?
Electron
shielding lessens the force of attraction between the nucleus and valence
electrons.
66) What is ionization energy, and how is it determined?
Ionization
energy is the amount of energy required to remove an electron from a gaseous
atom or ion. It can be determined by bombarding gaseous elements with electrons
of known energy.
67) Sodium reacts
with water to form a basic solution, as the following equation shows.
2
Na(s) +
2 H20(l) à 2 NaOH(aq) + H2(g)
Write the equation for the reactions of the following metals
with water:
a.
K
2 K (s) + 2 H20(l) à 2 KOH(aq) + H2(g)
b.
Li
2 Li(s) + 2 H20(l) à 2 LiOH(aq) + H2(g)
c.
Ca
Ca(s) + 2 H20(l) à Ca(OH)2(aq) + H2(g)
68) Use the
electron-sea model to explain the high electrical conductivity of metals.
The
electrons are mobile and can carry an electrical current (the valence electrons
can move easily between the metallic atoms.
69) Explain why fire involving an alkali metal cannot be put
out with water.
They
react with water and with oxygen in the air.
70) What is the most
reactive element?
Ans: fluorine
71) Briefly explain
how the potential energy changes as two hydrogen atoms approach each other.
As
the hydrogen atoms get closer, the attractions get larger and the potential
energy falls. At some point, the potential energy is at a minimum. When the atoms get closer still, the effect
of repulsions causes the potential energy to rise.
72) How does an ionic
bond form?
An
electron (or electrons) is (are) transferred from a metal atom to a nonmetal
atom and the resulting cation and anion are attracted
by electrostatic forces.
73) How does a
covalent bond form?
One,
two, or three electron pairs are shared by two atoms.
74) What is a crystal lattice?
The
structure representing the position that ions occupy in an ionic compound.
75) What types of elements usually form an ionic bond?
metals and nonmetals
76) What types of elements usually form a covalent bond?
nonmetals with other nonmetals
77) Classify bonds in
each of the following substances as nonpolar
covalent, polar covalent, or ionic.
a.
K2O
ionic
b.
BeO ionic
c.
NH3 polar
covalent
d.
KCl ionic
e.
CBr4 polar
covalent
f.
N2 nonpolar covalent
g.
CS2 nonpolar covalent
78) Show the partial
distribution in the following bonds:
a.
H—Br
b.
N—O
c.
Be—H
d+ d-
18. a. H—Br
d+ d-
b. H-C
d+ d-
c. C
– Cl
d- d+
d. N-H
d+ d-
e. N-O
d+
d-
f. Be-H
79) Describe the arrangement of the electron pairs for each
of the following molecules and predict their shape.
a.
CCl4
b.
SiH4
c.
AsH3
d.
H2Se
e.
CS2
Answers
for 79-80:
a. tetrahedral 4 bonded pair of electrons
b. tetrahedral 4 bonded pair of electrons
c. trigonal
pyramidal 3 bonded pair and one lone pair of
electrons
d. bent 2 bonded pair and 2 lone pair
of electrons
e. linear 2 bonded pair (double covalent
bonds around the
Polarity
of 79 a-e:
a. nonpolar
b. nonpolar
c. polar
d. polar
e. nonpolar
bonds, polar
80) Show the bond
dipoles and predict the molecular shape of each of the molecules in question
79.
Ans:
see above
81) On a molecular level, describe
dipole-dipole intermolecular forces.
Dipole-dipole forces
between molecules have medium strength.
They are formed when each molecule has a definite dipole (one end of the
molecule having a partial positive charge and one end of the molecule having a
partial negative charge). The partial
positive end of one molecule will be attracted to the partial negative end of a
different molecule, causing an attractive force between the two molecules.
82) Explain the origin of dispersion forces
between two molecules of nitrogen, N2.
Even though N2
is considered to be a nonpolar molecule (equal pull
of each N on the shared pairs of electrons), the electrons are constantly in
motion. At a given second of time, the
shared pair may be closer to one of the nitrogen atoms, creating a temporary
dipole (instantaneous dipole). At that
moment, one end will have a temporary partial positive charge while the other
end has a temporary partial negative charge.
For a moment, this creates at attractive force between one N2 molecule
and another N2 molecule.
83) Describe the relationship between the
polarity of individual molecules and the nature and strength of intermolecular
forces.
Polar molecules will
have much stronger intermolecular forces than nonpolar
molecules. In fact, the stronger the polarity
of the bonds within the molecule, the stronger the intermolecular forces will
be. For example, dipole-dipole forces
have medium strength, but molecules with much stronger polar bonds, such as H
attached to F, O, or N, will create a very strong dipole-dipole interaction
that we call hydrogen bonds (but not true chemical bonds).
84) How do dipole-dipole forces
differ from H bonds?
Dipole-dipole forces
are medium strength IMF. H bonds are
also dipole-dipole interactions, but on a much stronger level because of the
strong attraction of F, O, and N for the shared pair of electrons when bonded
with H.
85) Place the following in order
of strength from least strength to greatest strength: all chemical bonds, dipole-dipole, H bond,
Ans:
All intermolecular
forces (
86) Of the following types of
solids, order these from least structured to most structured: ionic, covalent molecular, covalent network,
amorphous, metallic.
amorphous, covalent
molecular, ionic, metallic, covalent network
87) What effect do intermolecular
forces have on the boiling point of a liquid?
The stronger the
intermolecular forces, the higher the boiling point. (The more attraction the molecules have for
each other, the less likely they will convert to gas, so more energy will have
to be added to the system to reach the boiling point).
88) What is the relationship
between boiling point, external pressure, and vapor pressure?
In order for boiling
point to be reached, the vapor pressure of the liquid must become equal to the
external pressure put on the system.
89) For a particular substance, why is the molar
heat of fusion less than the molar heat of vaporization?
The molar heat of
fusion is the energy required to convert a substance from solid to liquid. In order to do this, the IMF must be
lessened. The molar heat of vaporization is the energy required to convert a
substance from a liquid to a gas. Gas
particles have lots of space between them with basically no intermolecular
forces, so the energy required to take a substance from some intermolecular
force to none will be much greater than the energy required to just lessen the
intermolecular forces.
90) Why is the temperature of a substance
constant during melting?
The substance is
gaining energy in the form of potential energy so that it can change
phase. All molecules must have changed
state before the temperature will begin increasing.
91) What is the relationship between kinetic
energy and temperature?
Temperature is a
measurement of the average kinetic energy of a substance.
92) Why does increasing the
temperature increase the vapor pressure of a liquid?
As temperature is
increased, the motion of the molecules increases. As the motion of the molecules increases, the
molecules try to spread out more by breaking the intermolecular forces between
them, which increases the likelihood of the substance changing to gas phase,
thus increasing vapor pressure.
93) Explain why 12 g of steam at
100 degrees C can melt more ice than 12 g of liquid water at 100 degrees C.
12 g of water at 100
degrees Celsius is using some of the energy for phase change and melting,
whereas 12 g of steam at 100 degrees
Celsius is using all of the energy it contains to melt the ice.
94) What is the triple point of a
substance?
The point at which a
slight change in either pressure or temperature or both could force the
substance into one of three states—solid, liquid, or gas.
95) What is an indicator? A substance that
changes color depending on the acidity or basicity of
a substance
96) What does it mean for an acid or base to be
neutralized? The
acid or base is weakened, or brought closer to a neutral pH. It loses some of its characteristic
properties of being an acid or base.
97) On a heating curve, why do
some areas increase in temperature and other areas plateau? Areas that increase
in temperature are increasing in kinetic energy; areas that plateau are gaining
energy, but in the form of gaining potential energy that can be used to change
phase.
98) What
does pH measure? hydronium ion concentration (or hydrogen ion concentration)
99) What are the products of a
neutralization? salt and water
(acid + base à salt +
water)
100) What types of ionizing radiation exist? Know the characteristics of each.
alpha decay (a) --> a particle equivalent in mass and atomic to a He nuclide is released; lowest energy of radiation
beta minus decay (B-) --> an electron is released; middle energy of radiation
beta plus or positron decay (B+) --> a positive electron is released; middle energy of radiation
gamma radiation (g) --> highest energy of radiation--no mass as part of it
101) Complete the following
nuclear equation. 204 Pb à 200 Hg +
____.
0
-1e What type of decay is this?
102) Complete the following
nuclear equation. 234 Th à 234 Pa +
____.
4
2He What type of decay is this?
103) Complete the following nuclear
equation 8 B à 8
Be +
_____.
0
+1e What type of decay is
this?
104) What is the difference between fission and
fusion?
106) Describe the following nuclear equation:
106Ag + 0 -1e à 106Pd electron capture