200
Things to Know to Pass Chemistry
1. Protons are positively charged (+) with a mass of
1 amu.
Example: Which
has the greatest nuclear charge? Cl-35 Ar-40 K-39 Ca-40
2. Neutrons
have no charge and a mass of 1 amu.
3. Electrons are small and are negatively charged
(-) with a mass of almost 0 amu..
4. Protons & neutrons are in an atom’s nucleus (nucleons).
Which has the
greatest number of nucleons?
Sn-119 Sb-122 Te-128 I-127
5. Electrons are found in “clouds” (orbitals) around
an atom’s nucleus.
Where is most
of the mass of an atom found?
Where is most of
the size (volume) of an atom found?
6. The mass number is equal to an atom’s number of
protons and neutrons added together.
What is the mass number of an atom with 18
protons and 22 neutrons?
7. The atomic number is equal to the number of
protons in the nucleus of an atom.
Which has the
greatest atomic number?
S Cl Ar K
8. The number of neutrons = mass number – atomic
number.
Which correctly represents
an atom of neon containing 11 neutrons?
11Ne 21Ne 20 Ne 22Ne
9. In a neutral atom the number of protons = the number of
electrons.
10. Isotopes
are atoms with equal numbers of protons, but differ in their neutron numbers.
Two isotopes of the same
element will have the same number of
neutrons
and electrons, neutrons
and nucleons,
protons and nucleons, protons and electrons
11. Cations are positive (+) ions and form
when a neutral atom loses electrons.
They are
smaller than their parent atom.
Which of the
following will form an ion with a smaller radius that that of its atom?
Cl N Br Ba
12. Anions are negative ions and form when a neutral
atom gains electrons.
They are larger than their parent atom.
Which electron
configuration is correct for a fluoride ion?
2–7 2–8 2–8–1 2–6
13. Ernest Rutherford’s gold foil experiment
showed that an atom is mostly empty
space with a small,
dense, positively charged nucleus.
14. J.J. Thompson discovered the electron and
developed the “plum-pudding” model
of the atom.
+
- + - Positive
& negative
+
- + - + particles spread throughout
-
+ - + entire
atom.
15. Dalton’s model of the atom was a solid sphere of
matter that was uniform throughout.
16. The Bohr Model of the atom placed electrons in
“planet-like” orbits around the nucleus of an atom.
17. The current, wave-mechanical model of the atom
has electrons in “clouds”
(orbitals) around the
nucleus.
18. Electrons can be excited to jump to higher energy levels.
They emit energy as light when they fall
from higher energy levels back down to
lower (ground state) energy levels. Bright line spectra are
produced.
19. Elements are pure substances composed of atoms
with the same atomic number.
They cannot be decomposed.
A compound differs from an
element in that a compound
Has a
homogeneous
composition has one set
of properties
Has a heterogeneous
composition can be
decomposed
20. Binary compounds are substances made up of only two
kinds of atoms.
“Ternary”
compounds contain three (or more) kinds of atoms.
Which substance is a binary compound?
Ammonia magnesium potassium nitrate methanol
21. Diatomic molecules are elements that form two
atom molecules in their natural form at STP.
Which element is a diatomic
liquid at STP?
Chlorine fluorine bromine iodine
22. Use this diagram to help determine the number of
significant figures in a measured value…
Pacific 
Atlantic
Atlantic
If the decimal point is present,
start counting digits from the Pacific (left) side,
starting with the first non-zero digit.
0.003100
(….. sig. figs.)
If the decimal point is absent, start
counting digits from the Atlantic (right) side,
starting with the first non-zero digit.
31,400 (……sig. figs.)
23. When multiplying or dividing measurements, final answer must
have as many digits as the measurement
with the
fewest number of digits.
When adding or
subtracting, use place value.
What is the
density of the object measured in lab by the displacement of water according to
The data below:
Mass
of object: 23.6
g
Volume
of water: 15.0 mL
Volume
of water + object: 18.2 mL
24. Solutions are the best examples of homogeneous
mixtures. They have two
sets of properties.
25. Heterogeneous mixtures have discernable
components and are not uniform throughout.
Air is classified chemically as a(n)
Substance compound element mixture
26. A solute is the substance being dissolved; the solvent
is the substance that dissolves the solute.
NaCl (s) is
added to water.
The
solute is ….. the solvent is …… the solution is ……..
27. Isotopes are written in a number of ways: C-14 is also
Carbon-14, and is also
14C
6
atomic number = …….. mass number = ……..
28. The average atomic mass is the weighted average mass of all
the known isotopes of an element.
Find the average
atomic mass of lithium if 7.4 % are 6Li and 92.6% are 7Li.
29. The distribution of electrons in an atom is its electron
configuration.
30. Electron configurations are written in the bottom center of an
element’s box on the
periodic table in your reference tables. The outermost electrons are the
valence electrons.
2 = #
of electrons in ………..
8
= # of electrons in ………
3=
# of electrons in ……….
31. Use the mole map
to help you solve conversions
between moles, grams,
numbers of molecules/atoms, and liters of gases at STP.
Given the reaction CH4 + 2O2
--> CO2 + 2H2O,
what amount of carbon dioxide is produced by the reaction of 1 mole of CH4?
what amount of carbon dioxide is produced by the reaction of 1 mole of CH4?
1
gram 1 liter 1 mole 22 grams
32. An empirical formula is the simplest mole ratio among the
elements in a compound.
Use the mole map
to convert percent (mass) to moles.
Find the
empirical formula of a compound composed of 75% carbon and 25% hydrogen.
33. Electron dot model is a way of representing the
valence electron of an atom.
represents the electron-dot symbol of this
element C O B N
34. The kernel
of an atom includes everything in an atom except the atom’s valence
electrons.
The kernel of
this element contains 11 protons and 10 electrons
O F Ne Na
35. Polyatomic ions (Table E) are groups of atoms, covalently
bonded together, with an overall charge.
Nitrate: ………..., NH4+: ……..….., sulfite: ……..…..,
etc.
Which of the
following contains both ionic and covalent bonds?
NaOH CH3OH NaCl Cl2
36. Coefficients are written in front of the
formulas of reactants and products to balance chemical
equations. They give the ratios of reactants and
products in a balanced chemical equation.
……..Na
+ …….Cl2 à ………NaCl
37. Chemical formulas are written so that the charges of cations
and anions neutralize (cancel) one another.
calcium phosphate: Ca2+ PO43- = …………
38. When naming binary ionic compounds, write the name of the
positive ion (cation) first,
followed by the name of the negative ion
(anion) with the name ending in “-ide.”
CaCl2 …………….. MgS ……………..
39. When naming compounds containing polyatomic ions, keep the
name of the
polyatomic ion the
same as it is written in Table E.
NH4Cl ………. Dimercury
(I) nitrate ……….
40. Roman
numerals are used to show the positive oxidation number of the cation if it
has more than
one positive oxidation number
FeO: …………………….
Nickel (III) sulfate: ……………..
41. Physical changes do not form new substances.
They merely change the appearance of the
original material. (The melting of ice)
H2O (s) à H2O (l)
42. Chemical changes result in the formation of new
substances.
Which process is an example of a chemical change?
the melting of ice the electrolysis of water the boiling of water
43. Reactants are on the left side of the reaction
arrow and products are on the right.
44. Temperature is a measure of average
kinetic.
Which sample has the highest average
kinetic energy?
H2O
(l) at 0oC H2O
(s) at 0oC CO2
(g) at STP Mg (s) at 298K
45. Exothermic reactions release energy (energy is a product of the
reaction) while
Endothermic reactions absorb energy
and the energy is a reactant in the reaction.
Given the reaction: CH4 (g) + 2 O2 (g) →
2 H2O (g) + CO2 (g) + heat
What is the overall
result when CH4 (g) burns according to this reaction?
Energy is absorbed
and ∆H is negative. Energy
is absorbed and ∆H is positive.
Energy is released
and ∆H is negative. Energy
is released and ∆H is positive.
46. Only coefficients can be changed when balancing
chemical equations!
Given the unbalanced
equation: Al + O2 = Al2O3
When this equation is balanced using the smallest
whole numbers, what is the coefficient of Al?
1 2 3 4
47. Synthesis reactions occur when two or more
reactants combine to form a single product.
Na (s) + Cl2(g) à ……
48. Decomposition reactions occur when a single
reactant forms two or more products CaCO3(s)
à CaO(s) + ….…
49. Single replacement reactions occur when one
element replaces another element in a compound.
Which equation below represents a reaction classified
as a "single replacement" reaction?
2 H2 + O2
--> 2 H2O
Pb(NO3) 2
+ K2CrO4 --> 2 KNO3 + PbCrO4
HCl + NaOH -->
H2O + NaCl
Cu + Zn(NO3)
2 --> Zn + Cu(NO3) 2
CaCO3
--> CO2 + CaO
50. Double replacement reactions occur when two
compounds react to form two new compounds.
Potassium sulfide is mixed
with lead acetate. Which of the following products is expected?
PbSO4 K2S K3PO4 PbS Pb(C2H3O2)2
51. The masses (and energy and charge) of the reactants in a
chemical equation is always equal to
the masses (and energy and charge) of
the products. “Law of Conservation
of Mass(and Energy).”
52. The gram formula mass (molar mass) of a substance is the sum of
the atomic masses of all the atoms in it. H2SO4
= …….. g/mole
2 x H = 2 x
………g = ………g
1
x S = 1 x ………g = ………g
4
x O = 4 x………g = ………g
53. Know how to calculate the percentage composition of a
compound. (Formula is on Table T.)
Find the percent by mass of oxygen in CaCO3.
54. 6.02 x 1023 is called Avogadro’s number and
is the number of particles in 1 mole of a substance.
Equal volumes of gases contain an equal
number of molecules.
Under similar conditions, which
sample
contains the same number of moles of particles
as
1 liter of O2 (g)?
1
L Ne(g) 0.5 L SO2 (g) 2 L N2 (g) 1 L H2O(l)
55. Know how to convert an empirical formula into a molecular
formula.
A compound has the empirical formula NO2.
Find its molecular formula if the molar mass = 92g.
N2O NO2 N2O4
N2O
56. The kinetic molecular theory explains the behavior of matter
as particles with energy and motion.
57. The particles in a solid are rigidly held
together, closely packed in a lattice arrangement.
Which of the
following has a regular geometric arrangement at 298 K and 1.0 atm?
Br2
(l) CO2 (g) Mg (s) H2O (l)
58. Solids have a definite shape and volume.
In what region of the graph below would you only find molecules with definite shape and volume?
59. Liquids have closely-spaced particles that
easily slide past one another; they have no definite shape,
but have a definite volume.
60. Gases have widely-spaced particles that are in
random motion (collide with container to create pressure).
61. Gases are easily compressed and have no definite
shape or volume.
In what region of the graph below would you only find a sample with no definite shape or volume?
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62. Be able to read and interpret heating/cooling curves as pictured
below.
During which interval on the graph are
solid and
liquid in equilibrium?
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63. Substances that sublime turn from a solid
directly into a gas.
They have very weak attractive forces.
(examples include CO2 & I2)
64. As they evaporate, liquids become gases, which create vapor
pressure. (Reference Table H).
As temperature
increases, vapor pressure increases.
This liquid on
Reference Table H has the weakest attractive forces:
Propanone ethanol water acetic acid
65. “STP” means “Standard Temperature
and Pressure.” Reference Table B
These conditions define STP P = …..atm T = …..K
66. Degrees Kelvin = C
+ 273
Room temperature = 25oC = …….K Boiling point of helium = 4 K = ……….oC
67. Heat is a transfer of energy from a material at higher
temperature to one at lower temperature.
When an ice
pack is applied to a bruised arm, ……..
transfers from ……… to …………..
68. Use this formula to calculate heat absorbed/released by
substances. q = mcDt
q = heat
absorbed or released (Joules)
m =
mass of substance in grams
c =
specific heat capacity of substance (J/gC)
… for water it’s 4.18 J/g C.
Dt = temperature change in degrees Celsius
What is the total number of joules of heat energy
absorbed by 12 grams of water when it is heated from 30°C to 40°C?
69. The heat absorbed or released when 1 gram of a substance
changes between the
solid and liquid phases
is the substance’s heat of fusion. (Reference Table B: 334 J/g for water)
How many joules
are required to melt 15 g H2O (s)?
70. The heat absorbed or released when 1 gram of a substance
changes between the liquid and gaseous
phases
is the substance’s heat of
vaporization. (Reference Table B: 2260 J/g for
water)
How many joules
are required to boil 120 g H2O (l)?
71. Always use Kelvins for temperature when using
the combined gas law.
P1V1 = P2V2
T1 T2
Set up the equation to calculate
the volume of 50. mL of methane gas collected at STP
when the pressure rises to 2.4 atm and the temperature drops to 240 K.
72. As the pressure exerted on a gas increases, the volume
decreases proportionally.
25 L of a gas is
held at 1.2 atm pressure. Find the new volume if pressure drops to 0.80 atm at constant temperature.
73. As the pressure on a gas increases, temperature
increases.
A sample of gas exerts a pressure of 220. kPa at 373 K. Find the
pressure at 373 K at constant volume.
74. As the temperature of a gas increases, volume
increases.
15 mL of oxygen
gas is collected at 0oC. Find the volume at 50oC at
constant pressure.
75. Real gas particles have volume and are attracted
to one another.They don"t always behave like ideal gases.
Lighter gases
(with weaker attractive forces) are often most ideal.
Which of the
following is the most ideal gas?
He Ne Ar Kr
76. Real gases behave more like ideal gases at low pressures
and high temperatures.
77. Mixtures may be separated by several physical means:
Distillation separates
mixtures with different boiling points.
Fractional
distillation is a common method to separate and collect
Hydrocarbons Ionic solids Metals Precipitates
Filtration
separates mixtures of solids and liquids.
What
would collect in filter paper if a mixture of NaCl (aq) and CaCO3 (s)
were poured
through?
Chromatography
can also be used to separate mixtures of liquids and mixtures of gases.
78. The Periodic Law states that the properties of
elements are periodic functions of their
atomic numbers.
Elements are
arranged on the modern periodic table in order of increasing …………..
79. Periods are horizontal rows on the Periodic
Table.
In which energy
level are the valence electrons of the elements in Period 3 found?
80. Groups are vertical columns on the Periodic
Table.
Which group on the periodic table contains a solid, liquid, and gas(es)?
81. Metals are found left of the “staircase”
on the Periodic Table and at the bottom, nonmetals are
above it and at the
top, and metalloids border it.
Which of the following Group 14 elements has the greatest metallic character?
Carbon
silicon germanium tin
82. Complete and memorize this chart.
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Metals
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Malleable
and ductile
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All solids
except ……..
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Lustrous
|
Good
conductors of heat & electricity
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……. ionization
energy and electroneg.
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Tend to form
…… ions
|
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Nonmetals
|
Brittle when
solid
|
Mostly gases
at STP
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Dull
|
Good
insulators
|
……..
ionization energy and electroneg.
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Tend to form
….. ions
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83. Noble gases (Group 18) are unreactive and stable
due to the fact that their valence level
of electrons is completely filled.
84. Ionization energy increases as you go up and to
the right on the Periodic Table.
Which element
among the diagrams below has the lowest ionization energy?
85. Atomic radii decrease left to right across a period due
to increasing nuclear charge.
Which period 3
element among the diagrams below has the largest radius?
86. Atomic radii increase as you go down a group due to
increased electron energy levels.
Which alkali metal among the diagrams below has the
largest radius?
87. Electronegativity is a measure of an element’s
attraction for electrons.
Which of the following atoms has the greatest
tendency to attract electrons?
calcium carbon copper chlorine
88. Electronegativity increases as you go up and to the
right on the Periodic Table.
Which element
among the diagrams below has the greatest electronegativity?
89. The elements in Group 1 are the alkali metals;
those in Group 2 are the alkaline earth metals.
Which atom below
represents the alkali metal of period 2?
90. The elements in Group 17 are the halogens.
Which element
among the diagrams below is a halogen?
91. The elements in Group 18 are the noble gases.
Which element
among the diagrams below is a noble gas?
92. Use Table S to compare and look up the
properties of specific elements.
The freezing
point of phosphorus is ……..oC
93. Energy is absorbed when a chemical bond breaks. Energy
is released when a chemical bond forms.
The greater the energy, the more stable
the bond that forms.
Which of the
following, according to Reference Table I, is the most stable compound?
Ethane ethane ethyne hydrogen
iodide
94. The last digit of an element’s group number is equal to its number
of valence electrons.
Which contains
the greatest number of valence electrons?
Ca Ge Se Kr
95. Draw one dot for each valence electron when drawing an
element’s or ion’s Lewis electron dot
diagram.
Which dot
model would contain the fewest dots as valence electrons?
Ca Ge Se Kr
96. Metallic bonds can be thought of as a crystalline
lattice of kernels surrounded by
a “sea” of mobile
valence electrons.
Metallic bonding occurs between atoms of
sulfur sodium fluoride sodium carbon
97. Atoms are
most stable when they have 8 valence electrons (an octet) and
tend to
form ions to obtain
such a configuration of electrons.
Which of the
following atoms forms a stable ion that does not have an octet
structure?
Li F Na Cl
98. Covalent bonds form when two atoms share
a pair of electrons.
How many covalent
bonds are found in a nitrogen (N2) molecule?
99. Ionic bonds form when one atom transfers
an electron to another atom when
forming a bond with
it.
Which substance exhibits
ionic bonding rather than covalent bonding?
CO2 N2O4 SiO2 CaBr2 C6H12O6
100.
Dot models may be used to represent the formation of ions or covalent
molecules.
. Given the equation:
This
equation represents the formation of a
fluoride
ion, which is smaller in radius than a fluorine atom
fluoride
ion, which is larger in radius than a fluorine atom
fluorine
atom, which is smaller in radius than a fluoride ion
fluorine
atom, which is larger is radius than a fluoride ion
101. Nonpolar covalent bonds form when two atoms of
the same element bond together.
102. Polar covalent bonds form when the
electronegativity difference between two
bonding atoms is between
0.6 and 1.7.
Which of the
following combinations would form a polar covalent bond?
H and
H Na and N H and N Na and Br
103. Ionic bonds form when the electronegativity
difference between two bonding atoms is
greater than 1.7.
104. Substances containing mostly covalent bonds are called molecular
substances.
They are attracted to each other by weak
van der Waals or stronger hydrogen attractions
Which of the following is a molecular
substance?
Lithium chloride carbon monoxide sodium nitrate aluminum
oxide
105. Van der Waals attractive forces are the attractive
force between nonpolar molecules.
Nonpolar
molecules are molecules that have structural symmetry.
106. Van der Waals attractions become stronger with
increasing molar mass.
Which of the
following samples has the greatest forces of attraction?
F2 Cl2 Br2 I2
107. Polar molecules have stronger forces of attraction. The lack
structural symmetry.
Which of the
following is a polar molecule?
CO2 H2O C4H10 N2
108. Hydrogen bonds are attractive forces that form
when hydrogen bonds to the elements N, O, or F and
gives the compound
unexpectedly high melting and boiling points.
The strongest forces of attraction occur
between molecules of
HCl HBr HF HI
109. Substances containing mostly ionic bonds are called ionic
compounds.
They are made of metal and nonmetallic
ions. They are held together by electrostatic (ionic) forces.
110. Complete and memorize this table.
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Substance Type
|
Properties
|
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Ionic
|
Hard
(Low/high)
melting and boiling points
Conduct
electricity when molten or aqueous
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Covalent
(Molecular)
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Soft
(Low/high) melting and boiling points
Do not conduct
electricity (insulators)
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111. Remember: substances tend to be soluble in solvents with
similar molecular properties.
“Like dissolves like”
Pentane does not
dissolve in water because pentane is ………. and water is ………..
112. As temperature increases, solubility increases for most
solids.
For which
solid does increasing temperature have the least effect on solubility?
Potassium
nitrate ammonium chloride potassium chlorate sodium chloride
113. At low temperatures and high pressures solubility increases
for most gases.
Carbon dioxide
gas is least soluble in water at conditions of …. temperature and .… pressure.
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114. Use Table G to determine whether a solution is saturated,
unsaturated, or supersaturated.
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115. Use Reference Table F to predict soluble and insoluble
products of chemical reactions.
Which compound below would
"precipitate" if formed during a double replacement reaction?
AgNO3 K3PO4 Na2CO3 MgCl2 CaSO4
116. Molarity is a way to measure the concentration
of a solution.
Molarity is equal to the number of moles of solute divided by the
number of liters of solution.
(Reference Table
T).
What is the
molarity of an NaCl solution if 2.0 mol NaCl is present in 0.50 L solution?
117. Percent by mass
= (mass of the part / mass of the
whole) x
100%
A solution of
glucose is prepared by added 10. g glucose to 40. g water.
What is its
percent composition?
118. Parts per million (ppm) =
(grams of solute / grams of solution) x
1,000,000
A sample of
water is found to contain 0.010 g lead in 10. g solution.What is the
concentration in ppm?
119. Solutes raise the boiling points and lower the melting points of solvents.
Which of the
following will have the highest boiling point?
1 mol
NaCl in 100 g water 1 mole CH3OH
in 100 g water 1mol CaCl2 in
100 g water
120. Liquids boil when their vapor pressure is equal
to the atmospheric pressure. (Reference Table H)
Water will boil
at 90oC when the atmospheric pressure is …..…kPa.
121. The normal boiling point of a substance is the
temperature at which it boils at 1 atm pressure.
(Reference Table H)
What is the
normal boiling point of propanone?
122. Chemical reactions occur when reacting species collide
effectively.
123. Covalently bonded substances tend to react more slowly than
ionic compounds.
124. Increasing the concentration of reactants will increase
reaction rate.
Which sample
of HCl (aq) will react most rapidly with magnesium metal?
0.50
M HCl 1.0
M HCl 3.0 M HCl 6.0 M HCl
125. Reaction rate increases with an increase in temperature (and
pressure for gases).
126. Catalysts speed up reactions by lowering their activation
energies.
They are not changed themselves and can be reused many
times over.
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127. Be able to recognize and read potential energy
diagrams.
The heat content of the reactants of the forward reaction is about
…kilojoules.
The heat content of the
products of the forward reaction is about …kilojoules.
The heat content of the activated complex of the forward
reaction is about ….kilojoules.
The activation energy of the forward reaction is about
……kilojoules.
Add a dotted line to show the effect of a catalyst.
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128. The rates of the forward and reverse reactions are equal at
equilibrium.
A chemical reaction has reached equilibrium when the
reverse
reaction begins
reactants
are used up
rates
of the forward and reverse reactions are equal
concentrations
of products and reactants are equal
129. Adding any reactant or product to a system at
equilibrium will shift the equilibrium
away from the added substance.
130. Removing (taking out) any reactant or product
from a system at equilibrium will shift the
equilibrium point toward
that removed substance.
131. An increase in temperature shifts an
equilibrium system in the endothermic
direction.
132. A decrease in temperature shifts an equilibrium
system in the exothermic direction.
133. Increasing the pressure on a gaseous
equilibrium will shift the equilibrium point
toward the side with
fewer moles of gas (less gas
volume).
134. Decreasing the pressure on a gaseous
equilibrium will shift the equilibrium point
toward the side with
more moles of gas (greater gas
volume).
135. Catalysts have no effect on
equilibrium. It just establishes itself
more quickly.
Given the
reaction: H2 (g) + I2 (g) <=> 2 HI (g)
If a catalyst is added, the equilibrium
concentration of HI (g) produced …….
136. Enthalpy (H) is the heat energy gained or lost
in a reaction.
137. Entropy (S) is high in a highly unorganized
system, such as a gas, a messy room,
etc.
Which of the
following has the greatest entropy?
Na
(s) CO2 (g) H2O (l) N2 (g) + H2
(g)
138. A chemical reaction is most likely to occur (spontaneously)
in an exothermic reaction
with an increase in entropy.
In the reaction
below,
energy …(increases/decreases)…...… and
entropy ……(increases/decreases).…..….
N2
(g) +3 H2 (g)à 2 NH3 (g) + 91.8 kJ
139. Oxidation numbers can be assigned to atoms and ions.
What is the oxidation number of S in the sulfate ion?
140. Oxidation is the loss of electrons
by an atom or ion. The oxidation
number increases as a
result.
The electrons are usually on the right
side of the reaction arrow.
In the reaction Sn+4 + H2
(g) à Sn+2 + 2H+,
substance oxidized is
Sn+4 H2 Sn+2 H+
141. Reduction is the gain of electrons
by an atom or ion. The oxidation number
decreases (is
reduced!) as a result. The electrons are
on the left side of the reaction
arrow.
142. Redox reactions always involve the exchange of electrons.
Electrons lost = electrons gained.
143. Remember…. OIL RIG Oxidation
is loss of electrons Reduction
is gain of electrons
Identify the
element that gains electrons in the reaction: 2 Na + Cl2 à 2 NaCl
144. Identify redox reactions by looking for changes
in oxidation number.
Zn + 2HCl
à ZnCl2
+ H2
Write the oxidation and half reactions in the
above reaction.
145. Oxidizing agents are what get reduced in
a redox reaction.
Reducing agents are what get
oxidized in a redox reaction.
Identify the
oxidizing agent in the reaction: KMnO4
+ HCl + H2S à KCl + MnCl2 + S + H2O
146. Redox reactions can be balanced using the half-reaction
method
Balance the
equation in #145.
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147. Electrochemical cells produce electricity
with a
spontaneous
redox reaction.
In the
electrochemical cell shown at the right:
Electrons flow
from …….. to ………
The anode is
…..…..; the cathode is …………
…………..move through
the salt bridge
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148. The left electrode is usually the site of oxidation
in an electrochemical cell diagram.
149. Memorize this saying… “I have AN OX and a RED
CAT.”
In electrochemical cells, the ANode gets OXidized and REDuction occurs at
the CAThode.
In the chemical
cell reaction: Mg + Cu2+ à Mg2+ + Cu, the anode is …….
150. Use the Activity Series (Table J) to predict whether
or not a single replacement reaction will occur.
Which reaction will take place in a 1.0 molar aqueous solution?
Cu + Ag+
à Co + Zn+2 à Ag + Mn+2 à Sn + Fe+2 à
151.
Electrolytic cells use an applied electrical current to force a
nonspontaneous redox reaction to occur.
In
what kind of cell are redox reactions made to occur using an externally applied
electrical current?
galvanic cell chemical
cell electrochemical cell electrolytic cell
152. Electrolytic
cells are usually used for metal plating of objects.
When electroplating with
silver, the mass of the positive electrode
decreases increases remains the same
153. Acids
and bases are both good
electrolytes. Their solutions
conduct electricity well.
Which of the following is a
nonelectrolyte?
LiOH HBr CH3COOH C2H5OH
154. Weak acids
taste sour and react with metals.
155. Weak bases
taste bitter and feel slippery.
156. Acids and
bases turn indicators different colors. They’re listed on Table M.
Which
solution will change red litmus to blue?
HCl(aq) NaCl(aq)
CH3OH(aq) NaOH(aq)
157. pH is the
negative log (exponent) of the hydronium [H+] ion concentration.
What
is the pH of a 0.00001 molar HCl solution?
1 9
5 4
158. Acids have a
pH < 7. Bases have a pH > 7.
159. Every 1 pH
number decrease represents a ten-fold [H+] increase.
160. Tables
K & L list names and formulas of common acids and bases asked about
on the Regents.
161. The metals above H2 on Table J
will react with acids to make H2 gas bubbles.
Which of the following will react
with acid to produce hydrogen gas?
Au Cu Ag Zn
162.
Arrhenius model of acids and bases states:
“Acids give off H+ to form H3O+
ions in aqueous solution as their only (+) ion.”
“Bases give off OH- ions in aqueous solution as their
only (-).”
Which of the following is neither an
Arrhenius acid nor an Arrhenius base?
KOH CH3COOH CH3OH HNO3
163.
Brønsted model of acids and bases states:
“Acids donate protons.” “Bases
accept protons.”
Identify one Bronsted acid in the
reaction below:
H2O
+ NH3 ó NH4 + + OH-
164. Bronsted acids become Bronsted bases; Bronsted bases become
Bronsted acids; forming conjugate pairs.
Identify one conjugate acid-base pair from question #163
165. Acids and bases react in neutralization
reactions to make water and a salt.
Name the salt
produced when sulfuric acid is neutralized by potassium hydroxide.
166. Titrations are controlled neutralization
reactions used to find the concentration of an acid or base
sample.
Note the formula for it on Table T.
How many mL of a
0.25 M HCl solution are needed to neutralize 20. mL of a 0.40 M NaOH solution?
167. ALL organic compounds contain the element carbon and (usually) hydrogen.
Which of the following is an
organic compound?
CaCO3 KSCN CH3Cl CO2
168. Carbon
ALWAYS makes four covalent bonds
in molecules.
Which statement explains why the element
carbon forms so many compounds?
Carbon atoms
combine readily with oxygen.
Carbon atoms
have very high electronegativity
Carbon readily
forms ionic bonds with other carbon atoms
Carbon readily
forms covalent bonds with other carbon atoms
In a molecule of CH4, the hydrogen atoms are spatially oriented
toward the centers of a regular
pyramid tetrahedron square rectangle
169.
Saturated hydrocarbons have all single bonds within them
(alkanes).
Which compound is a saturated hydrocarbon?
ethane ethene ethyne ethanol
ethane ethene ethyne ethanol
170.
Unsaturated hydrocarbons have double or triple bonds in them
(alkenes & alkynes).
In which pair of hydrocarbons does each
compound contain only one double bond per molecule?
C2H2 and C2H6 C2H2 and C3H6 C4H8 and C2H4 C6H6 and C7H8
C2H2 and C2H6 C2H2 and C3H6 C4H8 and C2H4 C6H6 and C7H8
171. Hydrocarbons
contain ONLY the elements hydrogen and carbon.
They are nonpolar molecules, nonelectrolytes, and do not dissolve in water.
172. The homologous
series of hydrocarbons’ formulas are on Reference Table Q.
173. The functional
groups on organic molecules are listed on Reference Table R.
Which class of organic compounds
can be represented as R -- OH?
acids alcohols esters ethers
acids alcohols esters ethers
174.
Structural isomers of organic compounds have different
structural formulas but the same molecular formula.
Which compounds are isomers?
1-propanol and 2-propanol methanoic acid and ethanoic acid
1-propanol and 2-propanol methanoic acid and ethanoic acid
methanol
and methanal ethane
and ethanol
175. Number
the parent carbon chain in an organic molecule from the end closest to
the alkyl group(s).
Which molecule contains a total
of three carbon atoms?
2-methylpropane 2-methylbutane propane butane
2-methylpropane 2-methylbutane propane butane
176.
Combustion reactions occur when a hydrocarbon reacts with oxygen to
make CO2 and H2O.
177.
Organic substitution reactions occur when an alkane and a halogen (Group
17) reacts so that one or more hydrogen atoms on the alkane are replaced with
the halogen.
What type of reaction is CH3CH3 + Cl2 -> CH3CH2Cl + HCl?
an addition reaction a substitution reaction a saponification reaction an esterification reaction
an addition reaction a substitution reaction a saponification reaction an esterification reaction
178.
Organic addition reactions occur when an alkene or alkyne combine with
a halogen* to make one
product (halide).(the double bond between carbons becomes single;
triple bond becomes double).
The reaction CH2CH2 + H2 -> CH3CH3 is an example of
substitution addition esterification fermentation
substitution addition esterification fermentation
179 Esterification
occurs when an organic acid and an alcohol react to make water and an ester.
180.
Saponification occurs when an ester reacts with a base to make alcohol
and a soap.
181.
Fermentation reactions occur when yeast catalyze a sugar (C6H12O6)
to make carbon dioxide and ethanol.
The products of the fermentation
of sugar are ethanol and
water oxygen carbon dioxide sulfur dioxide
water oxygen carbon dioxide sulfur dioxide
182.
Polymers are long chains of repeating units called monomers.
What substance is made up of
monomers joined together in long chains?
ketone protein ester acid
ketone protein ester acid
183. Polymers
form by polymerization
reactions.
184.
Addition polymerization occurs when unsaturated monomers join in a long
polymer chain.
n(C2H2) à (C2H2)n
185. Condensation
polymerization occurs when monomers join to form a polymer by removing water.
Water is a product!
186. Natural polymers include
starch, cellulose, and proteins.
187. Synthetic polymers include
plastics such as nylon, rayon, and polyester.
188. Unstable
atoms that are radioactive are called radioisotopes. (Table N)
Which of the following represents
a stable nuclide?
Calcium-37 Potassium-42 Nitrogen-14 Phosphorus-32
189. Each radioactive isotope has
a specific mode and rate of decay (half-life).
Which sample will decay least
over a period of 30 days? [Refer to Reference Table N]
10 g of Au-198 10 g of I-131 10 g of P-32 10 g of Rn-222
190.
Radioisotopes can decay by giving off any of the particles/emanations listed in
Table O.
Which of the following decays by
positron emission?
Gold-198 Neon-19 Plutonium-239 Technetium-99
191. Alpha
particles (see Table J) are positively charged (+).
Beta particles (see
Table J) are negatively charged (-). Neutrons and gamma rays lack
charge.
Which particle cannot be
accelerated in a magnetic field?
alpha particle beta particle neutron proton
192. The sum of
the mass numbers and atomic numbers must be equal on both sides
of the reaction arrow for nuclear equations.
18 = 18
14N + 4He 17O + 1 H
9 =
9
193. When
radioactive nuclei decay, they undergo natural transmutation to form
new, stable atoms.
Complete the following decay
equation:
232Th à ……….. + …..…..
194. When
bombarded by radioactive particles, stable atoms undergo artificial
transmutation
Identify the element produced
when aluminum-27 is bombarded with an alpha particle.
(A neutron is also released).
27Al
+ 4He à 1n + …..…
195. Fission reactions split heavy
nuclei into smaller ones.
1n + 235U à 139Ba + 94Kr + 3 1n
+ Energy
196. Fusion reactions occur when
light nuclei combine to form a heavy nucleus and a lot of energy.
2H + 2H à 4He +
ENERGY
197. The half
life of a radioisotope is the length of time it takes for one
half of the
atoms in a sample to radioactively
decay. (Table N) (Table T).
Which sample
will decay least over a period of 30 days? [Refer to Reference Table N]
10 g of Au-198 10 g of I-131 10 g of P-32 10 g of Rn-222
198. Radioactive
isotopes have a variety of important uses.
Carbon-14, C-14, is used to determine the ages of organic material
up to 23,000 years old.
Uranium-238, U-238, is used
to determine the ages of rocks.
Iodine-131, I-131, is used to treat thyroid disorders.
Cobalt-60, Co-60, is used
to treat cancer tumors.
199. Radiation
can be used to kill bacteria on foods to slow the spoilage process.
200. Disposal of
radioactive waste is a problem associated with nuclear reactors
© 2005 John LaMassa
