Chemistry midterm study guide
Chapter 1
1. Be able to distinguish a scientific theory from a scientific law
a. Recognize that scientific hypotheses and theories are never proven true
2. Distinguish between different fields of chemistry
a. Organic chemistry
b. Inorganic chemistry
c. Analytical chemistry
d. Physical chemistry
e. Biochemistry
Chapter 2
3. Apply the classifications of matter
a. Matter is anything that has mass and takes up space
b. Matter is made of atoms
c. Matter is divided into pure substances and mixtures
i. Pure substances are elements or compounds
ii. Mixtures contain multiple elements or multiple compounds
1. Homogeneous mixtures
a. Solutions
2. Heterogeneous mixtures
a. Suspensions
b. Colloids
4. Distinguish between different types of matter based on physical properties like phase (gas,
liquid, solid), melting point, boiling point, density, etc.
a. Distinguish between physical and chemical changes
i. Physical changes do not create new substances
ii. Separations of mixtures into pure substances are physical changes
1. Chromatrography
2. Distillation
b. Know the difference between a chemical symbol and a chemical formula
5. Explain how chemical changes have reactants and products
a. Chemical changes create new substances
b. Chemical changes can usually be identified by things like changes in color, smell, the
production of gas, and changes in temperature.
6. Explain why all chemical reactions obey the law of conservation of mass
a. The law of conservation of mass says that in any chemical reaction the mass of the
products equals the mass of the reactants.
Chapter 3
7. Distinguish between qualitative and quantitative measurements
a. Qualitative measurements include things like color and odor, or anything that cannot be
measured with a number.
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b. Quantitative measurements are for things like mass, length, and density, i.e. things that
can be measured with a number and unit (cm, kg, etc)
Be able to multiply, divide, add, and subtract with scientific notation
Be able to make real measurements to the correct degree of precision depending on the scale of
the measuring tool.
Distinguish between accuracy and precision
a. Accuracy is how close to the actual value a measurement is
b. Precision is how repeatable a measurement is
Be able to calculate error and % error
Know and use the rules of significant figures
a. Every nonzero digit in a reported measurement is assumed to be significant
b. Zeros appearing between nonzero digits are significant
c. Leftmost zeros appearing in front of nonzero digits are not significant
d. Zeros at the end of a number and to the right of a decimal point are always significant
e. Zeros at the rightmost end of a measurement that lie to the left of an understood
decimal point are not significant if they serve as placeholders to show the magnitude of
the number.
f. There are two situations in which measurements have an unlimited number of
significant figures. They are things you can count with whole numbers, e.g. 14
baseballs, and definitions, like 1cm=10mm.
i. Know how to add and subtract with significant figures
ii. Know how to multiply and divide with significant figures
Know the SI units for length, mass, and volume
a. Know the SI prefixes
i. Nano, micro, milli, centi, deci, deca, kilo
Be able to use the question mass = density x volume
a. If given density and volume, find the mass
b. If given density and mass, find the volume
c. If given volume and mass, find the density
Know how to calculate specific gravity
Know how to convert degrees Celsius to Kelvin and Kelvin to degrees Celsius.
Chapter 4
17. Dimensional analysis
a. Be able to convert between SI units
i. Know how to cancel units
ii. Use cancelling units as a guide to setting up problems
b. Solve multistep chemistry problems
i. Example: gold has a density of 19.3g/cm3. What is the density in kilograms per
cubic meter?
Chapter 5
18. Know the composition of atoms
a. Protons and neutrons are in the nucleus
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i. Protons have positive charge, neutrons have no charge
b. Electrons orbit the nucleus
i. Electrons have negative charge
Understand and apply Dalton’s atomic theory
Be able to identify elements with physical properties like
a. Atomic number
b. Number of protons
c. Mass number
d. Number of electrons
e. Weighted atomic mass
f. Atomic mass units
Be able to calculate the weighted atomic mass of an element when given the mass of its
isotopes and their relative abundance.
Be able to calculate the relative abundance of isotopes based on their masses and the weighted
atomic mass of the element.
Describe how Mendel developed the periodic table
a. Mendel ordered the elements by increasing atomic mass, not atomic number
b. Mendel noticed that the chemical properties of elements repeated when ordered by
atomic mass.
c. Mendel predicted the existence of elements that had not yet been observed
d. Distinguish between groups and periods in the periodic table
Know the basic properties of metals like the alkali metals, the alkaline earth metals, and the
transition metals.
Know the basic properties of nonmetals, especially the halogens and the noble gasses.
Predict the number of valence electrons in an element based on its group in the periodic table.
Chapter 6
27. Distinguish between molecular compounds and ionic compounds
a. Nonmetals make molecular compounds
i. Molecular compounds have covalent bonds, where atoms share electrons
ii. Be able to predict the number of covalent bonds a nonmetal will make based on
its group in the periodic table
iii. Molecular compounds tend to have low melting points and are poor electric
conductors
b. Metals bond with nonmetals in an ionic bond
i. Ionic bonds form between positive and negative ions
ii. Metals make positive ions, or cations, and nonmetals make negative ions, or
anions.
iii. Ionic compounds tend to have high melting points and are good conductors of
electricity when dissolved in water.
28. Be able to name anions and cations, and predict the charge an element will have if it ionizes
based on its number of valence electrons.
a. Cations have the same name as the element
b. Anions change their ending to –ide, e.g. oxide, nitride, phosphide.
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Be able to write a chemical formula for molecular compounds and ionic compounds.
Understand and apply Dalton’s law of definite proportions
Understand and apply Dalton’s law of multiple proportions
Be familiar with the transition metals that can have multiple charges:
a. Cu + or 2+
b. Fe 2+ or 3+
c. Hg22+ or Hg 2+
d. Pb 2+ or 4+
e. Sn 2+ or 4+
f. Cr 2+ or 3+
g. Mn 2+ or 3+
h. Co 2+ or 3+
Know all of the polyatomic ions on page 147.
Write formulas for binary ionic compounds, and write their names when given the formula
Write formulas for tertiary ionic compounds, and write their names when given the formula
Name molecular compounds when given their formula, and write the formula for molecular
compounds when given their name
a. Know the Greek prefixes for naming molecular compounds
Know the names and formulas for common acids like hydrochloric acid (HCl), sulfuric acid
(H2SO4), etc.
Chapter 7
38. Know Avogadro’s number, 6.02 x 1023
a. Avogadro’s number is a counting number
b. Called a mole (mol) in chemistry calculations
c. 1 mol of anything = 6.02 x 1023 of that thing
39. Use the periodic table to
a. Find the gram atomic mass of an element
b. Find the gram molecular mass of a molecular compound
c. Find the gram formula mass of an ionic compound
40. Be able to convert between mass, moles, number of particles, and volume at STP.
a. 1 mol of any gas = 22.4L at STP
41. Use % composition to calculate the relative mass of an element in a compound
a. Use % composition to determine the empirical formula of a compound
b. Use % composition and molar mass to determine the molecular formula of a compound
Chapter 8
42. A chemical reaction can be concisely represented by a chemical equation
43. The substances that undergo a chemical change are the reactants. The new substances formed
are products.
44. Special symbols are written after formulas in equations to show a substance’s state. The
designations are for a solid, liquid, or gas are (s), (l), and (g) respectively. A substance dissolved
in water is designated (aq).
45. A catalyst is a substance that increases reaction rate without being used up by the reaction. If a
catalyst is used, its formula is written above the arrow.
46. In accordance with the law of conservation of mass, a chemical equation must be balanced. In
balancing an equation, coefficients are used so the same number of atoms of each element are
on each side of the equation.
47. In a combination reaction, there is always a single product. The reactants are two or more
elements and/or compounds that combine.
48. A decomposition reaction involves the breakdown of a single compound into two or more
simpler substances.
49. In a single-replacement reaction, the reactants and products are an element and a compound.
The activity series of metals can be used to predict whether single-replacement reactions will
take place.
50. A double-replacement reaction involves the exchange of cations (or anions) between two
compounds. This reaction generally takes place between two ionic compounds in aqueous
solution.
51. A combustion reaction always involves oxygen. The products of the complete combustion of a
hydrocarbon are carbon dioxide and water.
52. A complete ionic equation shows all dissolved ionic compounds as their free ions. It includes
spectator ions as well as ions involved in the reaction.
53. Both single- and double-replacement reactions can be written as net ionic equations in which
spectator ions are deleted from both sides of the equation.
54. The precipitate formed in a double-replacement reaction can be identified using a table of
solubilities.
Chapter 9
55. The coefficients in a balanced chemical equation tell the relative number of moles of reactants
and products.
56. Chemists use moles to do chemical arithmetic, or stoichiometry.
57. All stoichiometric calculations involving chemical reactions begin with a balanced equations
because mass is conserved in every chemical reaction.
58. The number and kinds of atoms in the reactants equal the number and kinds of atoms in the
products.
59. Stoichiometric problems are solved using conversion factors derived from a balanced chemical
equation.
60. A conversion factor called a mole ratio relates the moles a given substance to the moles of the
desired substance.
61. Units such as mass, volume of gases (at STP), and particles are converted to moles when
working stoichiometry problems.
62. Whenever quantities of two or more reactants are given in a stoichiometry problem, the limiting
reagent must be identified.
63. A limiting reagent is completely used up in a chemical reaction.
64. The amount of limiting reagent determines the amount of product formed in a chemical
reaction.
65. If there is a single limiting reagent in a reaction, all the other reactants are in excess.
66. A theoretical yield is the maximum amount of product that can be obtained from a given
amount of reactants in a chemical reaction.
67. An actual yield is the amount of product obtained when the reaction is carried out in the
laboratory.
68. A ratio of the actual yield to the theoretical yield, expressed as a percentage, is the percent yield
of a reaction.
Chapter 13
69. Rutherford pictured the atom as a dense nucleus surrounded by electrons.
70. In the Bohr model of the atom, the electrons move in fixed circular paths around a dense,
positively charged nucleus.
71. The energies of electrons in an atom are quantized in the modern quantum mechanical model
of the atom.
72. Current theory predicts the probability of finding an electron in terms of a cloud of negative
charge. The atomic orbital, or regions in which electrons are likely to be found, can be
calculated from a mathematical expression.
73. The ways in which electrons are arranged around the nuclei of atoms are called electron
configurations.
74. Correct electron configurations for atoms may be written using the aufbau principle, the Pauli
exclusion principle, and Hund’s rule.
75. For all electromagnetic waves, the product of frequency and wavelength always equals the
speed of light.
76. The concept of quantized electron energy levels in atoms grew out of the study of the
interaction of light and matter.
77. The line emission spectra of atoms are best explained by quantized energy levels.
78. The quantum concept developed in part from Planck’s studies of light radiation from heated
objects and from Einstein’s explanation of the photoelectric effect.
79. De Broglie proposed that all matter in motion has wavelike properties.
80. Schrodinger devised the most successful of these early quantum mechanical models.
81. In de Broglie’s equation, the wavelength of an object equals Planck’s constant divided by the
product of the mass and velocity of the object.
Chapter 14
82. Elements that have similar properties also have similar electron configurations and are members
of the same group.
83. The atoms of the noble gas elements have filled outermost s and p sublevels.
84. The outermost s and p sublevels of the representative elements are only partially filled.
85. The outermost s and nearby d sublevels of transition metals contain electrons.
86. The outermost s and nearby f sublevels of inner transition metals contain electrons.
87. Regular changes in the electron configuration of the elements cause gradual changes in both the
physical and chemical properties of the elements within a group and within a period.
88. Atomic radii generally decrease as you move from left to right in a given period because there is
an increase in the nuclear charge while the number of inner electrons, and hence the shielding
effect, remains constant.
89. Ionization energy, the energy required to remove an electron from an atom, generally increases
as you move from left to right across a period. Ionization energy decreases as you move down a
group.
90. Atomic radii generally decrease within a given group because the outer electrons are farther
from the nucleus as you go down the group. The attractive force of the increased nuclear
charge is unable to overcome the effect of the greater distance, which acts in opposition.
91. Ionic radii decrease for cations and anions as you move from left to right across a period and
increase as you move down a group.
92. Electronegativity measures the ability of a bonded atom to attract electrons to itself. It
generally increases as you move from left to right across a period. It decreases as you move
down a group.
Chapter 15
93. Atoms in compounds are held together by chemical bonds. Chemical bonds result from the
sharing or transfer of valence electrons between pairs of atoms.
94. Bonded atoms attain the stable electron configuration of a noble gas. The noble gases
themselves exist as isolated atoms because that is their most stable condition.
95. For the representative elements, the number of valence electrons is equal to the element’s
group number in the periodic table.
96. The transfer of one or more valence electrons between atoms produces positively and
negatively charged ions, or cations and anions, respectively.
97. The attraction between an anion and a cation is an ionic bond. A substance with ionic bonds is
an ionic compound.
98. Nearly all ionic compounds are crystalline solids at room temperature. They generally have high
melting points. The total positive charge in an ionic compound is balanced by the total negative
charge; thus ionic compounds are electrically neutral.
99. Ionic solids consist of positive and negative ions packed in an orderly arrangement. The
coordination number of an ion indicates the number of ions of opposite charge that surround
the ion in a crystal.
100.
When melted or in aqueous solution, ionic compounds can conduct electricity because
the ions can move freely when a voltage is applied.
Chapter 16
101.
Atoms form covalent bonds when they share electrons to form an octet.
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A shared pair of valence electrons constitutes a single covalent bond. Sometimes two or
three pairs of electrons may be shared to give double or triple covalent bonds
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Sometimes one atom may contribute both bonding electrons in a covalent bond. This
type of bond is called a coordinate covalent bond.
104.
Resonance structures help to visualize the bonding in molecules when more than one
electron dot formula can be written.
105.
Molecular orbital theory is a logical extension of the quantum mechanical description of
the electron structure of atoms. Covalent bonding is described in terms of sigma and pi bonds.
106.
The valence-shell electron-pair repulsion, or VSEPR, theory of molecular geometry
states that, as a general rule, molecules adjust their three-dimensional shapes so the valence
electron pairs are as far apart as possible.
107.
In some instances, molecular geometry is adequately described by simple overlap of
atomic orbitals. In others, a description of molecular shape that better fits experimental results
is obtained from hybridized atomic orbitals.
108.
When covalent bonds join like atoms, the bonding electrons are shared equally and the
bond is nonpolar. When the atoms in a bond have different electronegativities, the bonding
electrons are shared unequally and the bond is polar.
109.
Hydrogen bonds are attractive forces in which a hydrogen covalently bonded to a very
electronegative atom is also weakly bonded to an unshared electron pair of another
electronegative atom. Hydrogen bonds are strong relative to other dipole interactions.
110.
Weak intermolecular forces determine whether a covalent compound will be a solid,
liquid, or gas. Weak attractions between molecules are call van der Waals forces and include
dispersion forces, dipole interactions, and hydrogen bonds.
Chapter 20
111.
Acids taste sour, are electrolytes, react with active metals to produce hydrogen, react
with bases to form water and salts, and cause indicators to change color.
112.
Bases taste bitter, are electrolytes, react with acids to form water and salts, and cause
indicators to change color.
113.
Water molecules dissociate into hydrogen ions (H+) and hydroxide ions (OH-).
114.
On the pH scale, 0 is strongly acidic, 14 is strongly basic, and 7 is neutral. Pure water at
25 degrees C has a pH of 7.
115.
A Bronsted-Lowry acid is a proton donor, and a Bronsted-Lowry base is a proton
acceptor.
116.
An Arrhenius acid yields hydrogen ions in aqueous solution. An Arrhenius base yield
hydroxide ions in aqueous solution.
117.
A Lewis acid is an electron pair acceptor, and a Lewis base is an electron pair donor.
118.
A conjugate acid-base pair consists of two substances related by the loss or gain of a
single hydrogen ion.
119.
The strength of an acid or a base is determined by its degree of ionization in solution.
The acid dissociation constant (Ka) is a quantitative measure of acid strength.
120.
The base dissociation constant (Kb) is a quantitative measure of base strength.
Chapter 21.
121.
In the reaction of an acid with a base, hydrogen ions and hydroxide ions react to
produce a salt and water. This reaction, called a neutralization, is usually carried out by
titration.
122.
The end point in a titration is the point at which the indicator changes color. An
indicator is chosen so the solution reaches the end point very near the point of neutralization.
123.
An equivalent of an acid is the mass of the acid that provides one mole of hydrogen ions
in solution.
124.
A solution that contains one equivalent of an acid or base in a single liter of solution is a
one-normal (1N) solution.
125.
At the equivalence point of a titration, the number of equivalents of acid equals the
number of equivalents of a base.
126.
A salt consists of an anion from an acid and a cation from a base and forms when an acid
is neutralized by a base.
127.
Salts of strong acid—strong base reactions produce neutral solutions in water.
128.
Salts formed from weak acids or weak bases hydrolyze water and produce basic or
acidic solutions.
129.
Solutions that resist changes in pH are called buffer solutions. Within limits,
components of a buffer can react with hydrogen and hydroxide ions and minimize changes in
pH.
130.
The solubility product constant (Ksp) is the equilibrium constant for the reaction
between an ionic solid and its ions in solution.
131.
The solubility of a salt is decreased by the addition of a common ion.
Chapter 22
132.
Oxidation is the gain of oxygen or the less of electrons. Reduction is the loss of oxygen
or the gain of electrons.
133.
An oxidation process is always accompanied by a reduction process. The substances
that does the oxidizing is called an oxidizing agent. The substance that does the reducing is
called a reducing agent.
134.
Oxidation numbers help to keep track of electrons in redox reactions. An oxidationnumber increase is oxidation. An oxidation-number decrease is reduction.
135.
The oxidation number of an element in an uncombined state is zero. The oxidation
number of a monatomic ion is the same in magnitude and sign as its ionic charge. The sum of
the oxidation numbers of the elements in a neutral compound is zero.
136.
The oxidation-number—change method for balancing redox equations involves
determining the oxidation-number changes of the substances that are oxidized and reduced.
Coefficients are used to make the increase in oxidation number equal to the decrease.
137.
In the half-reaction method, which is another way to write a balanced equation for a
redox reaction, the net ionic equation is first divided into two half-reactions. One is for the
oxidation and the other is for the reduction. Each half-reaction is balanced independently for
atoms. H+, OH-, H2O are added as needed. The charge on both sides is balanced by adding
electrons. The half-reactions are then multiplied by factors to make the number of electrons
lost equal to the number of electrons gained. Finally, the half-reactions are added.