Download Unit Two Objectives

GHS Chemistry
Final exam Review Sheet 2014
Unit 1: Matter & Physical Thermodynamics
Part One: Sig Figs, Scientific Notation, SI Units, & Measurement
1. Significant Figures: know the rules of significant figures for:
a. Addition/subtraction: the answer can be no more significant than the weakest number
b. Multiplication/Division: the answer should have the same number of significant digits as the number
with the smallest number of sig figs.
2. Precision vs. Accuracy: Repeated measurements that
are in agreement are precise. A measurement that
matches the actual number is accurate.
3. Uncertainty in Measurement: When reading an instrument,
you are only able to estimate ONE DIGIT beyond the
smallest increment.
For example: On the ruler below, the measurement would be 4.50 cm, where the 0 is estimated as one digit
beyond the smallest increment.
4. Percent Error: Percent Error = ((Accepted – Experimental)/Accepted) x 100%
Note: Percent Error does not have a sign, since it is based on the absolute value of the difference.
5. SI Unit Conversions with Factor Labeling:
a. Know how to convert the SI Units for length, mass, & volume based on the following prefixes:
S.I. PREFIXES
Prefix
Symbol
Amount
Mega
M
1,000,000 (106)
Kilo
K
1,000 (103)
Hector
H
100 (102)
Deca
Da
10 (101)
Deci
D
0.1 (10-1)
Centi
C
0.01 (10-2)
Milli
M
0.001 (10-3)
Micro
U
0.000001 (10-6)
Nano
N
0.000000001 (10-9)
b. Convert Celsius, Fahrenheit, and Kelvin temperatures.
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GHS Chemistry
Final exam Review Sheet 2014
c. Compare/calculate mass vs. weight (with gravity)
6. Temperature Vs. Heat:
a. Temperature is a measure of the Average Kinetic Energy of the particles in a substance
b. Heat, conversely, is the SUM of energies of the ALL of the particles in a substance. More
about Heat later …
7. Density = Mass/Volume
 Be able to determine the Density, mass, and volume in calculations
 Determine the Density as the slope of a Mass (Y-axis) vs. Volume (X-axis) graph
Part Two: Matter and Its Changes
Chemistry: It is the study of Matter and its Composition, Properties, and the Changes that it
undergoes.
1. Matter: it is the stuff that occupies space, and has mass, and which can be changed by energy.
2. Composition: What is it matter made of?
A. PURE SUBSTANCE: contains only one kind of matter.
 Element: the fundamental substances of chemistry.
 COMPOUND: made up from the combination of 2 or more elements in a precise, well-defined
ratio.
B. MIXTURES: is a physical blend of two or more substances, where the composition of a mixture may
vary.
Mixtures can be divided into Two Types:
(1) A Heterogeneous Mixture: is one that is not uniform in composition. For example, a dinner salad,
or oil in water consists of several components that are not evenly distributed. CAN BE
PHYSICALLY SEPARATED into substances.
(2) A Homogeneous Mixture: is one that has completely uniform composition; its components are
evenly distributed throughout the sample (i.e. sugar in water)
Homogeneous mixtures are VERY important in chemistry. In the liquid phase, these homogeneous
mixtures are called SOLUTIONS.
.
The States of Matter & The Kinetic Molecular
Theory
Properties and Changes that Matter undergoes:
C. Physical Change: change in the physical
appearance without change in the chemical
composition. Changes of the states of
matter, from solid, liquid, to gas (shown to
the right) are PHYSICAL CHANGES.
D. Physical Properties: density, Boiling
Point, Melting Point, Freezing Point,
Conductivity. Physical properties are
characteristics that can be observed without a chemical reaction
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GHS Chemistry
Final exam Review Sheet 2014
These physical properties are INTENSIVE, which
means that the property does not depend on the
amount of material present. Physical/Intensive
Properties include Boiling Point, Melting Point,
Freezing Point, Color, Odor, State, malleability,
Ductility (metal can be drawn into a wire), Luster,
Solubility, Viscosity, Luminosity, Hardness, Density, &
Conductivity
By contrast, EXTENSIVE Properties depend on the
amount of material; i.e. mass, volume, energy.
E. Chemical Properties: describes how one substance reacts with another.
F. Chemical Change: the substance undergoes a change in chemical composition, in a Chemical Reaction.
A chemical reaction is given by a Chemical Equation:
A + B

C + D
Reactants
“Yields”
Products
Some example chemical reactions:
a. 2 H2O  2 H2 + O2
b. Combustion of a hydrocarbon: rapid combination of oxygen with other materials
i.e.
2 C8H18 + 25 O2  16 CO2 + 18 H2O
The Ten signs of a Chemical Reaction:
1.
2.
3.
4.
5.
6.
7.
8.
9.
Bubbles Appear
A Precipitate Forms
A color change occurs
The temperature changes.
Light is emitted.
A Change in Volume occurs.
A change in electrical conductivity occurs.
A change in melting point or boiling point occurs.
A change in smell or taste occurs.
10. A change in any distinctive chemical or physical property occurs.
Thermal Change: the change that occurs when matter absorbs or releases heat.
Remember:
1. Temperature is a measure of the average kinetic energy of particles in matter.
2. Heat is the total energy of all of the particles in the sample; it is the form of energy that flows
between two bodies when the bodies are at different temperatures; heat will flow from the hotter
body to the cooler body.
a. Endothermic Change: occurs when matter ABSORBS energy (in=endo)
b. Exothermic Change: occurs when matter RELEASES energy (exit=exo)
Separation of Mixtures: how can we separate mixtures into their components?
a. Physical methods: use of a magnet to separate a mixture of sulfur and iron? Why can these be
separated by a magnet?
i. Separation based on physical properties; i.e. magnetic
ii. Physical separation with tweezers
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GHS Chemistry
Final exam Review Sheet 2014
iii. Filtration of a solid that does not dissolve in a liquid, but is suspended, or precipitated
at the bottom of the flask.
iv. Solubilities: we can separate mixtures based on solubilities.
b. Distillation: separation of a liquid mixture based on the boiling points (physical, intensive property) of
each component.
c. Chromatography: separation based on solubility, and competition between attraction to paper and
eluting solvent (TLC)
Separation of Compounds into Elements: separation of compounds back to elements is accomplished only via
Chemical Change.
Phase Changes and Thermodynamics
Defining the Physical States of Matter.
1. Define kinetic and potential energy.
2. Describe the arrangement of particles in solids, liquids and gases in terms of particle motion (their
kinetic energy!!).
3. Define intermolecular forces.
States of Matter & The Kinetic Molecular
Theory
Assumes
shape &
volume of
container
Definite
Volume &
shape of
container
Definite
Shape and
Volume
GHS Chem
5. Define temperature as a measure of the AVERAGE KINETIC ENERGY of particles.
PART TWO: Measuring Energy Changes When Matter Changes.
1. Describe in words and with diagrams the
changes that occur in melting, freezing, boiling
and condensing.
The graph shows a pure substance which is
heated by a constant source of heat supplying
2000.0 joules per minute. Identify the area
4
Physical
Properties
4. Describe how particle organization (their kinetic
energy and intermolecular forces distinguish
solids from liquids from gases.
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GHS Chemistry
Final exam Review Sheet 2014
described in the questions below and complete the necessary calculations.
a.
b.
c.
d.
e.
f.
g.
h.
i.
j.
k.
being warmed as a solid ____UV_____
being warmed as a liquid ___WX____
being warmed as a gas _____YZ_______
changing from a solid to a liquid __VW___
changing from a liquid to a gas __XY____
What is its boiling temperature? 100
What is its melting temperature? 0 degrees
What segments represent a change in kinetic energy UV, WX, & YZ
Where on the curve do the molecules have the highest kinetic energy? Z
Where on the curve do the molecules experience a change in potential energy VW & XY
Moving on the graph from point X to point Y, the molecules experience An INCREASE in
POTENTIAL energy.
l. Moving on the graph from point V to point U, the molecules experience a DECREASE in kinetic
energy.
m. Which segment represents evaporation? XY Condensation? YX Melting? VW Freezing? WV
2. Explain how heat is a form of energy and how heat
changes accompany physical and chemical changes.
Thermodynamics
Heat Capacity & Specific Heat
3. Distinguish between specific heat capacity and heat
capacity and their SI units.
Heat Capacity: amount of heat needed to raise the temperature of
an object by 1 oC. This depends on the mass of an object.
4. Describe heat changes in terms of a system and its
surroundings.
The Specific Heat (Capacity): of a substance is the amount of heat
needed to raise the temperature of 1 gram of that substance by 1 oC.
5. Calculate the heat changes that accompany physical
and chemical changes.
What happens to the heat capacity as the mass is increased?
For WATER: 1 Calorie is the amount of heat needed to raise the
temp of 1 gram of water by 1 oC.
Specific Heatwater = 1 calorie /g
/goC)
since 1 calorie = 4.184 Joules
Specific Heatwater = 4.184 Joules /g
/goC)
GHS Chem
33
Here are some examples that you can use as a study guide.
The following values are for water:
Hfusion = 334 J/g, Hvaporization = 2260/g, Specific Heatliquid = 4.18 J/goC,
Specific Heatice = 2.03 J/goC, Specific Heatgas = 2.03 J/goC
1. What quantity of heat is released when 64 g of liquid water at 21.0 oC cools to 9.0 oC ?
Q = mcT = (64g) x (4.18 J/gC) x (21.0-9.0C)
= 3210 Joules
2. What is the mass of ice that will melt if 1250 Joules of heat are added?
Q = m x Hfusion so rearranging, m = Q/Hfusion
m = 1250J/334 J/g
m = 3.74 grams
3. 2.05 x 103 Joules of heat are added to 86g of water at 19.0oC. What will the final
temperature of the water be, assuming that it is still a liquid?
Q = mcT
2.05 x 103 Joules = (86g) x (4.18 J/gC) x (Tf – 19.0C)
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GHS Chemistry
Final exam Review Sheet 2014
Solving for Tf = 24.7 C
6. Construct equations that show the heat changes for physical and chemical processes.
7. Identify endothermic and exothermic changes by the heat energy term in a chemical equation.
PART THREE: Characteristics of Solids, Liquids and Gases
1. Define pressure as a force that is exerted on a unit of area, resulting from elastic collisions of gas
molecules with the inner surface of the container.
2. State the units of pressure: kiloPascal (kPa), atmosphere (atm), mm of Hg, along with their
relationships.
It VERY important to remember that Standard Atmospheric Pressure is 101.3 kPa,
or 1 atmosphere, or 760 mm Hg (torr). The boiling point for liquids AT THIS
PRESSURE is the NORMAL BOILING POINT.
3. Interpret a phase diagram of a substance at
any given temperature and pressure. The
TRIPLE POINT is where all three phases
coexist at the same time. Above the
CRITICAL POINT, a gas can longer be
changed to a liquid, regardless of how much
pressure is applied. The Normal BP and MP
are determined with a horizontal line at
standard atmospheric pressure of 1 atm, 740
mm Hg,, or 101 kPa.
4. Explain the vaporization of liquids using kinetic theory.
5. Describe what happens on a particle level at the boiling point of a liquid.
Molecules in the liquid phase gain more and more kinetic
energy, and finally have enough energy to escape from the
liquid phase, to the gas phase, at the boiling point of the liquid.
The pressure that these escaping molecules apply to the
atmosphere is the VAPOR PRESSURE. Boiling occurs when
the Vapor Pressure equals and then exceeds the Atmospheric
Pressure.
6. Determine the boiling point using a vapor pressure
curve.
Draw a horizontal line at the suggested pressure, and
determine the boiling point where the pressure line
intercepts the vapor curve. Alternatively, draw a line
from the desired temperature up to the vapor curve,
determine the pressure where the liquid would boil at
that temperature.
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GHS Chemistry
Final exam Review Sheet 2014
7. Compare and contrast the effect of intermolecular forces on the boiling point of substances using
the data on a vapor pressure curve. Simply stated, those liquids with higher intermolecular forces will
ave a vapor pressure curve that is shifted to the right of the graph. In the VP Graph above, water has the
higher intermolecular forces (Hydrogen bonds), so it’s VP curve is farthest to the right.
8. Define sublimation. Sublimation is the transition from a solid directly to the gas phase. This is a
ENDOTHERMIC reaction. The opposite transition, called Deposition, is Exothermic.
Unit 2: Atomic Structure /Electrons
1.
Describe the various models in the historical development of modern atomic theory:
a. Aristotle: Matter is made of Air, Fire, Earth, & Water.
b. Democritus: The first to say that matter is composed of atom, or “atomos.”
c. Dalton: Had five basic principles in his model of the atom
d. Thomson: discovered the charge of the electron by deflecting the flow of electrons through his
Cathode Ray Tube with magnetic and electrical fields, and theorized that the atom was a “plum
pudding” of electrons and positive charges.
e. Rutherford: Used the Gold Foil experiment to prove the Plum Pudding Model. He shot Alpha
particles (positively charged He nuclei) at Gold foil only to discover that the atom has a nucleus,
filled with protons, and other significant mass. Since most of the alpha particles passed through, he
concluded that the atom is mainly empty space. The small number of particles that bounced back
proved that the nucleus had mass, and positively charged protons.
f. Bohr: Electrons move about the nucleus in “orbits” in Quantum Energy levels. The maximum
number of electrons per energy level is 2n2, where n = the energy level number.
g. Quantum Mechanical Model: electrons spin around the nucleus in clouds of probabilities.
2.
Distinguish among protons, neutrons, and electrons in terms of their relative masses, charges, and location
with respect to the nucleus
3.
Infer the number of protons, neutrons, and electrons using the atomic number and mass number of an
element from the periodic table and symbol notation
i.e. 919F has 9 protons, 9 electrons (both determined by the Atomic Number), and 10 neutrons (Mass
Number – Atomic Number)
Explain how isotopes of an element differ: Isotopes are atoms of the same element with the same atomic
number but different Mass Numbers. Thus, isotopes have the same # of protons, but different numbers of
neutrons.
4.
5.
Explain why the atomic masses of elements are not whole numbers: Atomic Masses are “Weighted
Averages” of the naturally occurring isotopes.
6.
BE ABLE to calculate the average atomic mass of an element from isotope data. Predict which of the
naturally occurring isotopes is most abundant from average atomic mass.
7.
Differentiate between an atom and an ion. MAKE SURE that you understand that a “+” ion has LOST
electrons, while a “-“ ion has gained electrons … seems redundant … oh well?
8. Determine ion charge given proton, neutron, and electron data
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GHS Chemistry
Final exam Review Sheet 2014
9. Explain the atomic emission spectrum of an
element using Bohr’s model of the atom.
Electrons go from the ground state to the excited
state when energy is absorbed. Then the electron
moves from the excited state back to the ground
state, light is emitted!
The Electromagnetic Spectrum
Gamma Rays
0.001 0.01
X Rays
0.1
UV
nm
1
10
UVB
UVA
100
IR
Microwave
cm
1000 0.001 0.01 0.1
TV
1
10
1

VISIBLE

Near IR
V I B G Y O R
10. Calculate the Energy, frequency and wavelength of
electromagnetic radiation (light) using the
2x10 m
equations:
 E = hf, and c = f
12. Where h = Planck’s constant (6.63x10-34 J-s)
13. &
8
1. c = 3 x 10 m/s.
200
300
400
500
600
700
-7


800
900


9x10-7 m

Radio
meters
10
100 1000
R - Red
O - Orange
Y - Yellow
G - Green
B - Blue
I - Indigo
V - Violet
GHS Chemistry
14. What are the parts of the Wave? Predict the relationships between wavelength, frequency, and
energy for a wave.
15. Describe the properties of the different types of electromagnetic radiation.
16. Apply the Aufbau principle (filling of electron levels from low to high energy), the Pauli
Exclusion Principle (two electrons occupying the same orbital must have opposite spins), and
Hund’s rule (1 electron must be given to each of equivalent suborbitals; ie. 3 p and 5 d
suborbitals) to write electron configurations and orbital diagrams (1s2, 2s2, 2p6, 3s2, 3p6, …) of
elements. Do not forget the exceptions to the Aufbau principle, specifically the d4 and d9
exceptions.
17. Identify the electron configuration for an Excited Atom, where an electron has moved from
the Ground State to a higher energy level. For example, the electron configuration for the
ground state of 20Ca is 1s22s22p63s23p64s2. An possible electron configuration for an excited
2
2
6
2
5
2
1
20Ca atom may be 1s 2s 2p 3s 3p 4s 5s
18. Complete “Orbital Filling Diagrams” for an element’s electron configuration using the
Aufbau, Pauli, and Hund’s Rules.
19. Don’t forget the Quick equations:
b. Maximum number of electrons within an (n) energy level = 2n2
c. Maximum number of orbitals (i.e. s, py, pz, py, dzx, etc) within an (n) energy level = n2
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GHS Chemistry
Final exam Review Sheet 2014
Unit 3: Periodic Table
1. The Periodic Table: What is the Periodic Law? The Periodic Law states that when the elements
are arranged in order of increasing atomic number, there is a periodic repetition of their
chemical and physical properties.
a. The horizontal rows are called the periods. There are seven periods. Going across a period
from left to right, elements are filling that energy level’s “s & p” orbitals, eventually getting to a
full octet at the noble gas.
b. The vertical columns are called groups or families. Elements within the same group have the
same number of valence electrons, and thus have similar properties.
2. Where are the metalloids? Where are the Transition Metals? The Lanthaniides? The Actinides? The
Halogens? The Nobel Gases? The Alkali Metals? The Alkaline Earth Metals?
3. Describe the origins of the modern periodic table. Describe the organization of the periodic
table (periods, groups, periodic law, s block, p block, d block, & f block), and categorize
the elements as halogens, alkali metals, alkaline earth metals, noble gases, transition metals,
inner transition metals, and representative elements (all of the Group A elements). As a
refresher, the Periodic law states that when elements are arranged by their Atomic Number, there is
periodic repetition of properties within a group or family.
4. Contrast the physical and chemical properties of metals (LOSE ELECTRONS), nonmetals
(GAIN ELECTRONS), and metalloids (DO BOTH), and locate them on the periodic table
5. Explain the relationship between the electron configuration of an element, its position on the
periodic table, and its chemical properties. Simply put, be able to determine the electron
configuration using the Periodic Table, and predict the charge of the ion based on the Group A
number:
6.
i.e. You would expect a Group 2A metal, which has 2 valence electrons, to lose both of those
electrons and become a +2 ion.
_
within periods and groups of the Periodic Table.
+7.State the trends of properties of elements
st
nd
rd
Interpret the trend shown by atomic radii, ionic radius, electronegativities, and 1 , 2 , & 3
ionization energies, within the periodic table.
_
+
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GHS Chemistry
Final exam Review Sheet 2014
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GHS Honors Chem
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GHS Honors Chem
GHS Honors Chem
8. Determine the number of valence electrons &/or predict the stable ion formed by a
representative element, using the periodic table. Remember, the representative elements are
the GROUP A elements.
9. Predict Element location based on period and valence electrons. For example, a valence
configuration of 4s24p5 tells me that the element is in the 4th period, in Group 7A. The element is
therefore Br (Atomic No. 35).
Putting it all together……
Unit 4: Nomenclature/
Chemical Bonding


1. Predict the formation of cations & anions
from a representative atom, based on the
location on the Periodic Table; more
specifically, based on the atom’s Group and
the atom’s desire to obtain an OCTET of
electrons (or a noble-gas electron
configuration).







Solids have the least energy, and the
individual molecules are “frozen”
frozen” in
place by intermolecular forces.
In the Liquid State,
State, the energy is
increased, and the rigid structure of
the solid state is broken. Molecules
move past one another, but remain
relatively close due to intermolecular
forces.
forces.
Gases are formed when the energy in
the system exceeds all of the
attractive forces between molecules.
Thus gas molecules have little
interaction with each other beyond
occasionally bumping into one
another.
GHS Chem
Group 1A: loses s1 to become a +1
cation
Group 2A: loses s2 to become a +2 cation
Group 3A: loses s2p1 to become a +3 cation
Group 5A: gains 3 electrons to become s2p6, and a -3 anion
Group 6A: gains 2 electrons to become s2p6, and a -2 anion
Group 7A: gains 1 electron to become s2p6, and a -1 anion
2. Predict the formula unit of an ionic compound based on the
expected charge of the ions, and write the Lewis Dot
Structure for the Ionic compound.
3. What are the characteristics/properties of an ionic
bond? Ionic compounds occur when very reactive metals
transfer 1 or more electrons to reactive nonmetals. The
resulting positive metal ion and nonmetal negative ions are
highly attracted, and strongly bonded together. Ionic
10
18
Ionic Bonding
The Lewis Dot Struc ture for 3 Ca + 2 & 2 P3-
3Ca 2+ 2 P
Ca 3P2
GHS Honors Chem
3-
The Formula
Unit
GHS Chemistry
Final exam Review Sheet 2014
compounds exist as solids in a crystalline structure at room temperature, are dense, and have high
& specific melting points. Ionic compounds are conductive only when melted or dissolved in water.
4. Explain the physical properties of metals using the theory of
metallic bonding, where positively charged metal ions “float”
within a sea of mobile valence electrons. Metals are conductive,
malleable, and ductile due to this metallic bonding.
5. Describe the formation of a covalent bond
between two nonmetallic elements. These
bonds care either nonpolar covalent (for
diatomic molecules; i.e. H2, F2, Cl2), or polar
covalent (all others that are not ionic, or
nonpolar).
6. Identify bonds as ionic, polar covalent, nonpolar covalent using electronegativity values
If asked to judge the polarity of a covalent bond
based on the difference in electronegativity, use
the following chart as a guide:
7. Describe double and triple covalent bonds and draw Lewis structures to represent covalent bond
structures containing single, double, triple bonds, including exceptions to the octet rule. Those
exceptions covered in class were:
 Hydrogen (1 bond, 2 valence electrons only)
 Boron (3 bonds, 6 valence electrons only)
 Sulfur (can exceed the octet rule, and can have 6 bonding atoms, or 12 valence
electrons)
 Phosphorus (can exceed the octet rule, and can have 5 bonding atoms, or 10
valence electrons)
8. In the structure to the right, there are seven sigma bonds, and 3 pi
bonds.
9. Describe the shape of simple molecules using
orbital hybridization. Be able to describe a
molecule’s type (i.e. AX4, AX3E, etc), shape,
bond angle, hybridization, and overall
molecular polarity after building the
molecule’s Lewis Line Structure (like the
VSEPR worksheets and the Molecular Shapes
Worksheet – refamiliarize yourself with that
Molecular Shapes & Diagrams Worksheet that I
handed out. Here it is below).
AX Notation?
11
Hybrid?
GHS Chemistry
Final exam Review Sheet 2014
10. Identify examples, structure and properties of polar and non-polar molecules. Remember, all ionic
compounds are polar. Also, molecules that contain polar covalent bonds, BUT ARE
SYMMETRICAL, are nonpolar. SF6, PF5, and XeBr4 would all be examples of molecules with polar
covalent bonds, yet they are symmetrical, and thus, nonpolar molecules. If you’re not sure, then
draw the Lewis Structure to see it.
11. What are the attractions between molecules? How do they relate to attractive forces between
atoms? Describe the three Intermolecular forces covered in class from:
Weakest (Dispersion Forces) < Dipole-Dipole Interactions < Hydrogen Bonding < Covalent Bonds < Ionic Bonds (Strongest)
Inorganic Naming & Formula Writing
1.
2.
3.
4.
5.
6.
Distinguish among atoms, a molecular formula for covalent molecules, and a formula unit for ionic
compounds.
Distinguish between ionic compounds and covalent molecules.
Explain how a compound obeys the Law of Definite Proportions. In a sample of a molecule or
compound, the masses of the elements are always in the same proportions; i.e. water (H2O) is always 2
Hydrogens to 1 Oxygen.
Explain how two different compounds composed of the same elements obey the Law of Multiple
Proportions. The same atoms can combine in different ratios to form different molecules, with different
properties; i.e. water (H2O) and hydrogen peroxide (H2O2)
Memorize the charges of common monoatomic ions.
Memorize those metals that have multiple charges, and use both the Stock Naming System
(Roman numerals) and the Classical System for naming these metals in an ionic compound.
Those metals with Multiple charges;
Here are the Classical names that you are responsible
for:
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GHS Chemistry
Final exam Review Sheet 2014
Here’s a summary of monatomic ions:
7.
8.
9.
10.
11.
12.
13.
Distinguish between an ion and a
polyatomic ion.
Memorize the names, formulas and
charges of the common polyatomic ions.
Write chemical formulas for binary ionic
compounds (only 2 elements present).
Name binary ionic compounds when
given the chemical formula.
Identify by name and write the chemical
formulas for ternary ionic compounds
(with polyatomic ions).
Identify the name of a covalent molecule
from the formula. Write the chemical
formula for a covalent molecule given the
name.
Identify by name and write formulas for common acids.
For Simple Binary Acids:


If the anion attached to hydrogen ends in -ide, put the prefix hydro- and change -ide to -ic
acid
HCl - hydrogen ion and chloride ion becomes hydrochloric acid
For Oxyacids: If the anion has oxygen in it, the polyatomic ion’s name ends in -ate or -ite




change the suffix -ate to -ic acid (use no prefix)
H3PO4 Hydrogen and phosphate ions becomes Phosphoric acid
change the suffix -ite to -ous acid
H3PO3 Hydrogen and phosphite ions becomes Phosphorous acid
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GHS Chemistry
Final exam Review Sheet 2014
Unit 5 - Types of Reactions
1. Identify reactants and products in a chemical reaction.
Reactants  Products
2. Write chemical equations from a written description of a chemical reaction using appropriate symbols.
(s)  solid,(l)  liquid, (aq)  aqueous
3. Write word equations from chemical equations.
Zinc + Hydrochloric acid
 Zinc chloride + Hydrogen gas
Zn(s) + 2HCl(g)  ZnCl2(s) + H2(g)
4. Balance chemical equations when given the names or formulas of all reactants
and products in a chemical reaction. Follows the law of conservation of mass.
Number of atoms on left side (reactants) should be equal to number of atoms
on the right side (products).
2 NaBr + 1 Ca(OH)2  1 CaBr2 + 2 NaOH
5. Distinguish between these five types of reactions:
a) Combination:
 Use for Single Replacement
A + B  AB
Reactions.
reactant + reactant  1 product

The higher the metal on the
C + O2  CO2
series the stronger it is.
b) Decomposition:
AB  A + B
Reactant  Product + Product
2 H2O  2H2 + O2
c) Single replacement: (Use activity series)
A + XY  AY + X
Zn(s) +
HCl(g)  ZnCl2(s) + H2(g)
c) Double replacement: (Use solubility rules chart)
AB + CD  AD + CB
AgNO3(aq) + NaCl(s)  AgCl(s) + NaNO3(aq)
a. The following are one of the driving forces for a double replacement reaction:
i) Formation of a gas
ii) Formation of a precipitate
iii) Formation of water
d) Combustion of hydrocarbons: In a combustion reaction, a hydrocarbon (a compound containing
hydrogen and carbon) burns in oxygen (O2) and always produces carbon dioxide (CO2) and water (H20)
CH4 +
O2 
CO2 + H2O + heat
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GHS Chemistry
Final exam Review Sheet 2014
6. Describe the chemical properties of acids and bases
7. Interpret a pH scale and correctly use a variety of indicators to determine the pH of a solution.
8. Identify acids and bases from their chemical formulas and be able to name and chemical
formulas for common acids and bases.
9. Define: strong acid, strong base, weak acid, weak base; give at least one example of each.
(strong acids=hydrochloric, sulfuric, nitric; weak acid=acetic; strong base=sodium hydroxide,
potassium hydroxide, calcium hydroxide; weak base=ammonia / ammonium hydroxide)
10. Explain how the release of sulfur dioxide (SO2) into the atmosphere can form acid rain, and how
acid rain affects water sources, organisms and human-made structures.
11. Predict the products and balance simple combination (synthesis) and decomposition reactions, including
reactions of carbonates, chlorates, hydrogen peroxide and water.
12. General Prediction Rules for Decomposition Reactions:
Metal oxide  Metal + oxygen gas [2CuO 2Cu + O2]
Metal hydroxide  Metal oxide + water [Ca(OH)2  CaO + H2O]
Metal chlorate  Metal chloride + oxygen gas [2NaClO3 2NaCl + 3 O2]
Metal Carbonate  Metal oxide + carbon dioxide [(NH4)2CO3  (NH4)2O + CO2
Metal Bicarbonate  Carbonate + carbon dioxide + water
a. [2NaHCO3 (s) + heat  Na2CO3 + CO2 + H2O]
Hydrogen peroxide  water + oxygen gas [2H2O2  2H2O + O2]
Unit 6 – Energy Generation /Environment Impact
Unit 7- The Mole and Stoichiometry
Mole Conversions, Chemical
Equations, & Stoichiometry:
1. Moletown conversions:
2. Balancing Chemical
Reactions
3. Formula Weights:
a. Determine the
Empirical Formula from
% by weight
b. Determine the
Molecular Formula
from % Weight, and
actual Molecular
Weight
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GHS Chemistry
Final exam Review Sheet 2014
c. Determine the % Composition of each element from a molecular formula.
d. Stoichiometry, and More Stoichiometry. Know how to perform:
e. Standard Stoichiometry at STP (Standard Temp = -2730C or O0 K)
a. (Standard Pressure = 1 atm, 760 mm Hg, 760 torr, 101.3
kPa)
f. Limiting Reactant Stoichiometry (the reactant that produces the least
amount of product)
Unit 11 Gas Laws
1. Kinetic Theory of gases
2. Boyle’s law: P1V1 = P2V2
Boyles Law graph
3. Charles law:
V1 = V2
T1
T2
Charles Law Graph
4. Gay Lussac’s law:
P1 = P2
T1
T2
Gay Lussac’s
Law Graph
5. Remember, “Put Veggies on Table.” From the Combined Gas Law of P1V1/T1 =
P2V2/T2, you can derive the three other gas laws of Boyle’s (PxV), Charles’ (V/T),
and Gay Lussac (P/T). Remember, T must be in degrees Kelvin.
6. The Ideal Gas Law: PV=nRT. Remember, P in atm or kPa, V in L, and T in K.
7. Be able to solve simple gas-stoich problems.
8. Be able to describe real life applications of the gas laws.
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GHS Chemistry
Final exam Review Sheet 2014
Unit 9 - Solutions
1.
Define the terms solution (homogenous mixture), aqueous solution,
solute (substance being dissolved) and solvent (dissolving medium)
2.
Describe the role of solvation in the dissolving process and use the rule “like
dissolves like” to predict the solubility of one substance in another. Polar
substances will dissolve in polar and Non-polar substances will dissolve in nonpolar substances.
Ex. water (polar) will not dissolve in Oil (polar).
3.
List factors that determine how fast a soluble substance dissolves.
a) Stirred or shaken (will increase solubility)
b) Temperature (as temp increases, solubility increases)
c) Size of particle (smaller the size higher the solubility)
d) Nature of the solute / solvent
4.
Explain the difference among unsaturated( more solute
dissolves), saturated (no more solute dissolves), and
supersaturated (becomes unstable, crystals form) solutions.
5.
Apply information provided on a solubility curve.
a) Unsaturated (point below the curve)
b) Saturated (point ON the curve)
c) Super saturated (point above the curve)
6.
Calculate the Molarity of a solution.
Molarity = moles of solute
Liters of solution
7.
Prepare dilute solutions of given concentrations from concentrated solutions of known Molarity
using appropriate calculations (M1 V1 = M1 V1 )
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GHS Chemistry
Final exam Review Sheet 2014
Chemistry Formulas Reference Sheet
Formulas:
Q = m C ∆T
Q = m Hf or Q= m Hv
E = hf
c = λf
P1V1 = P2V2
T1
T2
PV=nRT
Constants
c = 3.00 x 108 m/s
h (Planck’s constant) = 6.626 x 10-34 J x sec
Hf = 334 J/g
Hv= 2260 J/g
C = Specific heat of ice = 2.06 J/g °C
C = Specific heat of liquid water = 4.18 J/g °C
C = Specific heat of steam = 2.02 J/g °C
R= 0.0821 liter-atm/mole-oK
8.314 liter-kPA/mole-oK
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