Chemistry Cram Sheet!
1.
2.
Laboratory Safety
 Always wear _____ !
 Never ____ chemicals!
 To smell a chemical ____!
Lab Equipment
A balance measures ____. Units include _____.
A graduated cylinder measures _____.


When mixing solutions ADD ____ to ____!
Always rinse chemicals off skin with ____!
A beaker and an Erlenmeyer flask measure _____. Units include ___.
A pipet measures ____. These are best when measuring small
amounts of liquid that need to be EXACT!
A crucible is used to ___ objects
because it can withstand very
high temperatures.
3. Scientific Method
 Observation
 Hypothesis
 Experiment


Results
Example: It takes me too long to get to school going
the way I go now. What can I do?
4.
Experimental Design
After studying about recycling, members of John’s biology class investigated the effect of various recycled products on plant growth.
John’s lab group compared the effect of different aged grass compost on bean plants. Because decomposition is necessary for release
of nutrients, the group hypothesized that older grass compost would produce taller bean plants. Three flats of bean plants (25
plants/flat) were grown for five days. The plants were then fertilized as follows: (a) Flat A: 450 g of 3 month old compost, (b) Flat B:
450 g of 6 month old compost, and (c) Flat C: 0 g of compost. The plants received the same amount of sunlight and water each day.
At the end of 30 days the group recorded the height of the plants (cm).
What is the hypothesis in this experiment?
What is the independent variable in this experiment?
What is the dependent variable in this experiment?
What are the controls?
5.
Theory vs. Law
A law is an ____. It tells ___ happens. Ex: Rocks fall on this guy’s head.
A theory is an ____. It tells ____ it happens. Ex: The earth’s pull causes loose rocks to fall on this guy’s head.
6.
Percent Error
Used to tell how “off” you are from the value you should have gotten. Used mostly in lab.
Ex: The specific heat capacity of iron is 0.45 J/gC. A student uses a calorimeter to experimentally determine the specific heat of iron
to be 0.60 J/gC. What is the student’s percent error?
(Observed-Actual)/Actual X 100
7.
Graphing
Direct Relationship
Indirect Relationship
8.
Scientific Notation
Ex: 2.5 x 10-3
If the exponent is ___ then the number in standard notation is ___ than 1
If the exponent is ___ then the number in standard notation is ___ than 1
Multiplying 2 numbers in scientific notation: ___ the bases and ___ the exponents
Dividing 2 numbers in scientific notation: ___ the bases and ___ the exponents
9.
Uncertainty and Significant Figures
When taking a measurement, always measure one decimal place past the scale of your
instrument. For instance, the graduated cylinder to the left is measured with a 0.1 scale.
The measurement recorded is 1.15 mL (1 place past the scale of the instrument). The
“5” is the digit we are uncertain about.
Significant Figures in Measurements



Non-zero digits are always significant.
Any zeros between two significant digits are significant.
A final zero or trailing zeros in the decimal portion ONLY are significant
How many significant figures does each number below contain?
123 ____
103 ____
0.001 ____
10300. ____
10300 ____
0.003010 _____
Addition and Subtraction
The answer cannot have more places after the decimal than your measurement with the fewest places after the decimal.
Ex: 2.59 + 2.3 = 2.9
4.506 cm + 2.9 cm = ____
2.5 g - .36 g = ____
Multiplication and Division
The answer cannot have more significant figures than your measurement with the fewest number of significant figures.
Ex: 2.500 x 2.0 = 5.0
6.5 x 3 = ___
100 / 4.00 = ___
10.
Precision vs. Accuracy
Precision - ___
Accuracy - ___
0.200 cm
0.190 cm
0.201 cm – accepted value
11.
How would you describe these results?
They almost always ask you to calculate percent error.
| accepted - observed| X 100%
accepted
Dimensional Analysis
What is the length (in cm) of a football field?
How many seconds have you been alive?
12.
Temperature Conversions
Celsius  Fahrenheit
C = 5/9(F – 32)
Fahrenheit  Celsius
F = 9/5C + 32
Celsius  Kelvin
K = C + 273
What is human body temperature in Celsius, Fahrenheit, and Kelvin?
13.
Density
D = mass/volume
Units = g/ml
Density determines whether or not an object will float.
If an object has a mass of 5.0 g and a density of 20.0 g/mL, what is the volume of the
object?
A graduated cylinder is filled to the 10.0 mL line with water. A cube of tin (density = 7.3 g/mL) is placed in the graduated
cylinder. The water level in the graduated cylinder rises to 20.0 mL. What is the mass of the cube of tin?
14.
Metric Conversions
1000 mL = 1 L
100 cm = 1m
1000 m = 1 km
15.
1000 mm = 1 m
Properties of the States Of Matter
16. Chemical and Physical Changes
Physical Changes –
Any change in the state of matter of a substance is a PHYSICAL change!
Solid  liquid = ___
Liquid  solid = ___
Liquid  gas = ___
Gas  liquid = ___
Solid  gas = ___
Chemical Changes –
What are four signs that a chemical reaction has occurred?
17.
Specific Heat Capacity
Specific heat capacity – J/oC g
If an object has a low specific heat capacity, it heats up quickly.
If an object has a high specific heat capacity, it heats up slowly.
Q = M x Cp x T
Heat gained/lost by the object = Mass of object x Specific heat capacity of object x change in temperature (in Celsius!)
A 5.0 g object is heated from 25 C to 45 C. If 4.5 J of energy was required to make this change in temperature, what is the
specific heat capacity of the object?
18.
Protons, Neutrons, and Electrons
Protons – found in _nucleus have charge of __+1__
Electrons – found outside of nucleus in the electron cloud and have charge of _-1___
Neutrons – found in _nucleus have no charge
The number of _protons__ always equals the number of __electrons__ in a neutral atom.
In a magnesium ion, there are 2 more __protons__ than __electrons__ giving the ion a total charge of _+2___.
In a phosphide ion, there are 3 more _electrons___ than _protons___ giving the ion a total charge of __-3__.
ONLY _electrons___ CAN BE LOST OR GAINED!!!
19.
Atomic Number
The atomic number is determined by the number of protons. It is the identity of the element. The atomic number
NEVER changes – no matter what!!!
20.
Atomic Mass (Mass Number)
The atomic mass is determined by the number of protons and the number of neutrons. The atomic mass can change
depending on how many neutrons are present in the atom.
21.
Isotopes/Ions/Atomic Structure Review
Isotopes – atoms of the same element with different numbers of neutrons.
Ions - _depends on number of protons and electrons
Symbol
Atomic Number
Atomic Mass
# protons
# neutrons
15
ClCa2+
37
17Cl
# electrons
Charge
22.
Average Atomic Mass
The average atomic mass is an average of all the isotopes of an element. (This is why the atomic mass on the periodic table
is a decimal. That should make sense – you can’t have .01 neutrons!)
Average Atomic Mass = (% abundance x mass number)1 + (% abundance + mass number)2 + …
There are two isotopes of chlorine, 35Cl which is 75% of the chlorine in the world, and 37Cl. What is the AAM of chlorine?
23.
Chemists and their Contributions
Bohr –
Thompson –
Dalton –
Rutherford –
Democritus –
Hund –
Pauli –
Heisenberg –
Milikan –
Chadwick –
Mosley
Mendeleev
24.
Electron Configurations
Noble Gas Core – [Ne] 3s1 = sodium
S,p,d,f blocks
What is the electron configuration for Cd? What is the configuration for the Cd 2+ ion?
25.
Orbital Diagrams
Draw an orbital diagram for nickel.
26.
Energy Levels, Sublevels, and Orbitals (the Electron Hotel) or quantum numbers
Energy level = n = 1, 2, 3,4…. Sublevel = l ( s= 0 p=1 d=2 f=3) m= magnetic quantum # = -l to +l
Spin = +/- ½
27.
Periodic Table Families
Group 1 – Alkali metals
Group 2 – Alkaline Earth Metals
Group 3-12 – Transition Metals
28.
Group 16 – Chalcogens or oxygen family
Group 17 - Halogens
Group 18 – Noble Gases
Periodic Table Trends
Period
Group
Atomic Radius
Ionization Energy
Metallic Character
Electronegativity
Shielding
Which has the bigger atomic radius Na or Cl?
Which has a larger first ionization energy Li or Cs?
Will Ca form an ion larger or smaller than the original atom? P?
29.
Oxidation Numbers (Charges)
Charge results when an atom lose or gain an electron.
Metals lose electrons, therefore become positive ions called cations. (Did you lose an electron? Yes, I’m POSITIVE!)
Nonmetals gain electrons, therefore become negative called anions.
30.
Valence Electrons
Valence electrons – Number of electrons in the outer energy level
31.
Ionic Bonds
Ionic bonds are formed when electrons are transferred between a metal and a nonmetal.
32.
Covalent Bonds
Covalent bonds are formed when electrons are shared between two nonmetals.
33.
How do I tell if the Compound is Ionic or Covalent or Both?
Check to see what the compound is made up of:
A metal and a nonmetal…It’s IONIC!
2 nonmetals…It’s COVALENT!
A polyatomic ion and another element…It’s BOTH! (The polyatomic ion is the covalent part, the whole compound
will be ionic.)
Polarity
Covalent bonds are when electrons are shared between two nonmetals. If the electrons are shared equally, it is a nonpolar
covalent bond. If the electrons are shared unequally (meaning they are pulled closer to the more electronegative element), it
is a polar covalent bond.
34.
To determine whether a bond is polar, nonpolar, or ionic, you must use a table of electronegativities. (This will be given to
you on the SOL if you are supposed to use it.) When you subtract the two values, if the difference is…
…between 0 and 0.4, the bond is nonpolar, meaning the electrons are shared equally between the two atoms
…between 0.4 and 1.7, the bond is polar, meaning the more electronegative element is pulling harder on the
electrons
…greater than 1.7, the bond is ionic, meaning the more electronegative element pulled so hard on the electrons, that
they came off one atom and were transferred to the other atom.
35.
Drawing Lewis Structures
http://misterguch.brinkster.net/lewisstructures.html
Great website to review the rules for Lewis Structures! Don’t forget Lewis Structures only use VALENCE Electrons!
Draw structures for H2O, CO2, CCl4, and NH3
36.
VSEPR Theory
Valence Shell Electron Pair Repulsion Theory – basically means that the electrons want to be as far away from each other as
possible.
Important shapes for the SOL:
Shape
Structure
Bond Angle
Example
Bent
Trigonal planar
Trigonal pyramidal
Tetrahedral
Linear
Naming Ionic Compounds
37.
1.
2.
38.
39.
40.
41.
Name the positive ion (listed first)
a. If the positive ion is a transition metal, use a roman numeral to indicate the charge
Name the negative ion (ending should be –ide)
NaCl
LiBr
FeCl3
Writing Formulas for Ionic Compounds
1. Find the charge on the metal (positive).
2. Find the charge on the nonmetal (negative).
3. Cross over.
Magnesium nitride
Lead (II) phosphide
Lithium oxide
Potassium fluoride
Copper (II) sulfide
Beryllium chloride
Naming Covalent Compounds
1. Count the number of atoms of each element in the compound.
2. Use a greek prefix to identify the number of atoms of each element.
a. Mono – 1 (do not use on the first atom)
f. Hexa – 6
b. Di – 2
g. Hepta – 7
c. Tri – 3
h. Octa – 8
d. Tetra – 4
i. Nona – 9
e. Penta – 5
j. Deca – 10
N 2O 5
H2O
CCl4
CO2
Writing Formulas for Covalent Compounds
1. The prefix tells you how many atoms of each element are present in the compound.
2. If there is no prefix on the first atom, there is only one atom of that in the compound.
44.
H2SO4
HBr
HF
H3PO4
Hydrobromic acid
Nitric acid
Phosphate
Phosphate
Hydroxide
Cyanide
Ammonium
Chlorate
Chlorite
Writing Formulas for Acids
1. The charge on H is always +1.
2. Find the charge on the anion in the acid.
3. Cross over.
Sulfurous acid
43.
PF5
N2O
Carbon disulfide
Sulfur hexafluoride
Dinitrogen tetraoxide
Oxygen monofluoride
Naming Acids
1. Acids have H in the front!
2. If the acid does not have a polyatomic ion in the formula, it will be named hydro ___ ic acid
3. If the acid has a polyatomic ion in the formula:
If the polyatomic ion ends in –ate, the acid will end in –ic
If the polyatomic ion ends in –ite, the acid will end in –ous
HNO3
HCl
HNO2
42.
K2O
PbO
Cu2O
Polyatomic Ions
Nitrate
Nitrite
Sulfate
Sulfite
Diatomic Elements
hydrogen
nitrogen
oxygen
fluorine
chlorine
bromine
iodine
45.
Writing Chemical Equations
REACTANTS  PRODUCTS
**If something is written over the arrow, it usually means that that substance is required for the reaction to occur.
Solid potassium chloride reacts with oxygen gas to yield solid potassium chlorate.
46.
47.
Types of Chemical Reactions
Synthesis
Decomposition
Single Replacement (Brad & Jen)
Double Replacement
Combustion
Acid/Base
Balancing Chemical Equations
Balance equations to satisfy Law of Conservation of Mass.
Write and balance:
Magnesium reacts with nitrogen to yield magnesium nitride.
48.
Moles
6.02 x 1023 units
1 mole of gas at STP = 22.4 L
1 mole =
How many atoms are found in 10.0 g of
sodium?
13 L of hydrogen at STP has a mass of
___g.
49.
Molar Mass
grams/mole
Find the molar mass of potassium
nitrate.
50.
Percent Composition
% composition =
Find the percent magnesium in magnesium
oxide. Take the atomic mass and divide by the molecular weight X 100%.
51.
Empirical & Molecular Formulas
Empirical Formula – formula that gives the lowest whole number ratio of atoms in the compound (reduced version of
formula)
Molecular Formula – the actual formula of the compound (the way it exists naturally)
To calculate E.F.:
1. Assume 100 g of the substance so the percentage equals the # of grams of each element
2. Change grams to moles
3. Divide by the lowest number of moles
4. The numbers you get are the subscripts in your formula.
To calculate M.F.:
Divide the molar mass (given) by the empirical mass. Multiply that number by the E.F.
A compound is 80% C and 20% H. If the molar mass of the molecular formula is 45 g/mol, find the empirical and molecular
formulas of this compound.
52.
Stoichiometry
*Must have a balanced equation to solve these problems!
Remember: grams to moles, mole ratio, moles to grams
H2 + O2  H2O
How many grams of water will be produced from 5.0 g of hydrogen?
53.
54.
Kinetic Molecular Theory
The Major Points
Temperature is related to kinetic energy. The hotter it is the faster the particles move.
Gas particles are in constant random motion
Gas particles have no volume
Collisions are perfectly elastic (no transfer of energy).
Gas Laws – Know how they are used. The formulas will be on the exam
Boyle’s
Charles’
Avogadro’s
Combined
55.
Dalton’s
Graham’s
Gay-Lussac’s
Gas Stoich.
Ideal Gas Law
PV = nRT
Remember: No change occurs!
P = pressure in atm or kPa
V = volume in L
N = Moles
R = constant (0.0821 L.atm/mol.K OR 8.314 L.kPa/mol.K)
T = temperature in K
56.
Endothermic Reactions
Heat is added. It appears on the left side of the equation. The quantity of heat will be in Joules. ∆H = +
57.
Exothermic Reactions
Heat is released. It appears on the right side of the equation. The quantity of heat will be in Joules. ∆H = -
58.
Activation Energy
The energy required to make a reaction take place. A catalyst lowers the activation energy.
59.
Reaction Progress Diagrams
60.
Phase Diagrams
61.
Heating Curves
Temperature does not change during a phase
change!
How much energy is required to melt 15.0 g of
ice if the heat of fusion for water is 6.02 J/g?
How much energy is required to raise the
temperature of 15.0 g of water from 10 C to 25 C?
62.
Spontaneity
Spontaneous Reactions – happen without putting
a lot of energy in (have low activation energy)
Non-spontaneous Reactions – must put in energy
in order for the reaction to occur ( have high
activation energy)
63.
Enthalpy
Heat!
*If is negative, the reaction is exothermic. If is positive, the reaction is endothermic.
Hrxn = Hproducts – Hreactants
64.
Entropy
Disorder!
*If S is a large positive number, the system is very chaotic. If S is a low number, the system is very organized.
S = Sproducts – Sreactants
65.
Free Energy
G =  - TS
66.
Kinetics
Kinetics - Study of the rate of a reaction
What are four things that affect the rate of a reaction? What is the collision theory?
Know how to write the Rate Law for a reaction based on the order of each component of the reaction. Determine order from
experimental data.
67.
Catalysts
*Increase the rate of a reaction by _____.
*Not used up in a reaction.
68.
Electrolytes
An electrolyte conducts electricity in solution.
STRONG ELECTROLYTES
WEAK ELECTROLYTES
NaCl
HCl
Dissociate 100%
69.
HC2H3O2
Weak acids and bases
Molarity
Molarity = moles of solute/L of solution
Calculate the molarity of a solution in which 15.0 g of NaCl is dissolved in 100 mL of water.
70.
Dilution
Molarity1 x Volume1 = Molarity2 x Volume2
What volume of a 4.0 M HCl solution should be used to make 100 mL of a 0.15 M HCl solution?
71.
Solubility Rules
You will not have to have the solubility rules memorized, they will most likely be on a chart in some form. Don’t forget:
Soluble – will dissolve (becomes aqueous in solution)
Insoluble – will not dissolve (stays in solid form in solution)
72. Solubility Curves – Just read the graph to determine the number of grams of a salt that will dissolve at a given temp.
This produces a saturated solution.
How many grams of NaNO3 will dissolve in 100 g of water at 20 C?
A supersaturated solution of KNO3 at 50 C would have more than ___ g of solute in
solution.
How many grams of KI will dissolve in 400 g of solution at 10 C?
73.
Precipitation Reactions
A reaction in which a precipitate forms. A precipitate is a solid.
Write the molecular, complete ionic, and net ionic equations for the following:
NaCl(aq) + AgNO3(aq) 
74.
Colligative Properties
Properties that depend on how much solute is present Know how to use the formulas
75.
Chemical Equilibrium
Equilibrium –
Reversible reactions –
Homogeneous equilibria –
Heterogeneous equilibria –
76.
Writing Equilibrium Expressions
Only include products and reactants that are aqueous or gaseous, solids are not included.
Write expressions for:
NO2(g)
NO(g) + O2(g)
Na+(aq) + Cl-(aq)
NaCl(s)
77.
LeChatelier’s Principle
A reaction at equilibrium wants to stay at equilibrium. To accomplish this, the reaction will shift to the left or right to
maintain equilibrium when a change is made.
78.
Acids
Properties of Acids
79. Bases
Properties of Bases
Arrhenius Bases Arrhenius Acids Bronsted-Lowry Bases Bronsted-Lowry Acids 80.
pH
On the SOL, the base will always be 1!!
If the [H+] is 1 x 10-5, the pH is 5
if the [OH-] is 1 x 10-4, the pH is 10
82.
Strong Acids
Strong Acids dissociate completely!
HCl
HBr
HNO3 H2SO4
All others are weak (therefore dissociate partially)
81.
pOH
On the SOL, the base will always be 1!!
If the [H+] is 1 x 10-5, the pOH is 9
if the [OH-] is 1 x 10-4, the pOH is 4
83.
Strong Bases
Strong Bases dissociate completely!
Any base with an alkali metal is strong. All others
are weak (therefore dissociate partially)
84. Titrations
Add acid to base to find the molarity of either the acid or the base.
An indicator changes color to show the endpoint of the titration.
85.
Half Life
A sample of element X has a half life of 8 days. If you start with 200 g of the sample, how much is left after 40 days?
86.
Organic Chemistry Organic molecules have carbon. You cannot be asked anything specific to organic molecules, however you will most
likely see organic molecules in other questions. You probably want to look at the functional groups.