Name:
Teacher:
Period:
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Metric System
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Temperature
Kelvin (K)
Miscellaneous Conversions
0K = - 273C
0K = - 460F
Amount of Substance
Mole (mol)
1 mol = 6.022 x 1023 (Avogadro’s number)
Volume
@ STP
22.4 L / mol of gas
Pressure
Energy
Pascal (Pa)
Joule (J)
K = C + 273
1 atm = 101.325 kPa = 760 mm Hg = 760 torr = 14.70 psi
1 calorie (cal) = 4.184 J
Formulas, Measurements, and Mathematics
1.
Conversion Factor: a ratio of equivalent measurements, including units, that is expressed as a fraction
2.
Percent Error = │experimental value – accepted value│ x 100
accepted value
3.
Temperature Conversions
4.
Density (D) = mass__
volume
5.
Average Atomic Mass: the weighted average mass of the naturally occurring isotopes of an element
6.
Percent Composition = mass of element x 100
mass of compound
7.
Mole Ratio: a conversion factor derived from the coefficients of a balanced chemical equation
8.
Empirical Formula: % to Mass, Mass to Mole, Divide by Small, Multiply til Whole
9.
Molecular Formula = molecular formula mass = multiplier
empirical formula mass
°C = K – 273
K = °C + 273
10. Percent Yield = actual yield___ x 100
theoretical yield
11. Combined Gas Law
P1V1T2 = P2V2T1
12. Ideal Gas Law
PV = nRT
Where R = Universal Gas Constant = 0.0821 L ∙ atm/mole ∙ K OR
Where n = moles
8.314 dm3 ∙ kPa/mol ∙ K
13. Dalton’s Law of Partial Pressure PTotal = P1 + P2+ P3 + …
14. Percent Water in a Hydrate =
mass of water_ x 100
mass of hydrate
15. Molarity (M) = moles of solute
liter of solution
16. Dilution
Molality (m) = moles of solute____
kilograms of solvent
M1V1 = M2V2
17. Titration
MAVANA = MBVBNB
18. Heat is represented by two values: q or ΔH
19. Specific Heat (C)
20. Acids/Bases
pH = -log [H+]
Where 1 = initial and 2 = final
ΔHsystem = - ΔHsurroundings or qsystem = -qsurroundings
q = mCΔT
[H+] = 10 –pH
Where A = acid and B = base
pOH = -log [OH-] [OH-] = 10 –pOH
Where ΔT = Tfinal - Tinitial
pH + pOH = 14
21. Equilibrium Constant (Keq): a ratio of the concentration of the products to the reactants where the coefficient if each
substance in the balanced equation becomes an exponent for that concentration
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Keq =
[products]x__
[reactants]x
x is the molar coefficients from a balanced equation
NUMBERS WRITTEN in SCIENTIFIC NOTATION
In chemistry, we often have to deal with either very large numbers, such as 602 300 000 000 000 000 000 000
000 (the number of things in a mole), or very small numbers, such as 0.000 000 000 000 000 000 000 000 000
911 grams (the mass of an electron). Such numbers are much more conveniently expressed as some number
between 1 and 10 times an exponential term. The exponential term contains an exponent which is a base. The
base is usually 10. For example:
1 = 1 x 100
0.2 = .2/101 = 2 x 10-1
1
30 = 3 x 10
500 = 5 x 102
0.007 = 7/103 = 7 x 10-3
When calculating with numbers written in scientific notation using a scientific calculator, one must be careful
entering the values. For example:
The number 6.022 x 1023 is entered as follows:
1. Clear your calculator display
2. Enter 6.022
3. Hit the KEY marked EE, EXP
4. The display will look like this: 6.022 00
5. Enter in the exponent, 23
6. The display will look like this: 6.02 23
7. If a negative exponent is needed, then
8. Hit the  key
9. The display looks like this: 6.022 -23
For adding, subtracting, multiplying, or dividing numbers in scientific notation and using a calculator, it is vital
that you DO NOT ENTER “TIMES 10” for the “10x” portion of the number in scientific notation. “Times 10”
means multiply by 10, and “10x” means multiply the base number by 10 as many times as indicated by the
exponent “x”.
Writing a number in this form is a “shorthand” method of expression. The notation 6.022 x 1023 tells us to
multiply 6.023 by 10 twenty-three times. The notation 9.11 x 10-28 tells us to divide by 10 twenty-eight times.
Examples of converting a scientific notation number into the expanded version.
1.23 x 102
6.78 x 10-3
Positive exponent shifts the decimal to the right
Negative exponent shifts the decimal to the left.
Rules for Counting Significant Figures
1. Nonzero integers. Nonzero integers always count as significant figures.
2. Zeros. There are three classes of zeros
a. Leading zeros are zeros that precede all the nonzero digits. These do not count as significant
figures. In the number 0.0025, the three zeros simply indicate the position of the decimal point. This number
has only two significant figures.
b. Captive zeros are zeros between nonzero digits. These always count as significant figures. The
number1.008 has four significant figures.
c. Trailing zeros are zeros at the right end of the number. They are significant only if the number
contains a decimal point. The number 100 has only one significant figure, whereas the number 1.00 x 102 has
three significant figures. The number one hundred written as 100. also has three significant figures.
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3. Exact numbers. Many times calculations involve numbers that were not obtained using measuring devices
but were determined by counting: 10 experiments, 3 apples, 8 molecules. Such numbers are called exact
numbers. They can be assumed to have an infinite number of significant figures.
RULES for SIGNIFICANT FIGURES in MATHEMATICAL OPERATIONS
1. For multiplication or division the number of significant figures in the answer is the same as the number in
the least precise measurement used in the calculation. For example, consider this calculation:
4.56 x 2.4 = 6.36
(corrected)  6.4


least precise
two significant figures
has two significant figures
allowed in answer
The product should have only two significant figures, since 2.4 has two significant figures.
2. For addition or subtraction the result has the same number of decimal places as the least precise
measurement in the calculation. For example, consider the sum:
12.11
18.0  limiting term has one decimal place
1.013
31.123 (corrected to)  31.1
\ one decimal place
The correct result is 31.1, since 18.0 has only one decimal place.
Note that for multiplication and division, significant figures are counted. For addition and subtraction,
the decimal places are counted.
In most calculations, you will need to round numbers to obtain the correct number of significant figures. The
following rules should be applied when rounding.
RULES for ROUNDING
1.
In a series of calculations, carry the extra digits through to the final result, then round.
2.
If the digit is to be removed
a) is less than 5, the proceeding digit stays the same. For example: 1.33 rounds to 1.3
b) is equal or greater than 5, the proceeding digit is increased by 1. For example: 1.36 rounds to 1.4
Although rounding is generally straightforward, one point requires special emphasis. As an illustration, suppose
that the number 4.348 needs to be rounded to two significant figures. In doing this, we look only at the first
number to the right of the 3.
4.348
\ look at this ‘4’ to round number to 2 significant figures
The number is rounded to 4.3 because 4 is less than 5. It’s incorrect to round sequentially. In other words,
use only the first number to the right of the last significant figure.
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Aufbau Diagram
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Naming Compounds
Inorganic Compounds
 Metal ion plus non-metal ion
With two ions: -ide ending
Example: Na+ and Cl- = NaCl
sodium chloride
*notice
the first
name
the compound
doesending
not change!
With
a that
variable
ion
andinanother
ion: -ide
Example: Fe2+ and S2- = FeS
iron (II) sulfide
*notice that the variable ion has its charge (II) indicated
With an ion and a polyatomic ion:
Example: Ca2+ and (SO4)2- = Ca(SO4)
calcium sulfate
*the name of the polyatomic ion does not change
With a variable ion and a polyatomic ion:
Example: Cu2+ and (CrO4)2- = Cu(CrO4)
copper (II) chromate
* the variable ion has its charge (II) indicated
Inorganic
* the nameMOLECULAR
of the polyatomiccompounds
ion does not change
 Two non-metal ions forma a molecular compound
 HYDROGEN is often considered a non-metal (ex.1)
 These compounds use prefixes instead of roman numerals (such as mono-, di-, tri-, tetra-, etc.)
 Mono- is not used when the first ion in the compound is singular (ex. 3 & 4)
 The second member of the compound ends in –ide
Example 1: H2O
dihydrogen monoxide
Example 2: N2O3
dinitrogen trioxide
Example 3: CO2
carbon dioxide
ACIDS:
Monoprotic Acids (1 available H+)
Hydrochloric Acid
Nitric Acid
Nitrous Acid
Acetic Acid
HCl
HNO3
HNO2
HCH3COO
Diprotic Acids (2 available H+)
Sulfuric Acid
Sulfurous Acid
Carbonic Acid
Hydrosulfuric Acid (hydrogen sulfide)
H2SO4
H2SO3
H2CO3
H2S
Tripotic Acids (3 available H+)
Phosphoric Acid
H3PO4
BASES:
Ammonia
Sodium Hydroxide
Potassium Hydroxide
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NH3
NaOH
KOH
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Electronegativity of Elements
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Activity Series List
Potential Cations
Lithium
Potassium
Rubidium
Cesium
Radium
Barium
Strontium
Calcium
Sodium
Lanthanum
Cerium
Neodymium
Samarium
Gadolinium
Magnesium
Yttrium
Americium
Lutetium
Scandium
Plutonium
Thorium
Neptunium
Beryllium
Uranium
Hafnium
Aluminum
Titanium
Zirconium
Manganese
Vanadium
Niobium
Selenium
Zinc
Chromium
Gallium
Tellurium
Iron
Cadmium
Indium
Thallium
Cobalt
Nickel
Molybdenum
Tin
Lead
Hydrogen
Copper
Mercury
Silver
Rhodium
Palladium
Platinum
Gold
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Potential Anions
Flourine
Chlorine
Bromine
Iodine
The most active
potential
cations/anions
are at the top of
each list.
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Solubility Rules
Chloride,
Bromide, Iodide:
Cl-, Br-, I-
Negative ion
(anion)
Plus
Positive Ion
(cation)
Forms a
compound
Which
is
1. Any ion
+
Alkali ions (Li+, Na+, K+, Rb+, or Cs+)
“
Soluble,
i.e.  0.1 mol/liter
2. Any ion
+
Ammonium, NH4+
“
Soluble
3. Nitrate, NO3-
+
Any cation
“
Soluble
4. Acetate, CH3COO-
+
Any cation except Ag+
“
Soluble
5. Chloride, Cl- or
Bromide, Br- or
Iodide, I-
+
+
+
Ag+, Pb2+, Hg2+, or Cu+
Any other cation
“
“
Not Soluble
Soluble
6. Sulfate, SO42-
+
+
Ca2+, Sr2+, Ba2+, Ra2+, Ag+, or Pb2+
Any other cation
“
“
Not soluble
Soluble
7. Sulfide, S2-
+
+
+
Alkali ions or NH4+
Be2+, Mg2+, Ca2+, Sr2+, Ba2+, or Ra2+
Any other cation
“
“
“
Soluble
Soluble
Not Soluble
8. Hydroxide, OH-
+
+
+
Alkali ions or NH4+
Sr2+, Ba2+, or Ra2+
Any other cation
“
“
“
Soluble
Soluble
Not Soluble
9. Phosphate, PO43+ or
Carbonate, CO32- or
Sulfite, SO32-
+
+
Alkali ions or NH4+
Any other cation
“
“
Soluble
Not Soluble
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