Download Example 1-2

1
Matter – Its Properties
and Measurement
Section Objectives
1. Write the names and chemical symbols of the more common elements.
This is a memorization task: to know the symbol that goes with the name and vice versa.
“Common elements” means different things to different people. A reasonable goal would be the
main group elements along with those in the first transition series (Sc through Zn) plus Ag, Au,
Cd, and Hg. These are elements with atomic numbers 1-38, 47-56, and 79-88. The atomic
number is the whole number in each box in the periodic table, often given the symbol Z. The
best way to learn names is by groups (vertical columns) in the periodic table. By learning them
together, you are also learning chemistry, not just memorizing symbols.
2. Know common units in the English system, and the relationships between them.
Common English units, their abbreviations, and the relationships between them follow. They
should be memorized if you do not know them already. Most scientists, and practically all
engineers, in the U.S. do.
Volume measure: Gallon (gal) = 4 quarts (qt) Pint (pt) = 16 fluid ounces (fl oz) = 2 cups (c)
Quart (qt) = 2 pt = 32 fl oz Tablespoon (T) = 2 teaspoons (tsp) = 1/2 fl oz
Linear measure:
Yard (yd) = 3 feet (ft)
Mile (mi) = 1760 yd
Foot (ft) = 12 inches (in)
Mass measure:
Pound (lb) = 16 ounces
Ton (t) = short ton = 2000 lb
(oz)
3. For the metric system, state the basic units of mass, length, and volume, and the
common prefixes.
At present there are three units to be learned: gram (g) for mass, meter (m) for length, and liter
(L) for volume. The following prefixes should be memorized.
Mega (M)
or
1,000,000
One million of
106
(k)
103
or
1000
One thousand of
Deci
(d)
10
-1
or
1/10
One tenth of
Centi
(c)
10-2
or
1/100
One hundredth of
Milli
(m)
10
-3
or
1/1000
One thousandth of
Micro
(μ)
10-6
or
1/1,000,000
One millionth of
Nano
(n)
10-9
or
1/1,000,000,000
One billionth of
Kilo
4. State the relationships between English and metric units.
There are many relationships, but they can be reduced to only three if you know the
interrelationships within each system. These three can be one from each column that follows.
Mass
453.6 g = 1 lb
1 kg = 2.205 lb
28.35 g = 1 oz
Length
2.5400 cm = 1 in
1 m 39.37 in
30.48 cm = 1 ft
Volume
0.9464 L = 1 qt
1 L = 1.057 qt
29.57 mL = 1 fl oz
5. Determine the number of significant digits in a numerical calculation.
6. Express the result of a calculation with the appropriate number of significant digits.
Here are the rules that govern significant digits.
a) Significant digits include all non-zero digits, zeros located between significant digits
(“captive zeros”), and zeros located after non-zero digits to the right of the decimal point
(“trailing zeros”).
b) Zeros that precede non-zero digits and zeros that end a number with no decimal point are
not significant (“placeholders”).
c) The result of a multiplication or division should have as many significant digits as the
factor with the fewest number of significant digits.
d) The result of an addition or a subtraction is rounded to the same number of decimal
places as the term with the fewest number of decimal places.
EXAMPLE 1-1
Express the result of the following calculation with the appropriate number of significant digits.
[(725.6 - 19.1)/760. 00]75 x 10-3
 ? The result is 2.9 x 10-3
(0.082057)( 293.2)
7. Express numbers in scientific notation.
8. Write a conversion factor from a relationship between two quantities, and use
conversion factors to solve problems.
Probably the most powerful problem-solving method you can learn is the conversion factor
method. The method is valuable no only in chemistry but in any numerical problem-solving
course.
EXAMPLE 1-2
How many teaspoons are in 5.00 gallons?
4 qt 32 fl oz
3 tsp
5.00 gal x
x
x
 3.84 x 103 tsp
1 gal
1 qt
0.5 fl oz
EXAMPLE 1-3
How many gallons are in 1.00 cubic foot?
3
3
1L
1 qt
1 gal
 12 in   2.5400 cm 
1.00 ft 3 x 
x
x
 7.48 gal
 x
 x
 1 ft   1 in
 1000 cm3 0.946 L 4 qt
9. Express and use density in the form of conversion factors.
Density is both a physical property of a substance and the means of interconverting mass and
m
volume of that substance. The defining equation ( d  ) has three variables: density, mass, and
v
volume.
EXAMPLE 1-4
An empty container weighs 206 g. Filled with 242 mL of liquid it weighs 938 g. What is the
density of the liquid?
(938 g  206 g)
d
 3.02 g/mL
242 mL
EXAMPLE 1-5
A 27.4 mL gold (19.3 g/mL) object has what mass?
19.3 g
27.4 mL x
 529 g Au
1 mL
EXAMPLE 1-6
A 75.2 g piece of zirconium (6.42 g/mL) has what volume?
1 mL
75.2 g x
 11.6 mL Zr
6.42 g
10. Express and use percent composition in terms of conversion factors.
Percent means part per hundred. Thus, 40.0% C in acetic acid means 40.0 g C in 100.0 g acetic
acid.
EXAMPLE 1-7
What mass of acetic acid contains 247 g C?
100.0 g acetic acid
247 g C x
 618 g acetic acid
40.0 g C
11. Solve algebraic equations that arise in the course of working chemistry problems.
Solving an algebraic equation generally means obtaining a new equation, with the symbol for
one variable isolated on one side and the remainder of the equation of the of the equation on the
other side.
EXAMPLE 1-8
Solve the following equation for P.
2

 P  an V - nb   nRT


V2 

P
nRT
an 2
V - nb  V 2
2
Atoms and the
Atomic Theory
Section Objectives
1. List the numbers of protons, neutrons, and electrons present in atoms and ions, using
the symbolism A
ZE.
The complete symbol for an atom or ion consists of the elemental symbol surrounded by
subscripts and superscripts.
a) The leading superscript (upper left) is the mass number. This is also the number of
nucleons; a nucleon is a proton or a neutron.
b) The leading subscript (lower left) is the atomic number or proton number.
c) The trailing superscript (upper right) is the charge or the number of protons (atomic
number) minus the number of electrons. The sign (+ or -) always must be included. The
number is zero for a neutral atom, but the zero is written only for emphasis.
Mass
Charge
number
235 U3
92
Atomic number
EXAMPLE 2-1
What is the atomic number, mass number, and charge of 19 F  ?
19 F  has an atomic number of 9; a mass number of 19; and a charge of –1.
2. Use the periodic table to predict the charges of ions of main group elements.
Elements in the same column have similar properties. Each column is referred to as a periodic
family or group. The horizontal rows are called periods. Elements on the right side of the
periodic table are nonmetals; they form anions, or negatively charged ions. Elements on the left
side of the periodic table are metals; they form cations, or positively charged ions. Elements
within the same group will form ions with the same charge.
3
Chemical Compounds
Section Objectives
1. Know and apply the conventions used in determining oxidation states.
Because oxidation state is a formal rather than an experimental concept, it is possible to devise a
rigid set of rules that work in all but the most unusual circumstances. One set of rules is given
below.
Method of applying the rules
1. Apply the rules from the top to the bottom of the list.
2. Search the list to find a rule that fits. Apply it.
3. Then start again at the top of the list to find the next rule that fits.
Oxidation State rules
a. The OS (oxidation state) of all uncombined elements = 0.
b. The sum of the OS in compound = 0.
c. The sum of the OS in an ion = ionic charge.
d. Alkali metals (group 1A) have OS = +1.
e. Alkaline earth metals (2A) have OS = +2.
f. F has OS = -1 and H has OS = +1.
g. O has OS = -2.
h. Cl, Br, I (in order) have OS = -1.
i. S, Se, Te (in order) have OS = -2.
j. N, P, As (in order) have OS = -3.
2. Know the names, formulas, and charges of ions in the following tables and be able to
write formulas and names of the compounds formed from these ions.
Al
3+
H-
Names, Formulas, and Charges of Some Common Ions
Aluminum
Iron (III) or ferric
Au3+ Gold (III) or auric
Fe3+
Hydride
Sn2+ Tin (II) or stannous
Co2+ Cobalt (II) or cobaltous
Mn2+ Manganese (II)
Ni2+ Nickel (II)
Sn4+
Pb
Lead (II) or plumbous
Cu
Zn2+
Zinc
Pb4+
Cu2+
Cd
Cadmium
Cr2+
Ag+
Silver
Cr3+
Lead (IV) or plumbic
Chromium (II) or
chromous
Chromium (III) or
chromic
Au+
Gold (I) or
aurous
Fe2+
Iron (II) or ferrous
2+
2+
Tin (IV) or stannic
Co3+
+
Cobalt (III) or cobaltic
Copper (I) or cuprous
Copper (II) or cupric
Mercury (I) or
Hg22+ mercurous
Mercury (II) or
Hg2+ mercuric
NH4
Names, Formulas, and Charges of Some Common Polyatomic Ions
Ammonium
Sulfate
Hypofluorite
SO42FO-
C2H3O2-
Acetate
HSO4-
Hydrogen sulfate
ClO-
CO32-
Carbonate
SO32-
Sulfite
ClO2-
Chlorite
-
Chlorate
+
-
Hypochlorite
HCO3C2O42-
Hydrogen carbonate
HSO3
Hydrogen sulfite
ClO3
Oxalate
S2O32-
Thiosulfate
ClO4-
CN
Cyanide
HS
-
OCN-
Cyanate
SCN-
Thiocyanate
Perchlorate
Hydrogen sulfide
BrO
OH-
Hydroxide
BrO3-
Bromate
O22-
Peroxide
BrO4-
Perbromate
Chromate
IO
-
Hypoiodite
Dichromate
IO3-
NO2
-
Nitrite
CrO4
NO3-
Nitrate
Cr2O72-
3-
-
2-
-
PO4
Phosphate
MnO4
HPO42-
Hydrogen phosphate
MnO42-
H2PO4-
Dihydrogen phosphate
Permanganate
IO4
-
Hypobromite
Iodate
Periodate
Manganate
The lists presented above may seem rather extensive, but they contain practically all the ions you
are likely to encounter in AP Chemistry. (You may get a few more in September, but this will
give you a good start!)
Naming an ionic compound is simple. Write down the name of the cation (positive ion),
followed with a space, and then the name of the anion (negative ion).
Writing formulas from names is not quite so simple. The formula contains more than just the
symbols for the cation and anion. The cation and anion symbols are multiplied so that the total
charge from the cations just balances the total charge of the anions. The total cation charge plus
the total anion charge equals zero.
Cations
M+
[NH4+]
-
X
MX
-
[F ]
2-
X
Anions
[SO42-]
M
2+
M
3+
[Al ]
[NH4F] M2X [(NH4)2SO4]
MX2 [CaF2] MX
[CaSO4]
MX3 [AlF3] M2X3 [Al2(SO4)3]
M
4+
[Ce4+]
MX4 [CeF4]
2+
[Ca ]
3+
MX2
[Ce(SO4)3]
X3-
[PO43-]
M 3X
[(NH4)3PO4]
M3X4 [Ce3(PO4)4]
M3X2 [Ca3(PO4)2]
MX
[AlPO4]
3. Be able to write formulas and names of simple binary covalent compounds and of
binary acids.
Covalent compounds are formed between nonmetallic elements. The names of binary covalent
compounds are obtained from the names of the two elements. The elements are named in the
same order as they appear in the formula. The first element name is unchanged; the ending of
the second becomes “-ide.” The element names have prefixes depending on the subscript of
that element in the formula, except that the prefix mono- (meaning one of) is rarely used for the
first element in a formula. Other prefixes are: di = 2, tri = 3, tetra = 4, penta = 5, hexa = 6, hepta
= 7, octa = 8, nona = 9, and deca = 10.
Binary acids consist of hydrogen and a nonmetal. HCl is a binary acid. The name of a binary
acid has the prefix “hydro-” and the suffix “-ic” surrounding the root name of the element. HCl
is hydrochloric acid. The binary acid names are used when the compound is dissolved in water,
that is, in aqueous solution. When the compound is not an aqueous solution the name is the same
as any ionic compound.
4. Use oxidation states to name oxoacids and oxoanions.
Salts and acids of chlorine oxoanions
Ox.
State
+1
Salt
Example
Acid
Example
Hypo- -ite
NaClO
Sodium hypochlorite
NaClO2
Sodium chlorite
NaClO3
Sodium chlorate
NaClO4
Sodium perchlorate
Hypo- -ous
HClO
Hypochlorous acid
HClO2
Chlorous acid
HClO3
Chloric acid
HClO4
Perchloric acid
+3
-ite
+5
-ate
+7
Per-
-ite
-ous
-ic
Per-
-ic
All oxoanions of the same family with the same oxidation state have similar names. Another
generality is that the –ate anion and the –ic acid endings are used when the oxidation state of the
central atom equals the periodic table family number. The only exceptions to this occur in the
halogens, where the –ate and the –ic endings correspond to a +5 oxidation state and the noble
gases where they correspond to +6.
5. Use solubility rules to predict products of reactions.
The rules on the next page are one form of the solubility rules. You are responsible for
learning these rules in some format. Each textbook gives a slightly different approach.
SOLUBILITY RULES
LEARN!!
The solubility of a solute is the amount that can be dissolved in a given quantity of solvent at a
given temperature. For example, the solubility of lead (II) nitrate is
56 g/100 mL at 20oC. The solubilities of ionic solids in water vary over a wide range of values.
For convenience, we divide compounds into three categories called soluble, slightly soluble and
insoluble. Insoluble is a relative term and does not mean that no solute dissolves! Compounds
are classified as insoluble if their solubility is less than
0.1 g/100 mL of water. On the other hand, soluble compounds are those whose
solubilities are greater than 1.0 g/100 mL of water. The following “solubility rules” summarize
the solubilities of various compounds in water at 25oC.
1. All Group IA salts are soluble (aq).
2. All ammonium salts are soluble (aq).
3. All salts containing nitrate, acetate, chlorate and perchlorate are soluble (aq).
4. All salts containing halides (chlorides, bromides, iodides, and fluorides) are
soluble (aq) EXCEPT silver, mercury(I) and lead (s). (Lead halides are soluble in
hot water.)
5. All sulfate salts are soluble (aq) EXCEPT barium, calcium, strontium, silver, mercury(I) and
lead (s).
6. All salts containing carbonates, phosphates, and chromates are insoluble (s) EXCEPT
for rules #1 and 2 (aq).
7. All sulfide salts are insoluble (s) EXCEPT for rules #1 and 2 and calcium, strontium, and
barium (aq).
8. All hydroxide salts are insoluble (s) EXCEPT for rules #1 and 2 and barium and strontium
(aq). (Calcium hydroxide is very slightly soluble.)
Note: Rule #8 is the one that varies from text book to text book and causes the most trouble for
people writing net ionic equations. Are the Group IIA hydroxides soluble or not? At best they
are only moderately soluble – barium and strontium are a little more soluble than calcium and
usually are called soluble. Calcium hydroxide is usually called insoluble.