1. business and industrial
2. chemicals industry
3. plastics and polymers

CHEMISTRY 11 Notes - Mrs. Jeffrey
This is a compilation of miscellaneous notes that I have assembled for use in the CHEM 11
course. I do not expect you to just print out every page in a wholesale fashion, as you can
always access this from your computer whenever the need arises.
I will not remove this file from my homework page before the end of the course. It will always
be accessible to you from the web (either at home or at school).
These pages appear in the order they are used in the course. Chapters may appear to be out
order. Study Worksheet for Chemistry Math Review
1. Completely read the chapter.
2. Define the Key Terms at the end of the chapter.
3. Be familiar with the scientific method (P. 5).
4. Know how to apply the rules for determining the number of significant figures (s.f.) in a number. Also know the
3 rules that aren't in the book: a) a mathematical definition or equivalence (e.g. exactly 12 in. = 1 ft), has
infinite s.f., b) an integer number (e.g. 3 atoms, or 6 people) also has infinite s.f., and c) all of the digits in a
scientific notation number (e.g. 3.30 x 103) are always significant.
5. Know how many significant figures should be in the answer to any mathematical calculation, based upon the
number of s.f. in the problem (one rule is used for x and /, and another rule is used for + and - ).
6. Be able to round off any number to any number of significant figures.
7. Be able to convert any common number into scientific notation and vice versa.
8. Know all of the metric prefixes in bold on P. 20, and what they mean, so you can change one into another, (e.g.
100 cm = 1 m, 100 mg = 1 dg, 1000 uL = 1 mL, 10 cm = .1 m, etc.), and can tell which of 2 numbers are the
largest and by how much (i.e. 1 cm is 10 times larger than 1 mm).
9. Know the SI unit and its abbreviation for the first 5 items in Table 2.2 on P. 20.
10. Know how to solve dimensional analysis problems (like on P. 23 - 29).
11. Always answer all numerical questions with both the correct scientific units and the correct number of
significant figures.
12. 1 cm3 = 1 mL , which is also equal to 1 g when describing water only.
1,000 cm3 = 1 L , which is also equal to 1 kg when describing water only.
13. Know the 7 terms used to describe phase changes: melt, boil, evaporate, vaporize, sublime, freeze & condense.
14. Convert any temperature into the other 2 systems using:
(9/5)C = F - 32
and
K = C + 273
15. Density = mass / volume. Solve for any unknown if you are given the other 2 values.
16. When 2 substances are mixed, the higher density substance will always sink to the bottom.
17. You don't need to memorize the numbers, but have a good qualitative appreciation for which substances are the
most and least dense in Table 2.5 (P. 35).
These notes pertain to the Pictorial Period Table that is attached to the wall in my
classroom.
Important Notes about the Period Table
1) All of the Noble Gases (Group VIIIA) are colorless gases. In these
pictures, an electric current is arcing through the tubes containing
these gases, which then give off these bright colors that are
characteristic of each element
(e.g. red-orange for Ne, lavender for Ar, etc.).
2) Most of the elements above U do not exist in nature.
They are man-made and all quantities must be synthesized in the lab
from other elements. So these elements can’t be produced in large
quantities and usually decay radioactively within seconds.
Consequently, there often isn't a sufficient amount available -- even
to just photograph. Therefore, some of these elements don’t have any
picture, while others instead show a picture of a compound (usually
an oxide) made from that element. Most of these elements are
metals, so if we had pictures of the pure elements, they would look
like any other silver metal.
3) The RaO sample is green only because it had previously absorbed
light energy and is now slowly re-radiating its characteristic green
light. It was formerly used in watch dials that glow in the dark.
Chemical Identity of Common Household Materials
Household items are typically impure mixtures, with the following as the main ingredient (other than H2O):
Item
Acetylene welding
Ammonia
Antifreeze (car)
Aspirin
Baking soda
Bleach
Brass
Cane sugar
Car battery acid
Citrus fruits & candy
Cream of Tarter
Dental fillings
Drano, lye, oven cleaner
Dry ice
Epsom salts
Grain alcohol (drinking)
Low sodium salt
Milk of Magnesia
Mothballs
Muriatic acid
Formula
C2H2
NH3
C2H4(OH)2
C9H8O4
NaHCO3
NaClO
Cu, Zn
C12H22O11
H2SO4
H2C6H6O7
KHC4H4O6
Ag, Hg
NaOH
CO2
MgSO4.7H2O
C2H5OH
KCl
Mg(OH)2
C10H8
HCl
Name
acetylene or ethyne
ammonia (actually ammonia water)
ethylene glycol
acetylsalicylic acid
sodium bicarbonate
sodium hypochlorite
copper and zinc mixture
sucrose
sulfuric acid
citric acid or 2-hydroxy-1,2,3-propanetricarboxylic acid
potassium hydrogen tartrate
silver and mercury mixture
sodium hydroxide
carbon dioxide
magnesium sulfate heptahydrate
ethyl alcohol
potassium chloride (substitute salt)
magnesium hydroxide
naphthalene
hydrochloric acid
Nail polish remover
Natural gas
Peroxide
Rubbing alcohol
Rolaids
Rust
Soft Drinks
Spot remover
Stainless steel
Sterling silver
Table salt
Teflon
Tums, chalk
Vinegar
Vitamin A
Vitamin B12
Vitamin C
Vitamin E
12 karat gold
CH3COCH3
CH4
H2O2
C3H7OH
Al(OH)3
FeO, Fe2O3
H2CO3
C2H3Cl3
Fe, Cr, Ni, C
Ag, Cu
NaCl
[-CF2-CF2-]~1000
CaCO3
HC2H3O2
C20H30O
C63H88CoN14O14P
C6H8O6
C29H50O2
Au, Ag, Cu
acetone
methane (95%)
hydrogen peroxide
isopropyl alcohol
aluminum hydroxide
iron II oxide and iron III oxide
carbonic acid (H2O + CO2)
trichloroethane
iron, chromium, nickel and carbon mixture
silver and copper mixture
sodium chloride
polytatrafluoroethylene
calcium carbonate
acetic acid
(very long name)
5,6-dimethylbenzimidazolylcyanocobamide
ascorbic acid
2,5,7,8-tetromethyl-2-(4',8'12'-trimethyltridecyl)-6-chromanol
gold, silver and copper mixture
Study Worksheet for Chapter 11 1. Define an orbital – a geometric figure resulting from quantum mechanics that predicts where an electron is most
likely to be with respect to the nucleus. Each orbital can hold up to 2 e-s.
2. Using this orbital chart, write out the full electron configuration for all elements up through Kr (36).
E.g. Br = 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5
3. What does 2p4 mean? There are 4 electrons in the 2p orbitals (and there are always 3 p orbitals – one in each of
the x, y & z directions). Two electrons are in the first 3p orbital, one electron is in the second 3p orbital, and one
electron is in the third 3p orbital.
4. For any element, identify how many electrons or in the highest energy level. E.g. Br = 4p5
5. Given an electronic configuration, identify the atom. (just add up the exponents)
E.g. 1s2 2s2 2p4 -- 2+2+4=8 electrons = 8 protons = Oxygen (atomic number = 8).
6. Given an electronic configuration, state whether its in its ground state, its excited state, or its illegal. E.g.
1s2 2s8 2p6 3s2 3p6 4s2 3d10 4p5
Illegal -- an s orbital can’t hold more than 2 electrons
2
2
6
2
10
5
1s 3s 3p 4s 3d 4p
Excited -- the entire 2nd principle level was left out (electrons promoted)
1s2 2s2 2p2 3s2 3p2 4s2 3d10 4p5
Excited -- the lower p orbitals aren’t fully filled (electrons promoted)
2
2
6
2
6
2
1s 2s 2p 3s 3p 3d
Excited -- the 4s orbital fills before the 3 d orbital (electrons promoted)
1s2 1p6 2s2 2p6 3s2 3p6
Illegal -- there is no such thing as a 1p orbital
1s2 2s2 2p6 2d10 3s2
Illegal -- there is no such thing as a 2d orbital
7. Given any group in the periodic table, identify what is common to the outermost electron shell of each element.
E.g. in Group IA, all elements have only 1 electron in the outer s orbital.
8. What do s orbitals look like? p orbitals?
9. The 4th principle energy level can hold up to how many electrons? 2 + 6 + 10 + 14 = 32.
10. Compare the major differences between the Thomsom model, the Rutherford model, the Bohr model and the
Quantum Mechanical model of the atom.
11. Define quantum – a small, discrete quantity of energy in an atom. One or more quanta units of energy are
required to move electrons between the various atomic energy levels.
12. In any particular principle energy level, how many s orbitals are there? (1), p orbitals? (3), d orbitals? (5), f
orbitals? (7) How many electrons can fit in any orbital? (2). Therefore, how many electrons can fit in the one s
orbital (2), the three p orbitals (6), the five d orbitals (10), the seven f orbitals (14)?
13. How does filling effect the stability of an orbital? Fully-filled and fully-empty orbitals are very stable. Halffilled are somewhat stable. Any other kind of partially-filled orbital is unstable.
14. Explain why atoms emit a color spectrum. Electrons moving from higher energy levels to lower energy levels
emit light (photons). The lines in emission spectra can be correlated to the movement of these electrons back to
their lower energy levels (ground states). Why is each atom’s spectra unique?
15. Give electron configurations for elements beyond Kr (36).
Principal Energy Levels, Sublevels and Orbitals - Chapter 11
•
•
•
Principal
Energy Level
1
2
Number of
Sublevels
1
2
3
3
4
4
5
5
6
6
7
7
8
Etc.
8
Etc.
Sublevel
Name
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
(5g)
6s
6p
6d
(6f)
(6g)
(6h)
7s
7p
(7d)
(7f)
(7g)
(7h)
(8s)
Etc.
# of Orbitals in
this Sublevel
1
1
3
1
3
5
1
3
5
7
1
3
5
7
(9)
1
3
5
(7)
(9)
(11)
1
3
(5)
(7)
(9)
(11)
(1)
Etc.
# of electrons in
this Sublevel
2
2
6
2
6
10
2
6
10
14
2
6
10
14
(18)
2
6
10
(14)
(18)
(22)
2
6
(10)
(14)
(18)
(22)
(2)
Etc.
Be careful to distinguish between the words: principal energy level, sublevel and orbital.
These words have different meanings and are not interchangeable.
The sublevels in parentheses are not used in the ground state of the present 109 elements.
When more elements are discovered (i.e., synthesized), they will then begin to use these sublevels.
Also, when electrons are excited, they may briefly occupy these levels, then return to their ground states.
Generally, energy increases as you go from the top line down through the lower lines, but not exactly.
Beginning at the 3d sublevel, the sublevels no longer fill in the exact order indicated.
The correct order can always be obtained from the Periodic Table.
Electron Configurations and Orbital Diagrams - Chapter 11
Z
El
Electron Configuration
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
H
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
1s1
1s2
1s2 2s1
1s2 2s2
1s2 2s2 2p1
1s2 2s2 2p2
1s2 2s2 2p3
1s2 2s2 2p4
1s2 2s2 2p5
1s2 2s2 2p6
1s2 2s2 2p6 3s1
1s2 2s2 2p6 3s2
1s2 2s2 2p6 3s2 3p1
1s2 2s2 2p6 3s2 3p2
1s2 2s2 2p6 3s2 3p3
1s2 2s2 2p6 3s2 3p4
1s2 2s2 2p6 3s2 3p5
1s2 2s2 2p6 3s2 3p6
1s2 2s2 2p6 3s2 3p6 4s1
1s2 2s2 2p6 3s2 3p6 4s2
1s2 2s2 2p6 3s2 3p6 4s2 3d1
1s2 2s2 2p6 3s2 3p6 4s2 3d2
1s2 2s2 2p6 3s2 3p6 4s2 3d3
1s2 2s2 2p6 3s2 3p6 4s1 3d5 (*)
1s2 2s2 2p6 3s2 3p6 4s2 3d5
1s2 2s2 2p6 3s2 3p6 4s2 3d6
1s2 2s2 2p6 3s2 3p6 4s2 3d7
1s2 2s2 2p6 3s2 3p6 4s2 3d8
1s2 2s2 2p6 3s2 3p6 4s1 3d10 (*)
1s2 2s2 2p6 3s2 3p6 4s2 3d10
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p1
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p4
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
Orbital Diagram
1s
[↑ ]
[↑↓] 2s
[↑↓] [↑ ]
[↑↓] [↑↓]
2p
[↑↓] [↑↓] [↑ ][ ][ ]
[↑↓] [↑↓] [↑ ][↑ ][ ]
[↑↓] [↑↓] [↑ ][↑ ][↑ ]
[↑↓] [↑↓] [↑↓][↑ ][↑ ]
[↑↓] [↑↓] [↑↓][↑↓][↑ ]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
[↑↓] [↑↓] [↑↓][↑↓][↑↓]
1s 2s
2p
3s
[↑ ]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
3s
3p
[↑ ][ ][ ]
[↑ ][↑ ][ ]
[↑ ][↑ ][↑ ]
[↑↓][↑ ][↑ ]
[↑↓][↑↓][↑ ]
[↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓]
3p
4s
[↑ ]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑ ]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑ ]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
[↑↓]
4s
{your textbook stops on this line}
3d
[↑ ][ ][ ][ ][ ]
[↑ ][↑ ][ ][ ][ ]
[↑ ][↑ ][↑ ][ ][ ]
[↑ ][↑ ][↑ ][↑ ][↑ ]
[↑ ][↑ ][↑ ][↑ ][↑ ]
[↑↓][↑ ][↑ ][↑ ][↑ ]
[↑↓][↑↓][↑ ][↑ ][↑ ]
[↑↓][↑↓][↑↓][↑ ][↑ ]
[↑↓][↑↓][↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓][↑↓][↑↓]
[↑↓][↑↓][↑↓][↑↓][↑↓]
3d
(*)
(*)
4p
[↑ ][ ][ ]
[↑ ][↑ ][ ]
[↑ ][↑ ][↑ ]
[↑↓][↑ ][↑ ]
[↑↓][↑↓][↑ ]
[↑↓][↑↓][↑↓]
4p
The * indicate elements where the electron configuration is slightly different than what you would guess. One
electron from the 4s sublevel is promoted up the 3d sublevel making it either half full, or entirely full.
Study Worksheet for Chapter 11 VSEPR Model - Chem 11
Extremely Important Note: After printing this page, you MUST draw in the missing dots on all 6
Lewis dot structures (its the horizontal pairs of dots can't be produced by MS Word).
Lewis dot structures only show which atoms are bonded to which other atoms, how many electrons are
involved in the bonds, and how many are lone pairs. It is a flat representation. It does not show the
molecule's true 3-D structure.
Steps to determine the compound's real 3-D structure:
1) First draw the Lewis dot structure.
2) Determine how many different groups of electrons surround each central atom.
3) Choose the appropriate geometric orientation that describes the placement of the atoms where the
electrons move the farthest apart from each other in 3-D space. This is different than how they would
move apart from each other in 2-D space when constrained to the plane of a sheet of paper. Note that the
geometric orientation of the atoms will be different than the geometric orientation of the electrons for
central atoms that have 1 or more lone pairs of electrons.
Formula
# of e- densities
around the central
Structure
atom
Geometric Shape
CO2
:O : : C : : O: 2
linear (2 e- densities around C try to get the farthest away)
BF3
:F:
3
trigonal planar (3 e- densities around B try to get the farthest away)
:F:
B
:F:
H2O
H:O:H
4
bent (4 e- densities around O, but 2 have no atoms attached to them)
So the e- densities form a tetrahedral, but the atoms form a bent angle.
CH4
H
H:C:H
H
4
tetrahedral (4 e- densities around C try to get the farthest away)
4
pyramidal (4 e- densities around N, but 1 has no atom attached to it)
So the e- densities form a tetrahedral, but the atoms form a pyramid.
NH3
H:N:H
H
:F: :F:
SF6
:F : S : F: 6
octahedral (e- densities around S) - this is not in your book
:F: :F:
Chapter 5 - Naming Inorganic Compounds
In order to correctly name an inorganic chemical compound from its formula, you must first determine the
type of compound it is, and then select the appropriate naming rule to use. Lower case is always used.
IONIC COMPOUND (a representative metal and one or more non-metals)
a) Binary compound where the metal has only 1 ionic charge state1
Name Rule: full metal name root form of the non-metal + -ide
Examples: NaCl = sodium chloride, MgH2 = magnesium hydride, AlN = aluminum nitride
b) Binary compound where the metal has multiple ionic charge states2
Name Rule: full metal name + (roman numeral) root form of the non-metal + -ide
Examples: CuF2 = copper (II) fluoride, Fe2O3 = iron (III) oxide, Hg2S = mercury (I) sulfide
c) Polyatomic compound where one ion consists of a group of atoms (usually the anion)
Name Rule: use (a) or (b) rule, and the ion table in your textbook for the anion name
Examples: KMnO4 = potassium permanganate, NH4NO3 = ammonium nitrate
MOLECULAR COMPOUND (two non-metals)
a) Binary compound consisting of only non-metals
Name Rule: prefix + element prefix + root form of last element + -ide
a) If there is only 1 atom of the 1st element, the mono prefix is dropped. b) Also, if the 2nd
element begins with a vowel and its prefix ends with a vowel, then the 1st vowel is dropped.
Examples: NO2 = nitrogen dioxide, CCl4 = carbon tetrachloride, N2O = dinitrogen monoxide
ACID (contains hydrogen at the beginning of the formula and reacts by losing that hydrogen atom)
a) A binary acid (hydrogen and only one other atom)3
Name Rule: hydro + root of anion + -ic acid
Examples: HCl = hydrochloric acid, H2S = hydrosulfuric acid, HF = hydrofluoric acid
b) A polyatomic acid (hydrogen and multiple other atoms) Two slight variations:
Name Rule: root of anion + -ic acid
(when the anion name ends in -ate)
Examples: H2SO4 = sulfuric acid, HClO3 = chloric acid, HNO3 = nitric acid
Name Rule: root of anion + -ous acid
(when the anion name ends in -ite)
Examples: H2SO3 = sulfurous acid, HClO2 = chlorous acid, HNO2 = nitrous acid
When determining a formula from the name (the reverse process), apply the above rules backwards. Write
out the symbols (if its an acid, write an H at the beginning). Always balance the formula by determining
the charge on each ion. (Note: the prefixes in molecular compounds tell you how to balance them).
Ex.: iron (II) phosphate = Fe3(PO4)2, oxygen difluoride = OF2, hydrobromic acid = HBr.
Footnotes:
1
Nearly all of the representative metals have just 1 charge state (a + charge equal to its group number).
Exceptions are: Sn2+ & Sn4+, Pb2+ & Pb4+, Bi3+ & Bi5+, Tl+ & Tl3+
2
Nearly all of the transition metals have multiple charge states (a roman numeral identifies the charge state).
Exceptions are: Ag+, Zn2+, Cd2+ (and several other less common ones you don't need to know).
3
Exception: cyanide (CN-) is considered a binary anion since it ends in -ide, so HCN = hydrocyanic acid
This nomenclature worksheet has two of every kind of naming rule and exception category on it.
Pretend this is a test and then compare your answers with those on the page after it.
Keep the next 2 pages synchronized (1st is the questions; 2nd is the answers).
Naming compounds from formulas and writing formulas from names - Chapter 5
Write the names for the following chemicals:
a) CaS
_________________________________
First determine whether each is ionic, (if so,
b) H3PO3
_________________________________
whether single or multiple charge states exist),
c) NH4Br
_________________________________
whether its a molecular compound, or
d) SF6
_________________________________
whether its an acid (if so, whether its a
e) FeCl3
_________________________________
binary or a polyatomic acid).
f) H2S
_________________________________
g) Ag2O
_________________________________
h) Cr(NO3)3
_________________________________
i) N2H4
_________________________________
j) HBr
_________________________________
k) Al(ClO3)3 _________________________________
l) PCl5
_________________________________
m) PbI2
_________________________________
n) Sn(SiO3)2 _________________________________
o) HClO4
_________________________________
Write the proper formulas for each of the following chemicals:
a) magnesium sulfide
b) diphosphorus pentoxide
c) tin (IV) phosphide
d) chloric acid
e) sodium dichromate
f) bismuth (III) nitrate
g) hydrocyanic acid
h) potassium oxide
i) silver sulfide
j) dinitrogen trioxide
k) lead (II) acetate
l) aluminum sulfite
m) zinc iodide
n) carbon tetrabromide
o) phosphorous acid
________________________ Don't forget to balance all formulas.
________________________
________________________
________________________
________________________ (careful)
________________________
________________________ (careful)
________________________
________________________
________________________
________________________
________________________
________________________
________________________
________________________
Write the names for the following chemicals:
a) CaS
calcium sulfide
ionic, binary, 1 charge state
b) H3PO3
phosphorous acid
acid, polyatomic (comes from phosphite anion)
c) NH4Br
ammonium bromide
ionic, polyatomic cation
d) SF6
sulfur hexafluoride
molecular
e) FeCl3
iron (III) choloride
ionic, binary, multiple charge states
f) H2S
hydrosulfuric acid
acid, binary (often called H-2-S, a slang name)
g) Ag2O
silver oxide
ionic, binary, 1 charge (a transition metal, but an exception)
h) Cr(NO3)3
chromium (III) nitrate
ionic, polyatomic, multiple charge states
i) N2H4
dinitrogen tetrahydride
molecular
j) HBr
hydrobromic acid
acid, binary (often called H-B-R, a slang name)
k) Al(ClO3)3 aluminum chlorate
ionic, polyatomic, 1 charge state
l) PCl5
phosphorous pentachloride
molecular
m) PbI2
lead (II) iodide
ionic, binary, multiple charge states (a rep metal, but exception)
n) Sn(SiO3)2 tin (IV) silicate
ionic, polyatomic, multiple charge states (rep metal, exception)
o) HClO4
acid, polyatomic (comes from perchlorate anion)
perchloric acid
Write the proper formulas for each of the following chemicals:
a) magnesium sulfide
MgS
binary, ionic, 1 charge state (balance: +2 and -2)
b) diphosphorus pentoxide P2O5
molecular (the name tells you how to balance it)
c) tin (IV) phosphide
Sn3P4
ionic, binary, multiple charges (balance: +4 and -3)
d) chloric acid
HClO3
acid, polyatomic, chlorate anion (balance: +1 and -1)
e) sodium dichromate
Na2Cr2O7
ionic, polyatomic, 1 charge state (+1 and -2), di is not a prefix
f) bismuth (III) nitrate
Bi(NO3)3
ionic, polyatomic, multiple charges (balance: +3 and -1)
g) hydrocyanic acid
HCN
acid, binary (exception), cyanide anion (balance: +1 and -1)
h) potassium oxide
K2O
ionic, binary, 1 charge state (balance: +1 and -2)
i) silver sulfide
Ag2S
ionic, binary, 1 charge state (balance: +1 and -2) exception
j) dinitrogen trioxide
N2O3
molecular (the name tells you how to balance it)
k) lead (II) acetate
Pb(C2H3O2)2 ionic, polyatomic, multiple charges (balance: +2 and -1) except
l) aluminum sulfite
Al2(SO3)3
ionic, polyatomic, 1 charge state (balance: +3 and -2)
m) zinc iodide
ZnI2
ionic, binary, 1 charge state (balance: +2 and -1) exception
n) carbon tetrabromide
CBr4
molecular (the name tells you how to balance it)
o) phosphorous acid
H3PO3
acid, polyatomic, phosphite anion (balance: +1 and -3)
Rules and Hints for Balancing Equations - Chapter 7
Always balance the number of atoms within a molecule with subscripts.
Always balance the number of molecules within an equation with coefficients.
When the number of atoms aren't an exact multiple of each other,
use the Least Common Multiple (LCM) method.
Until you feel comfortable doing it in your head, make a chart of all atoms. Ex:
C
H
O
reactants
1
4
2x2
products
1
2x2
2+2
CH4 + 2 O2  CO2 + 2 H2O
You can balance the elements in any order, but certain methods usually (but not
always) make it easier to do:
balance the metals first,
balance the non-metals next,
balance all the free elements last.
If its not obvious what the coefficient should be, use the algebraic equation
method to balance the very last free element.
Double check all equations after you are finished balancing them, because
frequently balancing one element unintentionally wrecks the balancing of another
previously-balanced element.
If you have to predict the products of a single-replacement reaction,
use the activity table on P. 155.
If you have to predict the products of a double-replacement reaction, use the
solubility chart on pages 470 and 471 and look for formation of gas, water or
unstable product. S = soluble (no precipitate); I = insoluble (a precipitate)
Mole Concept - Chapter 6
abbreviation = mol
a mole is a quantity of a chemical
1 mole = the molar mass of a substance expressed in the units of grams.
1 mole = the amount of an element that contains exactly 6.022 x 1023 atoms.
1 mole = the amount of a compound that contains exactly 6.022 x 1023 molecules.
Avagadro's number = NA = 6.022 x 1023 items / mol.
molar mass, molecular weight, atomic weight, formula weight, mass number
These terms all mean essentially the same thing, with slight variations, depending upon whether you are
referring to an atom, a molecule or an ionic substance. Units are amu (atomic mass units) or grams/mole.
1 mol of He =
4.003 g
1 mol of C =
12.011 g
1 mol of Au =
196.97 g
1 mol of H2 =
2.0158 g
1 mole of CH4 = 16.043 g
1 mole of NH3 = 17.031 g
=
=
=
=
=
=
6.022 x 1023 atoms
6.022 x 1023 atoms
6.022 x 1023 atoms
6.022 x 1023 molecules = 12.044 x 1023 atoms
6.022 x 1023 molecules = 30.110 x 1023 atoms
6.022 x 1023 molecules = 24.088 x 1023 atoms
1 mol of any element contains the same number of atoms (but it does not weigh the same).
1 mol of any compound contains the same number of molecules (but it does not weigh the same).
1 mol of H2O (18.015 g) contains 2 mols of H (2.0158 g) and 1 mol of O (15.999 g)
1 mol of NH3 (17.031 g) contains 1 mol of N (14.007 g) and 3 mols of H (3.024 g)
+
N2
picture:
OO
oo oo oo
oOo
o
# of atoms
# of molecules
# of molecules (x 2)
# of molecules (x 10)
# of molecules (x 1,000)
# of molecules (x NA)
# mols
molar mass
2
1
2
10
1,000
6
3
6
30
3,000
8
2
4
20
2,000
6.022 x 1023
1
28.014 g
28.014 g
34.062 g
3 H2

chemical equation:
=
≠
≠
≠
≠
23
18.066 x 10 ≠
3
≠
3 x 2.016 g =
6.048 g
=
=
2 NH3
12.044 x 1023
2
2 x 17.031 g
34.062 g
34.062 g
Stoichiometry - Chapter 8
÷ Avagadro’s number
oOo
o
x molar mass
# of atoms
(or molecules)
# of moles
x Avagadro’s number
# of grams
÷ molar mass
How much does 2.4 mols of Cu weigh?
2.4 mol 63.546 g
!
= 152.51 g = 150 g
mol
How many atoms are there in 5 mols of Fe?
5 mol 6.022 x 10 23 atoms
!
= 3.011 x 10 24 atoms = 3 x 10 24 atoms
mol
How many atoms are there in 3.50 g of Au?
3.50 g
!
mole
6.022 x 10 23 atoms
!
= 1.07 x 10 22 atoms
196.97 g
mol
How many molecules are there in 3.50 g of H2O?
3.50 g
!
mol
6.022 x 10 23 molecules
!
= 1.17 x 10 23 molecules
18.0148 g
mol
How much will 4.62 x 1025 molecules of H2O weigh?
4.62 x 10 25 molecules
!
mol
18.0148 g
!
= 1380 g
23
mol
6.022 x 10 molecules
How many O atoms are there in 2.00 g of KClO3?
2.00 g KClO 3
!
mol KClO 3 6.022 x 10 23 molecules
3 atoms O
!
!
= 2.95 x 10 22 atoms O
122.55 g
mol KClO 3
molecule KClO 3