Chapter 2: Names, Formulas and
Properties
p. 40-81
Chemical Names and Formulas p 42
The IUPAC (International Union of Pure and Applied Chemistry)
was founded in 1919 and has developed a method to name
chemicals based on their composition.
A. Naming Binary Ionic Compounds
First Name:
Use the full name of the cation
(usually a metal) first
Second Name:
Put the name of the anion (usually
a non-metal) last and change the
ending to –ide (eg. oxygen
becomes oxide)
Name the Following Ionic Compounds:
1.
2.
3.
4.
5.
6.
NaCl (s)
LiBr (s)
CaO (s)
SrCl2 (s)
BaF2 (s)
K2S (s)
Name the Following Ionic Compounds:
1.
2.
3.
4.
5.
6.
NaCl (s)
LiBr (s)
CaO (s)
SrCl2 (s)
BaF2 (s)
K2S (s)
sodium chloride
lithium bromide
calcium oxide
strontium chloride
barium fluoride
potassium sulfide
Using Ion Charges and Chemical Names to Write Formulas
1. Write the metal element symbol with its charge, and the
non-metal element symbol with its charge
2. Balance the ion charges
3. Write the formula indicating how many atoms of each
element are in it with a subscript.
* If there is only 1 atom of an element, no subscript is
needed
• For example: barium chloride
• Ba2+
Cl-
• Study Model Problem 1 and figure 2.3 on p.
45 and Do Practice Problems # 5-8
The Stock System:
As you look at your periodic table, you
may notice that many transitional
elements have more than one
possible ionic charge.
Ex. Cu+ or Cu2+
A German chemist, Alfred Stock,
developed a way to indicate which
cation is in a compound. The Stock
System uses a roman numeral to
indicate the charge on the metal or
cation.
Ex. CuCl = copper (I) chloride
CuCl2 = copper (II) chloride
• If the cation can have more than one possible
charge, you must specify which ion is being
used.
Example:
Name Fe2O3
Fe3+ and Fe2+ are both possible ions of iron
• If the cation can have more than one possible
charge, you must specify which ion is being
used.
Example:
Name Fe2O3
Fe3+ and Fe2+ are both possible ions of iron
Fe3+ + O2- = Fe2O3 Iron (III) ioxide
• If the cation can have more than one possible
charge, you must specify which ion is being
used.
Example:
Name Fe2O3
Fe3+ and Fe2+ are both possible ions of iron
Fe3+ + O2- = Fe2O3 Iron (III) ioxide
Fe2+ + O2- = FeO Iron (II) ioxide
• If the cation can have more than one possible
charge, you must specify which ion is being
used.
Example:
Name Fe2O3
Fe3+ and Fe2+ are both possible ions of iron
Fe3+ + O2- = Fe2O3 Iron (III) ioxide
Fe2+ + O2- = FeO Iron (II) ioxide
Name the following compounds:
HgF2
NiBr3
PbS2
Name the following compounds:
HgF2
Hg2+ Hg1+
NiBr3
PbS2
F-
Name the following compounds:
HgF2
mercury (II) fluoride
Hg2+ Hg1+
FNiBr3
PbS2
Name the following compounds:
HgF2
mercury (II) fluoride
Hg2+ Hg1+
FNiBr3 nickel (III) bromide
Ni2+ Ni3+
BrPbS2
Name the following compounds:
HgF2
mercury (II) fluoride
Hg2+ Hg1+
FNiBr3 nickel (III) bromide
Ni2+ Ni3+
BrPbS2 lead (IV) sulfide
Pb2+ Pb4+
S2-
• Study Model Problem 2 on p. 46 and do
Practice Problems #9,10
• Study Model Problem 3 and do Practice
Problems p.47 #11,12
Polyatomic Ions
• These ions consist of two or more different
atoms joined together by covalent bonds. As a
group, the bonded atoms have either a
positive or negative charge. See the
polyatomic ion table on the top of the periodic
table, and become familiar with the common
polyatomic ions in table 2.3 p51.
• Ex. Carbonate ion = CO32-
Naming Rules:
1st name: name the cation as usual
2nd name: name the anion (use the name given
from your periodic table)
**parentheses must be placed
around the polyatomic ion if more
than one are found in a chemical
formula
Ex. Na2SO4 = sodium sulfate
NH4Cl = ammonium chloride
(NH4)3PO4 = ammonium phosphate
Now name these Ionic Compounds:
(these examples contain polyatomic ions)
1.
2.
3.
4.
Li2CO3
KClO2
CaSO4
Ba(NO2)2
Now name these Ionic Compounds:
(these examples contain polyatomic ions)
1.
2.
3.
4.
Li2CO3
KClO2
CaSO4
Ba(NO2)2
lithium carbonate
potassium chlorite
calcium sulfate
barium nitrite
Study Model Problems 4 and 5 p.51 and do
Practice Problems on p. 52 #13-19
Getting to know the polyatomic ions:
• Many polyatomic anions contain oxygen and
are part of a “family”.
Anion name
# of oxygen
Example
Formula
Per----ate
1 more
pernitrate
NO4-
nitrate
NO3-
----ate
---------------
----ite
1 less
nitrite
NO2-
Hypo----ite
2 less
hyponitrite
NO-
Name the polyatomic ion
note: these are not compounds, they are
ions! They have a charge.
•
•
•
•
ClO3ClO2ClOClO4-
Name the polyatomic ion
note: these are not compounds, they are
ions! They have a charge.
•
•
•
•
ClO3ClO2ClOClO4-
chlorate
chlorite
hypochlorite
perchlorate
• Now try these:
–
–
–
–
PO43PO33PO23PO53-
• Now try these:
–
–
–
–
PO43PO33PO23PO53-
phosphate
phosphite
hypophosphite
perphosphate
• All right, show me what you know!
Name Li3BO2
• All right, show me what you know.
Name Li3BO2
Lithium borite
•
Write the formulas for the following Ionic Compounds:
1.
2.
3.
4.
5.
6.
7.
Potassium chloride
Calcium chloride
Iridium oxide
Zirconium nitride
Cobalt (II) chloride
Sodium hydroxide
Nickel (II) bromate
•:
Write the formulas for the following Ionic Compounds:
1.
2.
3.
4.
4.
5.
5.
6.
6.
7.
7.
Potassium chloride
Calcium chloride
Iridium oxide
Zirconium
Zirconium nitride
nitride
Cobalt
chloride
Cobalt (II)
(II) chloride
Sodium hydroxide
Sodium hydroxide
Nickel (II) bromate
Nickel (II) bromate
K+ ClCa2+ ClIr4+ O2Zr4+ N3Co2+ ClNa+ OHNi2+ BrO3-
KCl (s)
CaCl2 (s)
IrO2 (s)
Zr3N4 (s)
CoCl2 (s)
NaOH (aq)
Ni(BrO3)2 (s)
• Study Model Problem 3 and do Practice
Problems p.47 #11,12
• Do BLM 2-1 and BLM 2-5
• Study Model Problems 4 and 5 p.51 and do
Practice Problems on p. 52 #13-19
Hydrates
• some ionic compounds exist as hydrates.
• hydrated compounds have water molecules
attached and are named using one of two
methods:
• the new IUPAC method uses a dash after the
compound name and then the word water. The
ratio of compound to water is given in brackets.
ex. NiSO4•6H2O
nickel (II) sulfate – water (1/6)
• the older nomenclature uses a prefix to indicate
the number, and the word hydrate
ex. NiSO4•6H2O
nickel (II) sulfate hexahydrate
Naming Molecular
Compounds
• A binary molecular
compound forms
when atoms of nonmetals form a
covalent bond and
become a molecule.
• Many molecules are
known by their
common names,
such as water, H2O,
and ammonia, NH3
.
1. Use the full name of the first element (the most
metal-like goes first)
2. Put the name of the second element last and
change the ending to –ide
3. Use the correct prefix to indicate the number of
each element
– Exception: do not use the prefix mono when the
first element only has 1 atom
Number of Atoms
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
Prefix
mono
di
tri
tetra
penta
hexa
hepta
octa
nona
deca
So: To name molecules . . .
Prefix + First Element, Prefix + Second Element (with –ide ending)
Name the following molecular compounds:
1. CO2
2. NO3
3. N2O
4. NF3
5. N2O3
6. CO
Name the following molecular compounds:
1. CO2
carbon dioxide
2. NO3
nitrogen trioxide
3. N2O
dinitrogen monoxide
4. NF3
nitrogen trifluoride
5. N2O3
dinitrogen trioxide
6. CO
carbon monoxide
Remember: The number in subscript tells us
how many of each element are in the
compound.
• Do practice problems p. 44 # 1-4
• Some molecules have common names you will
have to memorize. These include:
water
H2O (l) or HOH (l)
ammonia
NH3 (g)
ozone
O3 (g)
methanol
CH3OH (l)
ethanol
C2H5OH (l)
propane
C3H8 (g)
methane
CH4 (g)
glucose
C6H12O6 (s)
sucrose
C12H22O11 (s)
hydrogen peroxide
H2O2 (l)
hydrogen sulfide
H2S (g)
Molecular Elements
• A number of elements are found in their
elemental form as molecular elements. These
elements should be memorized.
• The diatomic elements are H2, N2, O2, F2, Cl2,
Br2, and I2.
• Sulfur and selenium exist as S8 and Se8,
phosphorous exists as P4.
1A
2A
3A
4A
5A
6A
7A
8A
(1)
(2)
(13)
(14)
(15)
(16)
(17)
(18)
H
He
3B
4B
5B
6B
7B
—
8B
—
1B
2B
(3)
(4)
(5)
(6)
(7)
(8)
(9)
(10)
(11)
(12)
1
H
2
Li
Be
B
C
N
O
F
Ne
3
Na
Mg
Al
Si
P
S
Cl
Ar
4
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
5
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
6
Cs
Ba
La
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
At
Rn
7
Fr
Ra
Ac
Rf
Db
Sg
Bh
Hs
Mt
Ds
Rg
Uub
—
Uuq
—
—
—
—
6
Ce
Pr
Nd
Pm
Sm
Eu
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
7
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
Here is a great anagram to remember the
diatomic elements.
H - Horses
N - Need
O – Oats
F - For
Cl - Clear
Br – Brown
I- eyes (I's)
• Do BLM 2-3
• and Check Your Understanding p. 55 #1-4
2.2 Explaining Properties of Substances p. 56
• Physical properties – observable and measurable
qualities: melting and boiling points, conductivity,
appearance, state
• Chemical properties – reactive properties of a
substance
Properties of Ionic and Molecular Compounds
Property
Ionic compound
Molecular compound
State at room temp.
Solid (hard, brittle,
crystal lattice shape)
Solid, liquid or gas
Melting point
High
Low
Attraction between
molecules/ions
Strong crystal lattice
Weaker
Conductivity when solid
No,
molecules held rigidly in
crystal lattice shape
No
Conductivity when
dissolved in water or
melted?
Yes,
electrolyte
tend to be
non-electrolytes
• Do Investigation 2-B “Ionic or Molecular”
p. 58 (use BLM 2-2) and Q #1-6 p59
• Do Check Your Understanding #1,3,5 p62
2.3 Properties of Acids and Bases P.63
• Acidic and basic solutions are electrolytes
• The stronger the acid or base, the stronger the
electrolyte.
• Acids and bases react neutralizing each other.
The resulting solution is a salt dissolved into
water.
• Strong acids react with metals to produce
hydrogen gas.
Acids
• Taste: Sour
• Turns litmus red
• Phenolphthalein: colorless
• pH < 7
• Arrhenius definition: Release H+ ions when
dissolved in water.
• HCl(s)
 H+(aq) Cl-(aq)
Bases:
• Taste: Bitter
• Feel: Slippery
• Turns litmus: blue
• Phenolphthalein: pink
• pH > 7
• Arrhenius definition: Release OH- ions when
dissolved in water.
• NaOH–(s)
 Na+(aq) OH-(aq)
Naming Acids:
• Name the compound as an ionic compound. Then use the
rules on the periodic table to complete the acidic name:
If the ionic name was:
The acidic name is:
1. hydrogen _______ide
hydro_____ic acid
eg. H2S = hydrogen sulfide = hydrosulfuric acid
2. hydrogen _______ate
______ic acid
eg. HClO3 = hydrogen chlorate = chloric acid
3. hydrogen _______ ite
______ ous acid
eg. H2SO3 = hydrogen sulfite = sulfurous acid
• Acids and bases are aqueous solutions and
must use (aq) after the formula.
• H2S (aq) , hydrosulfuric acid, can also be called
aqueous hydrogen sulfide.
H2SO4 (aq) – hydrogen sulfate  sulfuric acid
H2S (aq) – hydrogen sulfide  hydrosulfuric acid
H2SO3 (aq) – hydrogen sulfite  sulfurous acid
• When working back to the formula, remember
to balance charges.
• Ex. Phosphoric acid  hydrogen phosphate
H+ PO43-  H3PO4 (aq)
• If the polyatomic ion ends in a COO-, the
hydrogen ion bonds at this end to become --COOH (this is called a carboxylic acid).
• For example acetic acid CH3COOH(aq)
• Do Practice Problems p.70 #20-23
• Do Find Out Activity “Acid or Base” p. 67 and
BLM 2-6 (optional)
• Do Check Your Understanding p. 71 #3,4,5,6,8
2.4 Why Water is Weird p. 72
• cells of most living organisms are 80-95% water
• virtually all chemical reactions for life occur in
aqueous solutions.
• water has many unique properties:
– It occurs in all 3 states on Earth
– Ice is less dense than liquid water
– Its melting and boiling points are higher than similar
substances (see table 2.14 p72)
– It has a very high surface tension
• Do Investigation 2-D “Physical Properties of
Water” p. 73 (use BLM 2-9 to record results)
and do Q #1-3,6 p74
• Water molecules have a bent shape
• Water is a polar or dipole molecule.
• As a result water molecules attract each other. The
negative oxygen attracts the positive hydrogen and forms a
hydrogen bond.
• This is an intermolecular bond, one between molecules
(versus intramolecular bonds, within molecules, like
covalent bonds which are stronger).
Hydrogen bonds explain water’s unique properties:
• high melting and boiling points:
-more energy must be added to break the hydrogen bonds
• High heat capacity:
- strong attraction between water molecules must be overcome
to increase the average speed of the molecules
• concave meniscus and capillary action
- polar molecules are attracted to the sides of the container
• solid water floats in liquid water
-hydrogen bonds force water molecules farther apart as
it freezes, making ice less dense than liquid water.
• high surface tension / cohesion
-hydrogen bonds pull water molecules close together
• Do Check Your Understanding p. 78 #1,2,4,6
• Do Chapter 2 Review p. 80 #1,2,4,6,8-10,
13,14,17-19,22,25