Honors Chemistry
Unit 5
 Chemical bonding
 Chemical formulas
 Chemical naming
1
We are learning to:
1.
2.
3.
4.
5.
Distinguish between an element and compound.
Distinguish between ionic, covalent, and metallic bonds.
Write the name/formula for covalent compounds.
Write the name/formula for ionic compounds including the use of polyatomic ions.
Write the name/formula for acids.
We are looking for:
1. Compounds are a chemical combination (bonded together) of 2 or more elements.
2a. Electronegativity differences between 2 elements determines the ionic character of a bond; 00.3 is nonpolar covalent, >0.3-1.7 is polar covalent, >1.7 is ionic.
2b. Metallic bonds occur between 2 or more metals.
2c. Describe the characteristics of substances with ionic, covalent, or metallic bonds (behavior of
the electrons, solubility, conductivity, state at room temperature, melting point).
3a. Use prefixes when naming or writing formulas for covalent compounds and change second element
ending to –ide.
3b. Do not reduce subscripts on covalent compound formulas.
4a. Do not use prefixes when naming ionic compound formulas/names; name the metal and then
change the nonmetal ending to –ide.
4b. Use oxidation numbers to write the correct balanced formula for an ionic compound (criss-cross
method).
4c. Use roman numerals for metals with multiple oxidation states when naming ionic compounds.
4d. Reduce subscripts for ionic compound formulas.
4e. Recall of polyatomic ions ending in –ate.
4f. Determine all other forms of polyatomic ions based upon the –ate ion and how the amount of
oxygen has changed.
5a. Use hydro- prefix when naming binary acid (an acid with H and a nonmetal) and change ending of
nonmetal to -ic and add acid on the end.
5b. Use oxidation numbers to write the correct balanced formula for binary and ternary acids (crisscross method).
5c. Naming ternary acids (an acid with H and a polyatomic ion) convert polyatomic ion to the correct
ending: -ate becomes –ic; -ite becomes –ous) and add acid on the end.
2
Chemical Bonding
Electrical attraction between nuclei of one atom and valance (outer shell) electrons of a different atom.
I. Ionic Bonding
_________________ of electrons from the ___________(cation) to the ____________ (anion)
Electrical attraction between large numbers of cations and anions. “clusters”
Cation = positively particle
Typically are metals
Loves to give electrons away
Low Ionization Energy
Low Electron Affinity
Low Electronegativity
Most metallic element = Francium
Anion = negatively charged particle
Typically are non metals
Loves to accept or take on electrons
Higher Ionization Energy
Higher Electron Affinity
Higher Electronegativity
Most nonmetallic element = Fluorine
Ionic compounds are ________________
Ionic bonds _________________________in water
- Fall apart into cations and anions
-Will conduct electricity
EN (electronegativity) difference is 1.8 – 4
Maximum value is 4 (Fluorine is a 4)
3
II.
Covalent Bonds
Formed when 2 _______________ atoms ______________electrons
A. ____________ Covalent
__________________ sharing of electrons
H
Cl
Electrons attracted MORE to higher EN atom, so density is greater around atom
with higher EN value.
EN of H = 2.1
EN of Cl = 3.0
EN difference = .9
EN difference range =
>_______– ________ for polar covalent bonds
Polar Covalent compounds
 dissolve in H2O
 don’t fall apart into ions….
 No conductivity!
B. _________________ Covalent Bonds
______________ sharing of electrons
Br
Br
Electrons are shared equally between two atoms.
EN difference for nonpolar covalent bonds is _____-_______.
Usually nonpolar covalent bonds are between diatomic molecules – two of the
same atoms.
Nonpolar covalent compounds
 do not dissolve in water
 no conductivity.
Can be a gas, liquid, or solid.
4
III. Metallic Bonds
Bond formed between 2 _____________atoms.
Larger outer electron shells – which overlap
Electrons are free to move within electron clouds of all metal ions =
___________________________
Electrons are ______________________ - they don’t belong to any one atom anymore.
5
Periodic Table with Electronegativities
1A
2A
3B
4B
5B
6B
7B
8B
1B
2B
3A
4A
5A
6A
1
H
2.1
7A
8A
2
He
3
4
Li Be
1.0 1.5
5
6
7
8
9 10
B
C
N O
F Ne
2.0 2.5 3.0 3.5 4.0
11 12
Na Mg
0.9 1.2
13 14 15 16 17 18
Al Si
P
S
Cl Ar
1.5 1.8 2.1 2.5 3.0
19 20 21
K Ca Sc
0.8 1.0 1.3
22
Ti
1.5
23
V
1.6
24
Cr
1.6
25
Mn
1.5
26
Fe
1.8
27
Co
1.9
28
Ni
1.9
29
Cu
1.9
30
Zn
1.6
31 32 33 34 35 36
Ga Ge As Se Br Kr
1.6 1.8 2.0 2.4 2.8 3.0
37 38 39
Rb Sr
Y
0.8 1.0 1.2
40
Zr
1.4
41
Nb
1.6
42
Mo
1.8
43
Tc
1.9
44
Ru
2.2
45
Rh
2.2
46
Pd
2.2
47
Ag
1.9
48
Cd
1.7
49 50 51 52 53 54
In Sn Sb Te
I Xe
1.7 1.8 1.9 2.1 2.5 2.6
55 56 57
Cs Ba La
0.7 0.9 1.1
72
Hf
1.3
73
Ta
1.5
74
W
1.7
75
Re
1.9
76
Os
2.2
77
Ir
2.2
78
Pt
2.2
79
Au
2.4
80
Hg
1.9
81 82 83 84 85 86
Tl Pb Bi Po At Rn
1.8 1.9 1.9 2.0 2.2 2.4
87 88 89 104 105 106 107 108 109 110 111 112
Fr Ra Ac Rf Ha Sg Ns Hs Mt Uun Uuu Uub
0.7 0.9 1.1
58 59 60 61 62 63 64 65 66 67
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho
1.1 1.1 1.1 1.2 1.2 1.1 1.2 1.2 1.2 1.2
68
Er
1.2
69
Tm
1.2
70
Yb
1.2
71
Lu
1.3
90 91 92 93 94 95 96 97 98 99 100 101 102 103
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
1.3 1.5 1.7 1.3 1.3 1.3 1.3 1.3 1.3 1.3 1.3 1.3 1.5
Keep in mind that electronegativities are approximate measures of the relative tendencies of these elements to
attract electrons to themselves in a chemical bond. The greater an atom's electronegativity, the greater its
ability to attract electrons to itself.
6
Calculating Bond Type
(using Electronegativity Values)
Electronegativity
Difference
Bond Type
Ionic Character
0.0 – 0.3
Non-polar covalent
0 – 5%
0.3 – 1.7
Polar covalent
5 – 50%
Over 1.7
Ionic
Higher than 50%
1. Find electronegativity values for each element.
2. Subtract the lesser value from the larger value (so the value is positive).
3. Find the difference value on the above chart; convert it to a bond type.
Determine the type of bond that will form between each pair of elements:
1) Ca ,Br ___________________________
2) O, H ____________________________
3) C, H ___________________________
4) Na, Cl __________________________
5) S, O ____________________________
7
Using Rule of Thumb To Predict Bond Type
Metal + Non metal  Ionic Bond
Metal + Metal  Metallic
Non metal + Non metal  Covalent
Noble Gas + any element  typically no bond
Elements
Predict using
Rule of Thumb
Au and I
Au =
I=
Ca =
F=
P=
Br =
Ag =
Cl =
Ti =
O=
Ca and F
P and Br
Ag and Cl
Ti and O
C6H12O6
C–H
C–O
Predict based on
Electronegativity values
(see page 6)
Au =
I=
Ca =
F=
P=
Br =
Ag =
Cl =
Ti =
O=
O–H
KOH
H2O
MgSO4
CaCO3
8
Polar versus Non Polar
Ionic versus Covalent
Substance
Formula
Physical
Appearance
Polar or
Non Polar
(Dissolve?)
Conducts
Electricity?
Ionic or
Covalent
1.
Baking
Soda
2.
Sugar
3.
Salt
4.
Splenda
5.
Milk of
Magnesia
6.
Flour
9
Substance
Formula
Physical
Appearance
Polar or
Non Polar
(Dissolve?)
Conducts
Electricity?
Ionic or
Covalent
7.
Baking
Powder
8.
Epsom
Salts
9.
Cream of
Tartar
10.
Alum
11.
Rolaids
12.
Chicken
Bouillon
10
Name______________________
Chemical Bonds Lab
Part 1. Conductivity
Introduction:
A compound is defined as a chemical combination of two or more elements. A chemical bond is the
"glue" holding together atoms of different elements. Two types of bonds are Ionic and covalent.
Ionic bonds generally occur between a metallic and nonmetallic atom. The bond results from the
transfer of one or more electrons from the metallic atom to the nonmetallic atom resulting in a
charge difference. The positively charged metal ion is then attracted to the negatively charged
nonmetallic ion. Covalent bonding involves the sharing of electrons.
In the following activity, you will be exploring one additional property displayed by the molecules
that contain these bonds. In this investigation, you will be observing three groups of substances
known as strong electrolytes, weak electrolytes, and nonelectrolytes



When placed in an aqueous solution (water) ionically bonded substances dissociate into
their ions and demonstrate the ability to conduct electricity. These are considered
__________________electrolytes.
Some covalently bonded substances also have the ability to conduct electricity at low
levels. These are considered ____________________ electrolytes.
Finally, there are many substances that, when placed in water do not dissociate into ions.
As a result, these substances fall to conduct electricity. These are considered
____________________ electrolytes.
Materials:
Solutions
Equipment




24-well microplate
Conductivity meter
Deionized Water-dH2O
(to rinse electrodes)
Beaker






Ethyl Alcohol (CH3CH2OH)
Hydrochloric acid (HCl)
Sodium Hydroxide (NaOH)
Hexane (C6H14)
Copper (II) Chloride (CuCl2)
Bleach (NaClO)







Tap H2O
dH2O
Ammonia (NH3)
Sodium Chloride (NaCl)
Glucose (C6H12O6)
Acetic Acid (CH3COOH)
Calcium Chloride (CaCl2)
11
Experiment:
1) In data table 1, record the substances’ names and formulas that will be tested. Calculate
the highest electronegativity difference between any 2 elements in the formula and
predict the type of bond contained in the substance (Ionic, NP Covalent, P Covalent).
2) In the diagram of the well plate below, put the name of the compound in the well number
that you intend to put that compound in. This will act as a key since the solutions appear
identical.
1
2
3
4
5
6
A
B
C
D
Procedure
1) Obtain a small sample of each substance and add it to the appropriate well. Add just
enough to cover the bottom of the well.
2) Place both metal probes of the conductivity meter into the well. Press and hold the low
button. If the meter lights up less than 10, record the number in the data table. If the
meter lights up 10, then press and hold the high button. You must add 10 to the meter
reading when on high when you record it in your data table.
3) Rinse the metal probes of the conductivity meter with the wash bottle of dH2O while
holding it over a beaker. Wipe dry.
4) Repeat steps 2 & 3 for each compound.
12
Data Table 1:
Formula of
Substance
Tap H2O
Largest
Electronegativity
Difference
Predicted Type
of Bond
Conductivity
Meter Reading
dH2O
NH3
NaCl
C6H12O6
CH3COOH
CaCl2
CH3CH2OH
HCl
NaOH
CuCl2
NaClO
C6H14
Conclusions:
1) Identify those compounds that showed an unexpected result (e.g. a covalent compound that had
a high electronegativity value) and discuss possible reasons why that result was obtained.
13
Part 2. Properties of Chemical Bonds
Introduction:
The physical properties of a substance, such as melting point, solubility, and conductivity, can be
used to predict the type of bond that binds the atoms of the compound. In this experiment, you
will test various compounds to determine these properties. Your compiled data will enable you to
classify the substances as either ionic or covalent compounds.
Materials:
Equipment







Candle
Ring stand
Conductivity tester
Thin-stemmed pipets, 2
Iron ring
Tin can lid
Matches
Solutions








Benzoic acid (C6H5COOH)
Calcium chloride (CaCl2)
Potassium iodide (KI)
Sodium chloride (NaCl)
Sucrose (C12H22O11)
Citric acid (C6H12O7)
Deionized Water (dH2O)
Ethyl Alcohol (CH3CH2OH)
Experiment:
Part A. Melting Point
1. Place a can lid on an iron ring attached to a ring stand. Position the ring so that it is just
above the tip of a candle flame, as shown in Figure A (next page). Light the candle for a
moment to check that you have the correct height.
2. Place a few crystals of sucrose, sodium chloride, benzoic acid, calcium chloride, citric acid,
and potassium iodide in separate locations on the lid, as shown in Figure B (next page). Do
not allow the samples of crystals to touch. Draw the position of each compound in the
diagram below.
Ring
Stand
Diagram 1.
14
3. Write a brief description of each of
the six substances in Data Table 2.
4. For this experiment, it is not
necessary to have exact values for the
melting point. The lid will continue to
get hotter as it is heated, so the order
of melting will give relative melting
points. Light the candle and observe.
Note the substance that melts first by
writing a “1” in the data table. Record
the order of melting for the other
substances.
Benzoic acid
5. After 3 min., record an “n” in your data table for each substance that did not melt.
Extinguish the candle flame. Allow the tin can lid to cool while you complete the remainder
of the experiment.
Part B. Solubility & Conductivity
1. Put a few crystals of each of the white solids in the top row of your microplate. Repeat
with the second row. Add 10 drops of water to each well in the top row. Record the
solubility of each substance in data table 2.
2. Add 10 drops of ethanol to each well in the second row of the microplate. Record the
solubility of each substance in your data table.
3. Test the conductivity of each water solution in the top row by dipping both electrodes
into each well of the microplate. Be sure to rinse the electrodes and dry them with a paper
towel after each test. If the bulb of the conductivity apparatus lights up, the solution
conducts electricity. Record your results in Data Table 2.
4. Clean the microplate by rinsing it well with water in the sink. Make sure all the wells are
clean and then dry each well with a paper towel. Clean the substances off of the can lid in
the sink. Be very careful not to cut yourself. Wash your hands thoroughly before you
leave the lab and after all work is finished.
15
Data Table 2.
Compound
Description
Melting
Solubility
Solubility
Conductivity
Point
in water
in ethanol
meter reading
Calcium
Chloride
Citric Acid
Benzoic
Acid
Potassium
Iodide
Sodium
Chloride
Sucrose
Conclusions:
1) Group the white substances into two groups according to their properties.
Group1
Group2
2) List the properties of each group.
Group1
Group2
16
3) Use your prior knowledge and your experimental data to determine which of the groups
consists of ionic compounds and which consists of covalent compounds.
4) Write a short paragraph to summarize the properties of ionic compounds and another
short paragraph to summarize the properties of covalent compounds. Include all properties
tested in this experiment.
17
18
Types of Bonds
Ionic
Examples
Na and Cl
Covalent
C and Cl
Metallic
Na and Na
Types of elements
(metal? nonmetal?)
Electronegativity
Difference between
elements
How are the electrons
involved?
Compound formed from
example above
Name for a single unit of
this compound
Bond Strength
State at room
temperature
Polarity
Solubility in H2O
Conductivity (light bulb
test)
19
Ionic
Covalent
20
Intermolecular bond occurs between 2 different molecules
Hydrogen Bonding
The intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is
attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule.
One molecule is polar and has hydrogen in a H-F, H-O or H-N bond
Other molecule has unshared pairs of electrons, usually on F, O, or N
Hydrogen forms a weak bond with the unshared pair of the other molecule
H – O:
H
H–O
H
δδ+
21
Van der Waals bonds
London dispersion forces result from the intermolecular attractions resulting from the constant motion of electrons and
the creation of instantaneous dipoles.
They can be very weak intermolecular forces between noble gas atoms and nonpolar molecules
Electrons of Noble Gas get shifted to one side causing it to attract a slightly positive atom that is nearby.
Dispersion forces are present between any two molecules (even polar molecules)
22
Bonding Review Sheet
Name the type of bond that will occur between the elements/compounds below based
on the periodic table.
1. Zn + Zn
2. Ca + Cl
3. I + I
4. Co + Ne
5. Mg + O
6. H2O + H2O
Below list all of the characteristics of the following words:
Ionic Bond
Van der Waals Bonds
Metallic Bond
Polar Colvalent
Hydrogen Bond
Nonpolar Covalent
23
Write the electronegativity difference for the compounds below and decide what type of
bond is occurring
Compound
Electronegativity
difference
Type of bond
CaCl2
Br2
CH4
C6H12O6
Using the elements below, draw a nonpolar electron cloud and a polar electron cloud
and describe why they are different. Label the type of bond (polar or nonpolar covalent)
H
Cl
Br
Type of bond ____________
Element
Nitrogen
Nuclear
Symbol
Mass
Number
Br
_____________
# of
Protons
# of
Neutrons
# of
Electrons
15
48
Charge
3-
22
4+
50
38
0
24
Making Ionic Compounds!!
(with a metal and a non-metal)
Make compounds from the following cations and anions
Bromine , Tin , Calcium , Nitrate, Phosphorus, Gallium, Ammonium,
Carbonate
Cations (write their symbol and charge)=
(Put them along the side of the table; top to bottom)
Anions (write their symbol and charge)=
(Put them across the top of the table)
-----------
------------
-------------
--------------
25
Naming Ionic Compounds
Binary Ionic Compounds (2 elements; metal + nonmetal)
1. Name metal (cation) first, use name as it appears on the periodic table.
2. If it is a transition metal, it might need a roman numeral.
(The roman numeral equals the charge)
The transition metals that don’t need a roman numeral are:
3. Name nonmetal (anion) second. Use an -ide ending on the element name if it
is an element off of the periodic table.
Ex. NaCl = Sodium Chloride
Exceptions:
Oxygen becomes oxide
Sulfur becomes sulfide
Hydrogen becomes hydride
Phosphorus becomes phosphide
***When given the name and you need to write the formula, you must determine
the charge for each ion first and then criss-cross.***
26
Ionic Compounds with Polyatomic Ions:
Name the polyatomic ion; do not change it.
Ex. NaNO3 is Sodium Nitrate
Fe(NO3)3 is Iron (III) Nitrate
Making More Polyatomic Ions
Change in Amount of Oxygen
1 oxygen higher than “ate” level
ending
per-
-ate
“-ate” level
-ate
1 oxygen lower than “ate” level
-ite
2 oxygens lower than “ate” level
hypo-
-ite
Common ions have "ate" ending
o
BrO3-, Bromate
If you lose 1 oxygen atom from "ate" ion, the form is "ite"
o Ex: Bromate is BrO3o Br02- = Bromite
o Charge of ion does not change as form changes!
If you lose 2 oxygens atoms from "ate" ion, the form is "hypo
ite"
o Ex: BrO3- is Bromate
o BrO- is hypobromite
o
If you gain 1 oxygen atom from the "ate" ion, the form is "per ate"
o Ex: BrO3- is Bromate
o BrO4- is perbromate
27
Metal or
Polyatomic
Ion
Nonmetal or
Polyatomic
Ion
Mg
F
Fr
S
Ga3+
P
Zn
PO4
NH4
O
Cd
BrO
Ionic Formula
Name
Scandium Bromide
Lead (IV) Carbide
Iron (III) Oxide
Silver Sulfate
Sr(OH)2
Sn(C2H3O2)4
Zinc Perchlorate
Cerium Nitrite
Sodium Phosphide
28
Naming Binary Covalent Compounds
(2 non-metals)
 Write the less electronegative element first
o Electronegative Trend
 From top to bottom – decrease
 From left to right – increase
 General order of nonmetals
o C, P, N, H, S, I, Br, Cl, O, F
A. Similar to naming ionic compounds

the first nonmetal has the name as given on the periodic table

the second nonmetal has the –ide ending
B. Different than naming ionic compounds

There are numerical prefixes that are used depending on how many of each
nonmetal are present in the compound.
Number
Prefix

The only time the prefixes are not used is when
there is only one of the first nonmetal- the term
mono is not used in this case.

1
mono-
2
di-
3
tri-
o NO2 = nitrogen dioxide
o NO = nitrogen monoxide
o N2O4 = dinitrogen tetroxide
4
tetra-
5
penta-
Drop the “a” from tetra before adding oxide
6
hexa-
7
hepta-
8
octa-
9
nona-
10
deca-
o CO= carbon monoxide
o CO2= carbon dioxide
29
Molecular
First non-metal
Second non-metal
2 Carbons
6 Fluorines
1 Phosphorous
5 Oxygens
1 Silicon
4 Chlorines
3 Nitrogens
7 Bromines
Formula
Name
S2I4
F10Cl6
O2Br5
CN3
Trisilicon nonafluoride
Carbon monoxide
30
Naming Various Compounds
Elements
Ionic or
Covalent
Ions
Chemical
Formula
Name of Compound
4 Phosphorus +
10 Oxygens
CS2
Calcium + Chlorine
Iron (III) + Bromine
Na2O
Potassium + Sulfur
Si3N4
Triantimony pentasulfide
FeS
Strontium Chloride
CuO
Nitrogen Dioxide
N2O
31
Tin (IV) + Fluorine
Cl2O7
Phosphorus Triodide
CCl4
Beryllium + Fluorine
Silicon monocarbide
PCl3
Al2O3
1 Carbon + 1 Oxygen
32
Names & Formulas of Acids & Bases
How to recognize a compound/formula as being an acid or base:
Acid: has H, hydrogen, at the beginning.
Base: contains OH, hydroxide.
Naming Acids
Binary Acids
 Contain H and one other element
 To name:
hydro + second element(change ending to –ic) + acid
HCl = Hydro chlorine ic + acid
Hydrochloric acid
HF = Hydro fluorine ic + acid
Hydrofluoric acid
H 2S =
HI =
*If you are given the name and must write the formula, make sure you check
charges and do the criss-cross!
33
Ternary Acids (Oxyacids)
 Contain H, O and one more element
(a polyatomic ion with O in it; SO4)
 To Name:
For –ate and per
-ate polyatomic ions, name the non H, non O element (this is
the 2nd element in the formula) and change the –ate ending to –ic.
1. HNO3 = nitric acid (was nitrate)
2. H2SO4 = sulfuric acid (was sulfate)
3. HC2H3O2 = acetic acid (was acetate)
(acetate is also CH3COO)
4. H2SO5 = persulfuric acid (was persulfate)
5. H3PO4 = phosphoric acid (was phosphate)
6. HClO3 =
7. HNO4 =
8. HIO4=
-ite and hypo -ite polyatomic ions become –ous.
1. HNO2 = nitrous acid (was nitrite)
2. HNO = hyponitrous acid (was hyponitrite)
3. H2SO3 =
4. H2SO2 =
5. HClO =
6. HIO=
34
Naming Bases:
Name the metal first + hydroxide
NaOH = sodium hydroxide
Ca(OH)2 = calcium hydroxide
Fe(OH)3 = iron (III) hydroxide
*If you are given the name and must write the formula, make sure you check
charges and do the criss-cross!
35
Name ____________________________________
Ionic
Covalent
Acid
Formula
Chemical Name
Base
1.
H2SO2
2.
Br2Cl4
3.
Sn3(PO4)2
4.
Ba(IO3)2
5.
H2O
6.
HF
7.
Silver hydroxide
8.
Zinc perchlorate
9.
Potassium sulfite
10.
hypochlorous acid
11.
Trioxygen pentabromide
12.
Hydrosulfuric acid
36
Formula
Acid(binary or ternary)/
Base/Ionic/
or Covalent
Name
1) NaOH
__________
_____________________
2) H2SO4
__________
_____________________
3) SbBr3
__________
_____________________
4) Li2O
__________
_____________________
5) H2CO3
__________
_____________________
6) Si2Br6
__________
_____________________
7) HClO4
__________
_____________________
8) Co2(CO3)3
__________
_____________________
9) HCl
__________
_____________________
10) SCl4
__________
_____________________
11) NH4CN
__________
_____________________
12) Ca(OH)2
__________
_____________________
13) B2Si
__________
_____________________
14) Cu(HCO3)2
__________
_____________________
15) H3PO4
__________
_____________________
16) H3BrO3
__________
_____________________
17) N2O3
__________
_____________________
18) IO5
__________
_____________________
19) HF
__________
_____________________
20) Na2SO3
__________
_____________________
21) CF4
__________
_____________________
22) HClO2
__________
_____________________
23) NiPO3
__________
_____________________
37
24) P2O5
__________
_____________________
25) NiSe
__________
_____________________
26) HNO2
__________
_____________________
27) Cu(CH3COO)2 __________
_____________________
28) H2SO3
__________
_____________________
29) P4S5
__________
_____________________
30) HNO3
__________
_____________________
31) V3(PO5)5
__________
_____________________
32) Al2S3
__________
_____________________
33) HBr
__________
_____________________
34) HNO
__________
_____________________
35) H3PO3
__________
_____________________
36) Sn(OH)2
__________
_____________________
37) FeP
__________
_____________________
38) HClO3
__________
_____________________
39) NaMnO4
__________
_____________________
40) MnF3
__________
_____________________
41) Be(NO2)2
__________
_____________________
42) CO2
__________
_____________________
43) AgBr
__________
_____________________
44) Zn3(PO2)2
__________
_____________________
45) Mn(CO3)2
__________
_____________________
46) Pb3N4
__________
_____________________
38
Name ____________________________________
Fill in the Blanks!
Ionic
Covalent
Acid
Formula
Chemical Name
Base
1.
LiOH
2.
Cl4P6
3.
Mo(BrO2)3
4.
SnO2
5.
HClO4
6.
HF
7.
Copper(II) hydroxide
8.
Zinc persulfate
9.
Barium carbonate
10.
Hypocarbonous acid
11.
Trinitrogen pentiodide
12.
Hydrochloric Acid
39
Name ________________________________________________
Review for Naming Compounds
Number
Ionic or
Name
Formula
Covalent
1
Strontium Phosphate
2
3
CS2
acid
H2CO3
4
Calcium Nitrate
5
Iron (III) Sulfide
6
Fe3P2
7
Germanium tetriodide
8
LiOH
9
Mg(BrO3)2
10
Silicon Monocarbide
11
Barium Cyanide
12
Aluminum Chromate
13
acid
HF
14
15
NH4Cl
acid
Phosphorous Acid
40