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Transcript
CHAPTER 3
Chemical Compounds
Objectives
You will be able to do the following.
1. Given a description of a form of matter, classify as an element, compound, or
mixture.
2. Write a description of the polar covalent bond between the hydrogen atom and
the chlorine atom in HCl. Your description should include a rough sketch of the
electron cloud that represents the electrons involved in the bond.
3. Write a description of the process that leads to the formation of the ionic bond
between sodium and chlorine atoms in sodium chloride.
4. Write a description of the sodium chloride crystal structure.
5. Write a description of the difference between a nonpolar covalent bond, a polar
covalent bond, and an ionic bond. Your description should include rough
sketches of the electron clouds that represent the electrons involved in the
formation of the each bond.
6. Given the names or formulas for two elements, identify the bond that would
form between them as covalent or ionic.
7. Given a table of electronegativities, do the following. (See pages 354-357 of the
text.)
a. Classify chemical bonds as nonpolar covalent, polar covalent, or ionic.
b. Identify which of two atoms in a polar covalent bond has a partial negative
charge and which atom has a partial positive charge.
c. Identify which of two atoms in an ionic bond has a negative charge and
which atom has a positive charge.
d. Given two bonds, determine which of the bonds would be expected to be
more polar.
8. Identify the most common number of covalent bonds that each of the following
elements form: hydrogen, Group 17 (halogens), oxygen, nitrogen, and carbon.
9. Convert between the description of the number of atoms of each element found
in a compound and its chemical formula.
10. Draw Lewis structures from chemical formulas for compounds that have all of
their atoms with their most common bonding pattern.
11. Describe the molecular geometry of CH4, NH3, and H2O molecules.
12. Write a description of the attractions between water molecules in liquid and solid
water.
13. Write a description of the structure of liquid water. Your description should
include a sketch of the particles in the liquid.
14. Convert between the names and chemical formulas for water, ammonia,
methane, ethane, propane, butane, pentane, and hexane.
15. Write or identify prefixes for the numbers 1-10. (For example, mono- represents
one, di- represents two, etc.)
31
32
Chapter 3
Chemical Compounds
16. Write or identify the roots of the nonmetals. (For example, the root for oxygen is
ox−.)
17. Convert between the systematic names and chemical formulas for binary covalent
compounds.
18. Convert between the complete name, the common name, and the chemical
formula for HF, HCl, HBr, HI, and H2S.
19. Write or identify the ionic charges that form for the following elements: group
17 (halogens), oxygen, sulfur, selenium, nitrogen, phosphorus, hydrogen, group
1 (alkali metals), group 2 (alkaline earth metals), group 3 elements, aluminum,
iron, silver, copper, and zinc.
20. Write an explanation for why the elements listed in the previous objective (except
for (iron, silver, copper, and zinc) form their ionic charges.
21. Convert between the names and chemical formulas for the monatomic ions.
22. Given the common name, write the chemical formula for the following ions:
ferric, ferrous, cupric, and cuprous.
23. Convert between the names and chemical formulas for the common polyatomic
ions hydroxide, ammonium, cyanide, acetate, dichromate, permanganate, oxalate,
sulfate, nitrate, phosphate, carbonate, chlorate, chromate, bromate, and iodate.
24. Convert between the names and chemical formulas for the per(root)ate, (root)ite,
and hypo(root)ite oxyanions that have the same element as the (root)ate ions.
25. Convert between the names and chemical formulas for the polyatomic anions
that are derived from the additions of H+ ions to anions with −2 or −3 charges.
For example, H2PO4− is dihydrogen phosphate.
26. Given the common name bicarbonate, bisulfate, or bisulfite, write its chemical
formula.
27. Convert between the names and chemical formulas for ionic compounds.
28. Convert between the name or formula of a (root)ate anion and the name or
formula of the (root)ic acid that forms from the (root)ate anion.
29. Write or identify a description of alcohols.
30. Convert between the correct name and chemical formula for each of the
following alcohols: methanol, ethanol, and 2-propanol.
31. Given one of the following common names for alcohols, write its chemical
formula: methyl alcohol, ethyl alcohol, and isopropyl alcohol.
32. Convert between the name and chemical formula the sugars glucose and sucrose.
33. Given a name or chemical formula, identify whether it represents a binary
ionic compound, an ionic compound with polyatomic ion(s), a binary covalent
compound, a binary acid, an oxyacid, a common alcohol, or a common sugar.
34. Convert between the name and chemical formula for binary ionic compounds,
ionic compounds with polyatomic ion(s), binary covalent compounds, binary
acids, oxyacids, common alcohols, and common sugars.
35. Given a description of a compound, identify it as either a molecular compound
or an ionic compound.
For objectives 36-42, see pages 409-416 of the text.
36. Write or identify a description of dipole-dipole attractions between polar
molecules.
37. Draw a sketch of hydrogen chloride molecules in the liquid form showing the
dipole-dipole attractions that hold the particles together.
33
38. Write or identify a description of hydrogen bonds between molecules of HF,
H2O, and NH3.
39. Write or identify a description of London dispersion forces between nonpolar
molecules. Your description should include mention of how the attractions form
beginning with nonpolar molecules.
40. Write a description of how London forces form between polar molecules.
41. Write an explanation for why larger molecules have stronger London forces.
42. Write or identify the nature of the particles and the nature of the attractions
between the particles in the liquid and solid form for metallic elements, the noble
gases, carbon in the diamond form, the other nonmetallic elements (hydrogen,
nitrogen, oxygen, sulfur, selenium, phosphorus, and the halogens), any ionic
compound, any hydrocarbon, hydrogen chloride, hydrogen fluoride, water,
alcohols, and ammonia.
43. Draw constitutional isomers of C4H10, C5H12, and C6H14.
44. Convert between names and formulas for alkanes, including branched alkanes
with alkyl groups in Table 3.5 of the text.
45. Write or identify the two abbreviations associated with atomic (unified) mass
unit.
46. Given a periodic table that shows atomic masses of the elements, convert between
mass of an element and moles of that element.
47. Given a formula for a molecular substance and a periodic table that includes
atomic masses for the elements, calculate the substance’s molecular mass.
48. Given enough information to calculate a molecular substance’s molecular mass,
convert between mass and moles of the substance.
49. Write a description of the similarities and differences between the term molecule
and the term formula unit.
50. Given a formula for an ionic compound and a periodic table that includes atomic
masses for the elements, calculate the compound’s formula mass.
51. Given enough information to calculate an ionic compound’s formula mass,
convert between mass and moles of the compound.
52. Given a chemical formula for a compound, write conversion factors that convert
between moles of compound and moles of element in the compound.
53. Make conversion between mass of compound and mass of element in the
compound.
54. Given a chemical formula for a compound, calculate the percentage of the
elements in the compound.
55. Convert between the definition and the term for the following words or phrases.
Skip sections 3.10 and 3.12 in the text.
34
Chapter 3
Chemical Compounds
Chapter 3 Glossary
Element A substance that cannot be chemically converted into simpler substances; a
substance in which all of the atoms have the same number of protons and therefore
the same chemical characteristics.
Compound A substance that contains two or more elements, the atoms of these
elements always combining in the same whole-number ratio.
Pure substance A sample of matter that has constant composition. There are two
types of pure substances, elements and compounds.
Mixture A sample of matter that contains two or more pure substances and has
variable composition.
Chemical bond An attraction between atoms or ions in chemical compounds.
Covalent bonds and ionic bonds are examples.
Polar covalent bond A covalent bond in which electrons are shared unequally, leading
to a partial negative charge on the atom that attracts the electrons more and to a
partial positive charge on the other atom.
Nonpolar covalent bond A covalent bond in which the difference in electronattracting ability of two atoms in a bond is negligible (or zero), so the atoms in the
bond have no significant charges.
Ion Any charged particle, whether positively or negatively charged.
Cation An ion formed from an atom that has lost one or more electrons and thus has
become positively charged.
Anion An ion formed from an atom that has gained one or more electrons and thus
has become negatively charged.
Ionic bond The attraction between a cation and an anion.
Ionic hydrate Ionic compounds with water molecules trapped within the crystal
lattice.
Water of hydration The associated water in ionic hydrates.
Electronegativity A measure of the electron-attracting ability of an atom in a chemical
bond.
Molecular compound A compound composed of molecules. In such compounds, all
of the bonds between atoms are covalent bonds.
Ionic compound A compound that consists of ions held together by ionic bonds.
Chemical formula A concise written description of the components of a chemical
compound. It identifies the elements in the compound by their symbols and
indicates the relative number of atoms of each element with subscripts.
Empirical formula A chemical formula that includes positive integers that describe
the simplest ratio of the atoms of each element in a compound.
Molecular formula The chemical formula that describes the actual numbers of atoms
of each element in a molecule of a compound.
Valence electrons The electrons that are most important in the formation of chemical
bonds. The number of valence electrons for the atoms of an element is equal to the
element’s A-group number on the periodic table.
Electron-dot symbol A representation of an atom that consists of its elemental symbol
surrounded by dots representing its valence electrons.
Lewis structure A representation of a molecule that consists of the elemental symbol
for each atom in the molecule, lines to show covalent bonds, and pairs of dots to
indicate lone pairs.
35
Double bond A link between atoms that results from the sharing of four electrons. It
can be viewed as two 2-electron covalent bonds.
Triple bond A link between atoms that results from the sharing of six electrons. It can
be viewed as three 2-electron covalent bonds.
Lone pair Two electrons that are not involved in the covalent bonds between atoms
but are important for explaining the arrangement of atoms in molecules. They are
represented by pairs of dots in Lewis structures.
Bond angle The angle formed by straight lines (representing bonds) connecting the
nuclei of three adjacent atoms.
Tetrahedral The molecular shape that keeps the negative charge of four electron groups
as far apart as possible. This shape has angles of 109.5° between the atoms.
Binary covalent compound A compound composed of two nonmetallic elements.
Organic chemistry The branch of chemistry that involves the study of carbon-based
compounds.
Organic compound A carbon-based compound.
Hydrocarbons Compounds that contain only carbon and hydrogen.
Alcohols Compounds that contain a hydrocarbon group with one or more -OH
groups attached.
Monatomic anions Negatively charged particles, such as Cl−, O2−, and N3−, that
contain single atoms with a negative charge.
Monatomic cations Positively charged particles, such as Na+, Ca2+, and Al3+, that
contain single atoms with a positive charge.
Binary ionic compound An ionic compound whose formula contains one symbol
for a metal and one symbol for a nonmetal.
Polyatomic ion A charged collection of atoms held together by covalent bonds.
Oxyanion A polyatomic ions with the general formula HaXbOcd−. (The a can be 0.)
Binary acid Substances that have the general formula of HX(aq), where X is one of
the first four halogens HF(aq), HCl(aq), HBr(aq), and HI(aq).
Oxyacids (or oxoacids) Molecular substances that have the general formula HaXbOc.
In other words, they contain hydrogen, oxygen, and one other element represented
by X; the a, b, and c represent subscripts.
Intermolecular attraction Attraction between molecules.
Dipole A molecule that contains an asymmetrical distribution of positive and negative
charges.
Dipole-dipole attraction The intermolecular attraction between the partial negative
end of one polar molecule and the partial positive end of another polar molecule.
Hydrogen bond The intermolecular attraction between a nitrogen, oxygen, or
fluorine atom of one molecule and a hydrogen atom bonded to a nitrogen, oxygen,
or fluorine atom in another molecule.
Metallic bond The attraction between the positive metal cations that form the basic
structure of a solid metal and the negative charge from the mobile sea of electrons
that surround the cations.
London dispersion forces, London forces, or dispersion forces The attractions
produced between molecules by instantaneous and induced dipoles.
Isomers Compounds that have the same molecular formula but different molecular
structures.
36
Chapter 3
Chemical Compounds
Constitutional isomers (also called structural isomers) Compounds with the
same molecular formula that differ in the order in which their atoms are bonded
together.
Alkanes Hydrocarbons (compounds composed of carbon and hydrogen) in which all
of the carbon-carbon bonds are single bonds.
Weighted average mass A mass calculated by multiplying the decimal fraction of each
component in a sample by its mass and adding the results of each multiplication
together.
Atomic mass unit 1/12 the mass of a carbon-12 atom. It is sometimes called a unified
mass unit. Its accepted abbreviation is u, but amu is sometimes used.
Atomic mass The weighted average of the masses of the naturally occurring isotopes
of an element.
Mole The amount of substance that contains the same number of particles as there
are atoms in 12 g of carbon-12.
Avogadro’s number The number of atoms in 12 g of carbon-12. To four significant
figures, it is 6.022 x 1023.
Molar mass The mass in grams of one mole of substance. (The number of grams in
the molar mass of an element is the same as its atomic mass. The number of grams
in the molar mass of a molecular compound is the same as its molecular mass.
The number of grams in the molar mass of an ionic compound is the same as its
formula mass.)
Molecular mass The weighted average of the masses of the naturally occurring
molecules of a molecular substance. It is the sum of the atomic masses of the atoms
in a molecule.
Formula unit A group represented by a substance’s chemical formula, that is, a group
containing the kinds and numbers of atoms or ions listed in the chemical formula.
It is a general term that can be used in reference to elements, molecular compounds,
or ionic compounds.
Formula mass The weighted average of the masses of the naturally occurring formula
units of the substance. It is the sum of the atomic masses of the atoms in a formula
unit.
37
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Figure 3.1
Elements Versus Compounds
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Figure 3.2
Ionic Bond Formation
38
Chapter 3
Chemical Compounds
Summary of Covalent and Ionic Bond Formation
When atoms of different elements form chemical bonds, the electrons in the bonds
can shift from one bonding atom to another.
The atom that attracts electrons more strongly will acquire a negative charge, and
the other atom will acquire a positive charge.
The more the atoms differ in their electron-attracting ability, the more the electron
cloud shifts from one atom toward another.
If there is a large enough difference in electron-attracting ability, 1, 2, or 3 electrons
can be viewed as shifting completely from one atom to another. The atoms become
positive and negative ions, and the attraction between them is called an ionic
bond.
If the electron transfer is significant but not enough to form ions, the atoms acquire
partial positive and partial negative charges. The bond in this situation is called a
polar covalent bond.
If there is no shift of electrons or if the shift is negligible, no significant charges will
form, and the bond will be a nonpolar covalent bond.
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Figure 3.3 Classifying Compounds
Table 3.1 Guidelines for Classifying Bonds as Covalent or Ionic
Combination
Type of bond
Examples
Metal-Nonmetal
Usually Ionic
NaCl or MgO
Nonmetal-Nonmetal
Covalent
HCl or CO
EXERCISE 3.1 - Classifying Compounds
Classify each of the following substances as either a molecular compound or an ionic
compound.
a. formaldehyde, CH2O (used in embalming fluids)
b. magnesium chloride, MgCl2 (used in fireproofing wood and in paper
manufacturing)
39
Figure 3.4
Electronegativities
1
1A
2
3
4
5
6
7
18
8A
2
2A
1
13
3A
2.20
H
0.98 1.57
Li
Be
0.93
1.31
Na
Mg
3
3B
4
4B
0.82 1.00
1.36 1.54
Sc
Ti
0.82 0.95
1.22
1.33
0.79 0.89
1.27
K
Ca
Rb
Sr
Cs
Ba
0.7
Y
Lu
5
5B
6
6B
7
7B
1.63 1.66 1.55
V
Cr
Mn
Zr
Nb
1.6
2.16
Mo
Tc
1.3
1.5
2.36
1.9
Hf
Ta
W
1.9
Re
17
7A
2.04 2.55
3.04
3.44
3.98
1.61 1.90
2.19 2.58
3.16
2.01 2.18 2.55
2.96
C
9
8B
10
8B
1.83
1.88
1.91 1.90 1.65 1.81
2.2
2.28 2.20 1.93 1.69 1.78 1.96
2.2
2.20 2.28 2.54 2.00 1.62
Ru
Os
Co
Rh
Ir
Ni
Pd
Pt
Cu
Ag
Au
12
2B
16
6A
8
8B
Fe
11
1B
15
5A
B
14
4A
Zn
Cd
Hg
Al
Ga
In
Tl
Si
N
P
Ge
O
S
F
Cl
As
Se
2.05
2.1
2.66
2.33 2.02
2.0
2.2
Sn
Pb
Sb
Bi
Te
Po
Br
I
3.00
Kr
2.6
Xe
At
0.9
Fr
Ra
TIP-OFF You wish to (1) classify a chemical bond as nonpolar covalent, polar covalent,
or ionic, (2) identify which element in a polar covalent bond is partially negative and
which is partially positive, (3) identify which element in an ionic bond forms the
anion and which forms the cation, or (4) identify which of two bonds is more polar.
GENERAL STEPS
•
Use the following guidelines to identify a chemical bond as ionic,
nonpolar covalent, or polar covalent. (If both atoms are nonmetals, the
bond is covalent.)
∆EN < 0.4 → Nonpolar Covalent
∆EN 0.4-1.7 → Polar Covalent
∆EN > 1.7 → Ionic
•
Use the following guidelines to identify which element in a polar covalent
bond is partially negative and which is partially positive.
Higher electronegativity → partial negative charge
Lower electronegativity → partial positive charge
•
Use the following guidelines to identify which element in an ionic bond
forms the anion and which forms the cation.
Nonmetal, which has a higher electronegativity → anion
Metal, which has the lower electronegativity → cation
•
Use the following guideline to decide which of two bonds is more polar.
The greater the ∆EN is, the more polar the bond.
Sample Study
Sheet 3.1:
Electronegativity,
Types of
Chemical Bonds,
and
Bond Polarity
40
Chapter 3
Chemical Compounds
Figure 3.5
Electronegativities
and Bond Type
EXERCISE 3.2 - Electronegativities and Bond Type
Classify the following bonds as nonpolar covalent, polar covalent, or ionic. If a bond
is polar covalent, identify which atom has the partial negative charge and which has
the partial positive charge. If a bond is ionic, identify which atom has the negative
charge and which has the positive charge.
a. N bonded to H
b. N bonded to Cl
c. Ca bonded to O
d. P bonded to F
EXERCISE 3.3 - Electronegativities and Bond Polarity
Which bond would you expect to be more polar, P-H or P-F?
41
Table 3.2
Electron-Dot Symbols and Usual Numbers of Bonds and Lone Pairs for Nonmetallic
Elements
Group 4A
4 valence electrons
Group 5A
5 valence electrons
Group 6A
6 valence electrons
Group 7A
7 valence electrons
X
X
X
X
4 bonds No lone 3 bonds 1 lone
pairs
pair
carbon-C
nitrogen-N
C
2 bonds
2 lone
pairs
oxygen-O
1 bond
3 lone
pairs
fluorine-F
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phosphorus-P
sulfur-S
chlorine-Cl
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selenium-Se
bromine-Br
��
��
iodine-I
�
EXERCISE 3.4 - Drawing Lewis Structures from Formulas
Draw a Lewis structure for each of the following formulas:
a. nitrogen triiodide, NI3 (explodes at the slightest touch)
b. hexachloroethane, C2Cl6 (used to make explosives)
c. hydrogen peroxide, H2O2 (a common antiseptic)
d. ethylene (or ethene), C2H4 (used to make polyethylene)
42
Chapter 3
Chemical Compounds
Figure 3.6
Three Ways
to Describe a
Methane
Molecule
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Figure 3.7
Three Ways to
Describe an
Ammonia
Molecule
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Figure 3.8
Three ways to
Describe a Water
Molecule
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Figure 3.9
Attractions Between
water Molecules
Figure 3.10
Liquid Water
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43
A Shockwave animation on our website will help you to visualize the structure
of water:
http://www.mpcfaculty.net/mark_bishop/water.htm
Table 3.3 Names to Memorize for Some Binary Covalent Compounds
Name
Formula
Name
Formula
water
H2O
ammonia
NH3
methane
CH4
ethane
C2H6
propane
C3H8
butane
C4H10
pentane
C5H12
hexane
C6H14
Follow these steps to write the names for binary covalent compounds.
•
If the subscript for the first element is greater than one, indicate the identity
of the subscript using prefixes from the table below. We do not write monoat the beginning of a compound’s name.
Example: We start the name for N2O3 with di-.
•
Attach the selected prefix to the name of the first element in the formula. If
no prefix is to be used, begin with the name of the first element.
Example: We indicate the N2 portion of N2O3 with dinitrogen.
•
Select a prefix to identify the subscript for the second element (even if its
subscript is understood to be one). Leave the “a” off the end of the prefixes
that end in “a” and the “o” off of mono- if they are placed in front of an
element whose name begins with a vowel (oxygen or iodine).
Example: The name of N2O3 grows to dinitrogen tri-.
•
Write the root of the name of the second element in the formula.
Example: The name of N2O3 becomes dinitrogen triox-.
•
Add -ide to the end of the name.
Example: The name of N2O3 is dinitrogen trioxide.
Table 3.4 Prefixes Used in the Names of Binary Covalent Compounds to Indicate the
Number of Atoms of each Element in the Formula
Number of
atoms
1
2
3
4
5
Prefix
mon(o)ditritetr(a)pent(a)-
Number
of atoms
6
7
8
9
10
Prefix
hex(a)hept(a)oct(a)non(a)dec(a)-
44
Chapter 3
Chemical Compounds
Table 3.5 Roots for the Nonmetals
Element
C
Root
carb-
Element
Root
N
nitrP
phosphAs
arsen-
Element
O
S
Se
Root
oxsulfselen-
Element
F
Cl
Br
I
Root
fluorchlorbromiod-
Table 3.6 Names for Binary Covalent Compounds with Atoms that Combine in Only
One Ratio
Formula
HF
HCl
HBr
HI
H2S
Complete Name
hydrogen monofluoride
hydrogen monochloride
hydrogen monobromide
hydrogen moniodide
dihydrogen monosulfide
(or dihydrogen sulfide)
Common Name
hydrogen fluoride
hydrogen chloride
hydrogen bromide
hydrogen iodide
hydrogen sulfide
The first step in writing formulas when given the systematic name of a binary covalent
compound is to recognize the name as representing a binary covalent compound. It
will have one of the following general forms.
prefix(name of nonmetal) prefix(root of name of nonmetal)ide (e.g.
dinitrogen pentoxide)
or (name of nonmetal) prefix(root of name of nonmetal)ide (e.g. carbon
dioxide)
or (name of nonmetal) (root of nonmetal)ide (e.g. hydrogen fluoride)
Follow these steps for writing formulas for binary covalent compounds when you
are given a systematic name. Notice that they are the reverse of the steps for writing
names from chemical formulas.
• Write the symbols for the elements in the order mentioned in the name.
• Write subscripts indicated by the prefixes. If the first part of the name has no
prefix, assume it is mono-.
EXERCISE 3.5 - Naming of Binary Covalent Compounds
Write names that correspond to the following formulas: (a) P2O5, (b) PCl3, (c) CO,
(d) H2S, and (e) NH3.
EXERCISE 3.6 - Writing Formulas for Binary Covalent Compounds
Write formulas that correspond to the following names: (a) disulfur decafluoride, (b)
nitrogen trifluoride, (c) propane, and (d) hydrogen chloride.
45
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Figure 3.11
Common Monatomic Ions
Table 3.7
Names of the Monatomic Anions
Anion
N3−
P3−
Name
nitride
phosphide
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sulfide
selenide
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Monatomic Anion Charges
46
Chapter 3
Chemical Compounds
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Figure 3.13 Monatomic Anion Charges
EXERCISE 3.7 - Naming Monatomic Ions
Write names that correspond to the following formulas for monatomic ions: (a) Mg2+,
(b) F−, and (c) Sn2+.
EXERCISE 3.8 - Formulas for Monatomic Ions
Write formulas that correspond to the following names for monatomic ions:
(a) bromide ion, (b) aluminum ion, and (c) gold(I) ion.
Figure 3.14 Sodium Chloride Structure
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47
Figure 3.15
Cesium Chloride Crystal Structure
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Table 3.8 Ionic Charges to Memorize
Element
Charge
Element
Charge
iron
2+ or 3+
copper
1+ or 2+
zinc
2+
cadmium
2+
silver
1+ (and very rarely +2)
Table 3.9 Naming Metals with Two Possible Charges
Ion
Fe2+
Systematic
(Preferred) Name
iron(II)
Nonsystematic
Name
ferrous
Fe3+
iron(III)
ferric
Cu+
copper(I)
cuprous
Cu2+
copper(II)
cupric
Table 3.10 Polyatomic ions with common names.
Ion
OH−
Name
hydroxide
Ion
NH4+
Name
ammonium
CN−
cyanide
C2H3O2−
acetate
C2O42−
oxalate
MnO4−
permanganate
Cr2O72−
dichromate
48
Chapter 3
Chemical Compounds
Table 3.11 Convention for Naming Oxyanions
Relationship
General Name
Example Name
Example
Formula
ClO4−
one more oxygen atom
than (root)ate
per(root)ate
perchlorate
(root)ate
chlorate
ClO3−
one less oxygen atom
than (root)ate
(root)ite
chlorite
ClO2−
two less oxygen atoms
than (root)ate
hypo(root)ite
hypochlorite
ClO−
Table 3.12 (Root)ate Polyatomic Ions
Ion
SO42−
Name
sulfate
Ion
NO3−
Name
nitrate
PO43−
phosphate
CO32−
carbonate
ClO3−
chlorate
CrO42−
chromate
BrO3−
bromate
IO3−
iodate
Table 3.13 Formula and Names for Polyatomic Ions Containing Hydrogen
Formula
HCO3−
HS−
Name
hydrogen carbonate
Formula
HPO42−
Name
hydrogen phosphate
hydrogen sulfide
H2PO4−
dihydrogen phosphate
Table 3.14 Systematic and Nonsystematic Names for Some Polyatomic Ions
Formula
HCO3−
Systematic (Preferred) Name
hydrogen carbonate
Nonsystematic Name
bicarbonate
HSO4−
hydrogen sulfate
bisulfate
HSO3−
hydrogen sulfite
bisulfite
Note: You should use the systematic name, but it will also be useful to know the nonsystematic names.
Table 3.15 Summary of the Ways that Cations are Named
Type of cation
General Name
Example
For metals with one possible charge
(Groups 1, 2, 3 - Al, Zn, and Cd)
name of metal
Mg2+ is
magnesium.
For metals with more than one
possible charge (The rest of the
metals)
name of metal(Roman
numeral)
Cu2+ is
copper(II).
Polyatomic cations
Ammonium is our only
example in this category.
NH4+ is
ammonium.
49
Table 3.16 Summary of the Ways that Anions are Named
Type of anion
General Name
Example
For monatomic anions
(root of nonmetal)ide
O2− is oxide.
For polyatomic anions
name of polyatomic ion
NO3− is nitrate.
EXERCISE 3.9 - Ionic Formulas to Names
Write the names for the following formulas: LiCl, CaSO4, MnF3, NH4F, Cr2S3,
Mg3(PO3)2, ZnCrO4, and AgBrO2
EXERCISE 3.10 - Ionic Names to Formulas
Write the formulas for the following names: aluminum oxide, ammonium chloride,
cobalt(II) sulfide, ferrous sulfate, silver chlorite, ammonium hydrogen phosphate,
and calcium bicarbonate.
Table 3.17
Two ways to name binary covalent hydrides
Formula
Named as Binary Covalent
Formula
Named as Binary acid
HF
or HF(g)
hydrogen monofluoride or
hydrogen fluoride
HF(aq)
hydrofluoric acid
HCl
or HCl(g)
hydrogen monochloride or
hydrogen chloride
HCl(aq)
hydrochloric acid
HBr
or HBr(g)
hydrogen monobromide or
hydrogen bromide
HBr(aq)
hydrobromic acid
HI
or HI(g)
hydrogen moniodide or
hydrogen iodide
HI(aq)
hydroiodic acid
or hydriodic acid
H2S
or H2S(g)
dihydrogen sulfide or
hydrogen sulfide
H2S(aq)
hydrosulfuric acid
50
Chapter 3
Chemical Compounds
Table 3.18 Relationship Between the (Root)ate Polyatomic Ions and the (Root)ic Acids
Oxyanion
formula
NO3−
C2H3O2−
SO42−
Oxyanion
name
nitrate
acetate
sulfate
Oxyacid
formula
HNO3
HC2H3O2
H2SO4
CO32−
PO43−
carbonate
phosphate
H2CO3
H3PO4
ClO3−
BrO3−
IO3−
C2O42−
CrO42−
chlorate
bromate
iodate
oxalate
chromate
HClO3
HBrO3
HIO3
H2C2O4
H2CrO4
Oxyacid name
nitric acid
acetic acid
sulfuric acid (Note that the whole name
sulfur is used in the oxyacid name.)
carbonic acid
phosphoric acid (Note that an the root
of phosphorus in an oxyacid name is
phosphor-.)
chloric acid
bromic acid
iodic acid
oxalic acid
chromic acid
Note: When enough H+ ions are added to the (root)ate polyatomic ion to completely neutralize the
charge, the (root)ic acid forms.
Table 3.19 Convention for Naming Oxyacids
Relationship
one more oxygen
atom than (root)ic
one less oxygen
atom than (root)ic
two less oxygen
atoms than
(root)ic
General Name
Example Name
per(root)ic acid
perchloric acid
Example
Formula
HClO4
(root)ic acid
(root)ous acid
chloric acid
chlorous acid
HClO3
HClO2
hypo(root)ous acid
hypochlorous acid
HClO
Table 3.20 Names of Common Alcohols (You should memorize them.)
Preferred Name
methanol
ethanol
2-propanol
Common Name
methyl alcohol
ethyl alcohol
isopropyl alcohol
Formula
CH3OH
C2H5OH
C3H7OH
Table 3.21 Names and Formulas for Common Sugars (You should memorize them.)
Name
glucose
sucrose
Formula
C6H12O6
C12H22O11
51
Table 3.22 Identifying Types of Compounds from Formulas
Type of Compound
General Formula
Example
binary ionic
MaAb
NaCl
ionic with an
oxyanion
MaHbXcOd or
(NH4)aHbXcOd
(b can be 0)
Li2HPO4 or NaNO3 or
NH4NO3
binary covalent
AaBb
N2O5
binary acid
HX(aq) or H2S(aq)
HCl(aq) or H2S(aq)
oxyacid
HaXbOc
H2SO4
hydrocarbon
CaHb
C2H6
alcohol
memorize examples
CH3OH, C2H5OH, C3H7OH
sugar
memorize examples
C6H12O6, C12H22O11
M = symbol of metal
A and B = symbols of nonmetals
X = some element other than H or O
a, b, c & d indicate subscripts
Table 3.23 Identifying Types of Compounds from Names
Type of Compound
Tip-off
Examples
binary ionic
-ide ending with metalnonmetal
sodium chloride
ionic with an oxyanion
-ite or -ate ending
sodium nitrate
a. with common names
recognize name
water or methane
b. hydrocarbons
There are other possibilities,
but those we will see will
include: (root)ane
n-pentane
c. Other
-ide ending with
nonmetal-nonmetal
nitrogen
trifluoride
hydro(root)ic acid
hydrochloric acid
per(root)ic acid
perchloric acid
(root)ic acid
chloric acid
(root)ous acid
chlorous acid
hypo(root)ous acid
hypochlorous acid
alcohol
name ends in -anol or
alcohol
methanol or
methyl alcohol
sugar
name ends in -ose
glucose or sucrose
binary covalent
binary acid
oxyacid
52
Chapter 3
Chemical Compounds
EXERCISE 3.11 - Nomenclature, Formulas to Names
Write names for the following formulas: P2O5, PCl3, CO, H2S(g), H2S(aq), NH3,
H3PO4, H3PO3, H3PO2, CH3OH, and C12H22O11.
EXERCISE 3.12 - Nomenclature, Names to Formulas
Write formulas for the following names: disulfur decafluoride, nitrogen trifluoride,
butane, hydrogen chloride, hydrochloric acid, carbonic acid, periodic acid, ethanol or
ethyl alcohol, and glucose.
53
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London Forces
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Larger Molecules, Stronger
London Forces
54
Chapter 3
Chemical Compounds
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Figure 3.19
Polar Molecules and
London forces
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Solid and Liquid Form
Table 3.24
Type of Substance
Elements
Metal
Noble Gases
Carbon (diamond)
Other Nonmetal
Elements
Ionic
Compounds
Molecular
Compounds
Nonpolar
Molecular
Polar molecules
without H-F, O-H,
or N-H bond
Molecules with
H-F, O-H or N-H
bond
Particles to
Visualize
Cations in a
sea of electrons
Atoms
Carbon Atoms
Molecules
Examples
Au
Type of Attraction
Between Particles in
Solid or Liquid
Metallic Bond
Xe
C(dia)
H2, N2, O2,
F2, Cl2, Br2,
I2, S8, Se8, P4
NaCl
London Forces
Covalent Bonds
London Forces
Hydrocarbons
London Forces
Molecules
HCl
Dipole-Dipole Forces
Molecules
HF, H2O,
alcohols, NH3
Hydrogen Bonds
Cations and
Anions
Molecules
Ionic Bond
55
EXERCISE 3.13 - Types of Particles and Attractions
Complete the following table by (1) writing the name for the type of particle viewed
as forming the structure of a solid, liquid, or gas of each of the following substances
and (2) writing the name of the type of attraction holding these particles in the solid
and liquid form.
Substance
Particles to Visualize
Type of Attraction
iron
MgO
iodine
CH3OH
NH3
hydrogen chloride
C (diamond)
lithium sulfate
Figure 3.19
Types of Attractions
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56
Chapter 3
Sample Study
Sheet 3.3:
Calculation of
Atomic Mass
Chemical Compounds
TIP-OFF - You are asked to calculate atomic mass, and you are given the percentage
abundance and mass of each isotope. (The tip-offs are more subtle for other types of
calculations.)
GENERAL STEPS
•
Write the memorized atomic mass equation with the decimal fraction and
mass of each isotope given in the problem. (The decimal fraction can be
calculated by dividing the percentage abundance by 100.)
for each
Atomic mass = ∑ (decimal fraction of isotope) (mass of isotope)
isotope
Decimal fraction of isotope =
•
percent abundance
100
Complete the calculation and report your answer.
If you are told the number of decimal positions to report, round your
answer accordingly.
If you are not given the number of decimal positions to report, use the
rules for significant figures to decide how to round off your answer.
Atomic masses can be reported with the unit u for unified mass unit (or
atomic mass unit), but they are often reported without a unit.
Sample Study
Sheet 3.4:
Molar Mass
Calculations
for Elements
TIP-OFF - You are converting between moles and mass of element.
GENERAL STEPS - The procedure involves unit analysis using the molar mass as a
conversion factor. (See Figure 3.1.)
•
The molar mass comes from the atomic mass for an element.
for each
Atomic mass = ∑ (decimal fraction of isotope) (mass of isotope)
isotope
•
Use the following general form for the conversion factor that comes from
molar mass.
Atomic Mass
Mass of element
Molar mass
(atomic mass) g element
1 mol element
moles of atoms
57
EXERCISE 3.14 - Element Molar Mass Calculation
An analysis of the element lithium shows that 7.5% of the lithium atoms are
lithium-6 atoms, and 92.5% are lithium-7 atoms. Each atom of lithium-6 has a mass
of 6.0151214 u, and each atom of lithium-7 has a mass of 7.0160030 u.
a. What is the atomic mass of lithium? (Report your answer to the third
decimal position, ±0.001.)
b. Write a conversion factor that will convert between grams of the element
lithium and moles of lithium.
c. How many moles of lithium are in a sample of lithium that has a mass of
7.249 pounds?
EXERCISE 3.15 - Element Molar Mass Calculation
Gold is often sold in units of troy ounces. (To four significant figures, there are 31.10
grams per troy ounce.) How many moles of gold, Au, are there is 1.00 troy ounce of
pure gold?
58
Chapter 3
Chemical Compounds
Figure 3.20
Formula Units
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59
TIP-OFF - You are converting between moles and mass for a molecular compound.
GENERAL STEPS - The procedure involves unit analysis using the molar mass as a
conversion factor.
•
Form molecular compounds, the molar mass comes from the molecular mass
of a molecular compound.
Molecular mass =
•
for each atom
∑
in a molecule
atomic mass
Sample Study
Sheet 3.5:
Molar Mass
Calculations
for Molecular
Compounds
Use the following general form for the conversion factor that comes from
molar mass.
Molecular Mass
Mass of molecular compound
Molar mass
moles of molecules
(molecular mass) g molecular compound
1 mol molecular compound
EXERCISE 3.16 - Molar Mass and Molecular Compounds
A typical 6.0 fluid ounce glass of wine contains about 16 g of ethanol, C2H5OH.
a. What is the molecular mass of C2H5OH?
b. Write a conversion factor that will convert between mass and moles of
C2H5OH.
c. What is the volume in milliliters of 1.0 mole of pure C2H5OH? (The density of
ethanol is 0.7893 g/mL.)
60
Chapter 3
Chemical Compounds
Sample Study
Sheet 3.6
Molar Mass
Calculations
for Ionic
Compounds
TIP-OFF - You are converting between moles and mass for an ionic compound.
GENERAL STEPS - The procedure involves unit analysis using the molar mass as a
conversion factor.
•
The molar mass comes from the atomic mass for an element.
Formula mass =
•
for each atom
∑
in a formula
unit
atomic mass
Use the following general form for the conversion factor that comes from
molar mass.
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EXERCISE 3.17 - Molar Mass and Ionic Compounds
A quarter teaspoon of a typical baking powder contains about 0.4 g of sodium hydrogen
carbonate, NaHCO3, which is often called bicarbonate of soda.
a. Calculate the formula mass of sodium hydrogen carbonate.
b. Write a conversion factor that could be used to convert between mass and
moles of NaHCO3.
c. How many moles of NaHCO3 are there is 0.4 g of NaHCO3?
61
TIP-OFF - When you analyze the type of unit you have and the type of unit you want, Sample Study
you recognize that you are converting between a unit associated with an element and a Sheet 3.7:
unit associated with a compound containing that element.
Converting
GENERAL STEPS
Between Mass
• Convert the given unit to moles of the first substance.
of Element
This step often requires converting the given unit into grams, after which
and Mass of
the grams can be converted into moles using the molar mass of the
Compound
substance.
Containing the
• Convert moles of the first substance to moles of the second substance
Element
using the molar ratio derived from the formula for the compound.
You either convert from moles of element to moles of compound or
moles of compound to moles of element.
•
Convert moles of the second substance to the desired units of the second
substance.
This step requires converting moles of the second substance into grams of
the second substance using the molar mass of the second substance, after
which the grams can be converted to the specific units that you want.
The following describes a shortcut for these problems.
Use the following general conversion factor in your unit analysis set-up.
Like all conversion factors, this conversion factor can be used in the form
described below or in the inverted form.
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EXERCISE 3.18 - Conversion Between Mass of Element and
Mass of Compound
Disulfur dichloride, S2Cl2, is used in vulcanizing rubber and hardening soft woods. It
can be made from the reaction of pure sulfur with chlorine gas. What is the mass of
S2Cl2 that contains 1.238 kg S?
62
Chapter 3
Chemical Compounds
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Figure 3.21
General Steps for Converting between the Mass of an Element and the Mass of a
Compound Containing the Element The calculation can be set up to convert from the
mass of an element to the mass of a compound (top to bottom) or from the mass of a
compound to the mass of an element (bottom to top).
EXERCISE 3.19 - Conversion Between Mass of element and Mass
of Compound
Vanadium metal, used as a catalyst and to make steel, is produced from the reaction
of vanadium(V) oxide, V2O5, and calcium metal. What is the mass in kilograms of
vanadium in 2.3 metric tons of V2O5?
EXERCISE 3.20 - Conversion Between Mass of element and Mass
of Compound
Calamine has two definitions. It is a naturally occurring zinc silicate that has the
equivalent of 67.5% zinc oxide, ZnO, and it is a substance that is used to make
the calamine lotion. The calamine used for the lotion is 98% ZnO. The naturally
occurring calamine is used to make zinc metal. What is the maximum mass in
kilograms of zinc in 1.347 x 104 kg of natural calamine that is 67.5% ZnO?