Download Chapter 3 Molecules, Compounds, and Chemical Equations

Chapter 3
Molecules, Compounds, and
Chemical Equations
1
Elements and Compounds
• Elements combine together to make an almost limitless
number of compounds.
• The properties of the compound are totally different from the
constituent elements.
• Compounds are made of atoms held together by chemical
bonds.
• Bonds are forces of attraction between atoms.
• The bonding attraction comes from attractions between
protons and electrons
2
Bond Types
• Two general types of bonding between atoms are
found in compounds, ionic and covalent.
• Ionic bonds result when electrons have been
transferred between atoms, resulting in oppositely
charged ions that attract each other.
9 generally found when metal atoms bonded to nonmetal
atoms Usually metal + nonmetal
• Covalent bonds result when two atoms share
some of their electrons.
9 generally found when nonmetal atoms bonded
together
3
3
1
Which of the following compounds exhibits
both ionic and covalent bonding?
a.
b.
c.
d.
e.
SO2
CF4
NaCl
Na2SO4
P4O10
4
5
5
Types of Formula: Empirical Formula
• Empirical Formulas describe the kinds of
elements found in the compound and the
ratio of their atoms.
9They do not describe how many atoms, the
order of attachment, or the shape.
9The formulas for ionic compounds are empirical.
The empirical formula for the ionic compound fluorspar is CaCl2.
This means that there is 1 Ca2+ ion for every 2 Cl− ions in the
compound.
The empirical formula for the molecular compound oxalic acid is CHO2.
This means that there is 1 C atom and 1 H atom for every 2 O atoms in
the molecule. The actual molecular formula is C2H2O46.
2
Types of Formula: Molecular Formula
• Molecular Formulas describe the kinds of
elements found in the compound and the
exact numbers of their atoms.
The molecular formula for oxalic acid is C2H2O4. This does
not tell you that the shape of the molecule.
7
Types of Formula: Structural Formula
• Structural Formulas describe the kinds of
elements found in the compound, the numbers of
their atoms, the order of atom attachment, and
the kind of attachment.
9 use lines to represent covalent bonds
9 Each line describes the number of electrons shared by
the bonded atoms.
¾ single line = 2 shared electrons, a single covalent bond
¾ double line = 4 shared electrons, a double covalent bond
¾ triple line = 6 shared electrons, a triple covalent bond
8
Example of all three types of structures
Structural Formula of Oxalic Acid
but Molecular Formula of Oxalic Acid
C2H2O4.
And Empirical Formula of Oxalic Acid (simplest ratio)
CHO2.
9
3
Clicker question #3-1
H
ethylene glycol
C O
What is the empirical formula for ethylene glycol?
a) C2H6O2
b) CHO
c) CH3O
d) C2H3O2
10
Molecular View of
Elements and Compounds
11
“Molecular” Elements
• Certain elements occur as two-atom molecules.
9 Rule of 7’s
• Other elements occur as polyatomic molecules.
9P4, S8, Se8
H2
7A
N2
7
O2 F2
Cl2
Br2
I2
12
4
Ionic vs. Molecular Compounds
Propane—contains
individual C3H8
molecules
Table salt—contains
an array of Na+ ions
and Cl– ions
13
Ionic Compounds
• Compounds of metals with nonmetals are
made of ions.
9Metal atoms form cations; nonmetal atoms
form anions.
• no individual molecule units, instead they
have a three-dimensional array of cations
and anions made of formula units
• many contain polyatomic ions
9several atoms attached together in one ion
examples; CO3-2 NO314
Write the formula of a compound made
from aluminum ions and oxide ions.
1. Write the symbol for the
+3
metal cation and its charge. Al column 3A
2. Write the symbol for the
3.
4.
5.
O-2 column 6A
nonmetal anion and its
charge.
Charge (without sign)
becomes subscript for other
ion.
Reduce subscripts to
smallest whole number ratio.
Check that the total charge
of the cations cancels the
total charge of the anions.
from
Chapter 2
Al+3 O2Al2 O3
Al = (2) × (+3) = +6
O = (3) × (–2) = –6
15
5
Practice—What are the formulas for
compounds made from the following ions?
potassium ion with a nitride ion
• K+ with N3–
calcium ion with a bromide ion
16
Clicker question #3-2
What is the formula for the compound
made from the following ions?
Cr 6+ and O 21.
2.
3.
4.
CrO
Cr6O2
Cr2O6
CrO3
17
*Important
Formula-to-Name
*Rules for Ionic Compounds*
• made of cation and anion
• Write systematic name by simply naming
the ions.
9if cation is:
¾metal with invariant charge = metal name
¾metal with variable charge = metal name (charge)
¾polyatomic ion = name of polyatomic ion
9if anion is:
¾nonmetal = stem of nonmetal name + -ide
¾polyatomic ion = name of polyatomic ion
18
6
Metal Cations
1) metals with
invariant charge
9metals whose ions can
have only one possible
charge
¾Groups 1A & 2A ,
and Al, Ag, Zn, Sc
9cation name = metal
name
19
2) Metals with variable
charges
metals whose ions
can have more than
one possible charge
determine charge of
cation by charge on
anion
name = metal name
with Roman numeral
charge in
parentheses
20
Naming Nonmetal Anion
• Determine the charge from position on the
Periodic Table.
• To name anion, change ending on the
element name to -ide.
4A = 4−
5A = 3−
6A = 2−
7A = 1−
C = carbide
N = nitride
O = oxide
F = fluoride
Si = silicide P = phosphide S = sulfide Cl = chloride
Br = bromide
I = iodide
21
7
REVIEW
Predictable Ion Charges (for group A
only)
C4 Si4 - P3-
22
22
Naming Binary Ionic Compounds for
Metals with Invariant Charge
Name CsF
1. Identify cation and anion.
Cs = Cs+ because it is in Group 1A
F = F– because it is in Group 7A
2. Name the cation Cs+ = cesium
3. Name the anion F– = fluoride
4. Write the cation name first, then the anion name.
23
Cesium
fluoride
cesium fluoride
Name the following compounds.
1.
KCl
2.
MgBr2
3.
Al2S3
potassium chloride
24
8
Naming Binary Ionic Compounds for
Metals with Variable Charge
1. Metal cation name is metal name
followed by a Roman numeral in
parentheses to indicate its charge.
9 determine charge from anion charge
9 common ions Table 3.4
2. nonmetal anion named by changing the
ending on the nonmetal name to -ide
25
Find the charge on the cation.
1.
TiCl4
2.
CrO3
4 Cl = 4−,
∴ Ti = 4+
26
Clicker question #3-3
Find the charge on Fe in the compound Fe3N2
1.
2.
3.
4.
5.
+1
+2
+3
-3
+6
27
9
Example—Naming Binary Ionic
with Variable Charge Metal CuF2
1. Identify cation and anion.
F = F− because it is Group 7
Cu = Cu2+ to balance the two (−) charges from 2 F−
2. Name the cation.
Cu2+ = copper(II)
3. Name the anion.
F− = fluoride
4. Write the cation name first, then the anion name.
copper(II)
Copper (II)fluoride
fluoride
28
Name the following compounds.
1.
TiCl4
2.
PbBr2
3.
Fe2S3
titanium(IV) chloride
29
Practice—What are the formulas for
compounds made from the following ions?
copper(II) ion with a nitride ion
iron(III) ion with a bromide ion
30
10
Compounds Containing
Polyatomic Ions
• Polyatomic ions are single ions that contain
more than one atom.
• Name and charge of polyatomic ion need
to be learned !
31
Some Common Polyatomic Ions
Name
Formula Name
Formula
acetate
carbonate
hydrogen carbonate
(aka bicarbonate)
hydroxide
nitrate
nitrite
chromate
dichromate
ammonium
C2H3O2–
CO32–
ClO–
ClO2–
ClO3–
ClO4–
SO42–
SO32–
HCO3–
OH–
NO3–
NO2–
CrO42–
Cr2O72–
NH4+
hypochlorite
chlorite
chlorate
perchlorate
sulfate
sulfite
hydrogen sulfate
(aka bisulfate)
hydrogen sulfite
(aka bisulfite)
phosphate
HSO4–
HSO3–
PO43-
32
Special Pattern for Oxygenated
halogens
• -ate ion
9chlorate = ClO3–
• -ate ion + one O ⇒ same charge, per- prefix
9perchlorate = ClO4–
• -ate ion – one O ⇒ same charge, -ite suffix
9chlorite = ClO2–
• -ate ion – two O ⇒ same charge, hypoprefix, -ite suffix
9hypochlorite = ClO–
33
11
Example—Naming Ionic Compounds
Containing a Polyatomic Ion Fe(NO3)3
1. Identify the ions.
2.
3.
4.
NO3 = NO3− a polyatomic ion
Fe = Fe3+ to balance the charge of the 3 NO3−
Name the cation.
Fe3+ = iron(III), metal with variable charge
Name the anion.
NO3− = nitrate
Write the name of the cation followed by the
name of the anion.
iron(III)
Iron
(III) nitrate
34
Name the following compounds
1.
NH4Cl
2.
Ca(ClO2)2
3.
Cu(NO3)2
NH4+ Cl-
ammonium chloride
35
Practice—What are the formulas for
compounds made from the following ions?
aluminum ion with a sulfate ion
Al3+ with SO42−
Al2(SO4)3
chromium(II) with hydrogen carbonate
36
12
Clicker question #3-4
All the following ions have the same charge ?
EXCEPT
1) Sulfide ion
2) Bromate ion
3) Oxide ion
4) Sulfate ion
5) monohydrogen phosphate ion
37
Hydrates
• Hydrates are ionic compounds
•
containing a specific number of waters
for each formula unit.
in formula, attached waters follow
9 CoCl2·6H2O
• in name, attached waters indicated by
prefix+hydrate after name of ionic
compound
9 CoCl2·6H2O = cobalt(II) chloride
hexahydrate
9 CaSO4·½H2O = calcium sulfate
hemihydrate
Prefix
No. of
Waters
hemi
½
mono
1
di
2
tri
3
tetra
4
penta
5
hexa
6
hepta
7
octa
8
38
Writing Names of Binary Molecular
Compounds (nonmetal + nonmetal)
1. Write the name of the first element in the
formula.
9 element furthest left and down on the Periodic Table
9 Use the full name of the element,
2. Write the name of the second element in the
formula with an -ide suffix.
3. Use a prefix in front of each name to indicate the
number of atoms.
a) Never use the prefix mono- on the first element.
39
13
Name the following.
NO2
nitrogen dioxide
PCl5
I2F7
40
Write formulas for the following.
dinitrogen tetroxide
N2O4
sulfur hexafluoride
diarsenic trisulfide
41
Acids
• Acids are molecular compounds that form H+
when dissolved in water.
9 To indicate the compound is dissolved in water, (aq) is
written after the formula. HCl (aq)
¾ NOTE: not
named as acid if not dissolved in water
• dissolve many metals
9 like Zn, Fe, Mg; but not Au, Ag, Pt
• formula generally starts with H
9 e.g., HBr (aq), H2SO4 (aq)
42
14
Acids
• contain H+ cation and
anion
9in aqueous solution
• Binary acids have H+
cation and nonmetal
anion.
or
HCl (aq)
H2SO4(aq)
•Oxyacids have H+ cation
and polyatomic anion.
43
Naming Binary Acids
•
•
•
•
Write a hydro prefix.
Follow with the nonmetal name.
Change ending on nonmetal name to -ic.
Write the word acid at the end of the name.
44
Example—Naming Binary
Acids: HCl(aq)
1. Identify the anion.
Cl = Cl−, chloride because Group 7A
2. Name the anion with an -ic suffix.
Cl− = chloride ⇒ chloric
3. Add a hydro- prefix to the anion
name.
hydrochloric
4. Add the word acid to the end.
hydrochloric acid
45
15
Naming Oxyacids
• If the polyatomic ion name ends in -ate,
then change ending to -ic suffix.
• If the polyatomic ion name ends in -ite, then
change ending to -ous suffix.
• Write word acid at the end of all names.
46
Example—Naming Oxyacids
H2SO4(aq)
1. Identify the anion.
SO4 = SO42− = sulfate
2. If the anion has -ate suffix, change it to -ic.
If the anion has -ite suffix, change it to -ous.
SO42− = sulfate ⇒ sulfuric
3. Write the name of the anion followed by the
word acid.
sulfuric
Sulfuricacid
acid
47
Name the following
H2S (aq)
hydrosulfuric acid
HClO3 (aq)
HNO2 (aq)
48
16
Clicker question #3-5
Choose the name-formula pair that does NOT
match.
1) ammonium hydrogen carbonate, (NH4)2CO3
2) sodium chlorite, NaClO2
3) calcium hydride, CaH2
4) nitric acid, HNO3
5) calcium hydroxide, Ca(OH)2
49
Writing Formulas for Acids
• when name ends in acid, formula starts with H
• Write formulas as if ionic, even though it is
molecular.
• Hydro prefix means it is a binary acid, no
hydro prefix means it is an oxyacid.
• for oxyacid, if ending is -ic, polyatomic ion
ends in -ate;
– if ending is -ous, polyatomic ion ends
in -ite
50
Practice—What are the formulas for the
following acids?
chlorous acid
H+ with ClO2–
HClO2
phosphoric acid
51
17
Formula Mass
• is the mass of an individual molecule or
formula unit
9also known as molecular mass or molecular
weight
• sum of the masses of the atoms in a single
molecule or formula unit
mass of 1 molecule of H2O
= 2(1.01 amu H) + 16.00 amu O = 18.02 amu
52
Molar Mass of Compounds
• The masses of molecules can be calculated from
atomic masses.
• since 1 mole of H2O contains 2 moles of H and 1
mole of O
Molar Mass = 1 mole H2O = 2(1.01 g H) +16.00 g O = 18.02 g
Molar Mass of H2O is 18.02 g/mole
NOTE!
You will need to more precise molar masses during
tests and assignments.
A more precise molar mass of water uses the
exact atomic weight from the peroidic table.
Molar Mass = 2(1.0079 g H) +15.9999 g O = 18.0148 g/mole
of H2O
53
Practice—How many moles are in 50.0 g of
PbO2?
Given: 50.0 g PbO2
Find: moles PbO2
Conceptual
Plan:
g PbO2
mol PbO2
Relationships: 1 mol PbO2 = 239.2 g
Solution:
Check: Since the given amount is less than 239.2 g,
the moles being <1 makes sense.
54
18
Example Find the number of CO2
molecules in 10.8 g of dry ice.
Given: 10.8 g CO2
Find: molecules CO2
g CO2
Conceptual
Plan:
Relationships:
mol CO2
molec CO2
1 mol CO2 = 44.01 g,
1 mol = 6.022 × 1023
Solution:
Check:
Since the given amount is much less than
1 mol CO2, the number makes sense.
55
Clicker question #3-6
How many molecules are in 1.00 g of mannose,
C 6H 12 O 6, which is a sweet tasting sugar that has a
bitter aftertaste?
1) 2.18 x 10 18
2) 3.34 x 10 18
3) 2.18 X 10 21
4) 3.34 x 10 21
5) 3.34 x 10 24
56
Percent Composition
• percentage of each element in a compound
9 by mass
• can be determined from
1. the formula of the compound
2. the experimental mass analysis of the
compound
57
19
Example 3.13
Find the mass
percent of Cl in C2Cl4F2.
Given:
Find:
C2Cl4F2
% Cl by mass
Conceptual
Plan:
Relationships:
Solution:
Check:
Since the percentage is less than 100 and Cl is much
heavier than the other atoms, the number makes sense.
58
Practice—Determine the mass percent
composition of the CaCl2
59
Clicker question #3-7
The mineral leadhillite, which is essentially
Pb4(SO4)(CO3)2(OH)2 (FW = 1079 amu), contains
____ % oxygen by weight.
1) 10.4
2) 11.9
3) 13.3
4) 14.8
5) 17.8
60
20
Empirical Formula
• simplest, whole-number ratio of the atoms
of elements in a compound
• can be determined from elemental analysis
by:
method 1.
percent composition
or
Method 2.
combustion analysis i.e., masses
of elements formed from decomposition or
reaction
61
1. Empirical Formula from % mass
1) Convert the percentages to grams.
- assume you start with 100.00 g of the compound
2) Convert grams to moles.
- use molar mass of each element
3) Write a pseudo-formula using moles as subscripts
4) Divide all by smallest number of moles.
- If result is within 0.1 of whole number, round to whole
number.
5) If required, multiply all mole ratios by number to
make all whole numbers.
- if ratio 0.5, multiply all by 2; if ratio 0.33 or 0.67, multiply
all by 3; if ratio 0.25 or 0.75, multiply all by 4; etc.
62
Example 3.16
• Laboratory analysis of aspirin determined
the following mass percent composition.
Find the empirical formula.
C = 60.00%
H = 4.48%
O = 35.53%
63
21
Strategy
The empirical formula is CxHyOz
and found by:
ggCC
molCC
mol
ggHH
molHH
mol
O
ggO
molO
O
mol
mole
pseudo- ratio
pseudoformula
formula
whole
number
ratio
empirical
empirical
formula
formula
Masses found in step #1 (next slide)
Step 1 Assume 100 g of sample; therefore
there are 60.00 g C, 4.48 g H, and 35.53 g O.
Step 2 Find moles
Step 3
write a pseudoformula
C4.996H4.44O2.220
find the mole ratio by dividing by the
smallest number of moles
Step 4 multiply subscripts by factor to give
whole number
{C2.25H2O1} x 4 =C9H8O4
22
Clicker question #3-8
Analysis of a hydrocarbon showed that it contained
14.4% hydrogen and 85.6% carbon by weight. What
is its simplest formula?
1) CH
2) CH 2
3) CH 3
4) C 2H 3
5) C 2H 5
67
Molecular Formulas
•
•
The molecular formula is a multiple of the
empirical formula.
To determine the molecular formula, you
need to know the empirical formula and
the molar mass of the compound.
68
68
Example 3.17 Find the molecular formula of
butanedione given the emp. formula = C2H3O
and the molar mass is 86.09g /mole
Check: The molar mass of the calculated formula is in
agreement with the given molar mass.
69
23
Empirical formula from
Combustion Analysis
• A common technique for analyzing compounds is to
burn a known mass of compound and weigh the
amounts of product made.
9 generally used for organic compounds containing C, H, O
• By knowing the mass of the product and
composition of constituent elements in the product,
the original amount of constituent elements can be
determined.
9 All the original C forms CO2, the original H forms H2O,
the original mass of O is found by subtraction.
• Once the masses of all the constituent elements in
the original compound have been determined, the
empirical formula can be found.
70
71
Example 3.19
• Combustion of a 0.8233-g sample of a
compound containing only carbon, hydrogen,
and oxygen produced the following:
CO2 = 2.445 g
H2O = 0.6003 g
Determine the empirical formula of the
compound.
72
24
Write a conceptual plan for example 3.19
g
CO2, H2O
mol
C, H, O
mol
CO2, H2O
mol
C, H
pseudo
formula
g
C, H
mol
ratio
g
O
mol
O
empirical
formula
Example 3.19 (continued)
Step 1 Calculate the moles of C and H.
Step 2 Calculate the grams of C and H
Step 3 Calculate the grams and moles of O.
Step 4 Write a pseudoformula using moles
found of each element
C0.05556H0.06662O0.00556
divide by the smallest number of moles.
75
25
Practice 1
A 47.95 mg sample of a compound containing only
C and H was combusted and produced 145.21 mg of
CO2 and 74.33 mg of H2O.
(a) What is the empirical formula of this
hydrocarbon compound?
(b) The molar mass of the hydrocarbon was found
to be 87.18 g mol-1. What is its molecular formula?
Answer is: a) C2H5
b) C6H15
76
Clicker question #3-9
A given hydrocarbon is converted completely to
carbon dioxide and water, and equal numbers of
moles of CO2 and H2O are produced. The
hydrocarbon could be
1) CH4.
2) C2H6.
3) C3H8.
4) C3H4.
5) C4H8.
77
Chemical Reactions
• Reactions involve chemical changes in matter
resulting in new substances.
• Reactions involve rearrangement and
exchange of atoms to produce new
molecules.
Reactants
→
Products
78
78
26
Chemical Equations
• shorthand way of describing a reaction
• provides information about the reaction
9formulas of reactants and products
9states of reactants and products
9relative numbers of reactant and product
molecules that are required
9can be used to determine weights of reactants
used and products that can be made
79
Combustion of Methane
CH4(g) + O2(g) → CO2(g) + H2O(g) (unbalanced)
• To show the reaction obeys the Law of Conservation
of Mass, the equation must be balanced.
9 We adjust the numbers of molecules so there are equal
numbers of atoms of each element on both sides of the
arrow.
CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g)
Check equal # of elements
1C + 4H + 4O
1C + 4H + 4O
80
Balancing Chemical Equations
CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g)
• CH4 and O2 are the reactants, and CO2 and H2O
are the products.
• The (g) after the formulas tells us the state of
the chemical.
• The number in front of each substance tells us
the number of those molecules in the reaction.
9 called the coefficients
81
27
Symbols Used in Equations
• symbols used to indicate state after
chemical
9(g) = gas; (l) = liquid; (s) = solid
9(aq) = aqueous = dissolved in water
• energy symbols used above the arrow for
decomposition reactions
9 Δ = heat
9 hν = light
ΔH (chapter 6)
82
Example 3.21 Write a balanced equation for
the combustion of butane, C4H10.
Write a skeletal
equation:
Balance atoms in
complex substances
first:
Balance free elements
by adjusting coefficient
in front of free element:
C4H10(l) + O2(g) → CO2(g) + H2O(g)
4⇐C⇒1×4
C4H10(l) + O2(g) → 4 CO2(g) + H2O(g)
10 ⇐ H ⇒ 2 × 5
C4H10(l) + O2(g) → 4 CO2(g) + 5 H2O(g)
13/2 × 2 ⇐ O ⇒ 13
C4H10(l) + 13/2 O2(g) → 4 CO2(g) + 5 H2O(g)
If fractional coefficients, {C4H10(l) + 13/2 O2(g) → 4 CO2(g) + 5 H2O(g)} × 2
multiply thru by
2 C4H10(l) + 13 O2(g) → 8 CO2(g) + 10 H2O(g)
denominator:
Check:
8 ⇐ C ⇒ 8; 20 ⇐ H ⇒ 20; 26 ⇐ O ⇒ 26
83
Practice 1
When aluminum metal reacts with air, it
produces a white, powdery compound—
aluminum oxide, Al2O3.
Write a balanced equation including proper
physical states.
84
28
Clicker question #3-10
Which of the following equations is (are) balanced?
→ PbCl2 + NaNO 3
a. NaCl + Pb(NO 3)2 ⎯ ⎯
→ N 2 + 4H 2O + Cr2O 3
b. (NH 4)2Cr2O 7 ⎯ ⎯
→ 3Fe + 3CO 2
c. Fe3O 4 + 3CO ⎯ ⎯
1) a only
2) b only
3) c only
4) a and b only
5) a, b, and c
85
29