Fundamentals
The Atom in Modern Chemistry
1
You will learn to look at chemistry from three
points of view
For example, when methane burns:
Macroscopic view
Particulate view
Chemical equation view
CH4 (g) + 2O2 (g) → 2H2O (g) + CO2 (g)
2
Macroscopic View
Chemistry is concerned with matter: its forms and
transformations!
Classification schemes for matter:
Properties – extensive / intensive
Changes – physical / chemical
Composition – pure substances / mixtures
3
Macroscopic View
PROPERTIES:
Extensive –
Depend on the amount of matter
Cannot be used to identify a substance
Examples:
Mass, volume, length, energy, electrical resistance
Intensive –
Inherent to the material
Do not depend on amount of matter
Can be used to identify a substance
Examples:
Density, boiling point, melting point, color, viscosity, flammability
Physical – Measured without changing substance’s identity, are either
extensive or intensive
Chemical – Only measured through the course of a chemical reaction,
intensive only
4
Macroscopic View
Physical Changes
SOLID
FREEZING
MELTING
LIQUID
EV
AP
OR
CO
AT
ND
IO
EN
N
SA
TIO
N
N
O
I
T
SI
O
P
DE
ON
I
T
MA
I
L
SUB
Transformations between
solid, liquid, and gas are
called “phase changes”
GAS
Physical Properties
Melting Point (ºC)
Boiling Point (ºC)
O2
-219
-183
ethanol
-114
78
Br2
-7
59
H2O
0
100
Cu
1085
2562
Macroscopic View
CHEMICAL CHANGES: A new substance is formed
4 Indicators of Chemical Change
- Heat or light is given off or taken in
- Color change (a new color)
- A precipitate is formed
- Gas is given off (not due to change of state)
http://beautifulchemistry.net/reactions.html
6
Macroscopic View
What do you think?
1.
Boiling of water
a)
Physical change
b)
Chemical change
2.
Burning of paper
a)
b)
Physical change
Chemical change
7
Macroscopic View
COMPOSITION
MATTER
MIXTURE
PURE
Substance
ELEMENT
HETEROGENEOUS
Chemical
change
HOMOGENEOUS
COMPOUND
GROSS
MIXTURE
SUSPENSION
COLLOID
SOLUTION
8
Macroscopic View
Mixtures: >2 elements / compounds mixed together
- Can be separated by physical properties.
- Dissolving in polar/nonpolar solvent, magnet, Gravity
(centrifuge),
b.p. (distillation), m.p., filtering
Example: iron filings and sulfur
Suspensions:
- Particles > 1 μm diameter, settle out more rapidly,
- Solute can be filtered out
Example: Muddy water
Colloids:
-- particles 1 nm – 1 μm diameter, settle out VERY slowly,
solute cannot be filtered
Examples: dust, fog, jello
9
Macroscopic View
Solutions:
- Particles < 1 nm (atoms, ions, molecules)
- Never settle out
- Single phase
- Solute cannot be filtered out
Examples: Saltwater, soda, Kool-Aid
Solvent -
> 50 %
Solute -
< 50 %
Mixtures of two or more gases are ALWAYS homogeneous and
ALWAYS solutions.
10
Macroscopic View
Alloy: A homogeneous mixture (solution) of two or more elements, at
least one of which is a metal, and where the resulting material has
metallic properties.
Steel:
99.5% Fe and 0.5 - 1.5% C
Stainless steel: 87 - 74% Fe and 13-26% Cr
Brass:
80% Cu and 20% Zn
Bronze:
90% Cu and 10% Sn
Solder:
50% Sn and 50% Pb
Sterling silver: 92.5% pure silver and 7.5% other metals, usually Cu
11
Macroscopic View
MATTER
MIXTURE
PURE
Substance
ELEMENT
Chemical
change
HETEROGENEOUS
HOMOGENEOUS
COMPOUND
GROSS
MIXTURE
SUSPENSION
COLLOID
SOLUTION
12
Macroscopic View
COMPOUNDS
- can be broken down into smaller constituent
parts through chemical reactions (chemical change).
- more than one atom connected together
ELEMENTS
- cannot be broken down into smaller constituent
parts through chemical reactions (chemical change).
- unique atoms
-- atoms are not connected
-- except for the diatomic seven
N, O, F, Cl, Br, I, and H
Atoms = Particulate view
13
Macroscopic à Particulate View
SOLID
I
Mixtures?
Elements (atoms)?
LIQUID
II
GAS
III
Pure Substances?
Compounds (molecules)?
IV
14
Macroscopic à Particulate View
Some elements have similar macroscopic properties.
Arranged in the Periodic Table during 1800s according to observed
chemical properties.
Chemical Reactivity
Ionization Energy
Electronegativity
Number of Valence Electrons
Oxidation Number
15
http://www.ptable.com
16
Macroscopic à Particulate View
Atom:
The smallest particle of an element that
retains the chemical identity of that element.
Democritus
460 - 370 B.C.
Greece
“According to convention there is a sweet and a bitter,
a hot and a cold, and according to convention there
is order. In truth there are atoms and a void.”
-- Democritus, 400 B.C.
Macroscopic View
1789:
Law of conservation of matter (mass):
Matter is neither created nor destroyed in
any process.
Lavoisier, Antoine
1743-1794
France
Particulate View
1808 - Atomic Theory:
1. Each element is made of extremely small, indivisible
particles called atoms. (Democritus)
John Dalton
1766 - 1844
England
2. Atoms of one element are the same (properties), but differ
from other elements.
3. Atoms can physically mix together or chemically combine
in simple whole number ratios to form compounds. (Proust)
4. Atoms are neither created nor destroyed in any chemical
reaction. (Lavoisier)
5. A given compound always has the same relative numbers
and kinds of atoms. (Proust)
Particulate View
Law of Combining volumes
The ratio of the volumes of any pair of gases in a gas
phase chemical reaction (at the same temperature and
pressure) is the ratio of simple integers
Joseph Louis Gay-Lussac
1778 – 1850
France
Example:
2 volumes of hydrogen + 1 volume of oxygen give 2 volumes of water vapor
1 volume of nitrogen + 3 volumes of hydrogen give 2 volumes of ammonia
20
Particulate View
1811 - Avogadro’s Hypothesis
Equal volumes of different gases (at the same T, P)
contain equal numbers of molecules.
Amedeo Avogadro
1776 – 1856
Italy
H2
H2O
Ar
21
Particulate View
1839
Proposed that atomic structure is related to
electricity.
Ben Franklin had earlier established that there
are two types of electric charge: positive and
negative.
Michael Faraday
1791 - 1867
England
Particulate à Chem. Eq. View
Compounds are made of different elements in fixed, small whole
number ratios.
Molecules are described by a molecular formula with atoms in
fixed, small whole number ratios.
H2O
CO
N2H4
but not CO1.7
CO2
23
Particulate à Chem. Eq. View
A sample of liquid N2H4 (hydrazine) is decomposed to give gaseous
N2 and gaseous H2.
N2H4 (l) → N2(g) + 2 H2(g)
Question: The two product gases are separated and the nitrogen
occupies 15 ml at room temperature and atmospheric pressure. What
volume of hydrogen will be produced under these same conditions?
A) 15 ml
B) 30 ml
C) 7.5 ml
D) ????
24
Chemical Equation View
Forensic scientists know that the mass ratio of C:H is 10.88:1 in heroin.
Question: When forensic scientists chemically analyze a sample of
white powder and discover that there are 2.00 grams of hydrogen
atoms, how many grams of carbon should they expect to find if it is
heroin?
A) 2.00 g
B) 10.88 g
C) 21.76 g
D) ????
25
What are the three parts of an atom?
Helium
Proton
+
+
Neutron
Electron
Not drawn
to scale
26
Particulate View
Relative sizes/masses
Nucleus
1 mm (barely larger than the
period at the end of a
sentence.
Electron
Mass is 1/1836 the mass of a
proton
Atom – 100 yd stadium
27
1
Atomic Number (Z) = # of protons
H
1.008
3
4
6
7
8
9
Li
Be
C
N
O
F
6.941
9.012
11
12
12.011
14
Na
Mg
Si
22.990 24.305
14.007 15.999 18.998
15
16
17
P
28.086 30.974
S
Cl
32.06
35.453
Protons each have a charge of +1
The nuclear charge (charge of the nucleus) is always the same as the
number of protons.
What’s the nuclear charge on:
H? C?
Ca?
U?
Chemical Equation View
In an electrically neutral atom:
1
# of electrons = # of protons
H
1.008
3
4
6
7
8
9
Li
Be
C
N
O
F
6.941
11
9.012
12
12.011
14
Na
Mg
Si
22.990 24.305
14.007 15.999 18.998
15
16
17
P
28.086 30.974
S
Cl
32.06
35.453
Chemical Equation View
Ions have different numbers of protons and electrons
Ions have positive or negative charge.
If there are more electrons than protons, they have a net negative
charge. (Anions)
If there are more protons than electrons, they have a net positive
charge. (Cations)
Macroscopic matter is net neutral. A negatively charged ion must be
paired with a positively charged ion.
30
Chemical Equation View
Ion Symbols
The chemical symbol (letter(s)) is determined by the number of protons.
For example, the fluoride ion can be written as FThe F indicates that it has 9 protons
The – indicates that it has one more electron than protons, so it
has 10 electrons.
Question: How many electrons are in Mg2+?
31
Chemical Equation View
What is the symbol for oxygen with one extra electron:
O-
Oxygen with two extra electrons:
O2-
Beryllium with one less electron:
Be+
Beryllium with two fewer electrons:
Be2+
Chemical Equation View
What can change with Chlorine, so it’s still Chlorine
and it’s still electrically neutral?
- Number of neutrons!
Isotope: Two or more forms of an element that have identical numbers
of protons but differ in the number of neutrons (and hence masses) but
have identical chemical properties. Also can be considered as a
radioactive form of the element.
Question: What changes when more neutrons are added to an atom?
Size or mass?
Chemical Equation View
How much does an atom weigh?
Proton =
Neutron =
Electron =
1.672649 x 10-24 grams
1.674954 x 10-24 grams
0.00091095 x 10-24 grams
Chemical Equation View
Atomic mass is measured relative to 1/12 the mass of 12C.
The mass of a single atom of 12C is set at exactly 12 u where “u” is an
atomic mass unit (amu).
This is also called 12 Da
The mass of other atoms can be measured relative to 12C using a mass
spectrometer.
35
Chemical Equation View
To facilitate the specification of atomic isotopes we
use the mass number which is nuclear particle count
Mass Number
Proton =
Neutron =
Electron =
1
1
0
Mass Number = # of protons + # of neutrons
Chemical Equation View
What is the Mass Number for Helium?
Protons = 2
Neutrons = 2
Electrons = 2
Mass Number = 2 + 2 = 4
What is the Mass Number for Fluorine?
Protons = 9
Neutrons = 10
Electrons = 9
Mass Number = 9 + 10 = 19
Chemical Equation View
What is the Mass Number for these Chlorine isotopes?
A) Protons = 17
Neutrons = 18 Electrons = 17
B) Protons = 17
Neutrons = 20 Electrons = 17
Mass Numbers = 35 and 37
Isotope names:
Isotope names:
Chlorine-35
Carbon-12
Chlorine-37
Carbon-13
Carbon-14
Clicker Question
How many neutrons are in 1000
atoms of Cl-37?
A.
B.
C.
D.
E.
37
37,000
20
20,000
cannot be determined
Chemical Equation View
An alternate way : Element – Mass number
What is the symbol for Helium-4 ?
Protons = 2
Neutrons = 2
Electrons = 2
Element symbol
Mass Number =
Number of
protons + neutrons
4
2
Number of protons
He
Chemical Equation View
How many neutrons are in Nitrogen-15?
Neutrons = 15 - 7 = 8
How many neutrons are in Magnesium-26?
Neutrons = 26 - 12 = 14
Chemical Equation View
“Valley of Stability”
1.4
N / P ratio
1.3
1.2
1.1
Chemical Equation View
Which is more stable?
108
47
Ag
(108 – 47) / 47 = 1.3
125
47
Ag
(125 – 47) / 47 = 1.7
Chemical Equation View
Chlorine has two isotopes:
A) Chlorine-35
34.969 amu
B) Chlorine-37
36.966 amu
Natural abundance = 75.77 % 24.23 %
Average Atomic Mass = (34.969 + 36.966)/2 = 35.968 amu
Weighted Average Atomic Mass =
(75.77% x 34.969 amu) + (24.23% x 36.966) = 35.453 amu
n
M = M 1 p1 + M 2 p2 + ... + M n pn = å M i pi
i =1
Mn is the relative atomic mass of each of the n isotopes and pn is their
fractional abundance.
Chemical Equation View
The masses on the Periodic Table have no units – they are relative to
1/12 the mass of 12C
1
Average Atomic Mass
H
1.008
3
4
6
7
8
9
Li
Be
C
N
O
F
6.941
11
9.012
12
12.011
14
Na
Mg
Si
22.990 24.305
14.007 15.999 18.998
15
16
17
P
28.086 30.974
S
Cl
32.06
35.453
Chemical Equation View
Naturally occurring boron (B) consists of two isotopes: 10B (atomic
mass 10.013) and 11B (atomic mass 11.009) . The atomic mass of
the isotope mixture found in nature is 10.811.
WITHOUT doing the calculation – which isotope is more abundant.
A) 10B
B) 11B
Practice – Ga-69 with mass 68.9256 amu has an abundance of 60.11% and Ga-71
with mass 70.9247 amu has an abundance of 39.89%. Calculate the atomic mass of
gallium.
Given:
Ga-69 = 60.11%, 68.9256 amu
Ga-71 = 39.89%, 70.9247 amu
Find:
atomic mass, amu
Conceptual Plan:
Relationships:
isotope masses,
isotope fractions
avg. atomic mass
Solution:
Check:
the average is between the two masses,
closer to the major isotope
Chemical Equation View
Molecular mass
Molecular mass (relative) is determined by adding the atomic masses of
all the atoms in the molecule (note: the atomics masses used are the
weighted averages listed in the periodic table!)
For example,
the molecular mass of water, H2O, is
2(H) + 1(O) = 2(1.00797) + 15.9994 = 18.0153
48
Weighted average atomic masses
What is the weighted Average Atomic Mass of a sample
of Krypton that is composed of 3 isotopes:
10% krypton-79
50% krypton-85
40% krypton-81
(.10 x 79) + (.50 x 85) + (.40 x 81) =
7.9
+
42.5
+
32.4
= 82.8 amu
49
Clicker Question
The mass spectrum of gallium, Ga, is shown below. The atomic
mass of Ga is 69.7 amu. Which of the following statements is
correct?
Relative Abundance (%)
Gallium Mass Spectrum
70
a)
68.9 amu, 60.1 %
60
50
70.9 amu, 39.9 %
40
30
b)
c)
20
10
0
67
68
69
70
Mass (amu)
71
72
All Ga atoms weigh 69.7
amu.
The atomic mass of Ga is
the average of 68.9 and
70.9.
The atomic mass of Ga will
be closer to 69 than 71
because there are more
atoms that weigh 68.9 amu.
Precision and Accuracy
SI Units
Dimensional Analysis – unit conversions
Moles – Avogadro’s number
51
Precision and Accuracy
Precise &
Accurate
Precise
Accurate
Neither
precise nor
accurate 52
SI units
SI units definitions
http://www.npl.co.uk/upload/pdf/units-of-measurement-poster.pdf
53
Common SI prefixes
SI System, you’ll find them in the back binder
Mega
kilo
deci
centi
milli
micro
nano
pico
For Example
M = 106
k = 103
d = 10-1
c = 10-2
m = 10-3
µ = 10-6
n = 10-9
p = 10-12
1 Megawatt (wind turbine)
kilometer (distance)
milliliter, syringes
micrometer, cell diameter
nanometer, Chemical
bond
picometer, atom diameter
Angstrom 1 Å = 10-10 meters
54
Dimensional Analysis
The method uses conversion factors. Conversion
factors must be written with a numerator and
denominator.
Example: Suppose we want to convert a volume of 1.7 qt into
liters.
56
Dimensional Analysis
The method uses conversion factors. Conversion
factors must be written with a numerator and
denominator.
Example: Suppose we want to convert a volume of 1.7 qt into
liters.
1.7 qt = ?? L
where:
1 qt = 0.946 L
57
Conversion Factors - Density
The density of a chemical is the ratio of its mass
to its volume.
density = mass/volume
d = ρ = m/V
- Mass and volume are extensive properties.
- Density is an intensive property
- Density of water (at 4°C ) is 1.000 g/mL
- Units are usually written as
• g/ml for liquids
• g/cm3 for solids (1ml = 1cm3)
• g/ml or g/L for gases
58
Conversion Factors - Density
Which one weighs more (has more mass)?
Which one floats? Why?
Density increases as temperature decreases (generally)
Density increases as pressure increases
Conversion Factors - Density
Density changes as temperature changes
-- generally as T the density
Conversion Factors - Density
• If the density of acetone is 0.7925 g/mL, what is the
mass of 2 liters of the acetone?
A) 792.5g B) 1585g C) 1.585g D) ????
61
The Mole – Avogadro’s Number
Moles
Avogadro’s number is the number of atoms in 12.000 g of
12C
and has the symbol NA.
NA = 6.022 141 3 + 0.000 000 3 x 1023
There are 6.0221413 x 1023 molecules in a mole.
The relative masses on the periodic table are also the molar
mass (grams per mole) for the atoms.
1 mole of different substances
63
The Mole – Avogadro’s Number
Currently, best experiments use X-rays to
measure the distance between atoms in a crystal
Ti crystal = 2 atoms per unit cell
Edge length = 330.6 pm.
Ti density = 4.401 g/cm3
Ti atomic mass = 47.88 g/mol
The Mole – Avogadro’s Number
How many atoms are in 3.6 moles of aluminum foil?
How many moles are in 3.02 x 1026 molecules of
water?
Molar Mass
1 mole of magnesium has a mass of 24.305 g
1 mole of carbon has a mass of 12.011 g
2 moles of carbon have a mass of
24.022 g
1
H
1.008
3
Li
6.941
4
Be
9.012
6
C
12.011
7
N
14.007
8
O
15.999
9
F
18.998
11
Na
22.990
12
Mg
24.305
14
Si
28.086
15
P
30.974
16
S
32.06
17
Cl
35.453
Molar Mass
1
H
1.008
3
Li
6.941
4
Be
9.012
6
C
12.011
7
N
14.007
8
O
15.999
9
F
18.998
11
Na
22.990
12
Mg
24.305
14
Si
28.086
15
P
30.974
16
S
32.06
17
Cl
35.453
1 mole of NaCl has a Molar Mass of _____?
What is the mass of 3.57 moles of NaCl?
How many moles of NaCl in 150.2 grams?
Clicker question
Which has more atoms, 10.0 g Mg or 10.0 g Ca?
A. Magnesium
B. Calcium
C. Both have the same number
of atoms.
Conversion Factors Summary
- To convert between moles and particles, use Avogadro’s
Number
1 mole = 6.02 x 1023 particles
- To convert between moles and mass, use the Molar Mass
(needs to be calculated from Periodic Table)
1 mole Na = 22.99 g
- To convert between mass and particles, use both NA and
Molar Mass (two conversion factors)
particles
Avogadro’s
Number
moles
Molar
Mass
mass
Clicker
Which sample represents the greatest number of
moles?
1. 44.01 g CO2
2. 1.0 moles C3H8
3. 6.022 x 1023 molecules C4H10
4. 18.02 g H2O
5. All of the samples have
the same number of moles.
Which of the following has the largest mass?
a)
b)
c)
d)
e)
10.0 g Li
10.0 moles of Li
100 g Na
10.0 moles of K
100 g Rb
Clicker - Conversion Factors
Iron (Fe) is biologically important in the transport of oxygen by red blood
cells. In the blood of an adult human, there are approximately 2.60x1013
red blood cells and a total of 2.90 g of iron atoms. On an average, how
many iron atoms are present in each red blood cell.
HINT: Use dimensional analysis.
A) 8.97x1012 B) 3.13x1022 C)1.20x109
D) ????
72
Compounds
Nomenclature of Compounds
Sections C and D in Fundamentals
73
Compounds
Compounds are electrically neutral consisting of two or more
different elements.
Compounds are classified as either organic or inorganic.
Organic compounds contain carbon and usually hydrogen.
There are millions of organic compounds, including fuels,
sugars, and most medicines.
Inorganic compounds are all the other compounds; they
include water, calcium, ammonia, silica, hydrochloric acid, and
many more.
74
The forces holding ionic solids or molecules together
Ionic solids comprise of
positively and negatively
charged atoms. These are not
discrete ions, instead a vast
sea of charges holds the ions
together in a crystal.
NaCl crystal
Molecules are discrete groups of atoms
connected by neutral bonds composed
of pairs of electrons
Dimethylamine
75
Positively charged ions are called cations, and negatively
charged ions are called anions.
cation
Na+ sodium cation
Ca2+ calcium cation
An example of a "polyatomic" (many-atom) cation is the
ammonium ion, NH4+
anion
A negatively charged chlorine atom is an anion and is denoted
CI-.
An example of a polyatomic anion is the carbonate ion, CO32-
76
Metals typically form cations
Nonmetals typically form anions.
cations
A pattern observed is that metallic elements typically form
cations by electron loss.
77
Metals typically form cations
Nonmetals typically form anions.
anions
Another pattern observed is
nonmetallic elements typically form
anions by gaining electrons.
78
Classifying compounds by the kind of element
Two nonmetals are molecular: Water (H2O) is
an example of a binary molecular compound
nonmetal-nonmetal
A metal and a nonmetal are ionic: sodium
chloride (NaCI) is an example of an ionic
compound.
cation anion
79
Polyatomic ions consist of three or more atoms. The ions
can have either a positive or negative charge.
Cyanide ion, CN- , is diatomic, and the ammonium ion,
NH4+,is polyatomic.
Common polyatomic anions with oxygen are called
oxoanions like
carbonate, CO32nitrate NO3phosphate, PO43sulfate, SO42Review tables D.1 through D.3 (in Fundamentals)
80
Give the name for:
AuCI3;
ionic, group 11 (variable
oxidation)
gold (III) chloride
CaS;
ionic group 2 (fixed oxidation
number)
calcium sulfide
Mn2O3;
ionic group 7 (variable oxidation manganese (III)
number)
oxide
HCN(aq) molecular acid ( ends in ide→
hydro_acid name_ic)
HOCN
IF5;
Molecular acid (ends in ate→ ic
acid)
molecular
hydrocyanic acid
cyanic acid
iodine pentafluoride
82
Formulas for ionic compounds have a different
meaning from those of molecular compounds.
Ionic compounds form
crystals due to the vast sea
of charges between anions
and cations.
Formal units are discrete, electrically neutral, units in a crystal.
NaCI is the formula for the crystal containing one Na+ ion for
each Cl- ion.
The crystal for the binary compound, CaCI2, is formed from Ca2+
and Cl- ions in the ratio 1:2.
83
Molecular formulas represent the composition in
terms of elements.
Subscripts show the relative numbers of atoms in the
smallest unit.
Estrone, a female hormone and testosterone, a male hormone,
84
differ by only a few atoms
CH4O verses CH3OH
molecular verses condensed
A molecular formula of methanol (wood spirit) is CH4O.
The condensed structural formula is CH3OH; indicating atom
arrangement.
Structural formulas indicate how atoms are linked together;
lines represent chemical bonds.
Methanol
CH3OH
85
Electrostatic potential surfaces show electric charge
distribution.
A red tint indicates a negative potential
due to the negatively charged electrons.
CH3CH2OH
A blue tint indicates a positive potential
86
due to a positively charged nuclei .
Which of the following would NOT be classified as a
molecular compound?
•
•
•
•
•
CO
C2H4
C6H12O6
NH4C2H3O2
NH3
Which of the following compounds exhibits both ionic
and covalent bonding?
•
•
•
•
•
SO2
CF4
NaCl
Na2SO4
P4O10
Chemical Formulas
Empirical and Molecular
Formulas
Percent Composition
Section F in Fundamentals
89
Chemical Formulas
Empirical Formula – The smallest whole-number ratio of
atoms in the compound.
Molecular formula – The actual number of atoms in the
molecule. A whole number multiple of the empirical formula
for that compound.
Empirical Formula = CH2
Molecular Formula = C2H4 Ethene
Molecular Formula = C3H6 Propene
Molecular Formula = C4H8 Butene
Chemical Formulas
Glucose
http://en.wikipedia.org/wiki/File:Glucose_Fisher_to_Haworth.gif
Empirical formula: CH2O
What is the formula mass?
Molecular formula: C6H12O6
What is the molecular mass?
What is the molar mass?
91
Empirical Formulas
A compound is found to contain 72.08 g carbon and 21.21g
hydrogen. Find the empirical formula.
72.08 g
= 6.002 mol C / 6.002 = 1.0 mol C
12.011 g/mol
21.21 g
1.0079 g/mol
C1H3.5
= 21.04 mol H / 6.002 = 3.5 mol H
C2H7
Steps needed to find Empirical Formulas
1. Find number of moles for each element
2. Divide all results by smallest number of moles
present
3. Find smallest multiplication factor needed to
make all numbers whole
4. Multiply all formula subscripts by this factor
93
Empirical formulas
• Laboratory analysis of aspirin determined the following
mass percent composition. Find the empirical formula.
C = 60.00%, H = 4.48%, O = 35.53%
HINT: Convert percentages to grams by assuming you start with 100 g
94
Information
Given: 60.00 g C, 4.48 g H, 35.53 g O
Find: empirical formula, CxHyOz
gC
mol C
gH
mol H
gO
mol O
mole
ratio
whole
number
ratio
pseudoformula
empirical
formula
95
Empirical formulas
Chemical analysis of a hydrocarbon (containing only
hydrogen and carbon atoms) shows that it is 81.8% carbon
by mass. What is its empirical formula?
HINT: figure out the ratio of the number of moles and convert to integers
A) C3H8
B) C2H9
C) C9H2
D) ????
96
An unknown compound contains the following percents
by mass: C: 60.86%, H: 5.83%, O: 23.16%, and N:
10.14%. Find the empirical formula.
•
•
•
•
•
C6H8O2N2
C7H8O2N
C6H8O2N
C8H8O2N
C8H8ON
Percent Composition
• Some common areas of confusion are:
– Ionic compounds use Formula Mass
• Distinguishable molecular units don't really exist
• The forces between adjacent units are
indistinguishable (e.g. NaCl).
– Hydrates
• CoCl2Ÿ6H2O
• All must be included in the formula mass.
98
Cobalt(II) chloride hexahydrate
99
Percent Composition
CaF2
Molar Mass = 40.08 + 2(19.00) = 78.08 g/mol
mass Ca
40.08 g
% Ca =
*100% =
*100% = 51.33%
total mass
78.08 g
mass F
2(19.00) g
%F =
*100% =
*100% = 48.67%
total mass
78.08 g
Percent Composition
SiO2
Molar Mass = 28.09 + 2(16.00) = 60.09 g/mol
28.09 g
mass Si
% Si =
*100% =
*100% = 46.75 %
total mass
60.09 g
2 (15.999) g
mass O
%O =
*100% =
*100% = 53.25%
total mass
60.09 g
101
What is the molar mass of calcium phosphate?
•
•
•
•
•
310.18 g/mole
230.02 g/mole
324.99 g/mole
135.05 g/mole
214.18 g/mole
Concentration
Dilution
Section G in Fundamentals
103
Solutions
Concentrations are measured in mol/Liter = M.
This is called molarity.
The water is called the solvent.
The solute is the species dissolved in the water.
You can determine the number of moles if you know the
concentration and the volume.
E.g.: 50 ml of a 2.0M solution contains
(2.0mol/L)(0.050L)=0.10 mol of solute.
104
Dilutions
You also often need to be able to dilute aqueous solutions
to a different (lower) concentration.
The number of moles in solution A can be determined by
MAVA
(mol/L X L) = mol !
If you dilute the solution, the number of moles does not
change.
So in the diluted solution, the product of the molarity and
volume must be equal to MAVA.
This is generally written as MAVA = MBVB
105
Clicker question
Question: If you have a 1.0 M solution of NaCl and you
need 100 ml of an 0.20 M solution, how should you make it?
A. 20 ml of the 1.0M soln diluted to a total volume of 100ml
B. 10ml of the 1.0M soln diluted to a total volume of 100ml
C. 2 ml of the 1.0M soln diluted to a total volume of 100ml
D. ????
106
Chemical Reactions & Stoichiometry
Balanced Chemical Equations
Electrolytes
Solubility
Ionic equations
Sections H and I in Fundamentals
107
Chemical Reactions
Conservation of Matter
Matter is conserved. It cannot be created or
destroyed, only rearranged.
The building blocks of molecules are the atoms.
Chemical reactions rearrange the atoms in
molecules into different molecules. (Chemical
Change)
108
Chemical Reactions
5 types of reactions:
- Putting things together
synthesis reaction
- Tearing things apart
decomposition reaction
- Rearranging things
single displacement reaction
double displacement reaction
combustion reaction
Chemical Reactions
1. SYNTHESIS A + B
A-B
2N2 (g) + 5O2 (g)
2. DECOMPOSITION
2 HgO (s)
3. COMBUSTION
CxHy + O2
C7H16 (g) + 11O 2 (g)
2 N2O5 (g)
A-B
A + B
2Hg(l) + O2(g)
CO2 + H2O
7CO2 (g) + 8H2O(g)
Chemical Reactions
4. SINGLE DISPLACEMENT REACTION
A-B + C
A-C + B
CuCl2 (s) + Mg (s)
MgCl2 (s) + Cu (s)
5. DOUBLE DISPLACEMENT REACTION
A-B + C-D
A-C + B-D
2NaOH (aq) + CaBr2(aq)
Ca(OH)2(aq) + 2NaBr (aq)
Chemical Reactions - Symbols
2 HCl(aq) + Mg(s)
H2 (g) + MgCl2 (aq)
Reactions can be written as chemical equations:
Reactants è Products (forward rxn)
Reactants ⇄ Products (equilibrium)
“react to produce”, “yield”, “decompose into”
heat
350ºC
“react when heated”
“react under pressure”
“react with a platinum catalyst”
press
5 atm
Pt
Chemical Reactions - Symbols
2 HCl(aq) + Mg(s)
H2 (g) + MgCl2 (aq)
(s)
solid
(l)
liquid
(g)
gas
(aq)
aqueous (in water solution)
evolves as a gas (PRODUCT ONLY)
precipitate (solid) forms (PRODUCT ONLY)
Chemical Reactions - Balancing
The number and identity of atoms in reactants must be
the same as in products. (Balanced Reaction)
The numbers before the molecules are called the
stoichiometric coefficients.
Pb(NO3)2(aq) + 2NaI(aq)
PbI2(s) + 2NaNO3(aq)
Can NOT change subscripts – why not??
If no number is given, the value is 1
5 MgCl2
coefficient
subscript
Chemical Reactions - Balancing
Hints for balancing equations:
- Bottom line – make an educated guess and check.
- Look first for elements that appear only once on each
side of equation
- Treat polyatomic ions as groups
- Leave lone elements until the end
- CHO for combustion reaction
Chemical Reactions - Balancing
H2SO4 + Al(OH)3
3 H2SO4 + 2 Al(OH)3
Al2(SO4)3 +
H2O
Al2(SO4)3 + 6 H2O
Clicker Question
Ammonia is prepared by reacting nitrogen and hydrogen gases
at high temperature according to the unbalanced chemical
equation below.
__ N2(g) + __ H2(g) à __ NH3(g)
What are the respective coefficients when the equation is
balanced with the smallest whole numbers?
A.
B.
C.
D.
1, 1, 1
1, 3, 1
1, 3, 2
2, 1, 2
What are the coefficients for the decomposition of
nitroglycerin?
__ C3H5N3O9 à __ N2 + __ CO2 + __ H2O + __ O2
a)
b)
c)
d)
e)
2,3,6,2,1
2,3,6,5,1
4,6,12,10,12
4,3,12,10,1
4,6,12,10,1
Chemical Equations
Solutions that conduct
electricity
Chemical reactions that
make precipitates
Chemical reactions that
make gases
Writing products of a
chemical reaction
What makes chemical reactions occur?
119
When in water some ionic and or molecular compounds
conduct electricity; some strongly and some not at all.
nonelectrolyte
weak electrolyte
strong electrolyte
Distinguishing between electrolytes is how well they conduct
electricity, hence strong, weak and non.
120
Strong electrolytes
Only solutions with ions can conduct electrical current since ions
have charge (so become charge carriers). Therefore compounds
that dissociate into ions when dissolved in solution are referred to
as electrolytes.
Notice, ions become free to move when the solid dissolves.
Free
Bound
Dissociation
121
Nonelectrolyte
A nonelectrolyte does not form ions in solution.
They can be solids or dissolve yet no ions are present.
In a nonelectrolytic solution,
molecules do not dissociate, they
remain intact.
122
Weak electrolyte
Weak electrolytes barely ionize
in solution; they mostly remain
intact.
Acetic acid is a weak
electrolyte.
Slight dissociation
Only a small fraction (1%) of
CH3COOH molecules separate
(slightly or dissociates
negligibly) into hydrogen ions,
H+, and acetate ions, CH3CO2.
123
Dissolving -versus- Dissociation: Sugar dissolves in water but
does not dissociate into ions. You can dissolve vodka in water but
it is a nonelectrolyte. Oxygen dissolves in your blood but it does
not ionize
No dissociation
nonelectrolyte
n
ion
Partial dissociatio Complete dissociat
weak electrolyte
124
strong electrolyte
Soluble versus insoluble
Soluble substances dissolve in a solvent. K2Cr2O7 is soluble in water;
sugar, vodka and oxygen also dissolve (soluble). Soluble implies in water
(aqueous) unless otherwise stated.
Soluble ionic compounds will both dissolve and dissociate.
Soluble molecular compounds will dissolve and may or may not dissociate.
Insoluble substance do not dissolve significantly in a solvent. A copper
wire is insoluble in water, like rubber or skin.
Soluble
Insoluble
125
Soluble verses insoluble
Soluble: are all in the same
phase eg. solid table salt
dissolves in water.
Soluble compounds can be
strong, weak and
nonelectrolytes; these may or
may not disassociate.
Insoluble: a different
phase i.e. solid in a liquid.
Insoluble compounds are
nonelectrolytes, and do
not disassociate.
126
When to use (aq) or (s)
Soluble (aq) [Dissolves]
only dissolves
Partially
disassociates
disassociates
Insoluble (s)
do not dissolve
nonelectrolyte weak
electrolyte
strong
electrolyte
nonelectrolyte
Sugar, vodka,
oxygen
KMnO4 and
many more
Skin, copper
wire, tires
Acetic acid,
most organic
acids
127
128
Not all soluble ionic compounds dissociate.
Here is a list of common weak electrolytes that only
partially dissociate;
•
•
•
they are soluble;
they do react;
but they do not fully break apart into ions.
Weak Acid
Electrolyte
Dissociate
HC2H3O2
H3PO4
Weak Base
Electrolyte
Dissociate
NH3 or NH4OH
Amines
Weak Salt
Electrolyte
Dissociate
Hg2Cl2, CdCl2
Pb(C2H3O2)2
129
Dissolution of ionic compounds
What happens at
the Particulate
level when salt
(NaCl) dissolves
in water?
http://www.youtube.com/watch?v=EBfGcTAJF4o
Chemical Reactions – Ionic Equations
The Particulate View of dissolving a salt in water:
Pb(NO3)2(aq) → Pb2+ (aq) + 2 NO3-(aq)
NaI(aq) → Na+ (aq) + I-(aq)
What if we start with 3 molecules
of Pb(NO3)2 and 4 molecules of NaI,
dissolved in water?
3 ions of Pb2+
6 molecular ions of NO34 ions of Na+
4 ions of IThe product will be PbI2(s)
131
Chemical Reactions – Ionic Equations
3Pb2+(aq) + 6NO3-(aq) + 4Na+(aq) + 4I-(aq)
→ 2 PbI2(s) +1Pb2+(aq) + 6NO3-(aq) + 4Na+(aq)
132
Chemical Reactions – Ionic Equations
2
3Pb2+(aq) + 6NO3-(aq) + 4Na+(aq) + 4I-(aq)
→ 2 PbI2(s) +1Pb2+(aq) + 6NO3-(aq) + 4Na+(aq)
2Pb2+(aq) + 4I-(aq) → 2PbI2(s)
Simplify
Pb2+(aq) + 2I-(aq) → PbI2(s)
Net ionic equation
Stoichiometric coefficients
Full molecular equation:
Pb(NO3)2(aq) + 2NaI(aq)
PbI2(s) + 2NaNO3(aq)
133
Putting the solubility rules to work
Starting with a molecular equation,
We can use the solubility rules to predict the outcome of the reaction.
(1) Identify species that are aqueous (i.e. dissolve and/or dissociate)
(1) Therefore, we have, AgNO3(aq), NaCI(aq) and NaNO3(aq)
(1) These are aqueous salts that will dissociate back into ions:
Ag+ (aq), NO3- (aq), Na+ (aq), and CI- (aq)
(4) And we are left with a non-electrolyte AgCl(s)
134
AgNO3(aq) + NaCl(aq) è AgCl(s) + NaNO3(aq)
The molecular equation in electrolytic form (dissociated), is called
the complete ionic equation:
Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) è
AgCl(s) + Na+(aq) + NO3-(aq)
We notice the same ions on each side of the equation, which means
they play no role in our chemical reaction. These are called
spectator ions.
Rewriting our complete ionic equation without spectators, gives
us the net ionic equation.
Ag+(aq) + Cl-(aq) è AgCl(s)
This net ionic equation is the net change taking place in the
reaction when Ag+ ions combine with Cl- ions to form the
135
precipitate solid silver chloride, AgCl.
Example: Write a molecular, complete ionic and net ionic
equation when concentrated aqueous solutions of barium nitrate
and ammonium iodate are mixed.
Ba(NO3)2 + NH4IO3 → Ba(IO3)2 + NH4NO3
136
Four common Gas Forming reactions
The following will steadily disassociate into their more stable
gas products.
1) H2CO3(aq) à CO2(g) + H2O(l)
2) H2SO3(aq) à SO2(g) + H2O(l)
3) NH4OH(aq) à NH3(g) + H2O(l)
4) Mmetal + HCl(aq) à H2(g) + MmetalCl
137
Chemical Reactions & Stoichiometry
Limiting Reagents
Percent Yield (Reaction Yield)
Fundamentals Parts L and M
138
Chemical Reactions – Stoichiometry
Convert moles of one substance into moles
of another substance
Mole Ratio:
3 mol Li2SO4
1 mol Mg3P2
Start with a balanced chemical equation.
3 Li2SO4 +
Mg3P2 à 3 MgSO4 + 2 Li3P
Mole ratio
A:B
MM of A
Grams A
Moles A
NA
Particles A
MM of B
Moles B
Grams B
NA
Particles B
139
Chemical Reactions – Percent Yield
Theoretical Yield – Maximum amount of product (mass) that
could be formed from a reaction.
- Calculated from a stoichiometry approach
Actual Yield – Mass of product, measured in the lab.
Percent Yield =
actual yield
theoretical yield
x 100 %
Can be done in grams or moles as long as they are the same
in numerator and denominator
140
Chemical Reactions – Percent Yield
What is the percent yield if 24.8 g of CaCO3 is heated,
producing 13.1 g of CaO?
CaCO3 à CaO + CO2
Given:
Molar mass of CaCO3 = 100.1 g/mol and CaO= 56.1 g/mol
Actual Yield
13.1 g CaO
Theoretical Yield
Percent Yield –
141
Chemical Reactions – Percent Yield
Maximizing percent yield generally requires skill in choosing
the best precursors, best conditions for the reaction, and
using the best techniques.
Can the percent yield be greater than 100%?
If we do observe it, what are the causes?
142
Chemical Reactions – Limiting Reactant
• The Limiting Reactant is the one that runs out first.
• Excess Reactants are the ones that don’t run out. (extra)
Approach
1. Convert reactant masses to moles and determine the
theoretical yield in moles from each reactant.
2. The reactant that can produce the least product is the
limiting reactant.
143
Chemical Reactions – Limiting Reactant
Pb(NO3)2(aq) + 2NaI(aq)
PbI2(s) + 2NaNO3(aq)
What is the limiting reactant if 1.5 g of NaI is reacted with
2.0g of Pb(NO3)2?
The molar mass of NaI is 149.9g/mol
The molar mass of Pb(NO3)2 is 331.2g/mol
A) NaI
B) Pb(NO3)2
C) ????
144
Chemical Reactions – Limiting Reactant
Aluminum and sulfur react to produce aluminum sulfide.
If 9.0 grams of each reactant is mixed together
a) Which is the limiting reactant?
b) Which is in excess?
c) How much (mass) excess is there?
145
Chemical reactions
Acid – Base reactions and titrations
Redox reactions
Fundamentals Parts J and K
146
Acids and bases
Acids
Tart
Litmus
Red
< 7 pH
Bases
Bitter
soapy
Litmus
Blue
> 7 pH
Acids and bases change the color of certain dyes
known as indicators.
147
Acid and base definitions
The Swedish chemist Svante Arrhenius, in 1884:
An acid is a compound that contains hydrogen and reacts with water
to form hydrogen (H+) ions.
HCl(aq)
Cl- (aq) + H+ (aq)
A base is a compound that produces hydroxide ions (OH-) in water.
NH3(aq) + H2O(l)
NH4+ (aq) + OH- (aq)
Limitations to the Arrhenius definition:
1. It is specific to one particular solvent, water.
2. Not all base reactions produce hydroxide ion, OH-.
3. The key process in an acid and base reaction is a proton (H+) transfer
(little to do with OH-).
148
Acid and base definitions
In 1923, Thomas Lowry in England and Johannes Brønsted in
Denmark, came up with a proton (H+) transfer idea.
An acid is a proton donor and a base is a proton acceptor.
HCl(aq) + H2O(l)
Cl- (aq) + H3O+ (aq)
HCl proton donor.
H2O proton acceptor.
HCI releases a hydrogen ion, H+, to water, producing
hydronium ions (H3O+ ) and chloride ions.
H2O accepts the hydrogen ion to form H3O+, water is acting as a
Brønsted base in this reaction.
149
Example
H+ transfer
CH3COOH(aq) + H2O(l)
CH3COO- (aq) + H3O+ (aq)
150
Classifying acids and bases
Brønsted-Lowry acids and bases are further categorized
based on their extent of deprotonation or protonation:
A strong acid is completely deprotonated in solution.
A weak acid is incompletely deprotonated in solution.
A strong base is completely protonated in solution.
A weak base is incompletely protonated in solution.
HCl(aq) + H2O(l) è Cl- (aq) + H3O+ (aq)
strong
100% ionization
CH3COOH(aq) + H2O(l) è CH3COO- (aq) + H3O+ (aq)
weak
About 1% ionization
151
Strong vs weak acids
You will never see HCl in water,
because it completely dissociates
(100%) into H+ and Cl- ions.
On the other hand, for a weak acid
like acetic acid, CH3COOH, you will
see mainly the acid and only about
1% acetate, CH3COO- .
Emphasis on this is seen again
chapters 12 and 13.
152
Clicker Question
Acids that ionize extensively in solution are referred to as
A.
B.
C.
D.
strong acids.
weak acids.
Arrhenius acids.
Brønsted-Lowry acids.
Neutralization reactions
An acid base reaction is called a neutralization reaction.
Neutralization reactions take place between a strong acid
and metal hydroxide:
Acid + metal hydroxide
salt + water
“Salt” is taken from ordinary table salt, sodium chloride.
HCI(aq) + NaOH(aq)
NaCI(aq) + H2O(l)
154
Neutralization reactions
Another acid and base reaction producing a salt and water.
2 HNO3(aq) + Ba(OH)2(aq)
Ba(NO3)2(aq) + 2 H2O(l)
The complete ionic equation:
2 H+ (aq) + 2NO3- (aq) + Ba2+ (aq) + 2OH- (aq)
Ba2+ (aq) + 2NO3- (aq) + 2 H2O(l)
And finally simplified into a net ionic equation.
H+ (aq) + OH- (aq) → H2O(l)
The net outcome of any strong acid base neutralization reaction
is the formation of water.
155
Writing a net ionic equation for a weak acid or weak base
*Molecular equation.
HC2H3O2(aq) + NaOH(aq)
H2O(l) + NaC2H3O2(aq)
*Complete ionic, weak electrolytes do not disassociate.
HC2H3O2(aq) + Na+(aq) + OH-(aq)
H2O(l) + Na+(aq) + C2H3O2- (aq)
*Net ionic equation.
HC2H3O2(aq) + OH- (aq)
H2O(l) + C2H3O2- (aq)
156
REDuction OXidation (REDOX) reactions
• Oxidation: Loss of electrons, or gain of Oxygen atoms
• Reduction: Gain of electrons, or gain of Hydrogen atoms
• Oxidation numbers increase with oxidation, decrease
with reduction
• Reducing agents reduce the species they react with
• Oxidizing agents oxidize the species they react with
Reduction Oxidation
reactions range from common:
combustion, corrosion, to
elaborate: photosynthesis,
metabolism and metal
extraction reactions.
157
Electron gain and electron loss occur together
2 Mg (s) + O 2 (g)
2 MgO (s)
Here Mg atoms lose electrons to form
Mg2+ ions, and the oxygen O2, gains
electrons to form O2-.
2 Mg (s) + O 2 (g)
Fe 2O3 (s) + 3 CO (g)
2 Mg 2+O 2- (s)
2 Fe(l) + 3 CO 2 (g)
Here Fe3+ gains electrons and is
consequently reduced to Fe0 metal.
The CO is oxidized to CO2.
158
Production of steel
Oxidation Numbers: Keeping Track of Electrons
The words “oxidation number” and “oxidation state” are
interchangeable.
Assigning oxidation number to elements use simple rules:
1) The oxidation numbers of uncombined elements are zero
(0) (free elements or atoms with themselves) i.e. H2, O2 F2,
Cl2, Li(s), U(s).
2) The sum of the oxidation numbers of all the atoms in a
species is equal to its total charge.
LiCl
MgCl2
MnO4-
Li+ClMg2+Cl2[Mn7+O42-]-
total charge = 0
total charge = 0
total charge = -1
159
Oxidation Numbers: Keeping Track of Electrons
Common oxidation numbers:
• Hydrogen is H+ unless it’s with a metal when it becomes H•
The reason for this is that most metals are more electropositive than H so H
takes on the negative charge.
• Groups 1 and 2 oxidation numbers are equal to their group
number.
• Halogens are -1 unless the halogen is in combination with
oxygen or another halogen higher in the group. E.g. ClF2
• Oxygen is O2- except when combined with fluorine. Less
common are peroxides (O22-), superoxides (O2-), and ozonides
(O3-).
160
Clicker question
Consider the conversion of SO2 to SO42Is it an:
A) Oxidation
A) Reduction
B) ???
161
Example: Find the oxidation numbers of sulfur, nitrogen, and
chromium in (a) SO3 ; (b) N2O ; and (c) Cr2O72-.
+6 -2
SO3
+1 -2
N2O
+6 -2
Cr2O72-
+6 + 3(-2) = 0
2(+1) + (-2) = 0
2(+6) + 7(-2) = -2
(a) The “-” side is -6, so what combines with -6 = 0 ? +6
“+6” (+6 -6 = 0) or S+6
(b) The “-” side is -2, so what combines with -2 = 0 ? +2
Since there are 2N atoms with total charge +2, each is N+1
(c) The “-” side is -14, so what combines with -14 = -2 ? +12
Since there are 2Cr atoms with total charge +12, each is Cr+6
162
Oxidizing and Reducing Agents
In the following what is oxidized and reduced ?
Zn(s) +
Cu2+
(aq) è
Zn2+
163
(aq) + Cu(s)
Definition: Oxidizing and Reducing Agents
Zn(s) + Cu2+ (aq)
Here zinc ( Zn
Zn2+ (aq) + Cu (s)
Zn2+ + 2 e- ) loses two electrons. These are taken by
the Cu2+ causing reduction to Cu. The role or purpose of the Zn is provide
e- to reduce Cu2+. Therefore Zn(s) is the reducing agent.
Here copper cations ( Cu2+ + 2 e-
Cu ) each gain two electrons. Cu2+
took the electrons from Zn, causing the oxidation of Zn to Zn2+. The role or
purpose of Cu2+ is to oxidize Zn. Therefore, Cu2+ is the oxidizing agent.
164
Identifying oxidizing agents and reducing agents
Identify the oxidizing agent and the reducing agents.
First, we’ll have to identify the oxidation number of each element.
+6
-2
+2
+1
+3
+3
+1 -2
Cr2O72- (aq) + 6 Fe2+ (aq) + 14 H+ (aq) → 6 Fe3+ (aq) + 2 Cr3+ (aq) + 7 H2O(l)
165