Download I. scientific notation. – a shorthand that scientists use when dealing

CHEM 101 LECTURE NOTES
Tuesday, August 30, 2001
Fall 2001
Chapter 1
Dr. Joy Heising
S 572-580
I. scientific notation. – a shorthand that scientists use
when dealing with very large or very small numbers.
a. 108 g silver = 602,000,000,000,000,000,000,000
silver atoms.
To simplify, we write 6.02 x 1023.
b. 1 silver atom = 0.000 000 000 000 000 000 000
179 gram
To simplify, we write 1.79 x 10-22.
c. 7,400,000 = 7.4 x 106
d. 0.000246 = 2.46 x 10-4
II. significant figures.
a. Exact number - no estimation is involved. Most
numbers are not exact.
b. Accuracy – describes how closely measured
values agree with correct values.
c. Precision – describes how closely individual
measurements agree with each other.
d. Significant figures – the digits believed to be
correct by the person making the measurements.
Ex. 10 mL cylinder is graduated in tenths. 4.73 mL
– how many significant figures? (3)
Ex. 2. population of a city is 2,678,342 people.
How many ‘significant’ figures? (7) Do you
believe this number?
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CHEM 101 LECTURE NOTES
Tuesday, August 30, 2001
Fall 2001
Chapter 1
Dr. Joy Heising
S 572-580
III. some rules for significant figures.
a. Zeroes i. ‘leading’ zeroes are never significant
ii. ‘trailing’ zeroes are sometimes significant –
definitely, if there is a decimal point in the
number
iii. zeroes in the middle of a number (4,309) are
significant
b. math involving significant figures.
i. Addition/subtraction
3.6923
+ 1.234
+ 2.02
= 6.95
ii. multiplication/division
2.7832
x 1.4
= 3.9
iii. rounding – in a multi-step calculation, carry
all digits (even non-significant ones) until
the end of the calculation. However, if you
are reporting the results of each step of the
calculation, report the correct number of
significant figures.
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CHEM 101 LECTURE NOTES
Tuesday, August 30, 2001
Fall 2001
Chapter 1
Dr. Joy Heising
S 572-580
IV. Unit Factor method (also called dimensional
analysis).
a. 1 foot = 12 inches. This relationship can be
converted into two unit factors. The numerator
and denominator describe the same distance:
ex.
1 foot
12 inches
12 inches =
1 foot = 1
9.32 yards = ? millimeters?
1 meter = 39.37 inches
9.32 yards
x
3 feet
1 yard
x
12 inches
1 foot
x
1 meter
39.37 inches
x
1000 mm
1 meter
= 8520 mm
627 mL = ? gallons?
1 gallon = 0.946 L
627 mL
x
1L
1000 mL x
1 gallon
0.946 L
= 0.663 gallons
V.
Density
density – mass per unit volume. D = M/V (g/mL)
742 g / 97.3 mL = 7.62 g/mL
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CHEM 101 LECTURE NOTES
Tuesday, August 30, 2001
Fall 2001
Chapter 1
Dr. Joy Heising
S 572-580
VI. Heat.
a. Temperature – the intensity of heat in a body,
that is, how hot or how cold the body is. Three
different temperature scales represent same temp:
i.Fahrenheit – bp H2O 212ºF
ii.Celsius – bp H2O 100ºC 5 x (ºF – 32)/9
iii.Kelvin - bp H2O 373 K (Celsius + 273)
b. the SI unit of heat and energy is the joule (J).
[another useful unit is the calorie – the amount of
heat required to raise the temperature of 1 g of
H2O from 14.5ºC to 15.5ºC. 1 cal = 4.184 J.]
c. specific heat – the amount of heat (usually
expressed in joules) required to raise the
temperature of one gram of the substance one
degree Celsius with no change in state.
d. Heat capacity – the amount of heat required to
raise the temperature of one mole of a substance
one degree Celsius with no change in state.
e. Specific heats of a few common substances
Substance
specific heat
Ice
Liquid water
Steam
2.02 J/gºC
4.18 J/gºC
2.03 J/gºC
Ex. calculate the amount of heat required to raise
the temperature of 200 g H2O from 10.0ºC to
55.0ºC.
55.0ºC – 10.0ºC
=
200.0g
x 45.0ºC
x
45.0ºC
4.18 J
gºC
= 3.76 x 104 J
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CHEM 101 LECTURE NOTES
Tuesday, August 30, 2001
Fall 2001
Chapter 2
Dr. Joy Heising
S 572-580
Chapter 2
Stoichiometry–describes the quantitative relationships
among elements
1. in compounds (composition stoichiometry)
2. in chemical changes (reaction stoichiometry)
Dalton’s Atomic Theory – summarized experimental
observations and interpretations in the nature of atoms:
1. an element is composed of extremely small,
indivisible particles called atoms
2. Atoms cannot be created, destroyed, or transformed
into atoms of another element
3. compounds are formed when atoms of different
elements combine with each other in small, wholenumber ratios.
4. the relative numbers and kinds of atoms are
constant in a given compound.
Atoms – (review, chapter 1)
Molecules - (review, chapter 1)
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CHEM 101 LECTURE NOTES
Tuesday, August 30, 2001
Fall 2001
Chapter 2
Dr. Joy Heising
S 572-580
Chemical formula – the chemical composition of a
substance.
Some examples of chemical formulas for elements:
Monoatomic elements – sodium (Na), copper (Cu), barium (Ba)
Diatomic elements – O2, N2, Br2
More complex molecules – S8, P 4
For compounds, the formula indicates
1. elements present
2. the ratio in which the atoms of the elements occur
Some examples:
HCl – 1 H atom, 1 Cl atom
H2O -2 H atoms, 1 O atom
NH3 -1 N atoms, 3 H atoms
C3H8 – 3 C atoms, 8 H atoms
1 chemical formula
1 chemical formula
1 chemical formula
1 chemical formula
(Fig. 2.6)
law of constant composition – (review, chapter 1)
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CHEM 101 LECTURE NOTES
Tuesday, August 30, 2001
Fall 2001
Chapter 2
Dr. Joy Heising
S 572-580
molecular compounds: (Table 2.2)
H2O2
CH3CH2OH
Methane
NH3
CO2
benzene
hydrogen peroxide
methanol
CH4
ammonia
carbon dioxide
C6 H6
Ion – an atom or group of atoms that carries an electrical
charge
Anion – negatively charged ions (Cl-)
Cation – positively charged ions (Na +)
Ionic compounds – extended array of ions in which the total
positive and negative charges are equal (Fig 2.7)
Formula unit – the simplest whole number ratio of ions in
the compound (NaCl)
Is PO43- a molecule? No.
polyatomic anion.
Compound names and formulas (table 2-3)
NaCl
Sulfuric acid
Molecule
sodium chloride
H2 SO4
?
NH4NO3
Na2SO4
ion
7
ammonium nitrate
sodium sulfate
CHEM 101 LECTURE NOTES
Tuesday, August 30, 2001
Fall 2001
Chapter 2
Dr. Joy Heising
S 572-580
Atomic weight – the relative masses of atoms of different
elements that are proportional to the actual masses of
atoms.
amu – units of atomic weight. (atomic mass unit)
1 amu – 1/12 the mass of carbon-12.
H 1 amu
Mg 24.3 amu
N
Ag
14 amu
107.8 amu
amu is not terribly practical ‘in real life’ for experimental
chemists, as the quantities are too small (most scientists
aren’t weighing individual atoms to do their experiments).
As a result, chemists use another larger unit called the
mole.
Mole – 6.022 x 1023 atoms, molecules or formula units
Avogadro’s number
**one mole of atoms of an element has a mass in grams
numerically equal to the atomic weight of the element**
H 6.022 x 1023 atoms
N 6.022 x 1023 atoms
Mg 6.022 x 1023 atoms
H2
5.022 x 1023 molecules, 1.204 x 1024 atoms
How many atoms are contained in 1.67 moles of Mg?
How many grams does 1.67 moles of Mg weigh?
= 1.00 x 1024 atoms
= 40.6 grams
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