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Unit 1 Introduction: Matter and Measurement

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Title: Unit 1 Introduction: Matter and Measurement


1
Unit 1IntroductionMatter and Measurement
CHM 1045 General Chemistry and Qualitative
Analysis
  • Textbook References
  • Module 1
  • Dr. Jorge L. Alonso
  • Miami-Dade College Kendall Campus
  • Miami, FL

2
Scientific Method
  • A systematic approach to solving problems.

Observation the detection of a phenom enon by
our sensory organs or their extensions
(instruments). Scientist
study CAUSE ? EFFECT Relationships Which factors
affect the behavior of gasses?
?,P, V, T
GasVariables
Hypothesis initial or tentative explanation of
the causes of a phenomenon.
Experiment carefully designed hypothesis
testing done by controlling all the variables
except your suspected CAUSE (independent
variable, x) which is manipulated in order to
observe its EFFECT (dependent variable, y).
3
Experiment carefully designed hypothesis
testing done by controlling all the variables
except your suspected CAUSE (independent
variable, x) which is manipulated in order to
observe its EFFECT (dependent variable, y).
OR
Dependent Variable
Boyles Law
Independent Variable
Law concise verbal or mathematical summary for a
variety of observations and experiences.
Theory a comprehensive explanation for natural
phenomena that has withstood repeated analysis
and experimentation.
Kinetic Molecular Theory
4
Matter Anything that has mass and takes up space.
What is the universe composed of?
  • Chemicals
  • Substances
  • Things

Chemistry
Energy the ability to perform an activity (work).
Physics
  • Kinetic (motion heat)
  • Electromagnetism (light, electricity
    chemical bonds)
  • Nuclear
  • Gravity

Matter with Energy
5
Describing Matter Physically and Chemically
Chemistry The study of matter and the changes it
undergoes.
(1) Physical Properties (3) Chemical
Properties (2) Physical Change (4) Chemical
Change
(reactions equations)
6
Lets get Physical!
Physical Properties The physical characteristics
(appearance) of matter. State or phase (s,l,g),
color, mass, volume, density, melting boiling
points, solubility, etc
H2O (g)
H2O (s)
H2O (l)
Physical Changes Changes in the physical
properties of matter. Changes of state (s l
g), density, temperature, shape, volume, etc.
Physical changes occur without changes in the
composition of matter.
Physical change
H2O (s) ? H2O (l) ? H2O (g)
Chemical change
2 H2O (s) ? 2 H2 (g) O2 (g)
7
Heat Energy Phase (State) Changes
(fusion)
Phase Changes Changes of State
8
Kinetic Energy States of Matter
Heat Kinetic Energy ? Temperature
c. pt. (condensation point)
f. pt. (freezing point)
m. pt (melting point)
b. pt. (boiling point)
Density (mass per unit volume)
For H2O
m.pt. f.pt 0OC
b.pt. c.pt 100OC
For Methanol
m.pt. f.pt -98OC
b.pt. c.pt 65OC
KineticMolecularTheory PhaseChange
9
Heating Curve Energy Phase Changes
Heating water vapor
Heat of Vaporization
Heating liquid water
Heating solid ice
Heat of Fusion
For H2O
m.pt. f.pt 0OC
b.pt. c.pt 100OC
For Methanol
m.pt. f.pt -98OC
b.pt. c.pt 65OC
(?Hv in kJ/mol)
10
Lets get Chemical!
What happens when you add Na to water? K?
Chemical PropertiesCan only be observed when a
substance reacts and is changed into another
substance. Does it react? With which
substance does it react? Flammability?
Corrosiveness?, etc.
Na K in H2O
Chemical Changes The changes that occur in the
process of producing new substances. Combustion,
oxidation, decomposition, etc.
Is this chemical or physical change?
Chemical change is always accompanied by physical
change!
11
Describing Chemical Change
Chemical Reaction the actual phenomenon that
occurs when chemicals change in
composition. Chemical Equation a symbolic
representation of a chemical reaction. Based on
the Atomic Theory.
Quiz Question Write the balanced chemical
equations for the reactions of (1) Sodium (Na)
Water (HOH) (2) Potassium
(K) Water
12
Chemical Reactions and Equations

-

2 Na 2 HOH ? 2 NaOH H2 2 K 2 HOH ? 2 KOH
H2
(Hint single displacement or replacement
reaction)
13
Matter and the Atomic Theory
  • Atoms are the building blocks of all matter.
  • Elements are made of the same kind of atom.
  • Compounds are made of two or more different
    kinds of atoms.
  • Mixtures are composed of different
    elements/compounds together.

14
Classification of Matter
Heterogeneous

Cu(NO3)2
Mixtures (Heterogeneous)
Physical Separation
Mixtures
Homogeneous
Solutions (Homogeneous)
Physical Separation
Cu(NO3)2 (aq)
Cu(NO3)2 (s)
Compounds
Chemical Decomposition
Pure Substances
Elements
15
methods of physically separating (purifying)
substances from mixtures and solutions into pure
substances.
Separatory Techniques
  • based on differences in physical properties of
    the substances present in the mixture/solution.
  1. Filtration by solubility vs. insolubility
  2. Metal Smelting Refining - by differences in
    melting point (ability to form a liquid)
  3. Distillation by differences in boiling points
    (ability to form a gas)
  4. Chromatography by differences in degree of
    solubility

16
(1) Filtration Separates insoluble solid
substances from liquids and solutions.
Mixture K2Cr2O7(s) NaNO3(s)
H2O

K2Cr2O7(s)
NaNO3(aq)
17
(2) Metal Smelting Refining
These techniques are used to differentially melt
mixtures of metals (alloys) by means of their
different melting points (ability to form a
liquid when heated).
The sweat furnace operates at a temp at which
one metal is selectively melted from a component,
leaving the metal with the higher melting point,
usually a ferrous metal as a recoverable solid.
Example mixture of Cu Zn, heated to 500C
18
(3) Distillation Separates homogeneous mixture
on the basis of differences in boiling points
(ability to form a gas).

Substance b.pt. Ethyl Alcohol
77oC Water 100oC Sodium Chloride 1413oC
Solution
Alcohol
19
Distillation of Hydrocarbons Petroleum Refinery
Towers
compounds composed of molecules arranged in a
long chain of carbon atoms with hydrogen atoms
attached to the carbon chain.
  • Name (b.pt. ?C) C Structural Formula
  • Methane (-162) 1 CH4
  • Ethane (-89) 2 CH3CH3
  • Propane (-42) 3 CH3CH2CH3
  • Butane (-0.5) 4 CH3CH2CH2CH3
  • Pentane (36) 5 CH3CH2CH2CH2CH3
  • Hexane (69) 6
    CH3CH2CH2CH2CH2CH3
  • Heptane (98) 7
    CH3CH2CH2CH2CH2CH2CH3
  • Octane (126) 8
    CH3CH2CH2CH2CH2CH2CH2CH3
  • Nonane (151) 9 CH3 CH2
    CH2CH2CH2CH2CH2CH2CH3
  • Decane (174) 10 CH3CH2CH2CH2CH2CH2CH2CH2CH2CH3

20
Distillation of Hydrocarbons Petroleum Refinery
Towers
0 ?C
120 ?C
200 ?C
250 ?C
300 ?C
21
Separation by differences in ability to form a
gas (boiling points)
Mixture / Solution or Pure Substance?
H2O vapor
Physical Separation
Cu(NO3)2(s)
  • Mixtures can be separated by differences in the
    physical properties of the substances they are
    composed of.
  • Pure substances cannot be separated by physical
    methods.

22
Separation by differential solubility, filtration
and evaporation
H2O vapor
Mixture CdS (yellow, insoluble substance),
Cu(NO3)2 (blue soluble substance), H2O(clear
liquid).
23
(4) Chromatography Separates substances on the
basis of their differences in their solubility in
a specific solvent.
Filter paper
Substance to be separated (black ink)
Solvent 50 50 Water Alcohol
Paper Chromatography
24
(No Transcript)
25
Classification of Matter
Heterogeneous

Cu(NO3)2
Mixtures (Heterogeneous)
Physical Separation
Mixtures
Homogeneous
Solutions (Homogeneous)
Physical Separation
Cu(NO3)2 (aq)
Cu(NO3)2 (s)
Compounds
Chemical Decomposition
Pure Substances
Elements
26
Chemical Decomposition of Pure Substances
  • Cannot be separated by physical means.
  • Composed of one substance only, which can be
    either an element or a compound.
  • Compounds can be broken down by chemical means,
    elements cannot.

Examples of pure substances Gold (Au), Oxygen
(O2), Water (H20), Methanol (CH3OH), Table salt
(NaCl) Each has its specific physical properties
(m. pt., density, etc.)
27
Compounds
  • can be broken down into more elemental particles
    (elements) by chemical decomposition reactions.

Electrolysis of Water
2 H2O (l) ? 2 H2 (g) O2 (g)
elect.
Electrolysis
28
How do we get pure Sodium?
2 NaCl (l) ? 2 Na (l) Cl2 (g)
elect
  • NaCl is electrolyzed in a Downs cell.
  • Gaseous Cl2 allowed to disperse
  • Molten Na siphoned off

29
Elements
  • cannot be broken down into more elemental
    particles by ordinary chemical means.

30
Classification of Matter
Heterogeneous
Physical Separation
Homogeneous
Mixture
Physical Separation
Solution
Chemical Decomposition
mixture vs. compound
Element
Compound
31
Units of Measurementlength (m)mass (g,
kg)volume (mL, L)temperature (oC, oK)time (s)
32
Metric System
SI Prefixes SI Prefixes SI Prefixes SI Prefixes SI Prefixes
Prefix Symbol Meaning Multiplier (numerical) Multiplier (exponential)
yotta Y septillion 1,000,000,000,000,000,000,000,000 1024
zetta Z sextillion 1,000,000,000,000,000,000,000 1021
exa E quintillion 1,000,000,000,000,000,000 1018
peta P quadrillion 1,000,000,000,000,000 1015
tera T trillion 1000,000,000,000 1012
giga G billion 1,000,000,000 109
mega M million 1,000,000 106
kilo k thousand 1,000 103
hecto h hundred 100 102
deka da ten 10 101
UNIT ONE 1 1 100
deci d tenth 0.1 10-1
centi c hundredth 0.01 10-2
milli m thousandth 0.001 10-3
micro ? millionth 0.000 001 10-6
nano ? billionth 0.000 000 001 10-9
pico ? trillionth 0.000 000 000 001 10-12
femto ? quadrillionth 0.000 000 000 000 001 10-15
atto ? quintillionth 0.000 000 000 000 000 001 10-18
zepto z (?) sextillionth 0.000 000 000 000 000 000 001 10-21
yocto y septillionth 0.000 000 000 000 000 000 000 001 10-24
When using dimensional analysis for metric
problems always consider the larger unit as
having a value of 1, then the smaller unit would
contain a large multiple of that unit.
X 1000
X 10
X 10
X 1000
Example 1 m compared to cm.
33
Atomic Dimensions
Atoms Tenth of a nanometer (10 -9 m) Nuclei of
atoms Hundredth of a picometer (10 -12
m) Protons Neutrons Fentometer (10-15
m) Quarks electrons Attometer (10-18 m)
34
Metric Conversions
SI Prefixes SI Prefixes SI Prefixes SI Prefixes SI Prefixes
Prefix Symbol Meaning Multiplier (numerical) Multiplier (exponential)
yotta Y septillion 1,000,000,000,000,000,000,000,000 1024
zetta Z sextillion 1,000,000,000,000,000,000,000 1021
exa E quintillion 1,000,000,000,000,000,000 1018
peta P quadrillion 1,000,000,000,000,000 1015
tera T trillion 1000,000,000,000 1012
giga G billion 1,000,000,000 109
mega M million 1,000,000 106
kilo k thousand 1,000 103
hecto h hundred 100 102
deka da ten 10 101
UNIT ONE 1 100
deci d tenth 0.1 10-1
centi c hundredth 0.01 10-2
milli m thousandth 0.001 10-3
micro ? millionth 0.000 001 10-6
nano ? billionth 0.000 000 001 10-9
pico ? trillionth 0.000 000 000 001 10-12
femto ? quadrillionth 0.000 000 000 000 001 10-15
atto ? quintillionth 0.000 000 000 000 000 001 10-18
zepto z (?) sextillionth 0.000 000 000 000 000 000 001 10-21
yocto y septillionth 0.000 000 000 000 000 000 000 001 10-24
Always convert PREFIXES to UNITS (not PREFIXES to
other PREFIXES)
Example Mm compared to pm.
meter, liter, gram
Factors, ratios, equivalences.
Example cm compared to ?m.
35
Metric Conversion Problems
LENGTH km hm dam METER (m) dm cm mm 2.54 cm 1 in. 1 mile 5,280 ft. 1 yd 36 in. 3 ft. 1 yard 12 in. 1 ft.
  • How many pm are there in 0.0023 cm?
  • Change 60. mph to km/s. Hint 1 mi. 1.6 km
  • How many m3 of water are there in 25 ft3 ?

3
3
3
36
Volume Liter (L) and the milliliter (mL)
10 cm
10 cm
  • A liter is a cube 1 dm3 10 cm long on each
    side.

10 cm
1 L dm3 (10 cm)3 (10 X 10 X 10) cm3 1000
cm3 1000 mL or 1mL 1/1000 L
Cubic centimeter
  • A milliliter (mL) is a cube 1 cm long on each
    side.

milliliter
37
Temperature measure of the average kinetic
energy (motion caused by heat) of the particles
in a sample.

K.E ? Temp
?T change in temp
As KE increases molecules vibrate more and their
volume expands (Temp).
  • 373
  • 273
  • 100

100 - 0 100
  • 212
  • 32
  • 180

?C (?F - 32) 1.8
?F 1.8(?C) 32
K ?C 273.15
38
Measured vs Exact Numbers
  • Measured Numbers (1) Accuracy Precision (2)
    Uncertainty (3) Significant figures
  • rounding-off
  • Exact Numbers from formulas, definitions
    counting

For sphere
1 mile 5,280 ft 1 km 1,000 m
39
Measured Numbers Accuracy versus Precision
  • Accuracy refers to the proximity of a
    measurement to the true value of a quantity.
  • Precision refers to the proximity of several
    measurements to each other.

40
Measured vs. Exact Numbers
  • Measured numbers are obtained when a measuring
    instrument (ruler, balance, thermometer) is used
    to determine a physical property of a substance.

13.7
0.1
uncertainty
The number of significant figures these
measurements contain depend on the accuracy of
the instrument being used.
7.63
0.01
uncertainty
41
Uncertainty in Measurements
  • Different instruments have different degrees of
    accuracy, uncertainty is 1 of estimated digit.

0.01
0.1
uncertainty
89.5 mL
2.65 mL
42
Measured vs. Exact Numbers
METRIC METRIC-ENGLISH ENGLISH
CONVERSIONS
Exact Numbers
Exact Numbers
LENGTH km hm dam METER (m) dm cm mm 2.54 cm 1 in. 1 mile 5,280 ft. 1 yd 36 in. 3 ft. 1 yard 12 in. 1 ft.
Measured Numbers
(1 in is exact, the 2.54 cm is measured)
How many km are there in a Marathon (26 miles)?
43
Significant Figures
  • Significant figures refers to digits that were
    accurately measured by an instrument.
  • Example 220g, 220. g, 220.5g, 220.50g,
    220.507g.
  • (all numbers above are measures of the same
    object, what is the difference?)

accuracy
44
Rules for determining the number of Significant
Figures
  1. All nonzero digits (NZD) are always significant.
  2. Zeroes between NZD are always significant. Ex
    103
  3. Zeroes to the left of NZD are never significant.
    Ex 0.0103
  4. Zeroes to the right of NZD are significant if a
    decimal point is written anywhere in the number.
    Ex. 0.01030

45
Rounding-off
  • Round-off your calculated numbers, to the correct
    number of significant figures, so we do not
    overstate the accuracy of our answers.
  • Example 23g 23.632g 46.632

47g
You cannot add an inaccurate measurement to a
accurate measurement and get and accurate answer.
46
Significant Figures in Addition Subtraction
  • When addition or subtraction is performed,
    answers are rounded to the least significant
    decimal place.
  • Example add the following numbers
  • 34
  • 231.678
  • 0.00354
  • 265.68154
  • 266

47
Significant Figures in Multiplication Division
  • Answers are rounded to the number of digits that
    corresponds to the least number of significant
    figures in any of the numbers used in the
    calculation.
  • Example (29.2 20.0) (1.79 x 105)
  • 1.39

(29.2 20.0) 9.2
Calculator answer 1.1847482 x 106
Correct answer 1.2 x 106
48
Uncertainty in Measurements
  • Piece of Black Paper with rulers beside the
    edges Determine the Area of Black Paper!

Lets look more accurately !
Area Length x Width
49
Uncertainty in Measurements
  • Piece of Paper Side A enlarged
  • How long is the paper to the best of your ability
    to measure it?

13.6 cm 0.1 cm
When using an instrument your last digit recorded
should be a significant digit estimated between
the two smallest measurement lines of your
instrument. Your precision would be 1 of that
digit.
50
Uncertainty in Measurements
  • Piece of Paper Side B enlarged
  • How wide is the paper to the best of your ability
    to measure it?

7.63 cm 0.01 cm
When using an instrument your last digit recorded
should be a significant digit estimated between
the two smallest measurement lines of your
instrument. Your precision would be 1 of that
digit.
51
Area of Paper
  • Area 13.6 cm x 7.63 cm 103.768 cm2 is the
    calculator answer.

104 cm2
52
Density
  • A Physical property of a substance, defined as
  • Amount of matter ( atoms) per unit volume
    compactness.
  • the mass divided by the volume.

Mass (g)
Volume (mL)
53
Density and Temperature
Density the mass of a substance divided by its
volume.
Temperature a measure of the amount of kinetic
energy (motion) an object possesses.
As the temperature increases the volume increases
due to the greater kinetic energy of the atoms or
molecules. The mass is not affected.
54
Density of water at various temperatures
C F D in g/cm³
0.0 32.0 0.9998425
4.0 39.2 1.0000000
15.0 59.0 0.9991026
20.0 68.0 0.9982071
25.0 77.0 0.9970479
37.0 98.6 0.9933316
50.0 122.0 0.9880400
100.0 212.0 0.9583665
55
Density Problems

For a Fe metal object whose density is 7.86
g/mL. (a) What is the mass (g) of a piece of this
metal if it displaces 12. mL of water in a
graduated cylinder?
(b) What is the volume in mL of 34 kg of this
same metal?
56
Density Problems

The density of Hg is 11.7 g/mL. What is it in
kg/m3?
3
57
Density of various substances
Density is directly proportional to the Molecular
Weight of a substance.
DBr ? MW
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