Transcript Introduction: Matter & Measurement

Introduction: Matter &
Measurement
AP Chemistry
Chapter 1
Chemistry
What is chemistry?
It is the study of
the composition of
matter and the
changes that
matter undergoes.
What is matter?
It is anything that
takes up space and
has mass.
Elements, Compounds & Mixtures
• A substance is matter that has a
definite composition and
constant properties.
• It can be an element or a
compound.
Elements, Compounds & Mixtures
• An element is the simplest form
of matter.
• It cannot be broken down
further by chemical reactions.
Elements, Compounds & Mixtures
• A compound can be separated
into simpler forms.
• It is a combination of two or
more elements.
Mixtures
• A mixture is a physical blend of two or
more substances.
1. Heterogeneous Mixtures
– Not uniform in composition
– Properties indefinite & vary
– Can be separated by physical methods
Mixtures
2. Homogeneous Mixtures
– Completely uniform in composition
– Properties constant for a given sample
– Cannot be separated by physical
methods (need distillation,
chromatography, etc)
– Also called solutions.
Separating mixtures
• Only a physical change- no new matter
• Filtration- separate solids from liquids with
a barrier.
• Distillation- separate different liquids or
solutions of a solid and a liquid using
boiling points.
– Heat the mixture.
– Catch vapor and cool it to retrieve the liquid.
• Chromatography- different substances are
attracted to paper or gel, so move at
different speeds.
Filtration
Distillation
Chromatography
Physical & Chemical Properties
• Physical property – characteristics of a
pure substance that we can observe
without changing the substance; the
chemical composition of the substance
does not change.
Physical & Chemical Properties
• Chemical property – describes the
chemical reaction of a pure substance
with another substance; chemical
reaction is involved.
Physical & Chemical Properties
Physical properties
• appearance
• odor
• melting point
• boiling point
• hardness
• density
• solubility
• conductivity
Chemical properties
• reaction with oxygen
(flammability)
• rxn with water
• rxn with acid
• Etc….
Intensive & Extensive Properties
Intensive properties
Extensive properties
• Do not depend on the • Depend on the
amount of sample
quantity of the
being examined
sample
–
–
–
–
–
–
temperature
odor
melting point
boiling point
hardness
density
– mass
– volume
• Etc….
Physical & Chemical Changes
Physical changes
Chemical changes
• The composition of • The substance is
the substance doesn’t transformed into a
change
chemically different
substance
• Phase changes (like
liquid to gas)
• All chemical
– Evaporation, freezing, reactions
condensing, subliming,
etc.
• Tearing or cutting the
substance
Signs of a Chemical Changes
1. permanent color change
2. gas produced (odor or
bubbles)
3. precipitate (solid) produced
4. light given off
5. heat released (exothermic) or
absorbed (endothermic)
Making Measurements
•A measurement is a number
with a unit.
•All measurements, MUST have
units.
Types of Units
Energy
Pressure
Joule
Pascal
J
Pa
Prefixes
1,000,000,000
109
• mega - M
1,000,000
106
•
•
•
•
•
•
•
1,000
0.1
0.01
0.001
0.000001
0.000000001
0.000000000001
103
10-1
10-2
10-3
10-6
10-9
10-12
• gigakilo decicentimillimicronanopico-
G
k
d
c
m
m
n
p
Measurements
There are two types of measurements:
 Qualitative data are words, such as
color, heavy or hot.
 Quantitative measurements involve
numbers (quantities), and depend on:
1. The reliability of the measuring
instrument.
2. The care with which it is read – this is
determined by YOU!
Accuracy & Precision
 Accuracy – how close a
measurement is to the true value.
 Precision – how close the
measurements are to each other
(reproducibility).
Precision and Accuracy
Neither
accurate
nor precise
Precise,
but not
accurate
Our goal!
Precise
AND
accurate
Which are Precise? Accurate?
Uncertainty in Measurements
Measurements are performed with
instruments, and no instrument can read to
an infinite number of decimal places
•Which of the balances below has the
greatest uncertainty in measurement?
1
2
3
Uncertainty
• Basis for significant figures
• All measurements are uncertain to
some degree
• Precision- how repeatable
• Accuracy- how correct - closeness to
true value.
• Random error - equal chance of being
high or low- addressed by averaging
measurements - expected
Uncertainty
• Systematic error- same direction each
time
• Want to avoid this
• Bad equipment or bad technique.
• Better precision implies better
accuracy.
• You can have precision without
accuracy.
• You can’t have accuracy without
precision (unless you’re really lucky).
Significant Figures in
Measurements
Significant figures in a measurement
include all of the digits that are
known, plus one more digit that is
estimated.
Sig figs help to account for the
uncertainty in a measurement.
To how many significant figures can
you measure this pencil?
What is wrong with this ruler? What is
it missing?
Rules for Counting Significant
Figures
Non-zeros always count as
significant figures:
3456 has
4 significant figures
Rules for Counting Significant
Figures
Zeros
Leading zeroes do not count as
significant figures:
0.0486 has
3 significant figures
Rules for Counting Significant
Figures
Zeros
Captive zeroes always count as
significant figures:
16.07 has
4 significant figures
Rules for Counting Significant
Figures
Zeros
Trailing zeros are significant only
if the number contains a
written decimal point:
9.300 has
4 significant figures
Rules for Counting Significant
Figures
Two special situations have an
unlimited (infinite) number of
significant figures:
1. Counted items
a) 23 people, or 36 desks
2. Exactly defined quantities
b) 60 minutes = 1 hour
Sig Fig Practice #1
How many significant figures in the following?
1.0070 m  5 sig figs
17.10 kg  4 sig figs
100,890 L  5 sig figs
3.29 x 103 s  3 sig figs
These all come
from some
measurements
0.0054 cm  2 sig figs
3,200,000 mL  2 sig figs
3 cats  infinite
This is a
counted value
Significant Figures in
Calculations
 In general a calculated answer
cannot be more accurate than the
least accurate measurement from
which it was calculated.
 Sometimes, calculated values need to
be rounded off.
Rounding Calculated Answers
 Rounding
 Decide how many significant figures
are needed
 Round to that many digits, counting
from the left
 Is the next digit less than 5? Drop it.
 Next digit 5 or greater? Increase by 1
Rules for Significant Figures in
Mathematical Operations
 Addition and Subtraction
 The answer should be rounded to
the same number of decimal
places as the least number of
decimal places in the problem.
Rules for Significant Figures in
Mathematical Operations
• Addition and Subtraction: The number
of decimal places in the result equals
the number of decimal places in the
least accurate measurement.
6.8 + 11.934 =
18.734  18.7 (3 sig figs)
Sig Fig Practice #2
Calculation
Calculator says:
Answer
3.24 m + 7.0 m
10.24 m
10.2 m
100.0 g - 23.73 g
76.27 g
76.3 g
0.02 cm + 2.371 cm
2.391 cm
2.39 cm
713.1 L - 3.872 L
709.228 L
709.2 L
1818 lb + 3.37 lb
1821.37 lb
1821 lb
2.030 mL - 1.870 mL
0.16 mL
0.160 mL
*Note the zero that has been added.
Rounding Calculated Answers
 Multiplication and Division
Round the answer to the
same number of significant
figures as the least number of
significant figures in the
problem.
Rules for Significant Figures in
Mathematical Operations
• Multiplication and Division: # sig figs in
the result equals the number in the least
accurate measurement used in the
calculation.
6.38 x 2.0 =
12.76  13 (2 sig figs)
Other Special Cases
• What if your answer has less significant
figures than you are supposed to have?
– Calculator Example: 100.00 / 5.00 = 20
• Add zeros!
– 20 is 1 sf
– 20. is 2 sf
– 20.0 is 3 sf
Sig Fig Practice #3
Calculation
Calculator says:
Answer
3.24 m x 7.0 m
22.68 m2
100.0 g ÷ 23.7 cm3
4.219409283 g/cm3 4.22 g/cm3
23 m2
0.02 cm x 2.371 cm 0.04742 cm2
0.05 cm2
710 m ÷ 3.0 s
236.6666667 m/s
240 m/s
1818.2 lb x 3.23 ft
5872.786 lb·ft
5870 lb·ft
1.030 g x 2.87 mL
2.9561 g/mL
2.96 g/mL
Dimensional Analysis
Using the units to solve problems
Dimensional Analysis
•
•
•
•
Use conversion factors to change the units
Conversion factors = 1
1 foot = 12 inches (equivalence statement)
12 in = 1 = 1 ft.
1 ft.
12 in
• 2 conversion factors
• multiply by the one that will give you the
correct units in your answer.
Examples
•
•
•
•
11 yards = 2 rod
40 rods = 1 furlong
8 furlongs = 1 mile
The Kentucky Derby race is 1.25 miles.
How long is the race in rods, furlongs,
meters, and kilometers?
• A marathon race is 26 miles, 385
yards. What is this distance in rods and
kilometers?
Examples
• Because you never learned dimensional
analysis, you have been working at a
fast food restaurant for the past 35
years wrapping hamburgers. Each hour
you wrap 184 hamburgers. You work 8
hours per day. You work 5 days a week.
you get paid every 2 weeks with a
salary of $840.34. How many
hamburgers will you have to wrap to
make your first one million dollars?
• A senior was applying to college and wondered how
many applications she needed to send. Her counselor
explained that with the excellent grade she received
in chemistry she would probably be accepted to one
school out of every three to which she applied. She
immediately realized that for each application she
would have to write 3 essays, and each essay would
require 2 hours work. Of course writing essays is no
simple matter. For each hour of serious essay writing,
she would need to expend 500 calories which she
could derive from her mother's apple pies. Every
three times she cleaned her bedroom, her mother
would made her an apple pie. How many times would
she have to clean her room in order to gain
acceptance to 10 colleges?
Temperature and Density
Temperature
• A measure of the average kinetic energy
• Different temperature scales, all are talking
about the same height of mercury.
• We make measurements in lab using the Celsius
scale, but most chemistry problems require you
to change the temperature to Kelvin before
using in an equation.
Converting ºF to ºC and vice versa
Fahrenheit to Celsius
(°F - 32) x 5/9 = °C
Celsius to Fahrenheit
(°C × 9/5) + 32 = °F
0ºC = 32ºF
0ºC 32ºF
0ºC = 32ºF
100ºC = 212ºF
0ºC 32ºF
100ºC 212ºF
Converting oC to K and vice versa
Celsius to Kelvin
K = oC + 273.15
Kelvin to Celsius
oC
= K - 273.15
Density
•
•
•
•
•
Ratio of mass to volume
D = m/V
Useful for identifying a compound
Useful for predicting weight
An intrinsic property- does depend on
what the material is.
Density Problem
• An empty container weighs 121.3 g. Filled
with carbon tetrachloride (density 1.53
g/cm3 ) the container weighs 283.2 g.
What is the volume of the container?
Density Problem
• A 55.0 gal drum weighs 75.0 lbs. when
empty. What will the total mass be when
filled with ethanol?
density 0.789 g/cm3
1 gal = 3.78 L
1 lb = 454 g