Transcript Chapter 1 measurements

Summary of the three States of Matter
ALSO CALLED PHASES,
HAPPENS BY CHANGING THE TEMPERATURE AND/OR PRESSURE OF A SUBSTANCE.
GAS: total disorder; mostly empty space; particles have
complete freedom of motion (vibrational, rotational, &
translational); particles are very far apart.
 Cool or compress (increase pressure) a gas to make a liquid
 Heat or reduce pressure of a liquid to make a gas
LIQUID: Disorder; particles or clusters of particles are
free to move relative to each other (vibrational &
rotational); particles are relatively close to each other.
 Cool or compress (increase pressure) a liquid to make a solid
 Heat or reduce pressure of a solid to make a liquid
SOLID: order ranges from amorphous(slightly
disordered) to crystalline (ordered); particles are
essentially in fixed positions (vibrational only); particles
are close to each other.
PHASE TRANSITIONS
Consider the following phase changes and properly fill-in
the schematic shown below:
1. condensation
2. evaporation
3. freezing
4. melting
5. sublimation
6. deposition
SOLID
LIQUID
GAS
Physical Changes:
The substance or mixture does not alter in atomic
composition. Some Physical Changes are boiling,
evaporation, condensation, freezing, melting,
sublimation, and deposition.
Associated with Physical Changes are Physical
Properties like boiling or freezing point, density,
hardness, and state of matter.
H2O (l)  H2O (g)
Chemical Changes:
The substance changes in its atomic composition, the
atoms are rearranged and new substances are formed.
2 H2O (l)  2 H2 (g) + O2 (g)
Matter
Elements
Vocabulary to Know:
Matter
Molecule
Compound
Heterogeneous Mixture
Intensive Property
Chemical Property
Chemical Change
Compounds
Mixtures
Atom
Element
Homogeneous Mixture
Extensive Property
Physical Property
Physical Change
ANALYSIS OF MATTER
MATTER
Is it uniform?
YES
NO
HOMOGENEOUS MIXTURE
HETEROGENEOUS MIXTURE
blood, soil
Can it be separated
by physical methods?
YES
NO
HOMOGENEOUS MIXTURE
saltwater, rubbing alcohol
PURE SUBSTANCE
Can it be decomposed into
simpler substances using
chemical methods?
YES
COMPOUND
water
NO
ELEMENT
carbon
LABORATORY APPLICATIONS
Define the following:
1.
Filtration
2.
Distillation
3.
Chromatography
4.
Extraction
5.
Crystallization
ELEMENTS to MEMORIZE
Aluminum
Antimony
Argon
Arsenic
Barium
Beryllium
Boron
Bromine
Calcium
Carbon
Cesium
Chlorine
Chromium
Cobalt
Copper
Fluorine
Gallium
Germanium
Gold
Helium
Hydrogen
Iodine
Iron
Krypton
Lead
Lithium
Magnesium
Al
Sb
Ar
As
Ba
Be
B
Br
Ca
C
Cs
Cl
Cr
Co
Cu
F
Ga
Ge
Au
He
H
I
Fe
Kr
Pb
Li
Mg
Manganese
Mercury
Neon
Nickel
Nitrogen
Oxygen
Palladium
Phosphorus
Platinum
Plutonium
Potassium
Radium
Radon
Rubidium
Selenium
Silicon
Silver
Sodium
Strontium
Sulfur
Tin
Titanium
Tungsten
Uranium
Xenon
Zinc
Zirconium
Mn
Hg
Ne
Ni
N
O
Pd
P
Pt
Pu
K
Ra
Rn
Rb
Se
Si
Ag
Na
Sr
S
Sn
Ti
W
U
Xe
Zn
Zr
SCIENTIFIC METHOD
1. FACT : An observable event; indisputable evidence which does
not explain but simply is.
2. HYPOTHESIS: A guess to try to explain an observation.
3. EXPERIMENT: A systematic exploration of an observation or
concept.
4. THEORY: An explanation of the facts; it can be proven by
experiment and it confirms an hypothesis.
5. LAW: A theory which has undergone rigorous experimentation
and no contradiction can be found.
Note:
MODEL: A visual or mathematical device or method used to demonstrate a
theory or concept.
SIX STEPS OF THE SCIENTIFIC METHOD
1. State a problem
2. Collect Observations
3. Search for scientific laws to state a
relationship between observed facts
4. Form a hypothesis or a temporary
observation for an observed fact
5. Develop a theory that provides a general
explanation for observations made over
time
6. Modify a theory to fit new facts
ACCURACY vs.
PRECISION
• Accurate & precise
• inaccurate & imprecise
inaccurate but
precise
PRECISION AND ACCURACY
1. Precision – refers to the degree of reproducibility of a measured
quantity.
2. Accuracy – refers to how close a measured value is to the
accepted or true value.
Precise (not accurate)
Accurate (not precise)
Both Precise/Accurate
MEASUREMENTS
Scientific Notation
Many measurements in science involve either very
large numbers or very small numbers (#).
Scientific notation is one method for
communicating these types of numbers with
minimal writing.
GENERIC FORMAT: # . # #… x 10#
A negative exponent represents a number less than 1 and a
positive exponent represents a number greater than 1.
6.75 x 10-3 is the same as 0.00675
6.75 x 103 is the same as 6750
MEASUREMENTS
Significant Figures
I. All nonzero numbers are significant figures.
II. Zero’s follow the rules below.
1. Zero’s between numbers are significant.
30.09 has 4 SF
2. Zero’s that precede are NOT significant.
0.000034 has 2 SF
3. Zero’s at the end of decimals are significant.
0.00900 has 3 SF
4. Zero’s at the end without decimals are either.
4050 has either 4 SF or 3 SF
SIGNIFICANT DIGITS WORKSHEET
1. Nonzero integers. Nonzero integers always count as significant digits.
1492 has ______ significant digits
2. Zeros. There are three classes of zeros:
A. Zeros that precede all nonzero digits are NOT significant.
0.00162 has ______ significant digits
B. Zeros between nonzero digits are significant.
4.007 has ______ significant digits
C. Trailing zeros at the right end of the number are significant only if the number
contains a decimal point.
200 has ______ significant digits
200. has ______ significant digits
200.0 has ______ significant digits
200 has ______ significant digits
D. When writing in scientific notation, all digits count.
2.370 x 10-3 has ______ significant digit
3. Exact numbers can be assumed to have an infinite number of significant figures.
The “2” in the circumference of a circle (2r) formula has ______ significant digits
MEASUREMENTS
Significant Figures & Calculations
Significant figures are based on the tools used to make the
measurement. An imprecise tool will negate the precision of the
other tools used. The following rules are used when measurements
are used in calculations.
Adding/subtracting:
The result should be rounded to the same number of
decimal places as the measurement with the least decimal
places.
Multiplying/dividing:
The result should contain the same number of significant
figures as the measurement with the least significant
figures.
WORKSHOP INVOLVING SIGNIFICANT DIGITS
1. For addition and subtraction, the result has the same number of decimal places as the
least precise measurement used in the calculation.
Example:
12.11
18.0
+ 1.013
2. For multiplication and division, the number of significant figures in the result is the same
as the number in the measurement with the fewest significant digits.
(a) 4.56 x 1.4 = ________
(b) (4.12 + 3.636) = _____
5.7
NOTE:
Rules for Rounding:
1. In a series of calculations, carry the extra digits through to the final result, then round
off.
2. If the digit to be removed is:
A. less than 5, the preceding digit stays the same. For example, 2.32 rounds to 2.3.
B. equal to or greater than 5, the preceding digit is increased by 1. For example, 3.46
rounds to 3.5.
DIMENSIONAL ANALYSIS
Unit Conversions
Common SI Prefixes:
Factor
Prefix
Abbreviation
106
Mega
M
103
Kilo
k
102
Hecto
h
101
Deka
da
10-1
Deci
d
10-2
Centi
c
10-3
Milli
m
10-6
Micro

10-9
Nano
n
10-12
Pico
p
MEASUREMENTS - METRIC
1. The mass of a young student is found to be 87 kg.
How many grams does this mass correspond to?
2. How many meters are equal to 16.80 km?
3. How many cubic centimeters are there in 1 cubic
meter?
4. How many nm are there in 200 dm? Express your
answer in scientific notation.
5. How many mg are there in 0.5 kg?
MEASUREMENTS
Since two different measuring systems exist, a scientist must be able to
convert from one system to the other.
CONVERSIONS
Length
Mass
Volume
 1 in = 2.54 cm
 1 mi = 1.61 km
 1 lb.... = 454 g
 1 kg = 2.2 lb....
 1 qt = 946 mL
 1 L = 1.057 qt
 4 qt = 1 gal
 1 mL = 1 cm3
MEASUREMENTS - CONVERSIONS
1. The mass of a young student is found to be 87 kg. How
many pounds does this mass correspond to?
2. An American visited Austria during the summer
summer, and the speedometer in the taxi read 90 km/hr.
How fast was the American driving in miles per hour?
(Note: 1 mile = 1.6093 km)
3. In most countries, meat is sold in the market by the
kilogram. Suppose the price of a certain cut of beef is 1400
pesos/kg, and the exchange rate is 124 pesos to the U.S.
dollar. What is the cost of the meat in dollars per pound
(lb)?
(Note: 1 kg = 2.20 lb)
TEMPERATURE CONVERSIONS
1. Fahrenheit – at standard atmospheric pressure, the melting
point of ice is 32 F, the boiling point of water is 212 F, and the
interval between is divided into 180 equal parts.
2. Celsius – at standard atmospheric pressure, the melting point
of ice is 0 C, the boiling point of water is 100 C, and the interval
between is divided into 100 equal parts.
3. Kelvin – assigns a value of zero to the lowest conceivable
temperature; there are NO negative numbers.
T(K) = T(C) + 273.15
T(F) = 1.8T(C) + 32
Introduction to Density
 Density is the measurement of the mass of an
object per unit volume of that object.
d=m/V
 Density is usually measured in g/mL or g/cm3
for solids or liquids.
 Volume may be measured in the lab using a
graduated cylinder or calculated using:
Volume = length x width x height if a box or
V = r2h if a cylinder.
 Remember 1 mL = 1 cm3
DENSITY DETERMINATION
1. Mercury is the only metal that is a liquid at 25 C. Given that
1.667 mL of mercury has a mass of 22.60 g at 25 C, calculate its
density.
2. Iridium is a metal with the greatest density, 22.65 g/cm3. What
is the volume of 192.2 g of Iridium?
3. What volume of acetone has the same mass as 10.0 mL of
mercury? Take the densities of acetone and mercury to be 0.792
g/cm3 and 13.56 g/cm3, respectively.
4. Hematite (iron ore) weighing 70.7 g was placed in a flask
whose volume was 53.2 mL. The flask was then carefully filled
with water and weighed. Hematite and water combined weighed
109.3 g. The density of water is 0.997 g/cm3. What is the density
of hematite?
BASIC MATH
used in Chemistry 101
The following slides are basic math review. Please look
the slides over to refresh your memory. I assume you
already know this material and will not cover it in
class.
Some general math equations:
(1) The Generic equation for percent:
% = ( portion / total ) 100
(2) The difference/change between two measurements:
D X = Xfinal - Xinitial
Mathematical Review
Fractions & Decimals
– A fraction represents division, the numerator
is divided by the denominator.
» 2/3 is read as 2 divided by 3
– Proper fraction: numerator is smaller than
denominator. Example: 3/5
– Improper fraction: numerator is larger than
the denominator. Example: 5/3
– A decimal is a fraction with the division
carried out.
– A decimal is a fraction expressed in powers
of 10.
–
0 . 0
0
1
–
ones . Tenths hundredths thousandths
Mathematical Review
Algebraic Equations
– Variables are the symbols used to represent a measurement.
» For example; T is the variable for temperature while t is the
variable for time.
– To isolate one variable of an equation remember to divide if the
unwanted variable is on top and to multiply if the variable is on
the bottom. An asterisk * represents multiplication.
» A = B / C
» to isolate C first rearrange the equation to it will read C=?
Do this by multiplying both sides by C (since it is on the
bottom of a fraction (denominator).
» C * A = B * C / C note: C/C = 1
» C * A = B
» Now to isolate C we need to divide by A (it is on top of a
fraction; A/1 = A)
» C * A / A = B / A
» Remember: what ever you do to one side you must do it to
the other side.
» C = B / A
Mathematical Review
Algebraic Equations
– When multiplication & division is mixed with adding & subtracting,
try the multiplication or division first.
» (A - D) / (C + F) = B
» to solve for C, first rearrange the equation to it will read
C=? Do this by multiplying both sides by C + F (since it is
on the bottom of a fraction (denominator).
» (A - D) * (C + F) / (C + F) = B * (C + F)
» (A - D) = B * (C + F)
» Now to isolate C we need to divide by B
» (A - D) / B = B * (C + F) / B
» (A - D) / B = C + F
» Now you can subtract F from both sides.
» [(A - D) / B] - F = C + F - F
» [(A - D) / B] - F = C
» which is the same as C = [(A-D) / B] -F
If A = 8, D = 2, B = 3, & F = 7
then C must = [(8-2) / 3] - 7 = -5
Mathematical Review
Exponents
– An exponent is a number written as a superscript.
• X2 is X-squared or “X to the power of 2”
– The base (X) is multiplied by itself the number of times
represented in the exponent(superscript, 2 in this example).
• 23 or two cubed (2 is the base and 3 is the exponent)
• 23 is 2 * 2 * 2 = 4 * 2 = 8
– A positive exponent represents a large number (greater than
one).
• 1 x 103 is 10 *10 *10 = 1000 thousand
– A negative exponent represents a small number (less than one).
• 1 x 10-3 is (1/10) * (1/10) * (1/10) = 0.001 thousandths
– When multiplying numbers written with exponents, add the
exponents. If dividing then subtract the exponents.
• x4 * x6 = x
(4+6)
= x10 or (2 x 103)(3 x 106) = 6 x 10(3+6) = 6 x 109
• 2x6/7x3 = 0.2857 x(6-3) = 0.2857 x3