Chapter 1: Matter and Measurements

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Chapter 1 Matter and Measurements
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What is Chemistry?

• Biology vs.
• Chemistry vs.
• Physics

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What is Chemistry?

• Biology
• Physics
• Chemistry

The study of living organisms
The study of forces motion
The study of matter and its reactions and
properties
4
What is Chemistry?

• Chemistry is the study of CHANGES in the stuff
  around us.
• (We formally define stuff as matter!)

5
What is Chemistry?

• Think, pair, share
• What are all the chemicals you use in your daily
  life?

6
Review Scientific Methods

• 1. Hypothesis
• Suggested solution to a problem
• 2. Experiment
• A controlled method of testing a hypothesis
• 3. Data
• Organized observations
• a. Data is always reproducible.

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Review Scientific Methods

• 4. Scientific Law
• Statement which summarizes results of many
  observations and experiment
• a. Scientific laws explain WHAT is observed.
• Example of a scientific law
• 5. Scientific Theory
• Explanation that supports a hypothesis and which
  has been supported with repeated testing
• b. Scientific theories explain WHY something is
  observed.
• Example of a scientific theory

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Review Scientific Methods

• 6. Steps of the Scientific MethodReview
• a.
• b.
• c.
• d.
• e.
• f.

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What is Chemistry?

• Biology
• Physics
• Chemistry

The study of living organisms
The study of forces motion
The study of matter and its reactions and
properties
10
What is Chemistry?

• Chemistry is the study of CHANGES in the stuff
  around us.
• (We formally define stuff as matter!)

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Matter
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Matter
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Elements

• Type of matter that cannot be broken down into
  simpler, stable substances and is made of only
  one type of atom

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Compounds

• A pure substance that contains two or more
  elements whose atoms are chemically bonded

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Compounds

• Fixed compositions
• A given compound contains the same elements in
  the same percent by mass

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Compounds

• The properties of a compound are VERY DIFFERENT
  from the properties of the elements they contain
• Ex.) Sodium Chloride (NaCl) vs. Sodium Chlorine
• Sodium http//www.youtube.com/watch?vRAFcZo8dTcU
  
• http//www.youtube.com/watch?v92Mfric7JUc

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Electrolysis
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Mixtures

• A blend of two or more kinds of matter, each of
  which retains its own identity and properties
• Homogeneous
• Heterogeneous

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Homogeneous Mixtures

• Composition is the same throughout the mixture
• Examples salt water, soda water, brass
• A.k.a. a solution
• Solute in a solvent (salt dissolved in water)

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Heterogeneous Mixtures

• Non-uniform composition varies throughout the
  mixture

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Separating Mixtures

• Filtration

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Separating Mixtures

• Distillation

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Separating Mixtures

• Chromatography

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Scientific Measurements

• Chemistry is a quantitative science.
• This means that experiments and calculations
  almost always involve measured values.
• Scientific measurements are expressed in the SI
  (metric) system.
• This is a decimal-based system in which all of
  the units of a particular quantity are related
  to each other by factors of ten.

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Types of Observations and Measurements

• We make QUALITATIVE observations of reactions
  changes in color and physical state.
• We also make QUANTITATIVE MEASUREMENTS, which
  involve numbers.
• Use SI units based on the metric system

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SI System

• Definition modernized version of metric system
  uses decimals
• All units derived from base units larger and
  smaller quantities use prefixes with base unit
• Must memorize prefixes from nano- (10-9) to
  tera- (1012)

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Prefixes (see handout Ebook)

• You will need to memorize all of the prefixes
  (factors, names and abbreviations from
• 109 (giga-) to 10-9 (nano)!
• One example of a memory device
•  

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INSTRUMENTS UNITS

• Use SI units based on the metric system
• Length
• Mass
• Time
• Temperature

Meter, m
Kilogram, kg
Seconds, s
Celsius degrees, C Kelvins, K
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Length

• The standard unit of length in the metric system
  is the METER
• which is a little larger than a YARD.
• USING THE PREFIXES WITH LENGTH
• cm often used in lab
• km
• Gm

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Length

• Base unit METER
• Conversions
• 1 km1000 m 1 cm 10 -2 m
• 1 Gm 106 m

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Units of Length

• 1 kilometer (km) 1000 meters (m)
• 10-2 meter (m) 1 centimeter (cm)
• 102 meter (m) 1 hectometer (Hm)
• 1 nanometer (nm) 1.0 x 10-9 meter

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Volume

• THE COMMON UNITS OF VOLUME IN CHEMISTRY ARE
• the liter (milliliter) and cubic centimeter
  (cm3)
• THE COMMON INSTRUMENTS FOR MEASURING VOLUME IN
  CHEMISTRY ARE graduated cylinder buret
• Note that 1 cm3 1 mL (We will use this exact
  conversion factor throughout the year, so you
  will need to memorize it!)

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Mass

• THE COMMON UNIT OF MASS IN CHEMISTRY IS
• the gram (g) often used in lab
• Mass IS A MEASURE OF THE AMOUNT OF MATTER IN AN
  OBJECT
• Weight IS A MEASURE OF THE GRAVITATIONAL FORCE
  ACTING ON THE OBJECT. CHEMISTS OFTEN USE THESE
  TERMS INTERCHANGEABLY.
• 1000 g 1 kg 1 Mg 10 6 g

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Temperature Scales

• Fahrenheit
• Celsius
• Kelvin

TEMPERATURE IS THE FACTOR THAT DETERMINES the
direction of heat flow.
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Temperature Scales


Celsius
Kelvin
Fahrenheit
Boiling point of water
Freezing point of water
Notice that 1 kelvin degree 1 degree Celsius
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Temperature Scales
100 oF 38 oC 311 K
oF
oC
K
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SI System

• English Units (inches, feet, degrees F, etc.) are
  NEVER used to take measurements in the lab!

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Calculations Using Temperature
Fahrenheit/Celsius T (F) 1.8 t (C) 32
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Calculations Using Temperature

• Some calculations are in kelvins (especially
  important for Ch 5!!)
• T (K) t (C) 273.15 (273)
• Body temp 37 C 273 310 K
• Liquid nitrogen -196 C 273 77 K

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• Problem
• Example 1L.1 A baby has a temperature of 39.8oC.
  Express this temperature in oF and K.

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SI System Base Units

• ESTABLISHMENT OF THE INTERNATIONAL SYSTEM OF
  UNITS (SI)
• SI UNITS AS ESTABLISHED BY THE SI
• LENGTH meter (m)
• VOLUME cubic meter (m3)
• MASS kilogram (kg)
• TEMPERATURE Kelvin (K)

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Time

• Base unit SECOND (sec)
• Conversions
• only non-decimal base unit
• 60 sec 1 min 60 min 1 hr

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Precision and Accuracy in Measurements

• Precision vs. Accuracy
• Definitions
• Precisionhow close answers are to each other
  (reproducibility)
• Accuracyhow close answer is to accepted (true)
  value (agreement to accepted value)

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Precision and Accuracy in Measurements

• Percent Error - a way to calculate accuracy in
  the lab
• Equation
• Error Accepted Value Exp. Value x
  100
• Accepted
  Value

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Precision and Accuracy in Measurements

• Ex1.9 A student reports the density of a pure
  substance to be 2.83 g/mL. The accepted value is
  2.70 g/mL. What is the percent error for the
  students results?
• Equation
• Error Accepted Value Exp. Value x
  100
• Accepted
  Value

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Scientific Notation

• Exponential (Scientific) NotationSee Worksheet

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Significant Figures Why are they Important?

• Numbers in math no units, abstract, no context,
  can read calculator output exactly for answer.
• vs.
• Numbers in chemistry measurements include
  units.
• SIG FIGS WILL BE IMPORTANT THROUGHOUT THIS
  COURSE!

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Graduated Cylinder Example
http//learningchemistryeasily.blogspot.com/2013/0
7/precision-of-measurement-and.html
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What are significant figures?(aka sig figs)

• Significant figures are all the digits in a
  measurement that are known with certainty plus a
  last digit that must be estimated.
• With experimental values your answer can have
  too few or too many sig figs, depending on how
  you round.

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How Rounding Influences Sig Figs

• 1.024 x 1.2 1.2288Too many numerals(sig figs)
• Too precise
• 1.024 x 1.2 1Too few numerals(sig figs)
• Not precise enough

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Why This Concept is Important

• We will be adding, subtracting, multiplying and
  dividing numbers throughout this course.
• You MUST learn how many sig figs to report each
  answer in or the answer is meaningless.
• You must report answers on lab reports
  tests/quizzes with the correct number of sig figs
  (/- 1) or else you will lose points!!

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How Do We Find the Correct Number of Sig Figs In
an Answer?

• First, we will learn to count number of sig figs
  in a number. You must learn 4 rules and how to
  apply them.
• Second, we will learn the process for rounding
  when we add/subtract or multiply/divide. We will
  then apply this process in calculations.

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Rules for Counting Sig Figs

• Rule 1 Read the number from left to right and
  count all digits, starting with the first digit
  that is not zero. Do NOT count final zeros
  unless there is a decimal point in the number!

3 sig figs 4 sig figs 5 sig figs
23.4 234 0.234 2340 203 345.6 3.456 0.03456 34560 3405 678.90 6789.0 0.0067890 67008 60708
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Rules for Counting Sig Figs

• Rule 2 A final zero or zeros will be
  designated as significant if a decimal point is
  added after the final zero.

3 sig figs 4 sig figs 5 sig figs
2340 23400 234000 2340000 2340. 2000. 20000. 23400.
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Rules for Counting Sig Figs

• Rule 3 If a number is expressed in standard
  scientific notation, assume all the digits in the
  scientific notation are significant.

2 sig figs 3 sig figs 4 sig figs
2.3 x 102 2.0 x 103 2.30 x 102 2.00 x 103 2.300 x 102 2.000 x 103
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Rules for Counting Sig Figs

• Rule 4 Any number which represents a numerical
  count or is an exact definition has an infinite
  number of sig figs and is NOT counted in the
  calculations.
• Examples
• 12 inches 1 foot (exact definition)
• 1000 mm 1 m (exact definition)
• 25 students 1 class (count)

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Practice Counting Sig Figs

• How many sig figs in each of the following?
• 1.2304 mm
• 1.23400 cm
• 1.200 x 105 mL
• 0.0230 m
• 0.02 cm
• 8 ounces 1 cup
• 30 cars in the parking lot

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Answers to Practice

• How many sig figs in each of the following?
• 1.2304 mm (5)
• 1.23400 cm (6)
• 1.200 x 105 mL (4)
• 0.0230 m (3)
• 0.02 cm (1)
• 8 ounces 1 cup (infinite, exact def.)
• 30 cars in the parking lot (infinite,
• count)

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General Rounding Rule

• When a number is rounded off, the last digit to
  be retained is increased by one only if the
  following digit is 5 or greater.
• EXAMPLE 5.3546 rounds to
• 5 (ones place)
• 5.35 (hundredths place)
• 5.355 (thousandths place)
• 5.4 (tenths place)
• You will lose points for rounding incorrectly!

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Process for Addition/Subtraction

• Step 1 Determine the number of decimal places
  in each number to be added/subtracted.
• Step 2 Calculate the answer, and then round the
  final number to the least number of decimal
  places from Step 1.

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Addition/Subtraction Examples
Example 1 Round to tenths place. Example 2 Round to hundredths place. Example 3 Round to ones place.
23.456 1.2 0.05 -------------- 24.706 Rounds to 24.7 3.56 - 0.14 - 1.3501 --------------- 2.0699 Rounds to 2.07 14 0.735 12.0 -------------- 26.735 Rounds to 27
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Process for Multiplication/Division

• Step 1 Determine the number of sig figs in each
  number to be multiplied/divided.
• Step 2 Calculate the answer, and then round the
  final number to the least number of sig figs from
  Step 1.

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Multiplication/Division Examples
Example 1 Round to 1 sig fig. Example 2 Round to 2 sig figs. Example 3 Round to 3 sig figs.
23.456 x 1.2 x 0.05 -------------- 1.40736 Rounds to 1 3.56 x 0.14 x 1.3501 --------------- 0.67288984 Rounds to 0.67 14.0/ 11.73 -------------- 1.193520887 Rounds to 1.19
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PracticeWrite the answers to the following
computations using the correct number of sig figs
a. 129.0 g 53.21 g 1.4365 g b. 10.00 m -
0.0448 m c. 23.456 4.20 0.010 d. 17
22.73
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Important Rounding Rule

• When you are doing several calculations, carry
  out all the calculations to at LEAST one more sig
  fig than you need (I carry all digits in my
  calculator memory) and
• only round off the FINAL result.

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Use of Conversion Factors

• Also known as dimensional analysis or
  factor-label method (or unit conversions)
• Dimensional analysis/ Use of conversion factors
• Definition technique to change one unit to
  another using a conversion factor
• Ex.) in original unit x new unit
  New in new unit
• 
  original unit
• 
  

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Using Dimensional Analysis

• Express the quantity 1.00 ft in different
  dimensions (inches, meters).
• Conversion factors

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Using Dimensional Analysis

• Example 1L.5 Calculate the following single step
  conversions
• a. How many Joules are equivalent to 25.5
  calories if 1 cal 4.184 joules?
• b. How many liters gasoline can be contained in
  a 22.0 gallon gas tank if 3.785 L 1 gal?

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Using Dimensional Analysis

• Example 1L.6 The following multiple step
  conversions can be solved, knowing that 1 in
  2.54 cm. Convert the length of 5.50 ft to
  millimeters.

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Using Dimensional Analysis

• Example 1L.7 The average velocity of hydrogen
  molecules at 0oC is 1.69 x 105 cm/s. Convert
  this to miles per hour.

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Using Dimensional Analysis

• Example 1L.8 A piece of iron with a volume of
  2.56 gal weighs 168.04 lbs. Convert this density
  to scruples per drachm with the following
  conversion factors
• 1.00 L 0.264 gal, 1.000 kg 2.205 lb, 1.000
  scruple 1.296 g, 1.000 mL 0.2816 drachm.

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Using Dimensional Analysis Area Conversions

• Example 1L.14 Express the area of a 27.0 sq yd
  carpet in square meters.
• Conversion factors needed

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Using Dimensional AnalysisVolume Conversions

• Example 1L.15 Convert 17.5 quarts to cubic
  meters. (1 L 1.057 qt, 1 ft3 28.32 L)

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Properties of Substances

• 1. Every pure substance has its own unique set
  of properties that serve to distinguish it from
  all other substances.
• 2. Properties used to identify a substance must
  be intensive that is, they must be independent
  of amount.
• Extensive properties depend on the amount.
• Classify the following as either intensive (I) or
  extensive (E)
• a. density
• b. mass
• c. melting point
• d. volume

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Properties of Substances

• Density is an INTENSIVE property of matter, which
  does NOT depend on quantity of matter.
• Contrast with EXTENSIVE properties of matter,
  which do depend on quantity of matter.
• Examples of extensive properties mass and volume.

Brick
Styrofoam
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Chemical and Physical Properties

• Chemical property Observed when the substance
  changes to a new one.
• Example of a chemical property
• Copper reacts with air to form copper (II) oxide.
• Physical property Observed without changing the
  substance to a new one.
• Example of a physical property
• Water boils at 100oC.

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Physical Changes

• Physical changes do not result in a new
  substance
• boiling of a liquid
• melting of a solid
• dissolving a solid in a liquid to give a SOLUTION.

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Physical vs. Chemical Change

• Another name for a Chemical change is a chemical
  reaction change that results in a new substance.

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Example

• Classify the following as either physical (P) or
  chemical (C) changes
• a. ice melting
• b. gasoline burning
• c. food spoiling
• d. log of wood sawed in half

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Density

• Density is an INTENSIVE property of matter, which
  does NOT depend on quantity of matter.
• Contrast with EXTENSIVE properties of matter,
  which do depend on quantity of matter.
• Examples of extensive properties mass and volume.

Brick
Styrofoam
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DENSITY ReviewDefinition ratio of mass to
volume for an object
2.7 g/cm3
13.6 g/cm3
21.5 g/cm3
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• Sample Problem A piece of copper has a mass of
  57.54 g. It is 9.36 cm long, 7.23 cm wide, and
  0.95 mm thick. Calculate density (g/cm3).

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Density as a Conversion Factor

• Density is a bridge between mass and volume, or
  vice versa
• Volume (cm3) x density g mass (g)
• cm3
• Mass (g) ? density cm3 Volume (cm3)
• g

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SAMPLE PROBLEM Mercury (Hg) has a density of
13.6 g/cm3. What is the mass of 95 mL of Hg in
grams? In pounds?
Solve the problem using DENSITY AS A CONVERSION
FACTOR.
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• Ex1L.9 What is the density of Hg if 164.56 g
  occupy a volume of 12.1cm3?

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• Ex1L.10 What is the mass of 2.15 cm3 of Hg?

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• Ex1l.11 What is the volume of 94.2 g of Hg?

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• Example 1L.12 Given the following densities
  chloroform 1.48 g/cm3 and mercury 13.6 g/cm3 and
  copper 8.94 g/cm3. Calculate if a 50.0 mL
  container will be large enough to hold a mixture
  of 50.0 g of mercury, 50.0 g of chloroform and a
  10.0 g chunk of copper.

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• Example 1L.13 How many kilograms of methanol (d
  0.791 g/mL) does it take to fill the 15.5-gal
  fuel tank of an automobile modified to run on
  methanol?

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Density of Water

• Density of water changes with temperature
• (As water temperature changes, volume changes)
• Maximum density of water is at
• 4oC 0.999973 g/cm3
• (often rounded to 1.00 g/cm3)

91
Derived Units

• Definition derived from base units
• Example m/sec (unit of speed)
• Divide meters by seconds
• Volume examples
• m3 (m x m x m) or cm3 (cm x cm x cm)