I. The Study of Chemistry

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I. The Study of Chemistry

• A. The Molecular Perspective of Chemistry
• Studying the properties and behavior of matter
• the physical material of the universe
• It is anything that has mass and occupies space
• It is comprised of combinations of only about
  100very basic substance called elements.
• Provides a background to understanding the
  properties of matter in terms of atoms
• the almost infinitesimally small building blocks
  of matter
• Atoms can combine to form molecules (chemical
  combination of atoms)
• First Assignment Learn the names and symbol for
  the first 48 elements by Friday.

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Oxygen
Ethanol
Water
Aspirin
Carbon Dioxide
Ethylene glycol
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• Why Study Chemistry?
• It provides an important understanding of our
  world and how it works
• Improvement of health care
• Conservation of natural resources
• Protection of the environment
• Increased food production
• Development of new materials
• It is, by its very nature the central science
• Astronomy, atmospheric science, biology, geology,
  environmental science, medicine, physics,
  material science, and polymers
• The language of chemistry is a universal
  scientific language

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• How Do I Study Chemistry?
• It takes lots of practice--homework, reviewing
  notes, reading the text.
• It is different than some other disciplines
• MICHELANGELO Buonarroti, Italian painter,
  sculptor and architect (1475-1564). If a block
  of marble were at the front of this room I
  suspect we would select Michelangelo to teach us
  this art form.
• Antoine Lavoisier (1743-1794guillotined) is
  called the Father of Modern Chemistry but he
  thinks waters formula is HO. He knows of about
  a dozen elements, nothing about polymers, nuclear
  chemistry, ceramics, etc.
• You must study differently for chemistry--study
  nearly every day. The best predictor of your
  final grade is your grade on the first exam.

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B. Classifications of Matter

• States of Matter
• a gas, a liquid, or a solid
• states of matter differ in some of their simple
  observable properties
• gases (vapors) have no fixed volume or shape.
  They can be compressed to occupy a smaller volume
  or allowed to expand to occupy a larger volume
• liquids have a distinct volume independent of the
  container that they occupy. They assume the shape
  of the portion of the container they occupy
• solids have a definite shape and a definite volume

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• Pure Substances
• Most forms of matter are not chemically pure
• air
• gasoline
• Sidewalks
• Pure substances have distinct properties and a
  composition that does not vary from sample to
  sample
• All substances are elements or compounds
• elements cannot be decomposed into simpler
  substances
• carbon, helium, iron, oxygen, chlorine, etc.
• compounds are composed of two or more elements
• Water (H2O) is composed of two elements, hydrogen
  and oxygen
• Mixtures are combinations of two or more
  substances in which each substance retains its
  own chemical identity

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Atoms of an element
Molecules of an element
Mixture of elementsand a compound
Molecules of a compound
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• Elements
• At the present time 114 elements are known
• Elements vary widely in their abundance

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The symbol for each element consist of one or two
letters, with thefirst letter capitalized. You
will need to know the symbols and names for the
first 100 elements
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• Compounds
• Most elements interact with other elements to
  form compounds
• Hydrogen burns in oxygen to form water (one O and
  two H atoms)

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• The observation that the elemental composition of
  a pure compound is always the same is known as
  the law of constant composition (law of definite
  proportions)
• Mixtures
• Most of the matter we encounter consists of
  mixtures of different substances
• Each substance in a mixture retains its own
  chemical identity and properties
• Mixture composition can vary
• a cup of sweetened coffee. chocolate chip
  cookies. water found in nature, rocks, wood,
  cement, steel, etc.
• Some mixtures are uniform throughout
• Homogeneous mixtures or solutions
• Some mixtures are not uniform throughout
• Heterogeneous mixtures

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C. Properties of Matter

• Every substance has a unique set of properties
• characteristics that allow us to recognize and
  distinguish one substance from another
• Properties of matter can categorized as either
  physical or chemical
• Physical properties can be measured without
  changing the identity and composition of the
  substance
• Color, odor, density, melting point, boiling
  point, hardness, etc.
• Chemical properties describe the way a substance
  may change or react to form other substances
• A common chemical property is flammability

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• Some properties, such as temperature, melting
  point, and density, do not depend on the amount
  of sample being examined intensive properties.
• used to identify substances
• Some properties, such as mass and volume, do
  depend on the amount of sample being examined
  extensive properties.

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• Physical and Chemical Changes
• During physical changes a substances changes its
  physical appearance, but not its composition
• ice melting to become water
• water evaporating to become steam
• All state changes are physical changes
• During chemical changes (chemical reactions) a
  substances is changed into a chemically different
  substance
• propane burning to form carbon dioxide and water
• scrambling an egg
• a change in state will not revert the substance
  back to its original form

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• Separation of Mixtures
• Mixtures can be separated into their constituent
  components
• mixture components retain their own properties
• Take advantage of the differences in the
  properties
• Heterogeneous Mixtures
• visual differences
• magnetic differences
• state differences
• Homogeneous Mixtures
• Boiling point difference
• Polarity differences

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D. Units of Measurement

• Many properties of matter are quantitative that
  it, they are associated with numbers
• To say that the length of a pencil is 17.5 is
  meaningless
• The units used for scientific measurements are
  those of the
• Metric System

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• 1. SI Units
• 1960 - international agreement specifying a
  particular choice of seven metric units for
  scientific measurements
• SI Systéme International dUnités

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• Prefixes are used to indicate decimal factions
  ormultiples of various units
• Exponential notation is used to avoid ambiguity
  with regard to value certainty, see section 1.8.
  Learn these prefixes in Table 1.3.

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• 2. Length and Mass
• SI base unit of length is the meter (m)
• 1 m 100 cm 39.37 inches (slightly longer than
  a yard)
• Mass is a measure of the amount of material in an
  object.
• Mass is different from weight, which depends upon
  gravity
• SI base unit of mass is the kilogram (kg)
• This base unit is unusual because it uses a
  prefix, kilo-, instead of the word gram alone.
• Other units of mass are obtained by adding
  prefixes to the word gram

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• 3. Temperature
• We sense temperature as a measure of hotness and
  coldness
• Temperature determines the direction of heat flow
• Heat always flows spontaneously from a substance
  at higher temperature to one at lower
  temperature.
• The temperature scales commonly employed in
  scientific studies are the Celsius and Kelvin
  scales
• The Celsius scale, the everyday scale of
  temperature in most countries throughout the
  world, was originally based on the assignment of
  0C to the freezing poing of water and 100C to
  its boiling point at sea level.

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• The Kelvin Scale is the SI temperature scale and
  the SI unit of temperature is the Kelvin (K).
• Zero on this scale is - 273.15C, once believed
  to be the lowest attainable temperature. Because
  of this, 0K is known as Absolute Zero.
• Both the Celsius and Kelvin scales have
  equal-sized units that is, a kelvin is the same
  size as a degree Celsius.
• K C 273.15
• The freezing point of water, 0C, is 273.15 K
• Notice that the degree symbol () is not used
  with temperatures on the Kelvin scale.

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• The common temperature scale in the US is the
  Fahrenheit scale
• Freezing point of water 32F
• Boiling point of water 212F
• C 5/9 (F - 32) or F 9/5 (C) 32

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• 4. Volume
• The volume of a cube is given by its length
  cubed, (length)3
• The basic SI unit of volume is the cubic meter
• Another common unitof volume is the liter
  (L),which equals a cubicdecimeter, dm3.
• It is slightly larger than aquart
• There are 1000 milliliters (mL)in a liter
• Each milliliter is the same volume as acubic
  centimeter

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Devices Used Most Frequently in Chemistry to
Measure Volume
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• 5. Density
• Widely used to characterize substances
• Amount of mass in a unit volume of the substance

• Densities of solids and liquids are commonly
  expressed in grams per cubic centimeter (g/cm3)
  or grams per milliliter (g/mL)
• The density of water is 1.00 g/ml
• Most substances change volume when heated or
  cooled Densities are temperature dependent
• Assume 25C unless otherwise noted

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E. Uncertainty in Measurement

• There are two kinds of numbers in scientific
  work
• Exact numbers (those whose values are known
  exactly)
• 12 eggs in a dozen, exactly 1000 g in a
  kilogram, and exactly 2.54 cm in an inch
• and inexact numbers (those whose values have
  some uncertainty)
• numbers obtained by measurement
• Uncertainties always exist for in measured
  quantities

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• 1. Precision and Accuracy
• Precision is a measure of how closely individual
  measurements agree with one another
• Accuracy refers to how closely individual
  measurements agree with the correct, or true
  value.

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• 2. Significant Figures
• Precision of a measured number is indicated using
  the concept of significant figures.
• Those digits in a measured number (or result of a
  calculation with measured numbers) that include
  all certain digits plus a final one have some
  uncertainty.

• Three measurements (9.12, 9.11, and 9.13 cm)
• Avg 9.12 First two digits (9.1) are
  certain, the next digit is estimated, so it has
  some uncertainty.

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• The greater the number of significant figures,
  the greater the certainty implied for the
  measurement
• In any measurement that is properly reported, all
  nonzero digits are significant. Zeros can be used
  either as part of the measured value or merely to
  locate the decimal point.
• Zeros between nonzero digits are always
  significant
• Zeros starting a number are never significant
• Zeros ending a number to the right of the decimal
  point are always significant
• Zeros ending a number to the left of the decimal
  point may or may not be significant
• Exponential notation is the solution

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• Significant Figures in Calculations
• Suppose that 0.0634 g of a compound
• will dissolve in 25.31 g of water. Howmany
  grams will dissolve in 100 g ofwater?

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Significant Figure Calculation Rules
Multiplication and Division. When multiplying or
dividing measuredquantities, give as many
significant figures in the answer as there arein
the measurement with the least number of
significant figures.
If the leftmost digit to be dropped is 5, round
the last significantfigure up, otherwise simply
drop the nonsignificant digits.
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Significant Figure Calculation Rules
Addition and Subtraction. When adding or
subtracting measuredquantities, give the same
number of decimal places in the answer asthere
are in the measurement with the least number of
significant figures.
If the leftmost digit to be dropped is 5, round
the last significantfigure up, otherwise simply
drop the nonsignificant digits.
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Dimensional Analysis A. This process is
called 1. Dimensional analysis (we will use this
name) 2. Factor-label method, Unit conversion
method, or the Unit Factor Method

• Dimensional analysis is a problem solving aid
  used to help ensure that the solutions to
  problems yield the proper units.
• It provides a systematic way of solving any
  numerical problems and of checking solutions for
  possible errors.
• The key to using dimensional analysis is the
  correct use of conversion factors to change one
  unit into another.
• This method will work for many chemical
  calculations. When it is not convenient you must
  still cancel the units to be certain your answer
  has the proper dimensions.

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• 1. A conversion factor is a fraction whose
  numerator and denominator are the same quantity
  expressed in different units.

• 2. Two important mathematical realities
• Doing the same thing to both sides of an equation
  does not change the relationship
• Multiplying one does not change the anything

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• 3. Since conversion factors are the number one
  in a different form, multiplying a measurement
  and its units by any number of conversion factors
  changes the value and the units but not the
  reality of the measurement itself.
• How many seconds in a century?

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B. Conversion Factors Constructed from any two
terms that are equal to each other. If two
quantities are equal and if one is then divided
by the other, the quotient equals one and it is
called a Unit Factor or a Conversion Factor.
Remember any quantity can be multiplied by one
without changing the result. Examples of
Conversion Factors, Memorize these. a. 1.00 inch
2.54 cm Length b. 454 g 1.00
lb Mass c. 1.00 L 1.057 qt Volume d. 1.00
mL 1.00 cm3 1.00 cc (learn these first four
conversions) e. 1.00 mol Mg 24.3 g of Mg f.
Many others that you will learn.
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• Volume, like other concepts that will be
  encountered throughout the semester, requires the
  use of relationships that necessitate the raising
  of numbers to a power, cubing in this specific
  case
• It is imperative to remember to raise both the
  number and the units to the appropriate power and
  not just the units
• 1 in 2.54 cm 1 in3 ? 2.54 cm3
• 1 in3 (1 in)3 (2.54 cm)3 16.39 cm3

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C. Recipe for Dimensional Analysis 1. On the far
right, write down the units of the answer. 2.
Analyze the information given or known (the
units) to select the proper starting quantity.
This information may be from the problem, the
periodic chart, some physical law, or something
you learned. 3. Analyze the dimensions (units)
of the answer and the dimensions (units) of the
starting quantity to determine the proper unit
factors (conversion factors) to convert the units
given into the units of the answer. 4. This may
require several unit factors. 5. Cancel the
units and do the numerical calculation.
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Examples 1. How many inches in 2.57
ft? DO ANS 2. How many nm in 3.72
yds? DO ANS
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3. A major contributor to global environmental
pollution is from coal burning power plants.
Carbon dioxide, a greenhouse gas, is the major
product and all coal contains various amounts of
sulfur that eventually contributes to acid rain.
A typical coal burning power plant burns 2500
tons of coal per day, with a sulfur content of
3, and the density of solid bituminous coal is
1346 kg/m3 (solid anthracite coal has a density
of 1506 kg/m3). a. What volume, in cubit
feet, of bituminous coal is burned by a typical
coal burning power plant each day? ANS
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• How many grams of sulfur are burned each day?
• ANS
• If a railroad car holds 100 tons of coal how many
  car loads are burned each day?
• ANS

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• 4. Determine the volume in ft3 of a piece of Pb
  that has a mass of 13.4 kg, if the density is
  11.2 g/cm3. DO
• ANS
• 5. A gallon of milk weighs 8.00 lbs.
• a. What is the mass of one pint in grams? DO
• ANS
• b. What is the density of milk in g/mL? DO
• ANS

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Douglas Isbell/Don Savage Headquarters,
Washington, DC Nov. 10,
1999 (Phone [phone])
Embargoed until 2 p.m. EST RELEASE 99-134
MARS CLIMATE ORBITER FAILURE BOARD RELEASES
REPORT, NUMEROUS NASA ACTIONS UNDERWAY IN
RESPONSE Wide-ranging managerial and
technical actions are underway at NASA's Jet
Propulsion Laboratory, Pasadena, CA, in response
to the loss of the 125 million Mars Climate
Orbiter and the initial findings of the mission
failure investigation board, whose first report
was released today. "The 'root cause' of the
loss of the spacecraft was the failed translation
of English units into metric units in a segment
of ground-based, navigation-related mission
software, as NASA has previously announced," said
Arthur Stephenson, chairman of the Mars Climate
Orbiter Mission Failure Investigation Board. "The
failure review board has identified other
significant factors that allowed this error to be
born, and then let it linger and propagate to the
point where it resulted in a major error in our
understanding of the spacecraft's path as it
approached Mars. WEB http//mars.jpl.nasa.gov/
msp98/news/mco991110.html