Introduction to Chemistry and Matter and Energy

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Introduction to Chemistryand Matter and Energy

• Summers over
• Hang tight
• Its going to be an exciting ride!

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What is Chemistry?

• What is Matter?
• What is Non-Matter?

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Why Study Chemistry?

• Central, fundamental science.
• Other sciences used chemistry as their backbone.
• Health care, conservation of natural resources,
  protection of the environment, food production,
  clothing, manufacturing, production of shelter,
  etc

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Scientific laws are the evidence used to support
a conclusion.  Scientific hypotheses and
theories are our best attempts at explaining the
behavior of the world, in ways that can be tested
by further experiment. 
We don't prove theories (and hypotheses) true. 
We just use the observations to convince
ourselves (and others) that we have a good idea. 
Scientists have a lot of confidence in scientific
theories, because they know there is a lot of
evidence to back them up.
Scientific law a generalized description,
usually expressed in mathematical terms, which
describes the empirical behavior of
matter. Scientific laws describe things.  They do
not explain them.
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Measurement and Scientific Notation

• Measurement define qualitative properties of a
  substance.
• Often in science, measurements require very large
  or very small numbers.
• Scientific notation a number between 1 and 10
  multiplied by 10 raised to a power.
• The number of places the decimal point has moved
  determines the power of 10. If the decimal point
  has moved to the ______then the power is _______,
  to the _____, ___________.
• e.g. 602,000,000,000,000,000,000,000.0
• e.g. 0.00524

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Convert the following

• Convert to scientific notation
• 24500
• 356
• 0.000985
• 0.222
• 12200

• Convert to non-scientific notation numbers
• 4.2 X 10-3
• 2.15 X 104
• 3.14 X 10-6
• 9.22 X 105
• 9.57 X 102

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Mathematics of SciencePrecision, Accuracy, and
Significant Figures

• No measurement of a physical quantity is
  absolutely certain.
• All measurements include a certain degree of
  uncertainty
• Causes of uncertainty

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• Precision
• Accuracy

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Consider three sets of data that have been
recorded after measuring a piece of wood that was
exactly 6.000 m long.

• Which set of data is the most accurate?
• Which set of data is the most precise?

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• Significant figures- measurements include one
  uncertain figure in addition to those known with
  certainty.
• Rules for Significant Figures
• 1.    All digits 1-9 are significant
• I.e.- 129
• 2.  Zeros between sig. Figs. are always
  significant
• I.e.- 5007
• 3.  Trailing zeros in a number are significant
  only if the number contains a decimal pt.
• I.e.- 1000.0
• 100
• 4.  Zeros in the beginning of a number whose only
  function is to place the decimal point are not
  significant.
• I.e.- 0.0025
• 5.  Zeros following a decimal sig fig are
  significant.
• I.e.- 0.000470
• 6.  A bar over a zero indicates significance

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• Atlantic Pacific Rule
• If a decimal is present, count from the Pacific
  side.
• If a decimal is absent, count from the Atlantic
  side.
• Start counting from the first nonzero digit you
  find, and count every digit including zero
  thereafter!

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Significant Figures Practice

• Determine the number of significant figures in
  the following
• 250.7
• 0.00077
• 1024
• 4.7 X 10-5
• 3400000
• 500.0
• 0.230970
• 0.03400
• 0.34030
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• Calculate the following to the correct number of
  sig. figs.
• 34.5 X 23.46
• 123/3
• 2.61X10-1 X 356
• 21.78 45.86
• 23.88887-11.2
• 6-3.0
• 32.559 X 34.555
• 4433-1187
• 1.2 X 4.3
• 8.08 21.98

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Rules for Calculations Using Significant Figures

• Multiplication and Division- limit and round to
  the least number of sig figs in any of the
  factors.
• I.e.- 144.6 X .0023 ?
• Addition and Subtraction Rule- limit and round to
  least number of decimal places in any of the
  numbers that make up the problem.
• I.e.- 5.42 g 131.1 g ?

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SI Units- preferred metric units used in science.
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Metric Conversion
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Unit Conversion Using Dimensional Analysis

• Write the term to be converted- both the number
  and the unit.
• 0.0342g
• Write the conversion formulas
• 1 g 1000 mg
• Make a fraction of the conversion formula such
  that
• a. if the unit in step 1 is in the numerator,
  the same unit in step 3 must be in the
  denominator
• b. if the unit in step 1 is in the denominator,
  the same unit in step 3 must be in the numerator.
• Note since the numerator and the denominator are
  equal, the fraction must be equal to 1.

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• 4. Multiply the term in step 1 by the fraction in
  step 3. Since the fraction equals 1, you can
  multiply it without changing the size of the
  term.
• 5. Check math by canceling your units.

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Convert the following quantities using the
following equivalence statements. Show work!

• 1 m 1.094 yd 1mile 1760 yd 1kg 2.205lbs
• 30.0 m to miles
• 1500 yd to miles
• 206 miles to m
• 34 kg to lbs
• 34 lb to kg

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Matter

• All matter is composed of 100 or so _____________
• A substance that cannot be separated into simpler
  substances by a chemical change simplest type of
  pure substance.
• The building block of matter is the _________
• The smallest particle of an element that retains
  the chemical identity of the element.
• Atoms can combine to form ___________

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Elements and Compounds Pure Substances

• Element

• Compound

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Mixtures? Mixtures can be

• Heterogeneous

• Homogeneous

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Classification of Matter
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Properties of Matter

• Physical
• Characteristics can be observed without altering
  the identity of the substance
• Volume
• Mass
• Maleability, ductility, conductivity etc

• Chemical
• Characteristics cannot be observed without
  altering the identity of the substance
• Flammability
• Tendency to corrode
• Reactivity
• Etc

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Changes Matter Can Undergo

• Physical Change
• Solid ? Liquid Melting
• Liquid ? Gas Boiling or Evaporating
• Gas ? Liquid _____________
• Solid ? Gas _____________
• Gas ? Solid _____________
• Liquid ? Solid Freezing, solidifying

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Changes Matter Can Undergo

• Chemical Change
• Rusting, rotting, burning, chemical reaction

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Distinguishing Chemical from Physical Change

• Did the change produce a different substance?
• Was there a color change?
• Is there a different density?
• Is there a melting or boiling point change?
• Did something precipitate out of solution?
• Did a gas or smoke form?

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EnergyRemember Matter- anything that has mass
and takes up space. Energy is the other stuff
of the universe.

• The capacity to do work (the ability to move or
  change matter)
• 1. Kinetic-
• 2. Potential-
• 3. Radiant/ electromagenetic- heat and light.
  We are mainly concerned with heat for this unit.

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Heat Vs Temperature

• Heat
• Energy due to _____________________
• Symbol ____
• Units ___________
• Does work by _____________________________________
  __________________________
• Flows from hot areas to cold areas
• Calorimetry

• Temperature
• A measure of _____________________________________
  ___________________________________
• Refers to the intensity of heat in an object
• Symbol T
• Units _______________
• Change in T Tf Ti D T
• NOT a form of energy but is a predictor of heat
  flow

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Heat Vs Temperature

• Keep in mind
• Objects can be the same temperature but have
  different amounts of heat energy
• Heat is dependent on MASS

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Temperature Scales

• 0 K ? absolute zero all molecular motion stops
• 0 K ? theroretical temperature not yet obtained
  (within a millionth of a degree)
• Closer to absolute zero ? atoms move more and
  more slowly much more difficult to remove heat
• Sig figs and temperature because the Celsius
  temperature is a continuum with both positive and
  negative values, a temperature measurement of 00C
  has 1 sig fig (0.10C 2 sig figs)

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Temperature Scale Conversions
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Significant Temperatures for Phases of Water
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Kinetic Molecular Theory

• 1.
• (atoms / molecules)
• The basic principles of KMT are theoretical and
  begin to break down under certain circumstances?
  KMT is better at describing matter in higher
  energy states (gases, for example)

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States/ Phases of Matter
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Calorimetry

• Physical and chemical changes are normally
  accompanied by energy changes.
• Energy changes in a laboratory setting are
  measured using a calorimeter.

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Types of Energy Changes

• If heat is consumed during the change, then the
  process/change/reaction is said to be
  ___________________.

• If heat is produced during a change, then the
  process/change/reaction is said to be
  ________________.

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Law of Conservation of Energy

• Within a closed system, energy transforms from
  one type to another.
• ______________________________________.
• Example electricity lights a bulb resistance
  builds up in the tungsen wire, it glows and gives
  off light and heat the total energy in the heat
  and light the energy in the electricity.
• Example when heat is added to water on a hot
  plate, that heat energy is absorbed by the water
  molecules, which move faster and faster
  (increased kinetic energy? higher temperature)

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Law of Conservation of Matter

• Matter can only be transformed during chemical
  and physical changes.
• ___________________________________________.
• Example when ice melts to make water during a
  phase change
• Example when two chemicals are mixed
• On our large scale, we see matter and energy as
  separate, but matter and energy interconvert at
  the subatomic level according to Einsteins
  Theory of Relativity Emc2)

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Calorie ProblemsTheoretical values for energy
changes during the heating or cooling of a
substance, or during a phase change, can be
calculated using three basic equations.

• DURING HEATING OR COOLING
• c specific heat for water 4.18 J/goC
• m mass of sample
• DT change in temperature of sample in oC

• DURING A PHASE CHANGE
• (freezing/ melting)
• (evap / condense)
• M mass of sample
• Hf heat of fusion (for water 334 J/g)
• Hv heat of vaporization (for water 2260 J/g)

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Why do we add propylene glycol (antifreeze) to
our cars radiators?

• The value of Q for any substance can be
  calculated, but note that each substance has
  unique values for specific heat capacity (c),
  heat of fusion (Hf), and heat of vaporization
  (Hv). Think about it its easier to raise the
  temperature of some substances than others.

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• High specific heat capacity (c) a large amount
  of energy must be added in order to increase the
  temperature.
• Water(l) 4.18 J/(gK)
• Low specific heat capacity (c) a small amount
  of energy must be added in order to increase the
  temperature.
• Iron(s) 0.129 J/(gK)

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Q mcDT

• How much heat is required to raise the
  temperature of 10.0 g of water from 5oC to
  25.0oC?
• What will be the temperature change if 418 J of
  heat are added to 25 g of water?

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• Q mHf
• How much heat is needed to melt 5.0 g of water?
• Q mHv
• How much water can be vaporized by 3135 Joules?