Pure substance by yaofenji

VIEWS: 0 PAGES: 53

									Chapter 3
  Matter and
   Energy
     Assigned Problems
Recommended
 Exercises: 1-27 (odd)

Required
 Problems: 29-85 (odd)

 Cumulative Problems: 87-105 (odd)

Optional
 Highlight Problems: 107-113 (odd)
        What Is Matter?
   Matter is any material that has mass and
    occupies space
   Matter is made up of small particles
      Atoms

      Molecules

   Includes all things (living and nonliving) such
    as plants, soil, and rocks and any material we
    use such as water, wood, clothing, etc.
   Classifications (of a sample of matter) is based
    on whether its shape and volume are definite
    or indefinite
           Classifying Matter According to Its State
   Solid
     Has a rigid, definite shape and definite volume
          Crystalline solids have a regular, internal long-range
           order of atoms, ions, or molecules
          Amorphous solids have no long-range order of atoms,
           ions, or molecules in their lattice structure
   Liquid
      Has an indefinite shape and a definite volume.

      It will take the shape of the container it fills

   Gas
      Has an indefinite shape and an indefinite volume.

      It will take the shape and completely fill the volume
       of the container it fills
      Gases are compressible
   Water is one of the few substances commonly found
    in all three physical states




             (a) Solid (Ice)   (b) Liquid (Water)   (c) Gas (Steam)
          Fig3_2
          Classifying Matter by Its Composition
   Matter can also be classified in terms of its chemical
    composition

                            Matter


        Pure Substance                   Mixture
       Pure Substances: Composed of only one atom or molecule
       Mixtures: Composed of two or more different atoms or
        molecules combined in various proportions
    Pure Substances
   Matter that has a definite and constant
    composition is a pure substance
   Composed of the same substance; no
    variation
   6 million pure substances have been isolated:
    112 are elements, the rest are compounds
   The two classifications of pure substances:
       Elements: e.g., a pure sample of copper or a pure
        sample of gold (one type of atom)
       Compounds: e.g., example, a pure sample of
        water or a pure sample of sucrose (one type of
        molecule)
        Pure Substances
   Elements
     Substances which can not be broken down into
      simpler substances by chemical reactions
     Fundamental substances

   Compounds
     Two or more elements combined chemically in a
      definite and constant ratio
     Can be broken down into simpler substances

     Most of matter is in the compound form
         Pure Substances
   Compounds
      Results from a chemical combination of two or more

       elements
      Can be broken down into elements by chemical processes

      Properties of the compound not related to the properties of

       the elements that compose it
          Water is composed of hydrogen and oxygen gases

           (combined in a 2:1 ratio)
     Mixtures
   Something of variable composition
   Result from the physical combination of two
    or more substances (elements or
    compounds)
   Made up of two or more types of
    substances physically mixed
   Each substance retains its identity because
    the substances are not chemically mixed
   Mixtures of the same components can vary
    in composition
        Mixtures
   Mixtures can be classified by the (visual)
    uniformity of the mixture’s components
   Homogeneous mixture:
      Same uniform composition throughout

      Not possible to see the two substances
       present
   Heterogeneous mixture:
      Composition is not uniform throughout the
       sample.
      It contains visibly different parts or phases
        Mixtures
   Homogenous mixtures
     A sugar solution

     14 karat gold, a mixture of copper and gold

     Air, a mixture of gases (oxygen, nitrogen)

   Heterogeneous mixture
     Oil and vinegar

     Raisin cookies

     Sand

   Pure substance
     e.g. copper (all elements are pure substances)
        Compounds vs. Mixtures
   Compounds are not mixtures
     Cannot be separated by a physical process

     Can be subdivided by a chemical process into two
      or more simpler substances
     Simpler substances have different properties from
      the compound

   Mixtures
     Unlike compounds, mixtures can be separated by
      a physical process
     Each substance in a mixture retains its own
      individual properties
       Classification of Matter

                             M atte r
                             Matter


      Pure Substance s
       Pure Substances                          Mixture
                                                Mixture


Ele me nts       Compounds          Homoge ne ous
                                     Homogeneous     He te r oge ne ous
                                                      Heterogeneous
Elements         Compounds
                                       Mixture
                                       M ixture           Mixture
                                                         M ixture


                                         Physical Methods
    Chemical Methods                       Pure Substance s
                                            Pure Substances
          Physical and Chemical Properties
   Various kinds of matter are differentiated by
    their properties
       Properties are the characteristics of a substance
        used to identify and describe it
   Two general categories:
     Physical Properties

     Chemical Properties

   Properties can be:
     Directly observable (physical)

     The interaction of the matter with other
      substances (chemical)
        Physical and Chemical Properties:
        Physical Properties
   A physical property is a characteristic of
    a substance that can be observed
    without changing a substance into
    another substance
     Characteristics of matter that can be directly
      observed or measured without changing its
      identity or composition
     Color, odor, physical state, density, melting
      point, boiling point
          Physical and Chemical Properties:
          Chemical Properties
   A chemical property describes the way a substance
    undergoes a change or resists change to form a
    new substance
   Properties that matter exhibits as it undergoes
    changes in chemical composition:
       Objects made from copper will turn green when
        exposed to moist air for long periods
       Gold objects will resist change when exposed to
        moist air for long periods
       Water will react strongly with sodium metal and
        produce hydrogen gas.
      Classifying Properties
    The boiling point of ethyl alcohol is 78 °C
     Physical property – describes an inherent
      characteristic of alcohol, its boiling point
   Diamond is very hard
     Physical property – describes inherent
      characteristic of diamond – hardness
   Sugar ferments to form ethyl alcohol
     Chemical property – describes behavior
      of sugar, ability to form a new substance
      (ethyl alcohol)
        Physical and Chemical Changes
   Changes in matter are regular occurrences:
      Food is cooked

      Paper is burned

      Iron rusts

   Matter undergoes changes as a result of the
    application of energy
   Changes in matter are also categorized as two types:
      Physical

      Chemical
          Physical and Chemical Changes
   A physical change is a process that alters the
    appearance of a substance but does not
    change its chemical identity or composition
       Folding aluminum foil sheets
       Crushing ice cubes
   No new substance is formed
     Most common is a change of a substance’s
      physical state
        The freezing of liquid water

        Evaporation of liquid water to steam
        Physical and Chemical Changes
   A chemical change is a process that changes
    the chemical composition of a substance
      Also called a chemical reaction

      (At least) one new substance is produced

      Wood burning, iron rusting, alka-seltzer
       tablet reacting with water
      During a chemical change, the original
       substance is converted into one or more
       new substances with different chemical and
       physical properties
        Classifying Changes
   Melting of snow
       Physical change – a change of state but
        not a change in composition
   Burning of gasoline
       Chemical change – combines with
        oxygen to form new compounds
   Rusting of iron
       Chemical change – combines with
        oxygen to form a new reddish-colored
        substance (ferric oxide)
        Classifying Changes

   Iron metal is melted
       Physical change – describes a state
        change, but the material is still iron
   Iron combines with oxygen to form rust
       Chemical change – describes how iron
        and oxygen combine to make a new
        substance, rust (ferric oxide)
   Sugar ferments to form ethyl alcohol
       Chemical change – describes how sugar
        forms a new substance (ethyl alcohol)
        Conservation of Mass
   During a physical change: No new substance is
    formed
   During a chemical change: At least one new
    substance is formed
   Whether it is a physical or chemical change, the
    amount of matter remains the same
   The law of conservation of mass states that the total
    mass of materials present after a chemical reaction is
    the same as the total mass before the reaction
   Matter is never created nor destroyed
          Energy
   Two major components of the universe:
       Matter
       Energy
   Energy is the capacity to do work or produce heat
     Electrical, radiant, mechanical, thermal, chemical,
      nuclear
     Nearly all changes that matter undergoes involves
      the release or absorption of energy
   Chemistry is the study of matter
     The properties of different types of matter
     The way matter behaves when influenced by other
      matter and/or energy
         Energy
   Energy is the part of the universe that has the ability
    to do work
   Energy can be converted from one form to another
    but it is neither created nor destroyed (the law of
    conservation of energy)

   Energy has two classifications
     Potential: Stored energy

     Kinetic: Motion energy



   All physical changes and chemical changes involve
    energy changes
           Energy
   Potential energy:
       Determined by an objects position (or composition)
       Chemical energy (also potential energy) is stored in the bonds
        contained within a molecule. It is released in a chemical reaction.
   Kinetic energy
       Energy that matter acquires due to motion
       Converted from the potential energy
   All forms of energy can be quantified in the same units
       Units of Energy
   The joule (J) is the SI unit of heat energy
   The calorie (cal) is an older unit used for
    measuring heat energy (not an SI unit)
     The amount of energy needed to raise the
      temperature of one gram of water by 1°C
         4.184 J = 1 cal        1 kcal = 1000 cal

   The Cal is the unit of heat energy in
    nutrition
                 1 Cal = 1000 cal = 1 kcal
           Energy: Chemical and Physical Change
   All physical changes and chemical
    changes involve energy changes which
    convert energy from one form to another
   In terms of a chemical reaction the
    universe is divided into two parts:
       The system (chemical reaction)
       The surroundings (everything else)
       The potential energy differences between the
        reactants and products determine whether
        heat flows into or out of a chemical system
       Whether a reaction is exothermic or endothermic
        depends on how the potential energy of the
        products compares to the PE of the reactants
           Energy: Chemical and Physical Change
   Chemical systems with low
    potential energy tend to change
    in order to increase their                             products

    potential energy by the
    absorption of heat
                                          reactants
       Chemical reactions that absorb
        heat are called endothermic
   Chemical systems with high
    potential energy tend to change
                                               reactants
    in order to lower their potential
    energy by the release of heat
       Chemical reactions that release
        heat are called exothermic                                    products
          Temperature
   Temperature is a number related to the
    average kinetic energy of the molecules of a
    substance
   In a substance, the temperature:
       measures the hotness or coldness of an object
       measures the average molecular motions in a
        system
       relates (directly) to the kinetic energy of the
        molecules
          Temperature
   Fahrenheit Scale, °F
     Used in USA

       Water’s freezing point = 32°F, boiling point = 212°F
   Celsius Scale, °C
     Used in science (USA) and everyday use in
      most of the world
     Temperature unit larger than the Fahrenheit

       Water’s freezing point = 0°C, boiling point = 100°C
        Temperature
   Kelvin Scale, K
     SI Unit
     Used in science

     Temperature unit same size as Celsius

     Water’s freezing point = 273 K (0 ºC),
      boiling point = 373 K (100 ºC)
     Absolute zero is the lowest temperature
      theoretically possible
     No negative temperatures
        Converting °C to °F
   Units are different sizes
     Fahrenheit scale: 180 degree intervals
      between freezing and boiling
     Celsius scale: 100 degree intervals
      between freezing and boiling
          180 F 9 F 1.8 F
                    
          100 C 5 C 1 C
             212ºF   100ºC             Boiling point


180                          100
Fahrenheit                   Celsius
degrees                      degrees


             32ºF    0ºC               Freezing point



1.8 F                       1.8 F
 1 C                         1 C
Fig2_9
       Converting °C to °F

   To convert from °C to °F
   Different values for the freezing points
         32 °F
                     add 32 to the °F value
          0 °C
   Different size of the degree intervals in each
    scale

               1.8 F
      T F           TC  32
                1 C
     Converting °C to K
   Temperature units are the same size
   Differ only in the value assigned to
    their reference points
        0 °C =
         273 °C
        K =K      + 273 273 to the °C value
                      add


   25°C is room temperature, what is the
    equivalent temperature on the Kelvin
    scale?
              25 ºC 273 = 298 K
             25 ºC + + 273 = 298 K
             25 ºC + 273 = 298 K
      Example
   A cake is baked at 350 °F. What is
    this in Centigrade/Celsius? In Kelvin?
         1.8 F
T F           TC  32
          1 C                      TC  176.6667 C
 1 C
       TF  32   TC 
1.8 F                                  176.7 273 =
     1 C
           350 F  32  TC             449.7 K
                                               450 K
    1.8 F
         Temperature Changes:
         Heat Capacity
   Heat is the total amount of energy in a system
     It is function of the amount of motion (kinetic

      energy) contained in molecules
     It is also a function of the potential energy of the

      molecules
     It involves the exchange of thermal energy caused

      by a temperature difference
        Heat vs. Temperature
   Within a quantity of matter:
   Heat has units of Joules and temperature has units in
    degrees
   Temperature relates only to kinetic energy within a
    molecule
   Heat is the total amount of energy in a molecule: It
    contains a kinetic and potential energy component
   Heat energy can be added or removed without a
    change in temperature (PE component only)
   As heat energy (KE component only) increases the
    temperature increases
    Temperature Changes: Heat Capacity

   Heat energy is the form of energy most often
    released or required for chemical and physical
    changes
   Every substance must absorb a different amount
    of heat to reach a certain temperature
   Different substances respond differently when
    heat is applied
   The amount of heat required to raise the
    temperature of a given quantity of a substance by
    1 ºC is called its heat capacity
        Temperature Changes: Specific Heat

   If 4.184 J of heat is applied to:
      1 g of water, its temperature is raised by 1 °C

      1 g of gold, its temperature is raised by 32 °C

      Some substances requires large amounts of heat

        to change their temperatures, and others require a
        small amount
      The precise amount of heat that is required to

        cause a given amount of substance (in grams) to
        have a rise in temperature is called a substance’s
        “specific heat”
          Specific Heat
     The amount of heat energy (q) needed to
      raise 1 gram of a substance by 1 °C
         Specific to the substance
         The higher the specific heat value, the less its
          temperature will change when it absorbs heat
         SH values given in table 3.4, page 71
         Only for heating/cooling, not for changes of state

          quantity of heat transf erred     q         J        cal
C                             heat              
                                           J (or cal)   C or g  C
                                         m  ΔT
                   SH 
   (grams of substance)  (temp. change)
                           grams  Δt
                                         
                                             g  C g 
    Specific Heat Expression with
    Calories and Joules

 1 cal is the energy needed to
  heat 1 g of water 1 °C
 1 cal is 4.184 J

   Make a conversion factor
    from the statements
                  1 cal    4.184 J
        C               
          water 1g  1 C 1g  1 C
          Specific Heat Equation
        The rearrangement of the SH equation gives the
         expression called the “heat equation”


     C     heat (q)                     heat (q)  C  mass ( g )  ΔT
          mass ( g )  ΔT

                   C
            J 
            g  C   m(g)  ΔT( C)
    q(J)                                              answ er in joules
                   
   q = heat
   C = specific heat (different for each substance)
   m = mass (g)
   ∆T = change in temperature (°C)
        Specific Heat Equation
   Energy (heat) required to change
    the temperature of a substance
    depends on:
     The amount of substance being
      heated (g)
     The temperature change (initial T
      and final T in °C)
     The identity of the substance
     Energy and T
             Heat (q)C  mass (g) Δt
              2×                    2×

   The amount the temperature of an object
    increases depends on the amount of heat added
    (q)
      If you double the added heat energy (q), the
       temperature will increase twice as much.
      When a substance absorbs energy, q is
       positive, temperature increases
      When a substance loses energy, q is
       negative, temperature decreases
          Energy and Heat Capacity Calculations
   Use same problem solving steps as before (Chapter 2)
      State the given and needed units

      Write the unit plan to convert the given unit to the final unit

      State the equalities and the conversion factors

      Set up the problem to cancel the units

   Pepsi One™ contains 1 Calorie per can.
    How many joules is this?
                   1 Cal = 1000 cal   4.184 J = 1 cal


        1 Cal 1000 cal 4.184 J
                                4184 J
               1 Cal 1 cal
       Energy and Heat Capacity Calculations

   The 4184 J from the Pepsi One™
    will heat how many grams of water
    from 0°C to boiling?
                          q
                    m
                       C  T



      4184 J 1 C1g            10 g  10 mL
             4.184 J 100 C
       Energy and Heat Capacity Calculations

   How many grams of water will the
    4184 J (from a can of Pepsi One™)
    heat from 0°C to boiling?
                              q
                        m
                           C  T

                                     dwater =1.0 g/mL
    4184 J 1 C  1 g            10 g 
                                           1.0 g
                                                    10 mL
             4.184 J 100 C0 C
                     100 C -               mL
                        Tf   Ti
        Energy and Heat Capacity Calculations
   How many grams of water would reach
    boiling if the water started out at room
    temperature (25°C)?
                                       q
                                 m
                                    C  T


        4184 J 1 C  1 g                
                 4.184 J         25
                            10075 C 
                              13.33 g  13.33 mL
      Energy and Heat Capacity Calculations

   If 50.0 J of heat is applied to 10.0 g
    of iron, by how much will the
    temperature of the iron increase?
     50.0 J g  C
                           11.11 C
            0.45 J 10.0 g
                              Solve for ΔT
                                     q
     q  C  m  ΔT          T   
                                    Cm
   end

								
To top