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					       Physical and Chemical Properties

• All substances have properties that we can use to identify
  them. For example we can identify a person by their face,
  their voice, height, finger prints, DNA etc.. The more of
  these properties that we can identify, the better we know
  the person. In a similar way matter has properties - and
  there are many of them. There are two basic types of
  properties that we can associate with matter. These
  properties are called Physical properties and Chemical
  properties:
• Physical properties:Properties that do not change the
  chemical nature of matter
• Chemical properties:Properties that do change the
  chemical nature of matter
• Examples of physical properties are: color, smell,
  freezing point, boiling point, melting point, infra-red
  spectrum, attraction (paramagnetic) or repulsion
  (diamagnetic) to magnets, opacity, viscosity and density.
  There are many more examples. Note that measuring each
  of these properties will not alter the basic nature of the
  substance.
• Examples of chemical properties are: heat of
  combustion, reactivity with water, pH, and electromotive
  force.
• The more properties we can identify for a substance, the
  better we know the nature of that substance. These
  properties can then help us model the substance and thus
  understand how this substance will behave under various
  conditions.
 Conservation of Mass in Chemical Reactions
• Democritus (460-370 BC) and somewhat later John
  Dalton (1766-1844) were the first to consider matter at its
  most microscopic form. They both came up with the
  concept of the "atom" as being the smallest unit of matter
  and thus being undivisible*. This observation has an
  important and fundamental consequence: mass is neither
  created nor destroyed during the course of a chemical
  reaction. How do we come to this conclusion? We know
  that chemical reactions take place at the atomic/molecular
  level. That is molecules and atoms interact with one
  another during a chemical reaction. If atoms are
  indivisible then they cannot be destroyed during a
  chemical reaction. If atoms cannot be destroyed then the
  mass of reactants must equal the mass of the products in a
  chemical reaction. e.g.,
               • Reactants -------> Products
         • Mass of Reactants = Mass of Products
• This can be visualized by considering the formation of
  water from oxygen and hydrogen molecules:




• Note that the hydrogen and oxygen atoms simply
  rearrange themselves but are not destroyed. Therefore
  mass is conserved.
  Iron + Oxygen -----> Rust
  100 g + ?g ------> 143g
  mass reactants = mass products
  mass products = 143g = mass reactants
  = 100 + mass of oxygen         mass oxygen = 43 g
     Elements, Compounds and Mixtures

• All substances have mass and therefore must be
  composed of atoms. These atoms and how they assemble
  themselves in the substance determines their chemical
  and physical properties. Substances can be classified
  according to how these atoms are assembled and is
  known as Classification of Matter: All matter falls into
  one of three categories: elements, compounds or
  mixtures. Furthermore, mixtures can be classified as
  homogeneous or inhomogeneous. The scheme looks
  something like the diagram next slide:
• This classification depends upon how we try and separate
  matter into its basic components. This separation is called
  the "process". There are two processes: a physical and a
  chemical process.
• Physical process: a process using physical properties
• Chemical process: a process using chemical properties
• If we have a sample of matter and can find a physical
  process such as evaporation, magnets, color etc. to
  separate it then the sample is a "mixture". Furthermore if
  the sample is a mixture of solids and liquids (e.g., sand
  and water) etc. or two or more liquids that don't mix (e.g.,
  oil and vinegar) then the mixture is "inhomogeneous".
  Otherwise the sample is a "homogeneous" mixture.
• If there is no physical process that will separate the
  sample then the sample is a "pure" substance. If a
  chemical process such as combustion or oxidation breaks
  the substance down to its constituent atoms then the
  substance is a "compound"(e.g., salt, sugar, water).
  Otherwise the substance is an "element" (e.g., copper
  penny, aluminum foil). Compounds are made up of
  molecules or salts. Elements are made up of single types
  of atoms.
                           Density

• Density is a physical property of matter. Most commonly
  density refers either to the mass per unit volume (mass
  density) or the number of objects (e.g., atoms, molecules)
  per unit volume (number density). We will focus out
  attention on mass density. The mass density has the units
  mass/volume. Since volume has the units length "cubed"
  then the SI unit of mass density is kg/m3. More common
  units of density are g/ml or g/l. Substances have different
  densities. In fact the density of a substance can often be
  used to help identify it. In the next slide there is a table of
  densities of common materials:
           Densities of Some Common Substances
               Substance        Density (g/mL)
               Ice (0 °C)           0.917
             Water (4.0 °C)         1.000
                  Gold              19.31
             Helium (25 °C)       0.000164
             Dry Air (25 °C)      0.001185
              Human Fat              0.94
                  Cork           0.22 - 0.26
              Table Sugar            1,59
              Balsa Wood             0.12
                  Earth              5,54

An important example is water. The above table states that
liquid water has a mass of 1 g in every ml. Thus 2 ml of water
has a mass of 2 g etc.. Table sugar is more dense than water
by about 60 percent. Density does not depend upon size. For
example the water in a swimming pool has the same density a
glass of that swimming pool water.
Calculations with density are straight forward and involve the
formula for density namely D=m/V, where D=density,         m=
mass and V = volume.
• Example 1: What is the volume of a nugget of gold that
  has a mass of 3.45 g? The density of gold can be looked
  upon as a conversion factor from mass to volume i.e.,


• Example 2: A light substance is found to weigh 23 g and
  to have a volume of 0.192 liters. What is the substance?




• Based upon this result we would guess that this substance
  might be balsa wood.
• Example 3: What is the mass of 1 liter of sugar?
                     Compounds

•   There are two basic types of compounds. They are
    distinguished by by the manner in which the atoms bind
    to one another in the compound. These two types are
    called "molecular" compounds and "salts" (or
    equivalently "ionic" compounds):
•   Molecular compounds:These compounds are made up of
    molecules whose atoms bind to one another through
    "covalent" bonds.
•   Salts:The atoms in salts are held together with "ionic"
    bonds. Unlike molecules, salts always form solids in a
    regular array called a "crystalline solid".
• A bond is the "glue" that holds atoms together. In
  compounds this glue can either be covalent or ionic.
• Covalent bonds:The electrons are shared between
  atoms. Therefore this sharing of electrons provides the
  glue.
• Ionic bonds:Ionic bonds occur due to the mutual
  attraction between atoms with positive and negative
  charges i.e., ions.
            Examples of Molecules




                               Acetaldeyhde
N-hexane (top)                 (top)

                              Taxol (left)
  An Example of a Salt




Sodium Chloride (NaCl)
      Energy and Chemical Reactions

• When matter undergoes transformations that
  change its chemical and physical properties then
  that transformation was brought about by a
  chemical reaction. On the other hand chemical
  reactions can only take place if there is sufficient
  energy to make the reaction proceed. Therefore
  energy is a prerequisite for chemical reactions.
• Energy can come in many forms e.g., heat, work, light, kinetic,
  potential, chemical etc.. Moreover, energy can itself transform
  among these various forms. For example a ball at the edge of a
  table has zero kinetic energy and positive potential energy. If
  the ball drops it will have zero potential energy and positive
  kinetic energy the instant it hits the floor. However the sum of
  the potential and kinetic energy is the same throughout the
  ball's dropping history. Therefore energy has neither been
  created or destroyed but has transformed from potential to
  kinetic energy.
• Molecular of chemical energy can mean several things:
  Chemical bonds are a source of energy, the movement of
  molecules in space is kinetic energy, the vibrations and
  rotations of molecules is another source of chemical
  energy. All of these forms of chemical energy contribute
  in one way or another to chemical reactions.
• The units of chemical reactions are straightforward and is
  given in the diagram below:




• There are many other units for energy including electron
  volt (ev), erg, kjoule (kJ) etc.
          Specific Heat and Heat Capacity
• Specific heat is another physical property of matter. All
  matter has a temperature associated with it. The
  temperature of matter is a direct measure of the motion of
  the molecules: The greater the motion the higher the
  temperature:




• Motion requires energy: The more energy matter has the
  higher temperature it will also have. Typically this energy
  is supplied by heat. Heat loss or gain by matter is
  equivalent energy loss or gain.
• With the observation above understood we can now ask the
  following question: by how much will the temperature of an
  object increase or decrease by the gain or loss of heat
  energy? The answer is given by the specific heat (S) of the
  object. The specific heat of an object is defined in the
  following way: Take an object of mass m, put in x amount of
  heat and carefully note the temperature rise, then S is given
  by;


• In this definition mass is usually in either grams or
  kilograms and temperatture is either in kelvin or degres
  Celcius. Note that the specific heat is "per unit mass". Thus,
  the specific heat of a gallon of milk is equal to the specific
  heat of a quart of milk. A related quantity is called the heat
  capacity (C). of an object. The relation between S and C is
  C = (mass of obect) x (specific heat of object).
• A table of some common specific heats and heat
  capacities is given below:

           Some common specific heats
           and heat capacities:
                                     C (J/°C)
             Substance  S (J/g °C) for 100 g
            Air         1.01        101
            Aluminum  0.902         90.2
            Copper      0.385       38.5
            Gold        0.129       12.9
            Iron        0.450       45.0
            Mercury     0.140       14.0
            NaCl        0.864       86.4
            Ice         2,03        203
            Water       4,179       417,9
• Consider the specific heat of copper , 0.385 J/g °C. What
  this means is that it takes 0.385 Joules of heat to raise 1
  gram of copper 1 degree Celsius. Thus, if we take 1 gram
  of copper at 25 °C and add 1 Joule of heat to it, we will
  find that the temperature of the copper will have risen to
  26 °C. We can then ask: How much heat will it take to
  raise by 1 °C 2g of copper?. Clearly the answer is 0.385 J
  for each gram or 2x0.385 J = 0.770 J. What about a
  pound of copper? A simple way of dealing with different
  masses of matter is to determine the heat capacity C as
  defined above. Note that C depends upon the size of the
  object as opposed to S that does not.
• Example 1: How much energy does it take to raise the
  temperature of 50 g of copper by 10 °C?
• Example 2: If we add 30 J of heat to 10 g of aluminum,
  by how much will its temperature increase?




• Thus, if the initial temperature of the aluminum was
  20 °C then after the heat is added the temperature will be
  28.3 °C.
               Dalton’s Atomic Theory
• Democritus first suggested the existence of the atom but it
  took almost two millennia before the atom was placed on
  a solid foothold as a fundamental chemical object by John
  Dalton (1766-1844). Although two centuries old, Dalton's
  atomic theory remains valid in modern chemical thought

. Dalton's Atomic Theory
  1) All matter is made of atoms. Atoms are indivisible and
  indestructible.
  2) All atoms of a given element are identical in mass and
  properties
  3) Compounds are formed by a combination of two or
  more different kinds of atoms.
  4) A chemical reaction is a rearrangement of atoms.
• Modern atomic theory is, of course, a little more involved
  than Dalton's theory but the essence of Dalton's theory
  remains valid. Today we know that atoms can be
  destroyed via nuclear reactions but not by chemical
  reactions. Also, there are different kinds of atoms
  (differing by their masses) within an element that are
  known as "isotopes", but isotopes of an element have the
  same chemical properties.
• Many heretofore unexplained chemical phenomena were
  quickly explained by Dalton with his theory. Dalton's
  theory quickly became the theoretical foundation in
  chemistry.
             Composition of the Atom
• Atoms have a definite structure. This structure determines
  the chemical and physical properties of matter. This
  atomic structure was not fully understood until the
  discovery of the neutron in 1932. The history of the
  discovery of atomic structure is one of the most
  interesting and profound stories in science. In 1910
  Rutherford was the first to propose what is accepted
  today as the basic structure of the atom. Today the
  Rutherford model is called the "planetary" model of the
  atom. In the planetary model of the atom there exists a
  nucleus at the center made up of positively charged
  particles called "protons" and electrically neutral atoms
  called "neutrons". Surrounding or "orbiting" this nucleus
  are the electrons. In elements the number of electrons
  equals the number of protons.
• The picture above greatly exaggerates the size of the
  nucleus relative to that of the atom. The nucleus is about
  100,000 times smaller than the atom. Nevertheless, the
  nucleus contains essentially all of the mass of the atom. In
  order to discuss the mass of an atom we need to define a
  new unit of mass appropriate to that of an atom. This new
  unit of mass is called the "atomic mass unit" or amu. The
  conversion between the amu and gram is
                   1 amu = 1.67x10-24 g
• The mass, in amu, of the three particles is given in the
  table below:




• Note that the mass of an electron is about 2000 times
  smaller than that of the proton and neutron. Also note that
  the mass of the proton and neutron is close to 1 amu. This
  is a useful fact to remember. If the number of electrons
  does not equal the number of protons in the nucleus then
  the atom is an ion:
    • cation: number of electrons < number of protons
     • anion: number of electrons > number of protons
 Rutherford’s Planetary Model of the Atom

• By 1911 the components of the atom had been
  discovered. The atom consisted of subatomic
  particles called protons and electrons. However,
  it was not clear how these protons and electrons
  were arranged within the atom. J.J. Thomson
  suggested the"plum pudding" model. In this
  model the electrons and protons are uniformly
  mixed throughout the atom:
• Rutherford tested Thomson's hypothesis by devising his
  "gold foil" experiment. Rutherford reasoned that if
  Thomson's model was correct then the mass of the atom
  was spread out throughout the atom. Then, if he shot high
  velocity alpha particles (helium nuclei) at an atom then
  there would be very little to deflect the alpha particles. He
  decided to test this with a thin film of gold atoms. As
  expected, most alpha particles went right through the gold
  foil but to his amazement a few alpha particles rebounded
  almost directly backwards.
• These deflections were not consistent with
  Thomson's model. Rutherford was forced to
  discard the Plum Pudding model and reasoned
  that the only way the alpha particles could be
  deflected backwards was if most of the mass in
  an atom was concentrated in a nucleus. He thus
  developed the planetary model of the atom which
  put all the protons in the nucleus and the
  electrons orbited around the nucleus like planets
  around the sun.
        Isotopes and Atomic Symbols


• Atomic Symbols:
  The atom of each element is made up of
  electrons, protons and neutrons. All atoms of the
  same neutral element have the same number of
  protons and electrons but the number of neutrons
  can differ. Atoms of the same element but
  different neutrons are called isotopes. Because of
  these isotopes it becomes necessary to develop a
  notation to distinguish one isotope from another -
  the atomic symbol.)
The atomic symbol has three parts to it:
  • 1. The symbol X: the usual element symbol
  • 2. The atomic number A: equal to the
    number of protons (placed as a left subscript)
  • 3. The mass number Z: equal to the number
    of protons and neutrons in the isotope (placed
    as a left superscript
• Examples 1:
• Consider two isotopes of gallium, one having the 37
  neutrons and the other having 39 neutrons. Write the
  atomic symbols for each isotope. Solution:




• Example 2:
• How many neutrons does the isotope of copper with mass
  number Z = 65 have?
  Solution: From the periodic table we see that copper has
  an atomic number of 29. Since Z is the number of protons
  plus the number of neutrons, then No. neutrons = 65 - 29
  = 36
                        The Mass

• The standard for every unit must be defined. Length is an
  example. The basic unit of length is the meter which was
  defined in 1983 as equal to the distance traveled by light
  in a vacuum in 1/299,792,458 of a second. Mass must
  also be defined. The definition of mass today is the amu
  (atomic mass unit). The amu is defined in the following
  way: the mass of one atom of the carbon-12 isotope is
  EXACTLY 12 amu.

          mass of one carbon-12 atom = 12 amu
• All other masses are measured relative to this carbon-12
  standard. For example, suppose we do an experiment and
  find that the isotope bromine-81 has a mass that is 6.743
  times that of carbon-12. Then the mass of bromine-81
  would be given by
                  Atomic Weights


• Most elements can be found on earth (with the exception
  of those elements that too unstable and thus must be
  synthesized in the laboratory). Since all elements have
  isotopes then we must consider how much of one isotope
  of an element exists versus another isotope of the same
  element. These are called the "natural" abundances on
  earth.
• Natural Abundances:
• Suppose we go to a cave and mine element "X".
  After careful analysis we find that in our sample
  of element X there exists three isotopes: Xa, Xb
  and Xc. Moreover, we find that out of every 100
  atoms the various isotopes are distributed as
  follows, and their masses are given.
               For Every 100       Isotope Masses of
                 atoms of X                 X
                         No. of                Mass 
            Isotope     atoms       Isotope    (amu)
                Xa         30          Xa        54
                Xb         60          Xb        56
                Xc         10          Xc        59
• Then the average mass (atomic weight) is given by:




• The atomic weight of each element is included along with
  the element symbol in the periodic table. It is important to
  note that no one atom has a mass equal to that of the
  atomic weight. Remember: the atomic weight represents
  that average mass of the atoms.

				
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