History of the Atomic Theory: by 3n8puy


									 TEACHER NOTES 2010                                                                                         UNIT 4 |1
                       UNIT 4 – ATOMIC THEORY AND STRUCTURE

You have a team of specialists longing to help you be able to…


       Identify metals, non-metals, and transition metalss on the periodic table.

       Locate the following groups on the periodic table: Alkali metals, Alkaline Earth Metals, Transition Metals,
   Halogens, Noble Gases.

      Compare and contrast physical properties of the following groups: Alkali metals, Alkaline Earth Metals, Transition
   Metals, Halogens, Noble Gases.

       Describe the contributions to atomic theory by Dalton, JJ Thompson, Rutherford, Bohr,

       Outline the cathode ray experiment.

       Explain the results of the gold foil experiment.

       Identify the three primary subatomic particles in terms of size, location, and charge.

       Compare and contrast atomic number, mass number, average atomic mass

       Compare and contrast the impact on an atom of changing the number of electrons, neutrons, and proton.

       Perform calculations involving numbers of protons and electrons and the ionic charge of a species.

       Define the terms cation and anion.

       Identify the isotopes of an element.

       Perform calculations involving atomic mass and % abundance of isotopes.
 TEACHER NOTES 2010                                                                                        UNIT 4 |2

             A. Organization

Using page 131 in your book, label the periodic table found on the next page according to the following guidelines: (NOTE:
Your key might not look exactly like the book!!!)

        a.   Color all metals blue
        b.   Color all non-metals yellow
        c.   Draw a dark line showing the transition between metals and non-metals (the line by the semi-metals)
        d.   Identify the transition metals
        e.   Label the halogens
        f.   Label the alkali metals
        g.   Label the alkaline earth metals
        h.   Label the noble gases
        i.   Number all groups and periods on the periodic table

                        Periodic Table
 TEACHER NOTES 2010                                                                                UNIT 4 |3
            B. Properties of selected groups:

                       Group 1           Group 2            Groups 3-12          Group 17           Group 18
     Name              ALKALI                              TRANSITION                                NOBLE
                                         EARTH                                 HALOGENS
                       METALS                                METALS                                  GASES
 Metal or non-
                       METALS             METALS             METALS           NON-METALS          NON-METALS
                                                                               F2, Cl2 gases
 State at 25oC                                             SOLID (except
                        SOLID              SOLID                                Br2 liquid           GASES
  and 1 atm                                                    Hg)
                                                                                  I2 solid
                    Y (not as much     Y (not as much
 Malleable and
                     as transition      as transition             Y                  N                   N
ductile (Y or N)
                        metals)            metals)
electricity (Y or          Y                  Y                   Y                  N                   N
                     Silvery color,
  Key physical      Soft, relatively
                                         Silvery color,     except for Cu.     F2 pale yellow
   properties.        low melting                                                                 Boil at very low
                                        higher melting        Very hard.       Cl2 pale green
 Include color,        point for a                                                                 temperatures.
                                       than Gr. 1, shiny     Readily form        Br2 brown
relative melting      metal, shiny                                                                   Very weak
                                        (dull surface is   alloys with each    I2 shiny black
points, density,    (dull surface is                                                                 attractions
                                        due to reaction      other. High        (sublimes to
 and luster (or     due to reaction                                                               between atoms.
                                         with oxygen)       melting points,    violet vapor).
 lack thereof!)      with oxygen).
                                                             high density
                     Low density
                                                            Less reactive                            Generally
                                                           than Gr. 1 & 2.       When pure,         unreactive or
                    Highly reactive.
                                                             Do not react     found paired up.    INERT. Kr can
                      Readily react
                                        Very reactive,     with water, but     F2 and Cl2 are       be forced to
                       with water,
 Key chemical                           but less than      some react with    poisonous, Br2 is      react with
  properties                           Gr.1. Barium is      acids. Often        toxic. Highly     fluorine and Xe
                       halogens –
                                        stored in oil.      form colored       reactive – love    can be forced to
                    almost anything!
                                                           solutions when          to steal          react with
                      Stored in oil
                                                             in a soluble         electrons!         fluorine or
                                                             compound.                                 oxygen


    1. Malleable

    2. Ductile

    3. Luster
 TEACHER NOTES 2010                                                                            UNIT 4 |4
                 PERIODIC TRENDS – Use these notes to help with your Inquiry!
The METALS are found to the LEFT of the staircase line, excluding hydrogen. Some of the metal groups
have specific names…
    Group 1 = Alkali Metals (part of the s-block)
    Group 2 = Alkaline Earth Metals (part of the s-block)
    Groups 3 – 12 = Transition Metals (the d-block)
    The two rows on the bottom = Lanthanides and Actinides (the f-block)

The NON-METALS are found to the RIGHT of the staircase line, as well as hydrogen. Some of the groups of
the non-metal groups also have specific names…
     Group 17 = Halogens (part of the p-block)
     Group 18 = Noble (Inert) Gases (part of the p-block)

GROUPS (aka Families) run VERTICALLY on the periodic table. Trends (such as reactivity and some
physical properties) tend to be consistent within a group.

PERIODS run HORIZONTALLY on the periodic table. As you work through the periods on the periodic
table, trends tend to repeat themselves from period to period (such as number of valence electrons).
There are some specific trends that can be summarized by looking at the periodic table as a whole. The
following contains a description of those trends. Realize with all of these trends, we are giving general
descriptions of how these trends work… as with anything in chemistry; there are exceptions to the rules!

   The radius is ½ of the distance measured
     from the center of one atom to the center
     of an adjacent identical atom. The bulk of
     this distance is the “electron playground”!
   Atomic size decreases across a period, due
     to the increased number of protons
     having a greater pull on the valence
     electrons more as you move across.
   Atomic size increases down a group, as
     you are adding more energy levels as you
     go down.
   This would mean that Fr (Francium)
     would theoretically have the largest
     atomic size. He (Helium) would have the smallest atomic size.
   EX: Which would have a larger atomic size – Ca, Br, or Ba?
 TEACHER NOTES 2010                                                                             UNIT 4 |5

      This refers to the amount of energy required to remove one outer-most (or VALENCE) electron.
      Since 8 valence electrons indicate a “full” outer shell (s & p sublevels), the closer the number of
       valence electrons is to 8, the more difficult it will be to pull off an electron - meaning it will take
       more ionization energy.
      Ionization energy increases across a period, because the atom is more tightly held together as the
       valence sublevels become full… meaning it will take more energy to pull an electron off.
      Ionization energy decreases down a group, because if you have more energy levels, the outer-most
       electrons are further away, meaning they aren’t held on as tight… meaning it takes less energy to
       pull them off.
      This would mean that He (Helium) would theoretically have the highest ionization energy. (Yes, he
       only has two valence electrons, but that’s all he can handle!) Fr (Francium) would have the lowest
       ionization energy.
      EX: Which would have a higher ionization energy – Ca, Br, or Ba?

    This refers to the attraction an atom has to on a nearby electron. The closer an atom is to 8 valence
      electrons (in the s & p sublevels), the higher the electronegatvitiy (because the nucleus has a
      stronger “+” charge!) However, would an atom that already has 8 valence electrons want to gain
    Electronegativity increases across a period, because as you move across, the atom has a greater pull
      to gain another electron (to achieve the 8 valence electrons).
    Electronegativity decreases down a group, because as the valence electrons get further from the
      nucleus, it has less of a pull to gain more electrons.
    Noble (Inert) gases have no electronegativity, because they already have 8 valence electrons (filled
      s & p sublevels), so they have no pull to gain any more! This is also part of the reason they are
      considered “INERT” – meaning unreactive.
    This would mean that F (Fluorine) would theoretically have the highest electronegativity. Fr
      (Francium) would have the lowest electronegativity.
    EX: Which would have a higher electronegativity – Ca, Br, or Ba?
 TEACHER NOTES 2010                                                                                  UNIT 4 |6

I. History of the Atomic Theory:

Atom - the smallest particle of matter which will exhibit the properties of that element. When broken down
smaller than an atom, the parts (protons, electrons, and neutrons) of different elements look exactly the same.
You cannot tell a proton in a gold atom from a proton in oxygen gas.

Atoms are very small—typically about 1 x 10-8 cm in diameter.

To give you an idea of how small this is: 1.0 gram of lead contains 2.9 x 1021 atoms of lead. By comparison, the
earth’s entire population is only 5 x 109 people.

Models – People use models to describe understanding of truth. Models, as with our understanding of science in
general, evolve over time as improved technology was used to study the atom. There were MANY scientists
involved in this process. We will focus on only a few.

        A. JOHN DALTON was an Englishman in the 19th century who was the first to develop and publish a
           theory about how atoms looked and behaved. He conceived of the atom as a solid sphere, much like a
           billiard ball. Hence, his theory was called the “Billiard Ball” theory.

                                     The following are statements of John Dalton’s ATOMIC THEORY:

                                     (1) Proposed: all elements are composed of very small particles called
                                     atoms which are indivisible.
                                     WRONG: Today we know that atoms can be divided into protons,
                                     electrons, neutrons and almost 200 other subatomic particles.

                                     (2) Proposed: All atoms of the same elements are identical.
                                     WRONG: We know that this is not true due to the presence of isotopes.
                                     Isotopes are atoms with the same number of protons (same element) but
                                     different numbers of neutrons (different masses). For example, “regular”
                                     carbon has 6 protons and 6 neutrons, but carbon-14 (radioactive carbon used
                                     in carbon “dating”) has 6 protons and 8 neutrons.

            (3) Proposed: Atoms of different elements are different. TRUE

            (4) Proposed: Atoms of different elements can combine with each other only in simple whole
            number ratios to form compounds. TRUE

            (5) Proposed: Chemical reactions occur when atoms are separated, joined or arranged. However,
            atoms of one element ARE NOT changed into atoms of another element by a chemical reaction.
            (Only by nuclear reactions) TRUE
 TEACHER NOTES 2010                                                                               UNIT 4 |7

           B. J. J. THOMSON discovered electrons by means of his CATHODE RAY TUBE EXPERIMENT

           (1) Proposed: Through the Cathode Ray Tube Experiment, proposed that electrons are negatively
           charged particles, abbreviated e‾. Electrons must be a subatomic particle because its mass was nearly
           2000 times less than a hydrogen atom (the smallest atom). TRUE

                                                                  Man - Don’t
                                                                 scientists ever


                                  + Magnetic field

              - Magnetic field

Electron Stream w/o applied electric field           Electron Stream after applying field, + on top, - on bottom

           Notice how the electron stream “bends” toward the + electrode (the cathode) and away from the –
           electrode (the anode). Thomson knew that electrons must be negative because they are attracted to
           the positive field! (Remember, positive attracts negative.)

            Final Note: Since the e‾ stream could be produced from any metal, Thomson suggested that all
           atoms had electrons and they were negatively charged.

            (2) Proposed: Thomson conceived of the atom as a plum pudding: the positive charge (he didn’t
           know about protons) and electrons were mixed together like a fruit salad (or plum pudding). The
           larger pieces of the “plum” were the electrons swimming in a pudding of positive charge. WRONG

           C. ERNEST RUTHERFORD worked for Thomson in his lab for a while. He performed the

           (1) Proposed: the atom as mostly space and that all of the positive charge was located in a very small
           central nucleus. TRUE
 TEACHER NOTES 2010                                                                                 UNIT 4 |8

        1. Most particles went through (which is what was expected)
        2. Some were slightly deflected
        3. Some were deflected back and missed the fluorescent screen

Why did some of the particles get deflected? Reason: the positive alpha particles were
hitting the positive, small, dense nucleus and getting deflected (since positive repels

Conclusions:      Dense positive nucleus (meaning plum pudding was incorrect).
Negative electrons were moving around the nucleus.

            D. NIELS BOHR – Danish scientist who originally worked for Thomson and Rutherford.

                          (1)        Proposed: Electrons were in energy levels. The further an electron was
                          from the nucleus, the higher its energy. This proposal provided a great starting point
                          for our current understanding of the atom – YEAH BOHR!! TRUE

                          (2)       Proposed: the “planetary model” of the atom in which the electrons orbit
                          the nucleus much as the planets orbit the sun. WRONG

                                                                                     Hmmm. I do look a
                                                                                     little “bohr”ed!
TEACHER NOTES 2010                                                                          UNIT 4 |9

        E. ERWIN SCHRODINGER – an Austrian physicist, along with Werner Heisenberg and Louis
                        de Broglie, postulated the quantum (wave) mechanical model of the atom
                        which we believe is true today – at least the math works! In this concept the
                        electrons do not actually “orbit” the nucleus, but are found only in definite
                        areas based on the amount of energy they have and move according to wave
                        functions. Key – we don’t really know how an electron moves – just the
                        probability of finding it in a particular region of space around the nucleus.

                                     We will talk WAY more about Schrodinger’s model very soon!


    A. PROTON- positively-charged particle which is found in the nucleus; it is considerable more massive
       than the electron (1837 times heavier). Shown as p+.

    B. ELECTRON- very light negatively-charged particle which
       is found somewhere outside the nucleus. Its mass is
       considered negligible when determining the mass of an atom. It
       weighs only 1/1837 that of a proton, but its negative charge is as
       powerful as the positive charge of a proton. Shown as e‾.

    C. NEUTRON- a particle found in the nucleus which is approximately
       the same mass as a proton but does not have an electrical charge
       associated with it – it is neutral. Shown as no.

    D. ATOMS ARE ALWAYS NEUTRAL PARTICLES—that is, they contain
       the same number of protons as electrons (and it doesn’t matter how many neutrons they have).
TEACHER NOTES 2010                                                                                U N I T 4 | 10

ATOMIC NUMBER – the number of protons found in the nucleus of an atom. On the Periodic Chart it is the
large whole number in the upper corner. Notice that the only thing which makes on element different from
another one is the number of protons it contains. An atom with 6 protons is carbon; an atom with 5 protons is
boron and an atom with 7 protons is nitrogen. ATOMIC NUMBER DEFINES ONE ELEMENT FROM


        E. MASS NUMBER – the number of protons and neutrons found in the nucleus of an atom, this DOES
           NOT show on the Periodic Chart. The atomic mass which does show on the chart is an average of all
           the different isotopes of that element (SEE BELOW).

                                                p + n = mass #

          We often use mass number to describe what isotope (same element, different # n0) we are talking
          about. For example, Carbon-14 means that we are talking about the carbon that has a mass
          number of 14.

          Example 4-1:How many neutrons are in Uranium-235?


        A. Changes in the number of neutrons – formation of isotopes

               Isotopes differ in the number of neutrons and hence the mass number
               Isotopes typically behave chemically the same.
               Isotopes are often designated with their mass number hyphenated after the element name:

                Example: Carbon-12              Carbon-13        Carbon-14
                Each has 6 protons and 6, 7, and 8 neutrons respectively.

             ATOMIC MASS – The atomic mass given on the Periodic Table for each element is really an
            average of the masses of all the isotopes of that element, weighted by their percentage of abundance.
            Every element on the Periodic Table has at least 3 isotopes; some of them have 20 or 30 or more

Avg. Atomic Mass (This is the number on the periodic table) = (fraction isotope #1) (mass of #1) +
(fraction isotope #2) (mass of #2) + (fraction of isotope 3) (mass of #3) +……. Continues on for however
many isotopes you have!!!

                              The fraction abundance is simply the percent ÷ 100
                                   NOTE: we express this in decimal form!
TEACHER NOTES 2010                                                                               U N I T 4 | 11

   Example 4-2:      Naturally occurring chlorine is 75.53% chlorine-35 and 24.47% chlorine-37. What is the
      average atomic mass which should be placed on the Periodic Table for the element?

Atomic Mass = (0.7553)(35) + (0.2447)(37) = 35.49

NOTE: since the mass number is a count of protons and neutrons, they aren’t considered in determining
significant figures.

   Example 4-3:        The element neon consists of three isotopes with masses of 19.99, 20.99 and 21.99 amus.
      These three isotopes are present in nature to the extent of 90.92%, 0.25% and 8.83% respectively.
      Calculate the atomic weight of neon.

Atomic Mass = (0.9092)(19.99) + (0.0025)(20.99) + (0.0883)(21.99) = 20.17

       B. Changes in the number of electrons – formation of ions

           IONS – particles which do not have the same number of protons and electrons so therefore they do
           have an electrical charge associated with them.

           IF: # p > # e─ THEN: there is an excess positive charge. CALLED: cation

           IF: # p < # e─ THEN: there is an excess negative charge. CALLED: anion

           # protons – # electrons = ionic charge

           NOTE: When adding and subtracting electrons you are dealing with adding a negative charge
           (making more negative) and subtracting a negative charge (making more positive).

                             Gain of electrons       Loss of electrons 

                     ─3        ─2       ─1        0      +1      +2     +3
   Example 4-4:          An element has 20 protons and 18 electrons. Identify the element and determine the ionic
      charge.        \

   Example 4-5:          The sulfide ion has a charge of ─2. How many protons and electrons does this ion have?
                                             16 protons, 18 electrons

     C. Changes in the number of protons – formation of different element
We will talk more about this phenomenon in Unit 14 – Nuclear Reactions
TEACHER NOTES 2010                                                       U N I T 4 | 12

IV. The NUCLEAR SYMBOL is a way of writing atoms or ions which gives you lots
of information.

                    massnumber 16              2-      

                    atomic  8
                     number              O

   Example 4-6:   Complete the following table

Protons Electrons Neutrons Mass # Atomic # Charge NUCLEAR SYMBOL
   5        5        6      11       5       0    11
                                                           5    B
  6         6        8         14         6           0    14
                                                            6   C
                                                                S 2
 16         18      16         32        16                32

                                                                Al 3
 13         10       14        27        13           +3   27

                                                                 Cr 2
 24        22        28        52         24          +2    52
                                                                P 3
 15         18      16         31        15                31

Hyphen notation
write the name of the element (dash) mass number
carbon-12 means the carbon atom that has 6 neutrons

carbon-13 means the carbon atom that has 7 neutrons

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