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Atomic Structure

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					Name ___________________________                     1                                    Atomic Structure

                        Atomic Structure and the periodic properties of elements

Concepts
  Electrons exist in atoms at discrete energy levels and with differently shaped probable locations. It is
    these energy and shape differences that result in the different chemical and physical properties of the
    elements.
  Light has both wave and particle characteristics.
  Matter has both wave and particle characteristics.
  Energy of light is proportional to frequency, inversely proportional to wavelength
  Much of atomic structure is not intuitive nor analogous to things that you're used to seeing, so you must
    make mental models to help you understand parts of atomic structure.
  Electrons occupy orbitals and do so according to rules
  Many properties of atoms are a reflection of the balance between electron repulsion and nuclear
    attraction.
  Like charges repel, opposite charges attract

Objectives
Student should be able to:
 Describe the basic structure of the atom using appropriate terminology
 Use the mole concept to convert between atomic and gram scales
 Interconvert among wavelength, frequency, and energy of light
 Identify the contributions of Dalton, Bohr, Planck, De Broglie, Einstein, and Schrödinger to our
   understanding of atomic structure
 Describe the shapes of each of the 4 atomic orbitals
 Rank the atomic orbitals in terms of energy
 Recognize and present the electronic configuration of atoms
 Utilize the electronic configuration of atoms to predict properties such as tendency to gain or lose
   electrons, charge on ions, and atomic radius

Videos, labs, websites, demos, and readings
  Spectral tubes, starburst demo
  Activity series lab, periodic properties charts, spectrophotometer lab
  Conceptual chemistry section 3.6, chapter 5
  Atomic structure video
  Chemthink website on atomic structure and ions; eths powerpoints on orbitals, periodic trends, and ions
Chemthink grading
Daily work - 3pts for completion of assigned tutorial
Quiz - 6 pts for completion of assigned quiz
Name ___________________________                  2                                  Atomic Structure

Glossary - as you find these terms in the outline or textbook, write their meanings here.
subatomic particle

atomic number

mass number

isotopes

atomic mass unit

nucleons

photons

wavelength

frequency

atomic spectrum

absorb

emit

quantized or quantum

principle quantum number

angular momentum quantum number

magnetic moment quantum number

electron spin quantum number

Atomic orbital

Max Planck

James Dalton

Neils Bohr

Albert Einstein

Erwin Schrödinger

Louis de Broglie
Name ___________________________                    3                                  Atomic Structure

Electron configuration

Aufbau principle

Pauli exclusion principle

Hund's rule

Shielding

Effective nuclear charge

Repulsion

Ionization energy



                 Describe the basic structure of the atom using appropriate terminology
What can you recall of Dalton's postulates about atoms? (see section 3.3 for help)




What do you recall of the structure of the atom - where is most of the mass of an atom found?



Where are electrons found?



Which subatomic particle is uncharged and where is it found?



Read section 3.6. What is the relative mass of the proton to the electron?



Which subatomic particle is negatively charged and where is it found?



What is the atomic number for the element manganese?



What does mass number tell us?
Name ___________________________                      4                                    Atomic Structure

What are isotopes? How do they differ from each other?



How do elements differ from each other?



Look at table 3.1. The mass of protons and neutrons are nearly the same, and it is given in kg, not g. There is
another common unit used to record atomic masses, rather than grams or kilograms, called the atomic mass
unit. At this point, the amu is defined as 1/12 the mass of 12C and corresponds within 1% to the mass of a
proton or neutron. Thus, for most purposes, the mass in amu of an atom will equal the sum of the protons
plus neutrons.

Practice: what is the mass, in amu, of an atom that has 32 protons and 38 neutrons?



What element does that atom correspond to?



What is the atomic number and atomic mass, in amu, of this element?



Which of the following pair(s) are isotopes and what elements are represented?
                  Z                   A                 Z                A                    Element(s)
 st
1 pair            6                   13                6                14

2nd pair          7                  14                   6                14

3rd pair          9                  19                   10               20



Isotopes exist for every element. In general, the number of neutrons is about equal to the number of
protons, but the ratio of neutrons to protons increases to about 1.5 as the atomic mass increases. More
neutrons are needed to keep the nucleus from blowing apart. Why might the nucleus blow apart?



                       The mole converts between atomic and gram scales!
Remember the mole? What was the number to which it corresponded?




Here's the basic mathematical conversion factor that you'll need here:
                              6.02 x 1023 atoms/mole and 6.02 x 1023 amu/gram
You also need to know how to determine molar mass, in grams, or atomic mass, in amu, as we did in the last
unit. For review:
What is the mass, in grams, of a mole of atoms that each has 32 protons and 38 neutrons?
Name ___________________________                          5                                 Atomic Structure

What is the number of moles of 5.0 g of atoms that each has 32 protons and 38 neutrons?




                                            56
What is the mass, in amu, of 25 atoms of         Fe? What is the mass in grams?




                                           56
What is the mass, in amu, of one mole of        Fe? What is the mass in grams?




How many atoms are in 2 moles of 133Cs?




How many atoms are in 2 x 10-5 moles of    133
                                                 Cs?




According to the Merck Index, an amount of HCN gas in the air at 1.5 x 10 -5 moles/L will kill you in less than
3 min. how many molecules of HCN are there in 1 L of air?




What is the molar mass of HCN?




How many grams of HCN in that 1 L of air?




Many biologically important molecules are found at very low concentration, about 1 x 10 -9 moles/L. Consider
insulin, which has a molar mass of 5823 g/mole. How many grams are present in 1 L of blood if the
concentration is 1 x 10-9 Moles/L?
Name ___________________________                        6                                     Atomic Structure

Human recombinant insulin costs $97.60 for 5 x 10 -2 g. How much goes 1 mole of insulin cost?




                       Interconvert among wavelength, frequency, and energy of light

Light gave us our first clues about the structure of atoms. I need to see what you know about light.
Examine the 2 statements below and select the one that best matches your understanding of light:

  Light is composed of electromagnetic waves. Light radiates out from a point in all directions much like a rock
                           thrown in a pond creates ripples that spread in all directions.

 Light is composed of small discrete packets called photons, which resemble particles of matter. A light beam is
                     like a river - instead of flowing water molecules, it is flowing photons.

Tell me what you like or dislike in these statements.




In the air all around us are electromagnetic waves of all manners of wavelengths. What are some forms of
electromagnetic waves that you can think of?



read section5.2.
how does wavelength vary going from blue to red light? (increase, decrease, stay the same)



how does frequency vary going from infrared to microwave? (increase, decrease, stay the same)



visible light corresponds to wavelengths between 400 and 700 nm. Can you see light with a wavelength of
650 nm? What kind of light does that wavelength correspond to?



Can you see light with a frequency of 8.15 x 10 15 Hz (or sec-1 - do you know what this symbol means?)



A relationship not in the book, but very important is this one: c =  or the speed of light = wavelength
times the frequency. The speed of light in a vacuum is 3.00 x 10 8 m/sec. Let's consider the speed of a car.
Wavelength is analogous to the distance that it takes for the tire to complete one revolution. Frequency is
how many revolutions the tire makes per unit time.
Then by dimensional analysis: distance/revolution x revolutions/time = distance/time or speed.
Name ___________________________                       7                                     Atomic Structure

Practice:
If a car tire is 2.0 m/revolution and the tire makes 8 revolutions/sec, how fast is the car moving?




If a wavelength of light is 3.75 x 10-6 m and has a frequency of 8 x 1013 sec-1, what is the speed of light?



If the wavelength is 6.50 x 10-8 m for light, what is the frequency?




When an atom of an element is excited by some energy source - electricity usually, then it emits light. H
atoms emit only particular wavelengths of light, regardless of how much energy was used. Na atoms emit
different wavelengths of light. These lines make up the atomic spectrum of each element. Let's look at a
few of these -

Looking through the spectroscope, you should be able to see some lines of color, which look different for
each different gas. Record here what you see:
                      Element                                              Colors of lines




Some researchers noticed that the frequencies of two lines could add up to that of a third line. Although
the scientists didn't know what that meant, you will shortly.

    Identify the contributions of Dalton, Bohr, Planck, De Broglie, Einstein, and Schrödinger to our
                                        understanding of atomic structure
In the early 1900s, Max Planck, Niels Bohr, and Albert Einstein were coming to realize that light energy is
 quantized, that is, it comes in discrete units, called photons. This means that light has characteristics of
   waves - wavelength and frequency, but also characteristics of particles - discrete energy amounts and
momentum. Planck found that the energy of light was proportional to the frequency in the equation E = h,
       where h is called Planck's constant and it has a value of 6.63 x 10 -34 kg*m2/sec or Joule*sec.

For light with a frequency of 1.4 x 1015 Hz, what is the energy in joules?
Name ___________________________                      8                                     Atomic Structure

As frequency increases, how does energy change? (increase, decrease, or remain the same)




As wavelength increases, how does energy change? (increase, decrease, or remain the same)




  This is an important relationship - energy is proportional to frequency and inversely proportional to
                                               wavelength.
                     Notice that light behaves both like a wave and like a particle!

Niels Bohr hypothesized that electrons absorb or emit light when they move between energy levels. The
light that you see then corresponds to the difference in energy between the levels. His model is often
depicted like the solar system - the further away the orbit, the higher the energy of the electron.

Look at a building with 5 floors. To go from the first to second floor requires some energy. If it takes 10 J
to go up one floor, how much does it take to go up 2 floors? How much energy is released when you go down 2
floors?




Draw an orbital model for an atom, with a nucleus in the center, and several electron orbitals around it. How
does the adding up of frequencies mentioned above make sense in this model?




For an atom, the difference in energy between orbitals isn't equal, as it would be in a building. The first
step is the biggest, and each subsequent step gets smaller until you reach the roof.



Read section 5.4.
Lots of things are quantized in the world. Name 3.
Name ___________________________                      9                                  Atomic Structure

One strange feature of this model is that the transition between orbits is instantaneous - the electron does
not exist in between! Can you find models for this in the 'real' world?




When an electron moves up hill (further away from the nucleus), it absorbs energy, but only enough to make
it to an existing energy level. Suppose that it took 1J to go up each step or orbital. How many steps would
the electron take if you gave it 1.5J? How about 3.5J? Suppose there were only 20 steps and you gave the
electron 21J. What would happen? To answer these questions, you can think about jumping up stairs.




Each step, or orbital, in Bohr's model is called a principal quantum number. The numbers start at 1 and
increase as the electron is found further from the nucleus. The symbol for the principal quantum number is
'n'.

Bohr's model described the energy levels and light bands for hydrogen atoms very well, but it did not
correctly describe the locations of electrons, nor did it work on other elements.

About the same time, Louis de Broglie observed that electrons could behave like light. They had wavelength
and frequency! Thus, energy and matter were interchangeable, behaved similarly at the atomic level, and the
interconversion was defined by Einstein's famous equation, E = mc2

Erwin Schrödinger and Werner Heisenberg described electrons as waves, and found that not only could they
get the same energy levels and light bands as for the Bohr model, but that they now could predict locations
of electrons that made sense of physical and chemical properties.

Read section 5.5 up to the middle of p.148.
What is one application of the wave character of electrons?




How fast do electrons appear to move in an atom?



what is the term used to denote the electron position?



How does that differ, if at all, from an atomic orbital?
Name ___________________________                      10                                    Atomic Structure




                          Describe the shapes of each of the 4 atomic orbitals
There are 4 numbers that describe the location/energy of electrons in an atom. The first is the principal
quantum number, n. It has values of 1, 2, 3, . . .

The second is called the angular momentum quantum number. It's symbol is a cursive lower case 'L', and it
has values between and including n-1 and 0. We won't use this, but will use the atomic orbital shape
designations. These are s, p, d, and f, corresponding to 4 different shapes of atomic orbitals.

The third is called the magnetic moment quantum number. It's symbol is mL, and it has values between L
and -L.

The fourth is called electron spin quantum number, and has values of 1/2 or -1/2.

Let's see how to use these numbers. Finsih reading section 5.5, starting at the middle of p. 148.
What is the difference between a Bohr orbit and an atomic orbital?




Look at the quantized whistle on p.151 - it is a good conceptual model for why an electron in an atom must
have only particular energies. Free electrons are not constrained to particular energies.

For n = 1, only an s orbital exists. It is spherical and can accommodate 2 electrons with opposite spins. It is
the lowest energy orbital (closest to the nucleus). Draw an s orbital here.




Now, draw p, d, and f orbitals (one of each). Notice that the number of lobes increases.




How many s, p, d, and f orbitals are there for any one value of n? you could predict this from the quantum
numbers, but it's easier to memorize.




we'll soon find an even easier way to know this - from the periodic table!
Name ___________________________                      11                                     Atomic Structure

                                 Rank the atomic orbitals in terms of energy

                   An important principle - opposite charges attract, like charges repel.
Read section 5.6.
There is a particular order to the energy of atomic orbitals. The more strongly the electrons in the orbital
are attracted to the protons in the nucleus, the lower the energy.
Let's devise a tool to help you memorize that order and show it here. Here's one tool. Use a square grid
(draw one below) and put 1, 2, 3, 4, 5, 6, 7 down the side, and s, p, d, and f across the top. Each box in the
grid represents a possible atomic orbital. Thus the upper left-hand box is 1s. the box directly beneath it is
2s. the box to the right of the 1s box is 1p, etc.




Now, draw a diagonal just to the right of 1s, 2p, 3d, 4f. All orbitals (boxes) to the right of that diagonal do
not exist.

S is always the lowest energy for any given number. Which is lower in energy, 3s or 3p?



P is lower energy than d, d lower than f.



Now for the tricky part. Draw diagonal lines from 2p to 3s, 3p to 4s, 3d to 4p to 5s, etc. Connect from the
low end of each diagonal to the high end of the next diagonal. Energy increases down these diagonals. Which
is lower in energy, 4s or 3d? 3d or 4p? 6f or 7p?




Remember the light bands for atoms? Those represent the energy difference between two of these levels.
Do you think that the energy difference would be different for different atoms? Why or why not?
Name ___________________________                     12                                    Atomic Structure




                        Recognize and present the electronic configuration of atoms
Okay, so what does this have to do with atoms? Let's start with hydrogen and work our way through the
periodic table. We draw a box for each orbital, and begin putting electrons in those boxes up to the number
of electrons available.
H has one atom. Where should it go? The rule is that electrons go into the lowest energy level available
(called the Aufbau principle).




Next is helium, with 2 electrons. Where should they go? Another rule is that each orbital can hold up to 2
electrons, so long as they have opposite spins (called the Pauli exclusion principle). We denote that by using
arrows, and showing them in opposite directions.




Next is lithium. What happens here?




Beryllium?




Boron?




Carbon? We've got a problem here - 2 electrons to put into 3 possible 2p orbitals. How do we do that? We
put the 6th electron in a distinct 2p orbital from that of the 5 th electron. Why?




Notice that the 2 electrons in the 2p orbitals could have the same, or different spins. It has been observed
that having the same spin is lower energy, so both electrons should be drawn pointing in the same direction
(called Hund's rule).

Now you do nitrogen, oxygen, fluorine, and neon.
Name ___________________________                      13                                    Atomic Structure




These arrangements of electrons are called an atom's electron configuration. Rather than using boxes
(which is a good tool to keep track of spins and orbital filling), we can use a shorthand system. See p. 155.
What is the electron configuration for potassium? For calcium?




Notice the choice of elements on p. 155. What kind of elements are these?



Notice that they all have in common an 's1' as the highest energy electron in their configuration. This
common structure accounts for their common properties! Now we can see how the periodic table has been
organized, and why elements within groups exhibit similarities.

What are the electron configurations of F, Cl, I, and Br.




What is the common feature of the electron configurations?




Now, let's color in the periodic table (last page of unit) based on the electron configuration of the outermost
electrons. Use blue for s1, orange for s2, green for d1 through d10, yellow for p1 through p4, red for p5, and
violet for p6. Label each of these groups.

Now, why are the rows grouped as they are. Remember that the rows are called periods. It turns out that
2s and 2p are closer in energy than 2p and 3s. likewise 4s, 3d, and 3p are closer in energy than 3p and 4s. so
the periods are grouped not only in order of filling up the electron orbitals from lowest to highest energy,
but also grouped by how close in energy they are.

Read section 5.7.
What other orbital(s) is 4f close in energy to?




how many electrons can each period hold?
Name ___________________________                    14                                    Atomic Structure




I don't like the shell model shown on p. 157 and figure 5.26 as a way to organize your thoughts about energy
levels, because it looks too much like the Bohr model of the atom, which doesn't correctly indicate electron
location. But if it helps you understand the periodic table, then feel free to use it.

Now, I've said that these electron configurations determine properties of the atoms. Section 5.8 begins to
explore this. Record below your thoughts on the starburst model.




   Utilize the electronic configuration of atoms to predict properties such as tendency to gain or lose
                                 electrons, charge on ions. and atomic radius

Read section 5.8.
Consider atomic size. It would seem that the more electrons you have, the larger the atom should be
(remember that the nucleus is so small that it affects size not at all). Is this true? Would potassium be
larger than sodium?



Thus a group trend is that size increases as you go down a group. However, within a period, the trend has to
take into account 2 factors - nuclear charge and shielding. Nuclear charge is proportional to the number of
protons in the nucleus, so that's easy. Shielding refers to electron repulsion. the more electrons between
the outer one and the nucleus, the more that electron is shielded, or repelled, from the nucleus. The sum of
those two factors is called effective nuclear charge (enc).

 This balancing act between nuclear attraction and electron repulsion determines chemical and physical
                                         properties of atoms.

Here are the trends that you need to know which arise due to this balancing act:
A. ionization energy or the tendency to lose electrons. The greater the ionization energy, the lower the
tendency to lose electrons.
B. electronegativity or the tendency to attract electrons.
C. due to A and B, certain atoms tend to gain electrons, others tend to lose electrons.
D. atomic size

We can use the periodic table to help us through each of these!

A. elements closest to Fr (francium) tend to lose electrons the most; elements closest to He (helium) tend
to lose electrons the least.
        1. as you go down a group, the tendency to lose electrons increases
        2. as you go across a period, the tendency to lose electrons decreases
        3. hydrogen is somewhere in the middle
Name ___________________________                         15                                   Atomic Structure

          4. noble gases do not tend to lose electrons

predict: which is more likely to lose electrons:
Na or Cl                                                 Ca or S



Fe or N                                                  Br or Ba



C or Pb                                                  Cl or I (does this explain what you saw in lab?)



[Ar]4s2 or [Ar]4s23d104p4                                1s22s22p1 or 1s22s22p5



[Ne]3s2 or [Xe]6s2                                       [Ar]4s23d104p3 or [Kr]5s24d105p3




Notice that metals have a greater tendency to lose electrons than non-metals, and that metalloids are in-
between!



B. elements closest to Fr (francium) tend to gain electrons the least; elements closest to fluorine (F) tend to
gain electrons the most.
        1. as you go down a group, the tendency to gain electrons decreases
        2. as you go across a period, the tendency to gain electrons increases
        3. hydrogen is somewhere in the middle
        4. noble gases do not gain electrons

predict: which is more likely to gain electrons:
F or Br                                                  Mg or Ca



Mg or Cl                                                 Al or S



C or P                                                   Mn or Si



[Ne]3s23p5 or [Ne]3s1                           [Ar]3s23p1 or [Ar]3s23p4



1s22s22p5 or [Ar]4s23d104p5                              [Ne]3s23p1 or [Xe]6s25d104f146p1
Name ___________________________                      16                                     Atomic Structure

C. These two trends only refer to one electron either gained or lost. There are other factors which affect
whether more electons are gained or lost.
       1. main group elements tend to gain or lose electrons until they have a full outer shell (transition
metals are much harder to predict)
       2. for metals, this means they lose electrons from their outer shell until it is empty
       3. for non-metals, this means they tend to gain electrons in their outer shell until it is full
       4. metalloids and hydrogen may gain or lose electrons until their outer shell is full

how many electrons do you think each element tends to gain or lose? How can you tell by looking at the
periodic table?
Oxygen                                              calcium



Potassium                                              beryllium



Carbon                                                 bromine



Aluminum                                               lead



[Kr]5s2                                                [Ar]4s23d104p1



1s22s22p4                                              1s22s1



Notice that neon, and all noble gases, already have a full outer shell. If neon were to gain an electron, where
would you put it?




Because of shielding, neon can't attract that extra electron strongly, and so it doesn't gain it.




How about chlorine - where would it put an extra electron, and would it be able to hold onto it easily?




If you gave chlorine two electrons, where would it put the second electron? Can it hold onto that electron?



D. atomic size is related to how many electrons you have and what period you put them in.
       1. atomic size increases as you go down a group
       2. atomic size decreases as you go across a period
Name ___________________________        17         Atomic Structure

which would you predict to be larger?
B or Ga                                 Ca or As



Li or Rb                                F or I



Sn or P                                 Ba or Be