Atomic Structure

Why does O₃ get broken down in the upper atmosphere and O₂doesn’t?

• Must look at chemical bonds in these molecules
  
• Must look at their atomic structures

Look at Periodic Table: there are two parts to it: on the left and right, we have "high rise apartments"; in the middle, "low rise". Today, we are looking at the "high rise" areas.

All elements have 2 numbers associated with them: a whole number and one with decimals.

Whole number = the atomic number . This is the defining characteristic of the element. This number also indicates the numbers of protons and electrons associated with it.

Decimal number = a weighted average of the different versions ( isotopes ) of the atom. This number is the atomic weight (or mass) of the element. (Ex: Carbon, which has an atomic weight of 12.0112 according to the periodic table, can exist as isotopes C¹³ and C_{14. But, there is much more C¹² than the other isotopes, so the atomic weight (or mass) = 12.0112)}

How to get the number of neutrons in an atom:

1) Round off the atomic weight (mass) to the nearest whole number
2) The atomic number has given you the number of protons/electrons 3) Subtract the atomic number from the atomic mass = number of neutrons

Ex: Fluorine (F)

Atomic # = 9
Atomic mass = 19
Mass # = 18.9988
Number of protons (electrons) = 9

Atomic mass (19) minus # protons (9) = 10. Therefore, fluorine has 10 neutrons.

Ex: Sodium (Na) Atomic # = 11
Atomic mass = 23
Mass # = 22.989
Number of protons (electrons) = 11

Atomic mass (23) minus # protons (11) = 12. Therefore, fluorine has 12 neutrons.

Isotopes are like "cousins" to atoms --- they can trace the same lineage. All isotopes of element have the same number of protons, but different numbers of neutrons. Because they have the same number of protons, they have the same chemical properties.

Ex --- Different isotopes of Hydrogen (H)

¹₁ H (this is protium)

where the superscript number = atomic mass, and the subscript number = atomic number

protons = 1
electrons = 1
neutrons = 0

Ex --- "Heavy Water": H₃O

This molecule is made up of deuterium, an isotope of hydrogen:

²₁ H

(again, the superscript number = atomic mass, and the subscript number = atomic number)

protons = 1
electrons = 1
neutrons = 1

Last example: "Tritium": H₃ (This isotope is radioactive!)

³₁ H

(again, the superscript number = atomic mass, and the subscript number = atomic number)

protons = 1
electrons = 1
neutrons = 2

A chemical reaction is an exchange of electrons. (If reaction occurs with protons or neutrons, result is an atomic explosion!)

The arrangement of electrons determines how an atom reacts. Electrons are found in "shells".

Shell #1 (innermost shell, closest to the nucleus) contains 2 electrons
Shell # 2 contains up to 8 electrons
Shell # 4, 5, etc. --- "Happy" to have 8 electrons, but can have more

This arrangement gives rise to the octet rule which states that an atom will do what is has to with its electrons (lose, gain or share) to have 8 electrons in its outermost shell.

Valence shell --- outermost shell of electrons

If we look at Group 1A on the periodic table, we can figure out how many electrons are in the out shell of each element, based on its atomic weight (same as the number of its electrons):

Na 11
# of electrons = 11
Inner shell - 2 electrons
Next shell - 8 electrons
Outer shell - 1 electron
Hydrogen 1
# of electrons = 1
Inner (outer also) shell - 1 electron
Li 3
# of electrons = 3
Inner shell - 2 electrons
Next (outer) shell - 1 electron

If the atom falls into Groups 1, 2 or 3 on the periodic table, it has 3 or less electrons in the outer shell and wants to give them away. If these electrons are "stripped", the atom becomes charged or electron poor.

Ex: Na looses one electron in its outer shell. It now has 10 electrons and 11 protons, giving it a single "+" charge, i.e., Na ⁺

Carbon has 4 electrons in its outer shell. It neither gives nor takes electrons, but rather shares. All elements that come after C in Groups 4-8 in the periodic table take in atoms into their outer shells, thus giving them a "-" charge for each additional electron. These atoms are called ions, and are electron rich.

Next time: The types of bonds in molecular structure (Lewis structure)











January 25, 1999 --- Dr. Bunce

Atmospheric Chemistry (Continued)


Reminders

Lab report #2 is due today, please hand in at the beginning of class.
Homework, Chapter 2 is due next Wednesday, 2/3
Lab report # 2 is due Monday, 2/1


Changes

The test will cover only Chapters 1 & 2. It will be in essay format and there may be problems to answer. A review sheet has been passed out to emphasize the importance of knowing and understanding concepts versus just memorization.

There will be no homework # 3 and no in-class assignment.

Today’s Schedule of Material

2:10-4:00:

• Lewis Structure
  
• Electron Magnetic Radiation
  
• O3/O2 Shield
  
• Chapman Cycle
  
• Preview of Lab # 2


Lewis Structures

Lewis structures are a way if illustrated the arrangement of the valence electrons in an element. All elements want eight electrons in their outer valence shell except for hydrogen and helium which only need two.In order to make these Lewis structures there are rules that need to be followed:

Rules of Lewis Structures for elements

1. Only show the valence electrons
2. Each dot represents 1 valence electron

For example: Chlorine (Cl)

First ask, how many valence electrons does chlorine have?
Answer: 7 (Found in row VII)

Now to write the Lewis Structure. Imagine a box around Chlorine.

Cl Now place one dot at each side going clockwise first, before doubling up of dots on one side.

Ionic Bonding Showing Lewis Structures

For example NaCl

. . .
Na : Cl :
. .

The only thing that holds these two elements together is a positive/negative attraction for one another. Why is the Na have a (+) charge and chlorine a (-) charge? Na is a (+) charged ion because it has 11 protons and 10 electrons. Therefore, it has an excess of one positive charge in the nucleus which gives it a (+ 1) charge. The reverse is for chlorine. Chlorine has 17 protons in its nucleus and 18 electrons in its outer shell so it has an excess of one (- ion) giving it a slightly negative charge.


Covalent Bonding

Covalent bonding is the sharing of electrons between elements. An example of a covalent bond is the bond between diatomic elements, such as Chlorine to make Cl2 . The difference between ionic bonding and covalent bonding is that in ionic bonding, electrons are notshared. Usually ionic bonds occur between elements that are on opposite sides of the periodic chart, such as rows 1 and 2 with rows 6 and 7. Covalent bonding occurs between elements that are closer to one another on the periodic table.

Additional rules for making Lewis Structures in Covalent bonding

1. Find the sum of all valence electrons
2. Draw a skeletal structure connecting elements with bonds. Each bond signifies two electrons
3. Subtract the number of electrons used in making the bonds.
4. Distribute the remaining electrons so that each element is surrounded by eight electrons, except of course, hydrogen and helium which are only surrounded with two.

Example # 1: Cl2: The sum of the valence electrons are 2 Cl x 7 e- = 14 valence electrons (remember to only count the ones in the outer shell. )

Skeletal Structure Cl--Cl

The bond between the two chlorines uses two electrons so we have 12 left to distribute around the chlorines.

.. ..
:Cl--Cl:
. . . .

Example # 2

H20 Total valence electrons 8

Connect the elements by the following additional rules:

1. If you have only one of a particular element, it is usually placed in the middle. 2. In almost every case, except for the structure of water, Oxygen is usually not in the middle. 3. After you have subtracted the electrons involved in the bonds, distribute the remaining electrons. If there is not enough electrons to put around each element, go back and a a double bond somewhere. If that isn’t enough, add another double bond. If there is still not enough remaining electrons, take awy the double bonds and make a triple bond.

Therefore, H--O--H

2 bonds, 4 electrons are used and there are 4 left over. Both hydrogens are surrounded by 2 electrons (which is all that they need.) The remaining four elctrons go around the oxygen.

.. ..
H--O—H

Example # 3: O2

Valence electrons 2 O x 6 electrons = 12

.. ..
Skeletal Structure: :O = O:

Now try yourself with Nitrogen, N2

Answer: :N=N:


The Lewis Dot Structures of O2 and O3, a comparision

.. ..
O2 = :O = O:

.. .. .. O3 = : O = O -- O : ..

The double bond in the Ozone molecule could be put on either side of the middle oxygen. at a result this is a hybrid which we represent in the following manner with a double arrow in both directions.

.. .. .. .. .. ..
:O = O -- O: :O --O = O:
.. ..

These are called resonance structures. The arrows indicate the flexibility in placement of the double bond. This resonance gives stability to the molecule.


Electromagnetic Radiation

Radiation is projected in the form of a wave. To understand how radiation is projected we need to look at the structure of a wave. The wavelength is the distance from one crest(top off a wave) to another. The greek letter lambda *. Every wavelength has a different frequency. The frequency is the number of waves that passa particular point per second. Something that has a long wavelength has a very low frequency and vice versa. The unit for frequency is the Hertz (Hz). Cosmic rays have high frequencies and a short wavelength. Television waves on the other hand have a low frequency but a very long wavelength.

[image]





Visible Light

When we see visible light we see a rainbow. The colors in the rainbow from the reds to the violets are always arranged in the same colored schematic order. Reds are on the lower end of the sepctrum, they have very long wavelengths but low frequencies. Violet on the other hand has a high frequency but a very short wavelength. Our sun enits electromagnetic light in the form of ultraviolet (UV) rays, infrared rays and visible light. The energy that is emitted by the son is not always continuous. The path is actually very discontinuous. At one point it was argued whether light travels in a particulate manner or as a wave. The great scientist Einstein tried to combine the two perspectives in his famous equation:

E= h *

Where E is equal to energy, h is a Planks constant (6.63x10-34 Joules/sec), and * is a given frequency.

Using the above equation we can compare an FM channel of 100 MHz to a UV photon that causes sunburn.

FM Channel	UV Radiation
E = hv
Convert MHZ to HZ = 106HZ
E = (6.63 x 10-36)(100 x 10
E = 6.63 x 10-26 J
		0

FM Channel UV Radiation E = h * E = h * Convert MHZ to HZ E =(6.63 x 10-34 Jsec)(1.00x1015) = 106HZ E = 6.63 x 10-18 J E = (6.63 x 10 -36) (100x106 HZ) E = 6.63 x 10 -26 J

We find that the only main difference between a photon of UV light and a photon of the radio is that the radio photon does not make as much energy. This is why we do not get a sunburn from listening to the radio but you can from going to the beach.

Pre-Lab Discussion for Friday

Acid-substance that releases H+
Base-substance that releases OH-

Acetic acid is a weak base with the chemical form CH3COOH

NaOH is a strong base

We are going to work with a 24 well plate, add H+ and a dye phenylthalein which is an indicator. It is usually clear but in the presence of OH it turns pink. Each time you add an OH- to an acidic solution the OH joins with a free H+. However, when there are no more free H+ groups the solution shifts to a basic environment and the indicator would turn pink. This first change of color is referred to as the endpoint. It is the first visible sign that you have just passed the neutralization point. The process of slowly adding OH to neutralize is called titration. You will be titrated and using two different types of vinegar on Friday. The calculation can be found at the end of lab 2. Focus your attention only on equation # 7.

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