UNIT 4 1

HONORS CHEMISTRY

HARVARD-WESTLAKE

UNIT 4

Atomic Structure,

Electron Configurations

And the Periodic Table,

Periodic Trends

Heisenberg is out for a drive when he's stopped by a traffic cop. The cop says: "Do you know how fast you were going?

Heisenberg replies: "No, but I know where I am".

UNIT 4 2

Reference Sheet

Speed of light (c) = 3.0x108 m/s Planck’s constant (h) = 6.63 x 10-34 J/Hz = 6.63 x 10-34 J•s

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Development of Modern Atomic Theory -- A Summary J.J. Thomson

discovered sub-atomic particles common to all elements in cathode ray tube experiments named negatively charged particles with very little mass electrons positively charged particles with much greater mass called protons

MODEL "plum pudding" model: positive charge of protons distributed throughout atom,

electrons embedded like raisins or nuts in pudding Ernest Rutherford et. al.

designed alpha-particle scattering experiment using radioactive source and gold foil observed that most particles passed through foil undeflected, some slightly deflected and

a very few reflected concluded that Thomson's model was incorrect

MODEL "nuclear" atom: positive, very small center of atom is nucleus containing protons;

electrons move around nucleus like planets around sun; atom is mostly empty space

Niels Bohr

worked out mathematical model of hydrogen atom which explained emission line spectrum of only certain wavelengths using Planck's quantum concept, E = hν

showed that calculations matched observed behavior of electrons as existing on only certain energy levels in the atom, "quantization"

MODEL Bohr atom: similar to Rutherford in structure but electron energy levels fixed at

certain amounts or distances from nucleus; electron transitions from higher to lower levels result in discrete wavelengths of emitted light

Erwin Schrödinger

combined some particle behavior with wave behavior as suggested by de Broglie and formulated mathematical model for hydrogen atom

derived equation which gives the probability of finding the electron at some point in a 3-dimensional space at any given instant; gives no information about the path the electron follows

solutions from equation yield information about probability maps or shapes on different energy levels

MODEL "wave/mechanical": basic layout of atom similar to Bohr and Rutherford, but

electrons do not follow simple orbits; their position can only be predicted in terms of probability

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Atomic Orbitals and Electron Configurations The BIG idea: four types of atomic orbitals emerge from the solutions to the Schrödinger equation. Each type is a different shape. Electrons fill these orbitals to build up the structure of an atom until the number of electrons is equal to the number of protons in the nucleus. The orbitals fill in a predictable way, following a pattern that can be traced on the periodic table. This is the essential information we need to understand why elements have various chemical properties. The ground rules: 1. Each row or period in the periodic table represents an energy level (1 through 7). These

levels are often designated by the letter n (called the principal quantum number).

Thus H and He have electrons in level n = 1. Li, Be, B, C, N, O, F and Ne have electrons in levels n = 1 and n = 2. Na, Mg, Al, Si, P, S, Cl and Ar have electrons in levels n = 1, n = 2 and n = 3. Get the idea?

2. Each orbital (regardless of type) may contain 1 or 2 electrons but no more. 3. The number of types of orbitals is equal to n. The total number of orbitals on an energy level

is equal to n2. 4. The types of orbitals are: s, p, d, f. These are listed in order of increasing energy within an

energy level (i.e., energy levels have sub-levels).

Thus on n = 1, there is only one s orbital. On n = 2, there is an s orbital and three p orbitals. On n = 3, there is an s orbital, three p orbitals and five d orbitals. On n = 4 and higher there is an s orbital, three p orbitals, five d orbitals and seven f orbitals. [the result of this progression is that there can only be one s orbital on a level, three p orbitals, five d orbitals, and seven f orbitals]

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The easier way: All of this information is summarized neatly in the periodic table if you learn to read it correctly. The first two columns of the table consist of s-orbitals. The last six are p-orbitals, the middle 10 are d-orbitals, and the bottom two rows are f-orbitals. The diagram below summarizes this information.

The electron structure surrounding the nucleus is of prime importance in understanding the chemistry of the elements. This structure can be read directly off the periodic table and expressed in abbreviated form. Some examples follow:

H 1s1 He 1s2 Li 1s22s1 Be 1s22s2 B 1s22s22p1 C 1s22s22p2 N 1s22s22p3 O 1s22s22p4 F 1s22s22p5 Ne 1s22s22p6 Na 1s22s22p63s1 OR [Ne]3s1

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Note: although the first d-orbitals appear on the periodic table in row 4, they are called 3d orbitals. Thus the entire d-section of the table is numbered one less than the row in which it appears. A similar offset occurs for the f-orbital block, but it is two numbers less.

1s 2s p d

"n" rules

each energy level is designated by the value n which is an integer from 1 (lowest energy or "ground state") on up

the number of types of orbitals possible on an energy level is also equal to n the maximum number of actual orbitals on an energy level is equal to n2 the maximum number of electrons in an orbital is equal to 2 the maximum number of electrons on an energy level is equal to 2n2

Summary Table

    

           

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To abbreviate a configuration, the previous noble gas may be used as a "core". Na 1s22s22p63s1 [Ne]3s1

Mg 1s22s22p63s2 [Ne]3s2

Al 1s22s22p63s23p1 [Ne]3s23p1

Si 1s22s22p63s23p2 [Ne]3s23p2

P 1s22s22p63s23p3 [Ne]3s23p3

S 1s22s22p63s23p4 [Ne]3s23p4

Cl 1s22s22p63s23p5 [Ne]3s23p5

Ar 1s22s22p63s23p6 [Ne]3s23p6

One last thing: the configurations we have been writing are ideal. As early on as elements like chromium and copper there are anomalies in the regular electron configurations. This gets much worse as the elements become heavier. You are not responsible for all of the exceptions, only the general principles.

Some notes on Periodic Properties: Mendeleev suggested the missing elements solely on what he thought he had discovered: the properties of elements are periodic functions of their atomic masses. There were errors in Mendeleev's table. Some of his atomic masses were incorrect. For example, Te and I along with Co and Ni were reversed in the original table. More importantly, had argon been known at the time, it and potassium would have been switched placing a very reactive metal among a family of inert gases. These problems were eventually corrected by Henry Moseley who also established the concept of atomic numbers. The table was eventually revised in order of atomic number and the law of chemical periodicity thus reads: PROPERTIES OF ELEMENTS ARE FUNCTIONS OF THEIR ATOMIC NUMBERS

Why is this?????? As the table is arranged, electron configurations repeat over and over. The outer electrons for each atom in a column are in the same arrangement. Could this be the factor?

Although we cannot "see" electron configurations there is experimental evidence for them and for their repetition through the periodic table. Some of the clearest evidence is the ionization energies of atoms.

Ionization energy (or, more properly, first ionization energy) is the energy needed to remove the outermost electron from a neutral atom:

X + energy X+ + e-

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Examination of the first ionization energies of the elements reveals that it is a periodic property, that is, it is a function of atomic number and repeats in a pattern that we can relate to chemical properties.

 The "size" of an atom is a curious property for something that does not have a hard surface! Chemists have agreed that atomic size should be considered as the 90% probability volume for the atom.

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Atomic radius continued:

Electron affinity is more or less the opposite of ionization energy:

X + e- X- + energy

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Electron Configuration, Shielding and Periodic Properties The BIG picture: properties such as atomic size (or radius), ionization energy, electronegativity, metallic character, etc. repeat regularly throughout the table because the electron configurations repeat regularly The LAST words on Shielding: inner spherical orbitals (s-orbitals) "shield" electrons on higher levels from the pull of the nucleus partially since they completely surround it. This is most readily observed in the great difference between the ionization energies of the noble gases (high) and the alkali metals (low) that immediately follow them. s-orbitals shield p and d orbital electrons on the same energy level to a much smaller extent. This effect is only really obvious when the ionization energy of an element like Mg is compared to that of Al. The poorer shielding as electrons are added on the same energy level and more protons are added in the nucleus accounts for the decrease in size across a period since the electrons feel the pull of the nucleus more. MORAL: periodic properties are best understood in terms of shielding Some of the major properties to know: 1. size (or radius): decreases from left to right (poor shielding) increases from top to bottom (good shielding)

atoms shrink as you advance across a period because added protons in the nucleus pull the added electrons more tightly in the absence of effective shielding atoms in the same column increase in size as you go down the family since electrons are being added on successively higher energy levels and are well-shielded by the previous completed electron energy levels

2. ionization energy: exact opposite of size

the outer-most electron becomes increasingly more difficult to remove as you advance across a period since shielding is very poor and the added protons continue to exert a stronger attraction the outer-most electron is easier to remove as you move down a column since each time you are one energy level farther from the nucleus, benefiting from that much more shielding

3. electronegativity/electron affinity: same trend as ionization energy

the ability to attract additional electrons increases as you advance across a period since shielding is largely ineffective and the pull of the increased protons in the nucleus is sufficient to influence electrons which approach the outer limits of the atom from other sources

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any attractive force of the nucleus is so diminished by effective shielding of electrons in lower levels that additional electrons become more difficult to hold onto as you move down a column

4. metallic character: same trend as size

metallic properties are related to the freedom of the outer electrons to move in a macroscopic sample and thus these properties are most pronounced at the left side of the table where shielding is at a maximum this means that metallic properties will also increase as you move down a family since shielding improves

Interpreting other facts using these principles: examples 1. The most reactive metal would be Fr. The most reactive non-metal would be F.

In reacting, metals typically form + ions. For this to happen, electrons must be lost. This will be favored by excellent shielding. In reacting, non-metals typically form - ions. For this to happen, electrons must be gained and held. This will be favored by poor shielding.

2. The noble gases are very unreactive. Although it might seem that the noble gases should be very electronegative, recall that they are at the end of their energy levels. Any attracted electrons would have to go into the next energy level, too far away to feel any permanent pull from the nucleus.

3. The common charges or oxidation states MAKE SENSE!

Alkali metals: +1 lone, well-shielded outer electron is removed and the remaining electrons are tightly held

Alkaline earth metals: +2 (as for alkali metals above) Group IIIA: +3 (as above) [why is +1 observed in some Group IIIA

elements?] Group VIA: -2 electrons are too tightly held to remove, but shielding is so

poor that high electronegativity is able to attract enough electrons to fill level (2 more)

Group VIIA: -1 similar to Group VIA

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UNIT 4 13

LAB: Emission Spectra The first evidence supporting the Bohr model of the hydrogen atom came from the striking spectrum produced by gaseous hydrogen atoms under conditions of electrical excitation. The series of visible lines known as the Balmer series (one of a set for hydrogen) is a kind of after-the-fact road map which indicates to what energy levels the single electron in the hydrogen atom has been excited. Electrons excited to higher energy levels give off energy when they 'relax' and return to lower levels. This energy is equivalent to a certain frequency of electromagnetic radiation, as illustrated by Max Planck:

∆E = hν

If this frequency, ν, happens to fall within the visible spectrum we perceive it as a color. In studying this phenomenon the diffraction grating is particularly useful because it enables us to split up what appears to be monochromatic light into any component colors (much the same way a prism splits up white light). How does it do that? The transmission grating you will be using consists of a sheet of transparent plastic that has thousands of tiny parallel groves ruled on it (around 5300 per cm). If you imagine light waves approaching the grating like waves in the ocean breaking on the shore, the lines on the grating act like little gates that the waves must pass through. Forcing the waves to pass through these "gates" causes a complex interference pattern (diagram below).

Where the crests (solid lines) from the "ripples of light" emerging from the gates are in the same place, the waves add to give a more intense wave (brighter). Where crests and valleys (dotted lines) meet, waves cancel and darkness results. If light of one wavelength (such as laser light) passes through a transmission grating, we see the undiffracted beam as well as a series of diffracted beams at angles depending on the wavelength ( ) of the light and the spacing of the rulings on the grating (the "gates"). In the diagram on the left, the wavelength of light is shorter than on the right. Note that the beams of light that emerge from the grating are spaced more closely together than on the right where the wavelength of light is longer. Thus when light composed of several wavelengths passes through the grating it is "broken up" into its component wavelengths since the different wavelengths emerge at slightly different angles from the gates. So when the light from excited hydrogen atoms (which appears to us as purplish) passes through the grating, four sharp lines of light emerge separated by darkness. The wavelengths (and therefore colors and spacing) of lines in a spectrum are characteristic of a substance and can be used to identify it. The science of spectroscopy is based on these principles.

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Preparing to experiment You will be provided with the following materials:

1. a grating spectroscope [see Technique section] 2. solid metal ions: Na+, Li+, Ca2+, Sr2+, unknown 3. gas discharge tubes containing three of the noble gases

Each solid sample is provided in a 24-well plate with moistened cotton sticks. Calibrate the spectroscope as described in the Technique section. Design an experiment to observe and record the visible lines of the metal and use this information to identify the unmarked solid. You will be helping the person who works next to you in lab (introducing the solid into the burner flame) but each student should record his/her observations. Finally, observe and record the spectra of the gas discharge tubes. (hint: begin with NaCl first since its single line near 600 nm is very persistent and easy to observe) Technique 1. The spectroscope The spectroscope has been constructed from an ordinary cardboard box according to the diagram below (the cover has been removed for clarity): How does it work?

Light enters through the slit (defined by the edges of electrical tape). It strikes the transparent grating at an angle and passes through it to the back of your eye, separated into its component wavelengths. As you look through the grating at the arbitrary scale on the opposite end of the box, the image in your eye is superimposed on the scale and you see the lines of the spectrum on the scale! 2. Calibration of the spectroscope To use your spectroscope to identify substances you will need to be able to distinguish one red line from another and so forth. That's what the scale on one end of the box is for. Look at a fluorescent light through the spectroscope. The overhead lights in the classroom or lab will not work because of the diffusing plastic covers. You can, however, use the lights over the counters in the lab. You should see three distinct lines superimposed over a dim continuous background spectrum. These lines are produced by transitions of electrons in the mercury (Hg) vapor that is inside the fluorescent tube. When high voltage passes through the mercury vapor it causes it to glow. The light bulb has a white coating on the inside which responds to

UNIT 4 15

this glow by giving off a mixture of light that is roughly "white" (thus the two spectra superimposed). Those mercury lines are: λ

Line 1: violet (closest to the slit) 436 nm Line 2: green 546 nm Line 3: yellow (sort of smeared) 580 nm Record the position of each line as its image appears on the scale. The dispersion of the grating (how much spreading of the different wavelengths occurs) is essentially linear. So if you make a plot of wavelength (in nm) vs. position of the line (in the arbitrary units of the scale) you will have a graph from which you can read off the wavelength of any line you see in the spectroscope! 3. Observing flame emission spectra You will need a flame that gives off very little of its own light and is of medium heat (small blue cone) and medium length (perhaps 7-10 cm). The hottest part of the flame is just above the blue cone but that will burn off your sample too quickly so you want to place the loop containing the solution in the lower outside part of the flame. This will tend to color the entire flame as the solution evaporates and make observation easier. Even so, these spectra are ephemeral at best. You will probably need to put new sample into the flame several times in order to observe and record each line for a substance. It will be easiest for you to observe the spectra if you have your back to the windows and cup your free hand around your eye so that you shield out stray light from the outside. Because the air/methane flames don't provide much energy, many lines that might be recorded in handbooks for the elements in this experiment will not be visible to you. But you should be able to get a unique "fingerprint" for each substance provided. You should also be aware that sodium is a contaminant in the air and therefore is likely to appear (if only dimly) in each spectrum. Analysis 1. Plot a calibration graph that extends (on the vertical axis) from 400 to 800 nm. Draw the best straight line among the three points you obtain from the fluorescent light data. 2. Prepare a table of wavelengths for the lines observed in each metal ion spectrum and for each discharge tube using your calibration graph. 3. Compare the wavelengths you get from your data with the wavelengths given on the table attached to this sheet. Remember, you may not have observed all lines. 4. Use the table prepared in #2 to identify the unmarked substance. Be sure to indicate the # of the substance! (see spectrum diagrams on following page) 5. Use the table prepared in #2 to match the discharge tubes with their correct noble gases. (see spectrum diagrams on following page) 6. Qualitative analysis is a series of techniques and procedures to determine the presence or absence of elements in a sample. You will do some of this later in the course. Most of the procedures involve the judicious use of solubility rules to precipitate, separate and identify elements. However, the alkali metals are almost always identified by flame tests such as you did in this experiment. Why does this make sense? (think about those solubility rules!!)

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[In the following spectra, dotted lines are variable and/or dim lines; gray patches are either broad bands or regions with many faint lines] Metals:

                                                                                                        800  780  760  740  720  700  680  660  640  620  600  580  560  540  520  500  480  460  440  420  400 

Sodium wavelength, nm   

                                                                                                        800  780  760  740  720  700  680  660  640  620  600  580  560  540  520  500  480  460  440  420  400 

Lithium wavelength, nm   

                                                                                                        800  780  760  740  720  700  680  660  640  620  600  580  560  540  520  500  480  460  440  420  400 

Calcium wavelength, nm   

                                                                                                        800  780  760  740  720  700  680  660  640  620  600  580  560  540  520  500  480  460  440  420  400 

Strontium wavelength, nm

Noble gases: 

                                                                                                        800  780  760  740  720  700  680  660  640  620  600  580  560  540  520  500  480  460  440  420  400 

Helium wavelength, nm  

                                                                                                        800  780  760  740  720  700  680  660  640  620  600  580  560  540  520  500  480  460  440  420  400 

Neon wavelength, nm  

                                                                                                        800  780  760  740  720  700  680  660  640  620  600  580  560  540  520  500  480  460  440  420  400 

Argon wavelength, nm  

                                                                                                        800  780  760  740  720  700  680  660  640  620  600  580  560  540  520  500  480  460  440  420  400 

Krypton wavelength, nm  

                                                                                                        800  780  760  740  720  700  680  660  640  620  600  580  560  540  520  500  480  460  440  420  400 

Xenon wavelength, nm

Example Data Table:

Color : Scaled Line

Gas Tube #1

Gas Tube #2

Gas Tube #3

UNIT 4 17

Instructions for Graphing with Microsoft Excel 1. Double-click the Microsoft Excel icon. A new spreadsheet will appear with vertical columns,

horizontal rows, and a tool bar along the top. In the first column, type your X (horizontal) axis data into each cell. In the second column, type in your Y (vertical) axis data.

2. Position the mouse pointer (it looks like a cross) on the first cell in the first column and press

the left mouse button. Hold this button down and drag the pointer diagonally to the bottom Y-value in the second column. The section will now be highlighted (white on black). The first cell is also selected but will not look like it.

3. From the tool bar, select Insert. From the Insert menu, select the Scatter plot option and

among the types of Scatter plot, select Scatter with only Markers as your plot of choice. Format your graph using the following steps:

a. Delete the legend (Series 1 box) by clicking on it and hitting delete on your keyboard.

b. Under the Chart Tools toolbox menu (only visible when you have clicked on the graph),

select the Layout tab. This will allow you to label your axes and title your graph.

To title your chart, select Chart Title under the Layout tab. Select the Above Chart option. Click and type directly in the chart title text box to re-title your chart. Recall: Graphs should always be titled in a Y vs. X format!

To title each axis, select Axis Titles under the Layout tab. For the x-axis, select Primary Horizontal Axis Title and then Title Below Axis. For the y-axis, select Primary Vertical Axis Title and then Rotated Title. Then, click and type directly into each axis title text box to re-title each x-axis.

c. To zoom in on your data points to present a nice, clear graph, right-click on any of the number values on the x-axis. Select Format Axis… From the pop-up menu, change your minimum and maximum axis options from Auto to Fixed. Then, choose minimum and maximum values that will nicely frame your x values. When done, hit the close button on the pop-up menu. Repeat this process with the y-axis.

d. To add a trendline, right-click on any of the data points on your graph and select Add Trendline… From the trendline options menu in the pop-up menu, select Linear. Also, check the boxes for Display Equation on chart and Display R-squared* value on chart. When done, hit the close button on the pop-up menu and reposition your equation so that it does not cover up any of your data points.

e. To move your graph to full page size, right-click on any white space on the chart outside of the graph and select Move Chart… From the pop-up menu, select New Sheet and hit OK. Your full-page graph will now have its own tab along the bottom of the worksheet.

4. Print your finalized graph by selecting the File tab in the top toolbar, followed by Print

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NOTE: the R2 value is an indication of how well your data conforms to the type of function selected; if you select Linear and your data is not very scattered, R2 should be close to 1.0 Excel is a very powerful piece of software and this introduction to the most rudimentary graphing skills barely scratches the surface of what is possible. If you are curious about what Excel can do, feel free to explore the Help menus or ask your teacher about specific things you'd like to be able to do. A.P. Chemistry students are also a good source of information since they have been trained in many uses of Excel.

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Name:___________________________ Per.:____ Date:_______________   

Investigating Periodic Properties

Purpose: to observe some reactions of alkali metals, alkaline earth metals, halogens, and period three elements and attempt to correlate the observations with known periodic properties such as shielding, atomic size, etc. Method: I will first observe a series of demonstrations by the instructor illustrating the behavior of Li and Na with water, chlorine and bromine, and some physical properties of those elements. I will then test the behavior of period 3 metals with distilled water. If there is no apparent reaction I will then try boiling water. I will also test the conductivity of the metals. I will test the remaining period three elements (except chlorine) for water solubility (and a possible acid/base reaction with water) and also for electrical conductivity. The halogens (chlorine, bromine and iodine) will be tested for water solubility and their strengths as oxidizing and reducing agents. Finally I will test the water solubility and acid/base character of the period 3 oxides. I will use these observations to write reactions where seen and to try and correlate the behaviors with various periodic properties. Data:

Table 1: behavior of alkali metals

with water with chlorine with bromine conductivity misc.

Li N/A N/A

Na

K N/A N/A

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Table 2: behavior of alkaline earth metals

with water with hot

water acidity? conductivity misc.

Mg

Ca

Table 3: behavior of remaining period 3 elements

with water with hot

water acidity? conductivity misc.

Al

Si

P N/A N/A N/A

S

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Table 4: behavior of period 3 oxides

water solubility acidity? physical state

Na2O

MgO

Al2O3

SiO2

P4O10

SO2

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Table 5: behavior of halogens

water solubility oxidizing/reducing

Cl

Br

I

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LAB: Investigating Periodic Properties The chemical and physical properties of elements (and their similar compounds) generally change gradually with atomic number. We say they are periodic because this change in properties repeats as a function of atomic number (every eight, not counting transition metals and inner transition metals). From your studies in class and your reading you will know (or soon know) that this regular change in properties is tied to the regular change in electron structure as one moves from left to right in a period of the table or as one moves from top to bottom in a group. But it is one thing to mention or memorize such properties and another to see them face-to-test tube, so to speak. So in this experiment you will have an opportunity to see some of them and try to explain why they change as they do. There are so many fascinating properties to choose from and there is only a limited amount of time and certain constraints of safety and economics. So we have chosen the following survey for you:

1. Properties within the alkali metal group* 2. Properties within the alkaline earth metal group 3. Properties within the halogen group 4. Properties of period 3 elements 5. Properties of period 3 oxides

*your instructor will demonstrate the behavior of Li and Na The properties to be examined for metals are the following:

1. reaction with room temperature water --- check acidity or alkalinity of resulting solution 2. if no reaction in (1), then with boiling water --- check acidity or alkalinity of resulting solution 3. electrical conductivity

In each case, only a small amount of metal is required and only a small amount of water. If no reaction occurs at first, the same sample of metal may be used for the next test. Gases that may be produced as metals react can be assumed to be hydrogen. Your instructor will demonstrate the detection of such gases with larger samples. The properties to be examined for the halogens are the following:

1. physical appearance 2. reactions with alkali metals* 3. solubility in water 4. oxidizing/reducing strength

*your instructor will demonstrate

The properties to be examined for the remaining period 3 elements include:

1. physical appearance 2. reaction with/solubility in water --- check acidity or alkalinity of resulting solution 3. electrical conductivity

The properties to be examined for the period 3 oxides are:

1. reaction with/solubility in water --- check acidity or alkalinity of resulting solution 2. physical state at room temperature

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You will notice that some of these test groups overlap. For example, Na is a metal, a member of the alkali metal family and a period 3 element. If you prepare your observation tables correctly, you may be able to eliminate duplicate tests, or just fill in the information rather than repeating the tests. Halogen generation and indicators adapted from: Small Scale Chemistry Laboratory Manual, Edward L. Waterman & Stephen Thompson, Addision Wesley, 1995 Preparing to experiment You will be provided with the following materials:

1. samples of Li†, Na

†, Mg, Ca, Al, Si, P

†, S

samples of Na2O‡ , MgO, Al2O3, SiO2, P4O10, SO2

*

2. bromthymol blue indicator 3. 2 M H2SO4 4. 0.5 M Na2SO3 5. 5% NaClO (household bleach) (to produce Cl2) 6. 5% KBrO3 + KBr (to produce Br2) 7. 5% KIO3 + KI (to produce I2) 8. 3 M NH3 9. redox indicating solutions [see Technique] 10. conductivity device 11. room temperature boiled distilled water 12. boiling distilled water 13. 24-well microplate 14. petri dish 15. hand lens

†to be displayed by instructor

‡ supplied already dissolved in water [see Chemicals]

*see Technique

Study the Technique section carefully and review the procedures recommended there before laying out your experiments. In your notebook you should plan out each section clearly. Tables will be helpful in organizing your observations. Technique 1. Acidity and alkalinity This is more of a reminder than anything else. Bromthymol blue is a 3-color indicator which is yellow in acid solution, green in neutral and blue in base. Solutions can be tested by simply adding a drop or two of the indicator. When gases are being tested (e.g., the gaseous oxides of period 3), it is easiest to place a drop of the indicator in the container where the gas will be generated. If any significant dissolving occurs and the acidity changes, the drop will change color. Thus a color change denotes both solubility in water and a change in acidity. 2. Redox indicators The redox indicators contain I- and starch (to indicate when oxidation has occurred) or I2 and starch (to indicate when reduction has occurred). When the I- is oxidized to I2 it forms a dark blue color with the starch present. The darker the color, the stronger the oxidizing agent. In the presence of a reducing agent the blue indicator fades or

UNIT 4 25

becomes colorless as the I2 is reduced to I-. As with the acid/base indicator, a color change indicates both solubility in water and the presence of an oxidizing or reducing agent agent. 3. Generating toxic gases safely in the lab--small scale In the past you have worked with chlorine and bromine in solution but for this experiment you will need them in gaseous form. A plastic petri dish (often used in microbiological work to grow bacteria) works well for this. Taking chlorine as an example the process works something like this:

a. place a drop of NaClO (the source of the Cl2) in the center of the bottom half of the dish b. to test solubility in water and redox activity place a few drops of redox indicator solution at various

places in the dish, but not touching the reactant c. hold the cover of the dish at an angle almost closing the dish and add a drop or two of 2 M H2SO4 to

the central drop of NaClO; close the lid quickly d. after observations have been made, open the lid slightly and add a drop or two of 3 M NH3 to the

central drop in order to neutralize the mixture and stop the production of gas; close the lid for a minute or two

e. clean out the dish with water and dry thoroughly before reuse

Similar techniques apply to the other gases. The reactants all include a source of H+ (2 M H2SO4) and the following:

1. for Cl2-----NaClO 2. for Br2-----KBrO3 + KBr 3. for I2------KIO3 + KI 4. for SO2-----Na2SO3 [SO2 is an oxide of period 3]

The Chemicals Lithium is a light metal in the alkali metal family which occurs to the extent of about 0.005% in the earth's crust. It is prepared by electrolysis of molten salts. Typical ores contain 3-10% lithium. The metal is silvery white, but tarnishes on exposure to air, acquiring a yellowish tint. Like some other alkali metals, lithium dissolves in liquid ammonia, giving a blue solution. It has been suggested that this color results from the "solvation" of the valence electron. It is unaffected by oxygen at room temperature but burns with a brilliant white flame when heated above 200oC, forming Li2O. Lithium and its compounds impart a carmine-red color to a bunsen flame. It is used in the manufacture of alloys, especially hardened bearing metals, in catalysts for the manufacture of some plastics and in fuels for aircraft. Lithium salts are used in porcelain enamels. Lithium compounds are also used in the treatment of manic psychosis, but their toxicity (especially if sodium intake is limited) is significant. Sodium was first prepared by Davy in 1807 by electrolysis of fused sodium hydroxide. It constitutes 2.83% of the crust of the earth but does not occur free. It is a light, silvery metal which is lustrous when freshly cut but tarnishes on exposure to air. Sodium is soft at ordinary temperatures. It reacts vigorously with oxygen, burning with a yellow flame and combines directly with most non-metals. Sodium compounds such as the cyanide, azide, peroxide, etc. have been used in the past to manufacture tetraethyllead--the former anti-knock and "octane" boosting additive in leaded gasoline. Sodium vapor is used in street lamps and it can be alloyed with potassium metal for use as a heat transfer coolant. Calcium is an alkaline earth metal, fifth most abundant in the crust of the earth (3.64%). Sea water contains about 400 g/ton. It can be produced by the electrolysis of molten calcium chloride. Like sodium, it is lustrous and white when first cut. Unlike sodium, it may be handled occasionally for short periods with bare (dry) hands in the laboratory. It burns, when finely divided, with a crimson flame. It is harder than sodium but softer than magnesium.

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Calcium can be used to harden lead for bearings and alloyed with cerium to make flints for cigarette and gas lighters. Aluminum is a tin-white metal with can take and hold a very high polish. It oxidizes superficially in moist air. Finely divided aluminum will burn and may cause explosions. The pure metal or its alloys are important structural materials where weight is an important factor, such as in aircraft. Aluminum is an excellent conductor and is used in some wiring. The coarse powder may be used for the thermite process and fine powder is used in flash photography. Aluminum has been used in skin protective pastes but it can result in contact dermatitis and even bronchial asthma. Recently some studies have implicated ingested aluminum as playing some part in the development of ailments such as Alzheimers. Silicon does not occur free in nature but is abundant in silica (sand, quartz, sandstone) and as silicate minerals. It is the second most abundant element on earth (27.6%). Only oxygen is present in greater proportions. It is black to gray and lustrous. Silicon is used in the manufacture of transistors, diodes, etc. It can also be alloyed with metals such as copper, iron, and tin. Silicon appears to be biologically inert, although prolonged exposure to its dust may cause pulmonary irritation. Sulfur is known by the ancient name of brimstone. It occurs in the free state as well as in sulfides and sulfates. It exists in several allotropic forms: two crystals and at least two amorphous and two liquid forms. It is used in the manufacture of sulfuric acid, insecticides, plastics, vulcanization of rubber, and the synthesis of dyes. Sulfur has low human toxicity but it may cause skin irritation and injury to the lungs from prolonged breathing of sulfur dust has been reported. Sodium oxide is a white amorphous substance which combines violently with water. Above 400oC it decomposes into sodium peroxide (Na2O2) and the metal. Aluminum oxide or alumina is used as an adsorbent, desiccant, and abrasive. It can also be used in making refractory materials for high temperature work. High concentrations of dust are harmful to the lungs. Some special grades of aluminum oxide can be used in column chromatography for the separation of mixtures. Silicon dioxide is used in the manufacture of glass, sodium silicate ("water glass"), refractory materials, abrasives and molds for casting. Tetraphosphorus decoxide or diphosphorus pentoxide (it actually exists in the dimer form, but is often labeled as the monomer) is an extremely deliquescent white powder that is used as a drying and dehydrating agent. It is a strong irritant and very corrosive. In any but small amounts it reacts violently with water and should be handled with extreme care. Sodium sulfite is used chiefly in photography in place of "hypo" (sodium thiosulfate) both in developers and fixer. It also has applications in bleaching. Sodium hypochlorite is an unstable solid, but its solutions are very stable, liquid Clorox being a good example. It can be used as an antiseptic and anti-fungal agent. Ingestion may cause corrosion of mucous membranes. Prolonged skin contact may result in irritation. Household bleach contains about 5% NaClO and should never be mixed with other cleaning agents since chlorine gas will be produced by acidic additions and chloramines may be produced by addition of ammonia. Both are toxic gases. Potassium bromate is a bread and flour improving agent ("bromated flour" is often sold as "bread flour"). Ingestion may cause vomiting, diarrhea, and renal injury. Potassium iodate is used as an oxidizing agent in chemical analysis and has been used as a topical antiseptic. In animal feed it is sometimes used as a source of iodine.

UNIT 4 27

Analysis 1. Write balanced net-ionic reactions for Li and Na with water.

a. What type of reaction is this? (precipitate, redox, etc.) b. What conclusion can you draw about the trend in reactivity for alkali metals based on your observations of

these reactions? c. Give at least one explanation for this trend based on some periodic property that is relevant to the type of

reaction you selected in (a) [hint: what is happening to the alkali metal in the reaction?] 2. Same as #1, but for Mg and Ca. 3. Now consider the horizontal trend in period 3 for the three metals examined (Na, Mg, and Al). Compare the

reactivity with water and explain the apparent trend. 4. Based upon the demonstrated reactions of Na with Cl2 and Br2, describe the reactivity trend within the halogen

family (at least as it relates to alkali metals). Explain this trend in light of some relevant periodic property [hint: what is happening to the chlorine or bromine during the reactions? check oxidation numbers!]

5. Write a balanced redox reaction for the production of Cl2 when ClO- is combined with H+. The skeleton half-

reactions are:

ClO- Cl2 ClO- ClO3

-

a. Based on your observations, are the halogens all equally soluble in water? b. Account for your observations in (a) based on at least one relevant periodic property. [hint: in order to

dissolve, the halogen molecules must "fit" into the empty spaces in the liquid water] c. Based on your observations, do the halogens all have the same strength as oxidizing agents? d. Account for your observations in (c) based on a relevant periodic property [hint: for a halogen to be an

oxidizing agent what must it be able to do?] 6. Account for your observations for the trend in conductivity of elements in period 3 with reference to at least one

periodic property. 7. Even chemical compounds may exhibit periodic properties if most variables remain unchanged. The oxides of

period three elements illustrate some periodic behavior that may not have been discussed in class. What trends do you observe in these compounds in terms of solubility in water and the acidity of their resulting solutions?

8. Chlorine forms several oxides, none of which were tested in this experiment since they are all dangerous to

handle. Based on your observations of the other period 3 oxides, what phase would you expect the oxides of chlorine to exist in at room temperature? Would you expect them to be soluble in water, and if so, would the solutions be acidic or basic?

UNIT 4 28

UNIT 4 29

Unit 4 Sample Test

The test will be similar in format to previous tests with five multiple choice questions, three required problems, a choice of two out of four chemical reactions to write and one essay question. Remember that on this test the first three solubility rules will be missing. Values for the speed of light (c) and Planck's constant (h) will also be given. The following are representative of typical multiple choice questions but do not necessarily indicate topics to be addressed on your actual test. _____1. The correct electron configuration for Mg is

a. 1s22s22p8 b. 1s22s42p6 b. 1s22s22p43s4 d. 1s22s22p63s2

_____2. As one moves down the periodic table, which of the following properties increases? Which decreases?

a. metallic character b. ionization energy c. electron affinity d. atomic radius

_____3. Which of the following statements are true? Which are false?

a. There are fewer metallic elements than non-metallic b. metals tend to have low electron affinities c. metalloids lie along a diagonal line in the periodic table which runs from upper

right to lower left d. non-metals are located toward the upper right corner of the table

_____4. Atoms form bright-line spectra when some of their electrons

a. are expelled from the atom b. move from ground to excited states c. fall from higher energy levels to lower energy levels d. travel in stationary orbits

_____5. Francium would be the most reactive element in group IA. This is so because

a. francium is rare b. francium has the lowest ionization energy c. francium ions are unstable d. francium has the largest atomic mass

UNIT 4 30

The next section consists of representative problems which might be found in the problem section. All students are expected to work on all 3 of the required problems. 6. In the Pfund series of lines for the excited hydrogen atom, a transition from n = 6 to a lower energy level has a frequency (ν) of 4.02 x 1013 Hz. Determine the wavelength (λ) of this EMR, and calculate the energy of the transition. On the diagram below, indicate with a properly oriented arrow the transition described above. ∞ ______________________________ 6 ______________________________ -6.05 x 10-20 J 5 ______________________________ -8.72 x 10-20 J 4 ______________________________ -1.36 x 10-19 J 3 ______________________________ -2.42 x 10-19 J

2 ______________________________ -5.45 x 10-19 J 1 ______________________________ -2.18 x 10-18 J 7. The following information concerning four consecutive--by atomic number-- (not necessarily in the same row) elements in the periodic table was collected as a result of laboratory experiments. The letters have been assigned arbitrarily: Element A reacts vigorously with water producing hydrogen gas. A also reacts with oxygen to form the compound A2O Element B is very unreactive and is monoatomic and gaseous at room temperature Element C exists as a diatomic gas at room temperature; reacts with A to form AC Element D reacts with C to form a compound DC2. It also reacts slowly with water to form hydrogen gas. ________________a. Assuming that not all these elements are in the same row, but are

consecutive, put them in order from lowest to highest atomic number. ________________b. Which of these elements is a halogen? ________________c. Which of these elements is an alkaline eart metal?

                                        Energy Level 

UNIT 4 31

8. Write out the electron configurations for these elements: (you may abbreviate)

Tc: ______________________________________________________________________ Ge: ______________________________________________________________________ V: ______________________________________________________________________ P: ______________________________________________________________________

9. Given: Element A is in the second period, Group I Element E is in the same row, Group V Elements J, D, M, and R are in period 3, in (respectively) Groups I, IV, VI, VIII Element L and G are in period 4, in Groups (respectively) III and VII

_____a. What would be the most probable formula for a compound of L and M? _____b. Select the element with the highest ionization energy. _____c. Select the element with the most metallic character.

10. When synthesized, the halogen in period 7 will have Z = ____.

At room temperature, this element will probably exist in the ____ (solid, liquid, gas) state. It will most likely have ________ (high, low) reactivity compared to other elements in its family.

UNIT 4 32

The next section consists of representative reactions to complete and write balanced net-ionic equations for. Note that some reactions do not occur in aqueous solution and thus molecular equations are all that would be needed. Each student is expected to choose two from this section. Recall that an additional solubility rule will be missing from the list this time. 11. For each of the following, complete the word equation, write a balanced net-ionic reaction, and tell what type of reaction it is. (reactions b and c do not occur in aqueous solution) a. potassium metal + water ____________________________________________________ ______________________________________________________________________________ ______________________________________________________________________________ type:__________________ b. chlorine + lithium _________________________________________________________ ______________________________________________________________________________ ______________________________________________________________________________ type:__________________ c. calcium + oxygen gas ______________________________________________________ ______________________________________________________________________________ ______________________________________________________________________________ type:__________________ *d. chloride ions and nitrate ions are combined in an acidic solution; nitrogen dioxide gas and chlorine gas are among the products ______________________________________________________________________________ ______________________________________________________________________________ ______________________________________________________________________________

The final section of the test will consist of one essay question selected from the following topics: --atomic models of Thomson, Rutherford, Bohr, and Schrodinger --predictions based on periodic trends

--electron configuration and periodic properties