Heater controls E.C.C. Fits: 2002 Volvo S40 4DRS W/O S.R | Jim Ellis Volvo of Marietta Parts

The objective of this experiment is to demonstrate and explore how chemical reactivity is affected by variations in these properties.
~Atomic radius, ionization energy, electron affinity, effective nuclear charge and electronegativity are a just a few such parameters.

Part A of this lab deals with the solubility of dissolved ions.
The two primary factors involved when an ionic solid is dissolved (dissolution) in water are dissociation and solvation.
--Dissociation, the separation of the ions from the crystalline solid, is related to the Lattice Energy..
This is an endothermic process as the stability of the ionic interaction forming the crystal is overcome when its constituent ions are energized into the gas phase. Lattice Energy essentially is a measure of how well the ions pack into a crystal – affected by ion charge, size and distribution.
--This is followed by hydration, or solvation of the ions with water:

alkali earth cations will be added to a variety of anions. By choosing large anions and alkali earth cations, the lattice energy of the potential precipitated salts will be relatively low (A). [Why are Lattice Energies low with large anions?] Additionally, anions have low heats of hydration. The formation of precipitates will be a function of the heats of hydration of the cations.

Part B will examine the interactions of halogens with halides.
The reactions will take place in water. Addition of another solvent, hexane, will allow you to monitor the reaction [(B) Why would hexane interact differently than water? How would each solvent interact with halogens? With halides?]
By examining several such reactions one may get a feel for how variations in the halogen (or halide) affect reactivity.

EXPERIMENTAL
Part A:

Atomic Number
In the modern periodic table, the elements are actually arranged in order of increasing atomic number--that's the number of protons in one atom of a particular element. An undisturbed atom is electrically neutral, so the number of electrons in it is the same as its atomic number.
Atomic weight almost always increases with atomic number, so Mendeleev's sequence of elements was almost exactly the same as the one used today, though there are a couple of weird exceptions. In general, it's correct to think of atoms getting heavier as you go down a column or to the right across a row.
That means that elements whose atomic weights are really close together can be very different, and some elements with far-apart atomic weights are very similar. As you move from lighter atoms to heavier ones, you keep periodically running across the same properties...
The periodic table is full of repeating patterns. Take atomic size, for instance: atoms get bigger as you move down a column, and smaller as you move to the right across a row, or period. There are two patterns to be explained: atoms get bigger as you go down a group, and smaller as you go to the right across a period. The reason for the first one shouldn't be so hard to see now; look again down the column of alkali metals.
Each time you move down, you add another primary level--lithium's highest electron is in a 2s state, for sodium it's 3s, and so on.
Exactly. And the higher an electron's energy, the farther from the nucleus it is.
So the atoms get bigger as you add electrons to higher energy levels--that makes sense. But why do they get smaller as you move to the right?
Well, you'll notice that within a period, the outermost electrons are all in the same primary level--that is, at (roughly) the same distance from the nucleus. But as you move to the right, the elements increase in atomic number; each element has one more proton than its left-hand neighbor. The more protons in the nucleus, the more strongly the valence electrons are pulled in...
...and so the atoms shrink! Also, I can see from the chart that the ionization energies get larger as you go to the right; that must be for the same reason.
Similarly, the ionization energies decrease as you move down a group.

You said that electron configurations are "the key to understanding why each element behaves the way it does." How does that work?
I'll give you an example: look again at that far left group with hydrogen and the alkali metals. Start at the top and go down, what do you notice about the electrons?
Um...well, in each one of these, the very top electron is starting a new colored row; it's all by itself in the s sublevel.
Very good! The chart will tell you that the ionization energy for that element is quite small. It wouldn't take much to send that one solitary electron sailing off into dizzying freedom--and that sort of thing, electrons leaving their home atoms, leads directly to chemical reactions.
So that's why the alkali metals react so violently--it's easy to set them off.

Take a look at the second column from the right in the periodic table--the one that starts with fluorine (F).
These elements have their highest electrons in p orbitals--five at the same energy. p has room for six, doesn't it?
That's right. These elements are called the halogens; they're highly reactive, too, but in a different way than the alkali metals. What sets them off is not losing one of their own, but picking up a stray electron, which fits perfectly into that "empty space" in the p sublevel.
In general, the arrangement of the outermost electrons, called valence electrons, tells you all about an element's chemical behavior.

Erroneous

The process of making an atom into an ion is called ionization--hence the term "ionization energy."-[ how much energy it would take to free that electron from the outermost electron].
¬¬Hmm...the electrons in the lower row have higher numbers listed. I guess that makes sense--the closer they are to the nucleus, the more strongly the electric force would be pulling them in. But that means the electrons in higher energy levels have lower numbers on the chart...
The terminology is a little confusing, I agree. Think of it this way: the more energy of its own an electron has, the less additional energy it needs in order to escape.
Okay, so why is that third electron in a higher level than the first two? Why not just add it to the lowest one?
Because the lowest level is "full"; it can't hold more than two electrons. The rule that's operating here is called the Pauli Exclusion Principle, Pauli guessed that two electrons can't be in the same "quantum state," In this context, it means that two identical electrons can't be in the same energy level in the same atom. Two electrons that are not identical. They differ in a characteristic called spin.

Crossover Energy Levels
What's happening here, roughly, is this: you can think of the series of primary levels as being based on the electrons' attraction to the protons; each successive row moves another "step" away from the nucleus. But once you start getting a lot of electrons around, they begin repelling one another like crazy and messing up this nice pattern.
How does the repulsion of other electrons mess up the energy levels?
Well...the fundamental difference between the various sublevels is that higher sublevels have more angular momentum. If you don't know what that means, don't worry; all you need to understand is that more angular momentum tends to fling an electron farther out from the nucleus. When there are many other electrons around, a screening effect occurs.
So that's why the different sublevels have different energies.
Yes. Furthermore, when the number of electrons becomes large, this screening effect becomes so strong that it actually begins to overlap the next primary level. An added electron will then prefer to enter that next level rather then go to the orbital where it "should" be.

a higher primary level usually, not always, means a higher energy. The energy levels aren't always as well-behaved as one might like; it sometimes happens that the first orbitals in a "higher" primary level actually have less energy than the top orbitals of the level below.
So in potassium, the 4s orbitals end up with less energy than the 3d states--that's why potassium starts a new row in the periodic table.
All those elements from scandium through zinc (Zn) are just filling the ten green d spaces. Then gallium (Ga) goes back up and starts the 4p orbital.

Exceptions to the General Pattern of First Ionization Energies
The figure below shows the first ionization energies for elements in the second row of the periodic table. Although there is a general trend toward an increase in the first ionization energy as we go from left to right across this row, there are two minor inversions in this pattern. The first ionization energy of boron is smaller than beryllium, and the first ionization energy of oxygen is smaller than nitrogen. These observations can be explained by looking at the electron configurations of these elements. The electron removed when a beryllium atom is ionized comes from the 2s orbital, but a 2p electron is removed when boron is ionized.
Be: [He] 2s2
B: [He] 2s2 2p1
The electrons removed when nitrogen and oxygen are ionized also come from 2p orbitals.
N: [He] 2s2 2p3
O: [He] 2s2 2p4
But there is an important difference in the way electrons are distributed in these atoms. Hund's rules predict that the three electrons in the 2p orbitals of a nitrogen atom all have the same spin, but electrons are paired in one of the 2p orbitals on an oxygen atom. Hund's rules can be understood by assuming that electrons try to stay as far apart as possible to minimize the force of repulsion between these particles. The three electrons in the 2p orbitals on nitrogen therefore enter different orbitals with their spins aligned in the same direction. In oxygen, two electrons must occupy one of the 2p orbitals. The force of repulsion between these electrons is minimized to some extent by pairing the electrons. There is still some residual repulsion between these electrons, however, which makes it slightly easier to remove an electron from a neutral oxygen atom than we would expect from the number of protons in the nucleus of the atom.

Colloids and Suspensions

Describe the following terms:
a) True solutions b) Colloids (Tyndall effect) c) Suspensions. a) True solutions Normally light passes through true solutions or they transparent to visible light. Because of small solute particle sizes they do not scatter light.. True solutions have solute particles with diameters less than 1 x103 pm. Colloids, on the other hand have solute particles with diameters in the range 1 x103 to 1 x105 pm. Suspensions have larger particle diameters which are greater than 1 x105 pm.
b) Colloids (Tyndall effect) Colloids have solute particle diameters in the rage 1 x103 to 1 x105 pm. Colloids scatter light and the solution become opaque and relative degree of opacity depends on the sizes and amount of the particles. The scattering of light by colloidal particles is called Tyndell effect which has been used to distinguish between true solutions and colloids. This effect will increase with increasing solute particle diameter of a colloid. For example, milk is opaue because of the higher diameters of the solute particles such as proteins in the milk.
c) Suspensions. Suspensions have particle diameters greater than 1 x105 pm. The particles in a suspension are lage enough to be affected by gravity and settle to the bottom of the mixture with time. For example, muddy water.
Surfactants
Surfactants constitute the most important group of detergent components. Generally, these are water-soluble surface-active agents comprised of a hydrophobic portion, usually a long alkyl chain, attached to hydrophilic or water solubility enhancing functional groups. The hydrophilic end, which is either polar or ionic, dissolves readily in water. The hydrophobic, or non-polar, end, however, does not dissolve in water. In fact, the hydrophobic, or "water-hating" end will move as far away from water as possible. This phenomenon is called the hydrophobic effect. Because one end of a surfactant resists water and the other end embraces it, surfactants have very unique characteristics.
Emulsifiers
A surfactant is also called an emulsifier. Emulsifiers do help oil and water remain in stable emulsions. Aggregates of oil and emulsifiers form micelles and stay dispersed in water solution. Examples of emulsifiers include sodium dodecyl sulfate.
All surfactants possess the common property of lowering surface tension when added to water in small amounts, as illustrated in the plot below. The characteristic discontinuity in the plots of surface tension against surfactant concentration can be experimentally determined. The corresponding surfactant concentration at this discontinuity corresponds to the critical micelle concentration (CMC). At surfactant concentrations below the CMC, the surfactant molecules are loosely integrated into the water structure (monomer see figure below). In the region of the CMC, the surfactant-water structure is changed in such a way that the surfactant molecules begin to build up their own structures (micelles in the interior and monolayers at the surface).
When a small concentration of surfactant is added to water, the hydrophobic end will immediately rise to the surface. The surfactant is then stable. The polar end is happily immersed in the polar solvent while the non-polar end rises above the surface. This configuration is called a monolayer. If more surfactant is added and there is not enough room for all of the hydrophobic ends to stick out of the water, a bilayer will form. However, if the concentration of surfactant is large enough, even the bilayer configuration will not be stable. At this point, a micelle forms, in which a group of hydrophilic ends surround their hydrophobic "tails" and shield them from the water. The micelle, therefore, consists of a hydrophilic shell and a hydrophobic core.
A very important effect of surfactants in cleaning products is the wetting effect. Because of the reduced surface tension, the water can be more evenly distributed over the surface and this improves the cleaning process. The emulsifying effect of surfactants is important for both cleansing and washing of textiles. Due to the hydrophobic and hydrophilic parts, surfactants can sorb to non-polar and polar materials at the same time. During cleansing and washing, the non-polar materials are kept in emulsions in the aqueous solution and removed by rinsing. By varying the hydrophobic and hydrophilic part of a surfactant, a number of properties may be adjusted, e.g. wetting effect, emulsifying effect, dispersive effect, foaming ability and foaming control.
Soap the most common surfactant
One of the organic chemical reactions known to ancient man was the preparation of soaps through a reaction called saponification. Natural soaps are sodium or potassium salts of fatty acids, originally made by boiling lard or other animal fat together with lye or potash (potassium hydroxide). Hydrolysis of the fats and oils occurs, yielding glycerol and crude soap. In the industrial manufacture of soap, tallow (fat from animals such as cattle and sheep) or vegetable fat is heated with sodium hydroxide. Once the saponification reaction is complete, sodium chloride is added to precipitate the soap. The water layer is drawn off the top of the mixture and the glycerol is recovered using vacuum distillation.
The crude soap obtained from the saponification reaction contains sodium chloride, sodium hydroxide, and glycerol. These impurities are removed by boiling the crude soap curds in water and re-precipitating the soap with salt. After the purification process is repeated several times, the soap may be used as an inexpensive industrial cleanser. Sand or pumice may be added to produce a scouring soap. Other treatments may result in laundry, cosmetic, liquid, and other soaps. Soap is an anionic surfactant.
Detergent
The term "detergent" is used for a surfactant material which cleans (or is used for cleaning), but in this definition soap is not included. Even so, this is still a wide definition, because, of course, it can refer to the active ingredient, or the solid, liquid, paste or powder compounded from this active matter. Surfactants are grouped according to their ionic properties in water:
• Anionic surfactants have a negative charge
• Nonionic surfactants have no charge
• Cationic surfactants have a positive charge
• Amphoteric surfactants have positive or negative charge dependent on pH A list of surfactants which are commonly used in biochemistry is presented in the table below.
Structures of common surfactants: Chemical or Trade Name Sodium dodecylsulfate (SDS) Sodium cholate Sodium deoxycholate (DOC) N-Lauroylsarcosine Sodium salt Lauryldimethylamine-oxide (LDAO

Cetyltrimethylammoniumbromide (CTAB)

Bis(2-ethylhexyl)sulfosuccinate Sodium salt

CHEMISTRY LAB Trends in the Periodic Table Names: _____________________ _____________________ _____________________ _____________________ Purpose Determine the trends, if they exist, for atomic size and ionization energy in the Periodic Table. Materials Graph paper Procedure
1. Use the information from the section of the periodic table. Be sure to give each graph a title and to label each axis.
2. For elements 3-20, make a graph of atomic radius as a function of atomic number. Plot atomic number on the X axis and atomic radius on the Y-axis.
3. For elements in Family 1A (1) and Family 2A (2), graph period number vs. atomic radius. Use a different color or symbol for each line.
4. For elements 3-20, make a graph of ionization energy as a function of atomic number. Plot atomic number on the X-axis and ionization energy on the Y-axis.
5. For elements in Family 1(1A) and Family 2 (2A), graph period number vs. ionization energy. Use a different color or symbol for each line. IA (1) IIA (2) IIIA (13) IVA (14) VA (15) VIA (16) VIIA (17) VIIIA (18) 2 3
Li
1.23
124 4
Be
0.89
215 5
B
0.80
191 6
C
0.77
260 7
N
0.70
335 8
O
0.66
314 9
F
0.64
402 10
Ne
0.67
497

3 11
Na
1.57
119 12
Mg
1.36
176 13
Al
1.25
138 14
Si
1.17
188 15
P
1.10
242 16
S
1.04
239 17
Cl
0.99
299 18
Ar
0.98
363

4 19
K
2.03
100 20
Ca
1.74
141

5 37
Rb
2.16
96 38
Sr
1.91
131 8
O
0.66
314 Atomic
Symbol
Atomic
Ionization number radius energy 6 55
Cs
2.35
90 56
Ba
1.98
120

Analysis
1. What happens to the atomic radius as the atomic number increases across a period? Down a family?

2. What happens to the ionization energy as the atomic number increases across a period? Down a family?

Conclusion
1. Why does atomic radius change as it does?

2. Why does the ionization energy change as it does?

*Important fact you may want to write down* The Modern Periodic Table is based on atomic number and electron configuration, not atomic mass.
When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.
• Atomic Radius: No definite shape, but scientist can get a rough measure.
• 12. Atomic Size It doesn’t depend on Weight Watchers
• 13. Ionization Energy The energy required to remove an electron from an atom in the gas phase. A (g) + Energy  A + (g) + e - Ionization energy measures how tightly an electron is held in the atom.
• 14. Ionization Energy (cont.) More and more energy is required to move each electron from an atom Metals generally have low IE. Nonmetals have high IE. IE increases as you move across a period and decreases as you go down a group or family.
• 15. Electron Affinity EA – It’s NOT a clothing company The energy used or released for a gaseous atom to gain an electron. A (g) + e -  A - (g) + Energy
• 16. Electron Affinity (cont.) In general. . . EA increases (becomes more negative) as you go across a period and decreases as you go down a group or family. The greater the electron affinity, the greater the IE. Metals have lower EAs Nonmetals have higher EAs
• 17. Electronegativity A comparative scale relating the abilities of elements to attract electrons when their atoms are combined. Active metals (IA) have the lowest Ens Most nonmetals have the highest Ens