Science is based on empiricism, that is, it is based on what is measureable Often direct measurement is not possible. Mathematical computation must take the place of direct measurement. In this lab, you will measure objects using the geometric principle of similar triangles.

The Principle:

Two triangles are similar if the only difference in them is size. You may have to flip them or turn them, however to see that they are the same. Here are some examples of similar triangles:

For similar triangles, corresponding sides always have the same ratio. You can use this fact to figure out distances and heights. Here are some examples:

Since the triangles have the same shape, if you know how far the man is from the mirror (say 6 feet) and you know how far the the mirror is from the stop light (say 24 feet) then you know the light is 24/6 or 4 times higher than the height of the man. It is a simple ratio.

Here is another example:

In this case the tree is 84/12, or 7 times closer to the end of the shadow than the woman is, so the tree is 7 times higher than she is.

Materials:

· Meter stick

· Ruler or other substitute

· Pencil

· String

· Text Box: PROCEDURE:
Build a device like the one shown in the picture, and measure how much higher the top of the stick is from the surface of the meter stick. Record that height difference.
To use your new invention,
Place it flat.
Draw the string tight.
Move the string up and down the meter stick until you see the top of the object through the straw. Tape

· Straw

Suggestions:

1. Take extra care that the pencil keeps the sticks at a right angle to each other.

2. You want to be really careful that you take measurements from level surfaces.

3. It would be possible to measure widths of distant objects by placing your device flat against a wall.

4. Improve the device. (Each unique and beneficial modification is worth 1/3 of a letter grade extra credit.)

Report:

NOTE: Each object you measure must not have been measured by other students.

Assignment	Description	Distance to object	Meter stick reading	Estimated size
Something in the room
Something on the second floor
Something far away
Something very small
Your Choice
Your Choice
Your Choice

Questions:

1. What is the accuracy of your device?

2. How could you improve this device?

3. Using this principle, how could you measure the distance the moon is from the earth?

Chemistry Lab

2 Glassware Fabrication

Materials:

· Glass rods

· Bunsen burner

· Sparker

· Eye protection

· Protective paper towels

Directions:Read any handouts about glassware fabrication and watch the videos, then manipulate the glass tubing to produce the objects listed in the table below and have them initialed when you finish for credit.

Sign off	Task
	Cut glass tubing using a triangular file
	Fire polish
	90® Bend
	Insert glass tube in stopper
	Remove glass tubing from stopper
	Glass U shape
	Draw a pipette
	Glass T

Chemistry Lab

3 Molecular Models 1

Using your phone and your text book, look up the chemical model or “configuration” for each of the following chemicals and build it using one of the modeling kits. Build each molecule with a different partner.

	Formula	Name	Chem Name	Pts	Partner Name	Signature
1	H₂O	Water	Dihydrogen Oxide	1
2	CH₄	Methane gas	Carbon Tetrahydride	1
3	CH3COOH	Vinegar	Acetic Acid	6
4	O₂	Oxygen gas	Dioxygen	1
5	NaCl	Salt	Sodium Chloride	1
6	CH₃CHOHCH₃	Rubbing Alcohol	Isopropyl Alcohol	6
7	NaClO	Bleach	Sodium Hypochlorite	3

4 The Borax Bead Test to Discriminate Metal Ions

From http://www.brainyresort.com/en/borax-bead-test/

The borax bead test is based on the creation of a glassy bead that can variously change color when put under a Bunsen flame, potentially revealing a whole series of compounds, depending on their chemical (electronic) features.

First of all a loop, eyelet, is made in the end of a Nichrome wire. Then it's loaded with borax or phosphorus salt and heating in the flame you get a clear, colorless glassy sphear, the so called "bead". The bead will be our reagent to identify specific cationic components (borax bead test) or cationic and anionic components (phosphorus salt bead test).

In the present article we're going to see how to carry out identification tests with this interesting method.

Bead Preparation

When we talk about "borax" we're referring to the compound with formula Na₂B₄O₇ · 10H₂O, IUPAC name sodium tetraborate decahydrate but also known as sodium tetraborate.As support to create the bead is commonly used a platinum wire that is firmly attached to the end of a glass stick.Before creating the pearl, it would be good practice to clean the wire moistening it with6N HCl, to solubilize impurities such as any residual crust of previous uses. The hydrochloric acid is necessary to transform the various compounds eventually present in chlorides, usually volatile and thus eliminated in the flame.Residue of other substances could indeed alter the outcome of the test.Alternatively, you can clean the wire with borax itself, loading the loop with borax and heating in flame.The impurities will be adsorbed within the melted bead, which is then eliminated.

Now, we're going to consider from a chemical (and physical) viewpoint what happens when the bead is heated in flame.

The first thing we notice is the so-called "popcorn" effect, namely the swelling of the substance caused by the loss of crystallization water:

Na₂B₄O₇ · 10H₂ONa₂B₄O₇ + 10H₂O ↑

For further warming, tetraborate decomposes to give sodium metaborate and boric oxide.

Na₂B₄O₇2NaBO₂ + B₂O₃

This is our pearl before the reaction with the unknown substance.It should appear clear and colorless.If it wasn't we must continue heating up to fusion and reloading with more borax.Alternatively, as mentioned earlier, we unload the bead onto the work surface, and we create a new one.

Recognition of the cationic component: bead test

We bring in our flame (Bunsen-Burner) the glass-like bead. We heat it up to incandescence and then we remove it from the flame. Then we let it cool slightly and then, with the bead (that's on the top of the wire, retained by the loop), we touch a small amount of unknown substance.

The real reagent of this type of assay is boric oxide,B₂O₃ .This reacts with the oxides of certain metals to give the corresponding metaborates.These impart to the bead particular color features that may be indicative about the cationic composition of our unknown substance.

In addition, the test may have different outcomes depending on which zone of the flame we select.In fact, in oxidizing flame we assist both to the formation of oxidized compounds (and then eventually with different colors) and tothermochromisms a physical phenomenon consisting in color change given by heating, whereas in the reducing flame we assist to the formation of reduced compounds (and then again eventually with different colors).

Let's see what are the reactions involved step by step:

The salts or compounds present in the unknown substance for prolonged heating or action of the oxidizing flame itself, give the corresponding oxides (some example):

CuCO₃ CuO + CO₂↑

CuNO₃ CuO + NO₂ + 1/2 O₂↑

Cu(OH)₂ CuO + H₂O↑

etc-etc.

These oxides react then with boric oxide to give corresponding metaborates.

Oxidizing flame: CuO + B₂O₃→Cu(BO₂)₂ cupric metaborate green (hot), blue-green (cold)

Reducing flame (the reducing agent is elemental carbon powder): 4NaBO₂ + 2Cu(BO₂)₂ + CCO₂ + 2Na₂B₄O₇ + 2Cu metallic copper → red-brown

According to the different color we can recognize different cations, and that's is the aim of this old but interesting assay.

Expedients

When you are loading the borax bead with the unknown substance it is good to do so it is adsorbed by the bead just a minimal amount of unknown substance.An excess of the reagent may not react and therefore remain as an unreacted dark mass that does not allow to observe well the coloration assumed by the small bead.

Conclusions

This test is a dry assay.It allows to recognize the following substances: Cr, Mn, Ni, Co, Cu, Fe.The real test reagent is boric oxide (B₂O₃) which reacts with the oxides of these metals to give colored metaborates.It may be useful to bring the pearl both in oxidizing flame, to observe highest oxidation states coloration of metals and eventual thermochromisms, and in reducing flame, where the metaborates of the metals react with carbon powder (C) and are reduced to lower oxidation states if not, in some cases, the same metal in the elemental state (as copper in the example above).

Scheme of possible outcomes (1 or 2)

Text Box:
Bunsen burner: leftmost: reducing flame, rightmost: oxidizing flame

Oxygen rich butane torch flame

Fuel rich butane torch flame Oxidizing and reducing flames

From Wikipedia, the free encyclopedia

In various burners, the oxidizing flameis the flameproduced with an excessive amount of oxygen. When the amount of oxygen increases, the flame shortens, its color darkens, and it hisses and roars.^[1] With some exceptions (e.g., platinum soldering in jewelry), the oxidizing flame is usually undesirable forweldingand soldering, since, as its name suggests, it oxidizesthe metal's surface.^[1]The same principle is important in firing pottery — seeReducing atmosphere.

The reducing flameis the flame with low oxygen. It has a yellow or yellowish color due to carbonor hydrocarbons ^[2]which bind with (or reduce) the oxygen contained in the materials processed with the flame.^[1] The reducing flame is also called the carburizing flame, since it tends to introduce carbon into the molten metal.

The neutral flameis the flame in which the amount of oxygen is precisely enough for burning, and neither oxidation nor reduction occurs.^[1]A flame with a good balance of oxygen is clear blue.

The reducing and neutral flames are useful in soldering and annealing.^[1]

Table of Reactions Obtained with Borax

From http://webmineral.com/help/BoraxBead.shtml#.Waf5bHQpCvM

This table of flux fusion reactions with borax is based largely on the book "Determinative Mineralogy and Blowpipe Analysis" by Brush & Penfield, 1906. The reactions are observed by fusing a crushed (and roasted in the case of sulfides) sample of the mineral in a borax lead embedded on a loop of platinum wire. The color is observed in the bead after heating the sample with the oxidizing flame and then the reducing flame of the torch (blowpipe).

Oxidizing Flame		Amount of Material	Chemical Element(s)	Reducing Flame
HOT	COLD	Amount of Material	Chemical Element(s)	HOT	COLD
Colorless	Colorless	Little or Much	Si, Al, and Sn	Colorless	Colorless
Colorless	Colorless or opaque White depending on the degree of saturation	Little or Much	Ca, Sr, Ba, Mg, Be, Zn, Y, La, Th, Zr, Ta, and Nb	Colorless	Colorless or opaque White depending on the degree of saturation
Pale Yellow	Colorless or White	Much	Pb, Sb, and Cd	Pale Yellow	Colorless
Pale Yellow	Colorless or White	Much	Bi	Gray	Gray
Pale Yellow	Colorless or White	Much	Mo	Brown	Brown
Pale Yellow	Colorless or White	Medium	W	Yellow	Yellow to Yellowish Brown
Pale Yellow	Colorless or White	Medium	Ti	Grayish	Brownish Violet
Pale Yellow	Nearly Colorless	Little	Fe and U	Pale Green	Nearly Colorless
Yellow	Pale Yellow	Little	Ce	Colorless	Colorless
Yellow	Yellowish Green	Little	Cr	Green	Green
Yellow	Yellowish Green, almost Colorless	Little	V	Dirty Green	Fine Green
Deep Yellow to Orange-red	Yellow	Medium to Much	Ce	Colorless	Colorless
Deep Yellow to Orange-red	Yellow	Medium to Much	U	Pale Green	Pale Green to nearly Colorless
Deep Yellow to Orange-red	Yellow	Medium to Much	Fe	Bottle Green	Pale Bottle Green
Deep Yellow to Orange-red	Yellowish Green	Medium to Much	Cr	Green	Green
Green	Blue	Little to Medium	Cu	Colorless to Green	Opaque Red with much Oxide
Green	Yellow, Green, or Blue of various Shades	Medium	Various Mixtures of Fe, Cu, Ni, and Co	(?) See Below	(?) See Below
Blue	Blue	Little to Medium	Co	Blue	Blue
Violet	Reddish Brown	Little to Medium	Ni	Opaque Gray	Opaque Gray
Violet	Reddish Violet	Little	Mn	Colorless	Colorless
Pale Rose	Pale Rose	Much	Nd	Pale Rose	Pale Rose

Assignment:

Use the Borax Bead Test to determine the metal ions in the following unknowns:

	Oxidizing Flame Color		Reducing Flame Color		Metal Ion in Sample
	Hot	Cold	Hot	Cold	Metal Ion in Sample
1
2
3
4
5

Chemistry Lab

5 Spectroscopy - Flame Test Lab

Background:

The normal electron configuration of atoms or ions of an element is known as the “ground state.” In this most stable energy state, all electrons are in the lowest energy levels available. When atoms or ions in the “ground state” are heated to high temperatures, some electrons may absorb enough energy to allow them to “jump” to higher energy levels. The element is then said to be in the “excited state.” This excited configuration is unstable, and the electrons “fall” back to their normal positions of lower energy (ground state). As the electrons return to their normal levels, the energy that was absorbed is emitted in the form of electromagnetic energy. Some of this energy may be in the form of visible light. The color of this light can be used as a means of identifying the elements involved. Such analysis is known as a flame test.

To do a flame test on a metallic element, the metal is first dissolved in a solution and the solution is then held in the hot, blue flame of a Bunsen burner. This test works well for metal ions, and was perfected by Robert Bunsen (1811 – 1899). Many metallic ions exhibit characteristic colors when vaporized in the burner flame.

Quick Test:

	Name for the most stable state for electrons
	Name one thing that can energize electrons
	What is an element in a high energy level called?
	How do chemicals make light?
	What is light?

Purpose:

The purpose is to observe the characteristic colors produced by certain metallic ions when vaporized in a flame and then to identify an unknown metallic ion by means of its flame test.

Materials:

Set of metal chloride salts (NaCl, CuCl₂, KCl, CaCl₂, SrCl₂, LiCl, CoCl₂, BaCl₂)

Bunsen Burner

Nichrome wire

Unknown solution (for each student)

Safety: Be sure to wear goggles and an apron at all times

Procedure:

Light the Bunsen burner and adjust it so that it has a hot blue flame.
Using a clean nichrome wire loop, wet the loop and dip it into one of the salts then hold it in the hottest part of the burner flame. Observe the color of the flame. Carefully record your observations in the data table. Be accurate here - your description of the color must be accurate enough to distinguish this metal ion from the other ions tested.
Clean the nichrome loop for each of the other salts, and check the color of their flame tests. Record your observations for each.

4. When you have tested all the known solutions and can distinguish the color of each metal ion, obtain unknown solutions and determine which metal ions are present by performing a flame test and comparing this data to your previous data.

Data table:

Metal ion	Color of Flame
barium
calcium
cobalt
copper
lithium
potassium
sodium
strontium

Unknown # ____
Unknown # ____

Based on your observations, identify the two unknowns you examined:

Unknown # _____ is ________________________________

Questions:

1. State at least three problems that may be involved when using flame tests for identification purposes.

2. Which ions produce similar colors in the flame tests?

3. What purpose did the cobalt glass serve?

4. Explain how the colors observed in the flame tests are produced.

5. How could this test be made more accurate?

6. What would happen if a chemical makes more than one color because its electrons jump to two or more levels?

Chemistry Lab

6 Hydrates

LABORATORY 6.6: DETERMINE THE FORMULA OF A HYDRATE[1]

Many ionic compounds exist in two or more forms. The anhydrous form of the compound contains only molecules of the compound itself. The hydrated form or forms of the compound contains molecules of the compound and one or more molecules of water loosely bound to each molecule of the compound. These water molecules are referred to as water of hydration or water of crystallization, and are incorporated into the crystalline lattice as the compound crystallizes from an aqueous solution.

Because these water molecules assume defined positions within the crystalline lattice, the proportion of water molecules to compound molecules is fixed and specific. For example, copper sulfate exists as an anhydrous compound (CuSO₄) and in hydrated form as the pentahydrate (CuSO₄ 5H₂0). Copper sulfate does not exist in the tetrahydrate (CuSO₄ 4H₂0) or hexahydrate (CuSO₄ • 6H20) forms, because the physical geometry of the crystalline lattice does not permit four or six water molecules to associate with one copper sulfate molecule. The number of molecules of water in a hydrate is usually an integer, but not always. For example, some hydrates exist in the form X2. 5H₂0, where each molecule of the compound X is associated with a fractional number (in this, case 2.5) molecules of water.

Some compounds, including copper sulfate, have only one stable hydrated form. (Monohydrate and trihydrate forms of copper sulfate are known. but are difficult to prepare and tend to spontaneously convert to the more stable anhydrous or pentahydrate forms by absorbing or giving up water molecules.) Other compounds have two or more common hydrated forms. For example, sodium carbonate exists in anhydrous form (Na₂CO₃), monohydrate form (Na₂CO₃ • 1H₂0), heptahydrate form (Na₂CO₃ • 7H₂0), and decahydrate form (Na₂CO₃ • 10H₂0). Many anhydrous compounds are hygroscopic, which means they absorb water vapor from the air and are gradually converted to a hydrated form. Such compounds, such as calcium chloride (CaCl₂), are often used as drying agents. (Some of these compounds absorb so much water vapor from the air that they actually dissolve In the absorbed water, a property called

REQUIRED EQUIPMENT AND SUPPLIES

1. goggles, gloves, and protective clothing

2. balance and weighing papers

3. crucible with cover and tongs

4. gas burner

5. ring stand, ring, and clay triangle

6. copper sulfate pentahydrate (-5 g)

deliquescence ,) Conversely. the water molecules in some grated compounds are so loosely bound that the compound spontaneously loses some or all of its water of hydration if left ina dry environment, a property called efflorescence. some compounds may be either hygroscopic or efflorescent. depending on the temperature and humidity of the environment. For example. anhydrous copper sulfate exposed to a humid atmosphere gradually absorbs water vapor and is converted to the pentahydrate form, and copper sulfate pentahydrate exposed to warm. dry air gradually loses water and is converted to the anhydrous form. Because the water of crystallization in a hydrate can be driven off by heating the hydrate, a hydrate is actually a mixture of an anhydrous salt with water rather than a separate compound. (Recall that a mixture is a substance that can be separated into its component parts by physical means. such as heating, as opposed to a substance that can be separated into its component parts only by using chemical means.) Because the number of water molecules in a hydrate are in fixed proportion to the number of molecules of the compound, it's possible to determine that fixed proportion by weighing a sample of a hydrate, heating the compound to drive off the water of crystallization, weighing the resulting anhydrous compound, and using the mass difference between the hydrated and anhydrous forms to calculate the relationship.

In this laboratory, we'll heat hydrated copper sulfate to drive off the water of crystallization and use the mass differential to determine how many molecules of water are associated with each molecule of hydrated copper sulfate. We could have used any number of common hydrates, but we chose copper sulfate because the hydrated and anhydrous forms have distinctly different appearances. In hydrated form, copper sulfate forms brilliant blue crystals; in anhydrous form, copper sulfate is a white powder (see Figure 6-7).

CAUTIONS This laboratory uses strong heat. Use extreme care with the heat source and hot objects. A hot crucible looks exactly like a cold crucible. Wear splash goggles, gloves, and protective clothing.

PROCEDURE.

1.If you have not already done so, put on your splash goggles. gloves. and protective clothing.

2.Set up your ring stand. support ring, clay triangle. and burner. Heat the crucible and cover gently for a minute or two to vaporize any moisture.

3.Remove the heat and allow the crucible and cover to cool to room temperature, which may require 10 or 15 minutes.

4.Weigh the crucible and cover. and record their mass to 0.01 g on line A of Table 6-7.

5.Transfer about 5.0 g of copper sulfate pentahydrate to the crucible. (The copper sulfate pentahydrate should be in the form of fine crystals. If it is in the form of large lumps, use your mortar and pestle to crush It into finer crystals.)

6.Reweigh the crucible, cover, and contents and record the mass to 0.01 g on line B of Table 6-7.

7. Subtract the mass of the empty crucible and lid from the mass of the crucible with the copper sulfate and record the initial mass of the copper sulfate pentahydrate on line C of Table 6-7.

8. Place the crucible and cover on the heat source and begin heating them gently. As the crucible warms up, increase the heat gradually until it is at its highest setting. Continue heating the crucible for at least 15 minutes.

9. Remove the heat and allow the crucible. cover, and contents to cool to room temperature, which may require 10 or 15 minutes.

10.After you are sure the crucible has cooled. reweigh crucible, lid, and contents. Record the mass to 0.01 g on line D of Table 6-7.

11. Subtract the initial mass of the crucible and lid (line A) from this value, and record the mass of the anhydrous copper sulfate to 0.01 g on line E of Table 6-7.

Text Box: Draw a picture of what anhydrous CuSO4 would look like. 12. Subtract the mass of the anhydrous copper sulfate (line E) from the mass of the copper sulfate pentahydrate (line C). and record the mass loss to 0.01 g on line F of Table 6-7.

	Table 6-7	Result
A	Mass Crucible
B	Mass: Crucible + sample
C	Mass of sample CuSO₄
D	Mass after heating
E	Mass of anhydrous salt
F	Mass loss
G	Molar mass CuSO₄
H	No. Moles of CuSO₄
I	Molar mass H₂O
J	No. Moles of H₂O
K	Ratio CuSO₄ to H₂O

7 Molar Solution of a Solid Chemical

In this lab, we make up 100 mL of a stock solution of copper (II) sulfate, which is used in many of the other lab. Although we won't standardize[2] the solution, we will make every effort to achieve an accurate concentration by weighing masses carefully and measuring volumes carefully.

1. Calculate the molar weight of copper sulfate pentahydrate.

1		= Molecular weight of Cu, copper		6	= Molecular weight of H, hydrogen
2		= Molecular weight of S, sulfur		7	= Molecular weight of O, oxygen
3		= Molecular weight of O, oxygen		8	= Molecular weight of H₂
4		= Molecular weight O₄ – four oxygen		9	= Molecular weight of H₂O
5		= Molecular weight of CuSO₄		10	= Molecular weight of 5 H₂O

11			= Molecular weight of CuSO4 × 5H₂O
12	g/mol		= Molar weight of CuSO4 × 5H₂O

2. Looking up copper sulfate in a reference book, we find: its solubility at 20°C is 317 g/L.

3. Dividing 317 g/L by your answer in box 12will tell you how many moles are in a saturated solution. This will protect us against the possibility that the solution will crystalize if the temperature drops.

= Moles in a saturated solution of CuSO₄

4. Calculate the number of grams you need to weigh out to make 100 ml of 0.1 molar CuSO₄. (Hint, because 0.1 molar solutions are one tenth as strong, you need much less CuSO₄. Again, because molarity is are measured as grams per liter, you are going to need much less again.)

= Grams of CuSO₄× 5H₂Oneeded to make up 100 ml of 0.1 molar CuSO₄.

5. Measure out 100 ml of H₂O.

6. Tare a scale.

15		= Mass of container and H₂O
16		= Mass of container + H₂O + CuSO₄× 5H₂O (from box 14)

7. Add the water to a container and weigh them, then carefully add the copper sulfate.

8.Transfer the 0.1M copper sulfate solution to a bottle and label it for use in later labs.

Chemistry Lab

Text Box: Notes, Evidence of ionic reaction:
1. Color change
2. Effervescence
3. Precipitate
4. Temperature
5. Smell
6. Light emission 8 Ionic Interaction

Using the Ionic Interaction Kittest for ionic reactions and post the results in the table below. Do this lab as a whole class.

1. Place 3 drops of the cation acetate in the top two rows.

2. Add 3 drops of each anion starting with Pb in sequence so all cations are tested against acetate.

3. Record the results below for each anion and cation.

4. Go to step 1 and repeat it for bromide and all the rest of the anions.

Cations

Pb²⁺

Ag⁺

Hg²⁺

Cu²⁺

Cd²⁺

Ni²⁺

Co²⁺

Mn²⁺

Zn²⁺

Al³⁺

Cr³⁺

Fe³⁺

Ba²⁺

Sr²⁺

Ca²⁺

Aions

Acetate

Bromide

Carbonate

Chloride

Chromate

Hexacyano

Ferrate(II)

Hexacyano

Ferrate(III)

Hydroxide

Iodide

Oxalate

Phosphate

Silicate

Sulfate

Ionic Interaction Lab

File:Acetic-acid-3D-balls-B.png Acetate C₂H₃0₂^- COOH^-

Bromide Br^-

Carbonate -

Hexacyanoferrate (II) C₆FeN₆^4-Ferricyanide

HexacyanidoferratIII 2.svg

Hexaxanoferrate (III) C₆FeN₆^3-Ferrocyanide

Hydroxide OH^-

Image result Iodide I^-

Oxilate C2O4, (COO)2

Image result for Phosphate ion

Phosphate PO₄^3-

Image result

Silicate SiO₃^4-

Image result for so4 2- ion

Sulfate SO₄^2-

Pb²⁺

Ag⁺

Hg²⁺

Cu²⁺

Cd²⁺

Ni²⁺

Co²⁺

Mn²⁺

Zn²⁺

Al³⁺

Cr³⁺

Fe³⁺

Ba²⁺

Sr²⁺

Ca²⁺

Text Box: Notes:
V:Sum up the Valence electrons
S:Make skeleton connecting all atoms
E:Add electrons starting with most electronegative and outermost
P:Map electron pairs
R:Review charges Chemistry Lab

9 Molecular Modeling-2

Using the Deluxe Molecular Modeling Kit and VSEPR(“vesper”) build 3-D models of each of the chemicals listed in the table.

1. Fill out the table.

2. Build a model.

3. Have the instructor check it off.

	Formula	Name	Elect rons	Skeleton	Lewis Dot Drawing	Configuration[3]
1	NH3	Ammonia
2	CO2	Carbon Dioxide
3	H2O	Water
4	CH4	Methane
5	PCl5	Phosphorus Pentachloride
6	C2H6O	Ethanol
7	C8H18	Octane
8	Yours

A Brief Tutorial on Drawing Lewis Dot Structures[4]

We will use three molecules (CO₂, CO₃^2- and NH₄⁺) as our examples on this guided tour of a simple method for drawing Lewis dot structures. While this algorithm may not work in all cases, it should be adequate the vast majority of the time.

Procedure for Neutral Molecules (CO₂)

1. Decide how many valence (outer shell)electrons are possessed by each atom in the molecule.

2. If there is more than one atom type in the molecule, put the most metallic or least electronegative atom in the center. Recall that electronegativity decreases as atom moves further away from fluorine on the periodic chart.

Arrangement of atoms in CO₂:

3. Arrange the electrons so that each atom contributes one electron to a single bondbetween each atom.

4. Count the electrons around each atom: are the octets complete? If so, your Lewis dot structure is complete.

5. If the octets are incomplete, and more electrons remain to be shared, move one electron per bond per atom to make another bond. Note that in some structures there will be open octets (example: the B of BF₃), or atoms which have ten electrons (example: P in PF₅) or twelve electrons (example: S in SF₆).

6. Repeat steps 4 and 5 as needed until all octets are full.

7. Redrawthe dots so that electrons on any given atom are in pairs wherever possible.

Procedure for Negatively Charged Ions (CO₃^2-)

Use the same procedure as outlined above, then as a last step add one electron per negative charge to fill octets. Carbonate ion has a 2- charge, so we have two electrons available to fill octets.

Using the procedure above, we arrive at this structure:

The two singly-bonded oxygen atoms each have an open octet, so we add one electron to each so as to fill these octets. The added electrons are shown with arrows. Don't forget to assign formal charges as well! The final Lewis structure for carbonate ion is:

Procedure for Positively Charged Ions (NH₄⁺)

Use the same procedure as outlined above, then remove one electron per positive charge as needed to avoid expanded octets. When using this procedure for positively charged ions, it may be necessary to have some atoms with expanded octets (nitrogen in this example). For each unit of positive charge on the ion remove on electron from these expanded octets. If done correctly, your final structure should have no first or second period elements with expanded octets.

Using the basic procedure outlined above, we arrive at a structure in which nitrogen has nine valence electrons. (Electrons supplied by hydrogen are red; electrons supplied by nitrogen are black.) Removal of one of these valence electrons to account for the 1+ charge of ammonium ion solves this octet rule violation.

Molecular Configurations (Shapes)[5]

The steric number is the number of atoms bonded to the central atom plus the number of non-bonding pairs, and thus the steric number for water is 4. With this information available, together with the table below, you can predict the 3-dimensional shape of the molecule.

Chemistry Lab

10 Ionic Salts 1: Separation of Dissolved Liquids

Why Do Some Solids Dissolve in Water?

The sugar we use to sweeten coffee or tea is a molecular solid, in which the individual molecules are held together by relatively weak intermolecular forces. When sugar dissolves in water, the weak bonds between the individual sucrose molecules are broken, and theC₁₂H₂₂O₁₁ molecules are released into solution.

diagram It takes energy to break the bonds between the C₁₂H₂₂O₁₁molecules in sucrose. It also takes energy to break the hydrogen bonds in water that must be disrupted to insert one of these sucrose molecules into solution.Sugar dissolves in water because energy is given off when the slightly polar sucrose molecules form intermolecular bonds with the polar water molecules. The weak bonds that form between the solute and the solvent compensate for the energy needed to disrupt the structure of both the pure solute and the solvent. In the case of sugar and water, this process works so well that up to 1800 grams of sucrose can dissolve in a liter of water.

diagram Ionic solids (or salts) contain positive and negative ions, which are held together by the strong force of attraction between particles with opposite charges. When one of these solids dissolves in water, the ions that form the solid are released into solution, where they become associated with the polar solvent molecules.

	H₂O
NaCl(s)		Na⁺(aq)	+	Cl^-(aq)

We can generally assume that salts dissociate into their ions when they dissolve in water. Ionic compounds dissolve in water if the energy given off when the ions interact with water molecules compensates for the energy needed to break the ionic bonds in the solid and the energy required to separate the water molecules so that the ions can be inserted into solution.

Solubility Equilibria

Discussions of solubility equilibria are based on the following assumption: When solids dissolve in water, they dissociate to give the elementary particles from which they are formed. Thus, molecular solids dissociate to give individual molecules

	H₂O
C₁₂H₂₂O₁₁(s)		C₁₂H₂₂O₁₁(aq)

diagram and ionic solids dissociate to give solutions of the positive and negative ions they contain.

	H₂O
NaCl(s)		Na⁺(aq)	+	Cl^-(aq)

When the salt is first added, it dissolves and dissociates rapidly. The conductivity of the solution therefore increases rapidly at first.

	dissolve
NaCl(s)		Na⁺(aq)	+	Cl^-(aq)
	dissociate

The concentrations of these ions soon become large enough that the reverse reaction starts to compete with the forward reaction, which leads to a decrease in the rate at which Na⁺ and Cl^-ions enter the solution.

			associate
Na⁺(aq)	+	Cl^-(aq)		NaCl(s)
			precipitate

Eventually, the Na⁺ and Cl^- ion concentrations become large enough that the rate at which precipitation occurs exactly balances the rate at which NaCl dissolves. Once that happens, there is no change in the concentration of these ions with time and the reaction is at equilibrium. When this system reaches equilibrium, it is called a saturated solution, because it contains the maximum concentration of ions that can exist in equilibrium with the solid salt. The amount of salt that must be added to a given volume of solvent to form a saturated solution is called the solubility of the salt.

Solubility Rules

There are a number of patterns in the data obtained from measuring the solubility of different salts. These patterns form the basis for the rules outlined in the table below, which can guide predictions of whether a given salt will dissolve in water. These rules are based on the following definitions of the terms soluble, insoluble, and slightly soluble.

A salt is soluble if it dissolves in water to give a solution with a concentration of at least 0.1 moles per liter at room temperature.
A salt is insoluble if the concentration of an aqueous solution is less than 0.001 M at room temperature.
Slightly soluble salts give solutions that fall between these extremes.

Solubility Rules for Ionic Compounds in Water

Soluble Salts

1. The Na⁺, K⁺, and NH₄⁺ ions form soluble salts. Thus, NaCl, KNO₃, (NH₄)₂SO₄, Na₂S, and (NH₄)₂CO₃ are soluble.

2. The nitrate (NO3-) ion forms soluble salts. Thus, Cu(NO₃)₂ and Fe(NO₃)₃ are soluble.

3. The chloride (Cl-), bromide (Br-), and iodide (I-) ions generally form soluble salts. Exceptions to this rule include salts of the Pb²⁺, Hg₂²⁺, Ag⁺, and Cu⁺ ions. ZnCl₂ is soluble, but CuBr is not.

4. The sulfate (SO42-)ion generally forms soluble salts. Exceptions include BaSO₄, SrSO₄, and PbSO₄, which are insoluble, and Ag₂SO₄, CaSO₄, and Hg₂SO₄, which are slightly soluble.

Insoluble Salts

1. Sulfides (S2-) are usually insoluble. Exceptions include Na₂S, K₂S, (NH₄)₂S, MgS, CaS, SrS, and BaS.

2. Oxides (O2-) are usually insoluble. Exceptions include Na₂O, K₂O, SrO, and BaO, which are soluble, and CaO, which is slightly soluble.

3. Hydroxides (OH-) are usually insoluble. Exceptions include NaOH, KOH, Sr(OH)₂, and Ba(OH)₂, which are soluble, and Ca(OH)₂, which is slightly soluble.

4. Chromates (CrO42-) are usually insoluble. Exceptions include Na₂CrO₄, K₂CrO₄, (NH₄)₂CrO₄, and MgCrO₄.

5. Phosphates (PO43-) and carbonates (CO32-) are usually insoluble. Exceptions include salts of the Na⁺, K⁺, and NH₄⁺ ions.

Solubility: How solubility is measured

The amount of solute that can be dissolved in a solvent is used as the measure of solubility. The conventional reference for solubility is the number of grams of solute that can dissolve in 100 mL of solvent. Sometimes the solubility is in grams of solute per 100 grams of solvent. The table below gives typical solubility data for some common inorganic compounds.

Solubility of Common inorganic compounds in grams solute per 100 mL of water

Substance	0^oC	10^oC	20^oC	30^oC	40^oC	50^oC
KI, potassium iodide	127.5	136	144	152	160	168
KCl, potassium chloride	27.6	31.0	34.0	37.0	40.0	42.6
NaCl, sodium chloride	35.7	35.8	36.0	36.3	36.6	37.0
NaHCO₃, sodium bicarbonate	6.9	8.15	9.6	11.1	12.7	14.45
NaOH, sodium hydroxide	-----------	-------------	109	119	145	174
MgSO₄• 7 H₂O, epsom salts magnesium sulfate heptahydrate	------------	23.6	26.2	29	31.3	-------------

These values are the amount of solute that will dissolve and form a saturated solution at the temperature listed. A saturated solution is one where there is an equilibrium between undissolved solute and dissolved solute.

NaCl(s) <---> Na¹⁺(aq) + Cl^1-(aq)

The solvent cannot dissolve more solute at that temperature.The solubility can be increased if the temperature is increased. The table shows that solubility usually increases with increasing temperature. Clearly there are exceptions such as Ce₂(SO₄)₃

Reading the Solubility Plot

A saturatedKCl solution at 10^oC will have 31 grams of KCldissolved in 100 grams of water. If there are 40 grams ofKCl are in the container, then there will be 9 grams of undissolvedKClremaining in the solid.

Raising the temperature of the mixture to 30^oC will increase the amount of dissolvedKCl to 37 grams and there will be only 3 grams of solid undissolved. The entire 40 grams can be dissolved if the temperature is raised above 40^oC.

Cooling the hot 40^oC solution will reverse the process. When the temperature decreased to 20^oC the solubility will eventually be decreased to 34 gramKCl. There is a time delay before the extra 6 grams of dissolvedKClcrystallizes. This solution is "supersaturated" -- a temporary condition. The "extra" solute will come out of solution when the randomly moving solute particles can form the crystal pattern of the solid. A "seed" crystal is sometimes needed to provide the surface for solute particles to crystallize on and establish equilibrium.

Separation of Dissolved Liquids

Goal:

Purify 70% alcohol by mixing it with salt. Then measure the purity of your solution by calculating its density.

Theory:

Alcohol is usually purified by distillation. Distillation is the process of separating components by careful temperature controlled boiling and condensation. In theory, because alcohol boils at a lower temperature than water, heating a mixture will boil the alcohol off and leave only the water. But, alcohol isazeotropic at 4%. An azeotrope is a mixture of two or more liquids whose proportions cannot be altered or changed by simple distillation. This happens because when an azeotrope is boiled, the vapor has the same proportions as the unboiled mixture.That means that a 4% to 96% alcohol-water mixture actually boils at a lower temperature than pure alcohol. We are going to see if we can produce a purer alcohol solution using ionized salts. First we will mix a measured amount of NaCl (noniodized salt) into the isopropyl alcohol. The salt will dissolve in the water. Because salt water is heavier than alcohol, it will be drawn out of the alcohol mixture and form a layer at the bottom of the container.

Procedure:

1. Measure out about 50 mL of isopropyl alcohol and record its exact volume →
2. Find out how many grams of salt it takes to make a saturated solution at 20⁰, using the internet, or the graph above. →
3. Calculate how much salt you need by solving this equation →
4. Combine the alcohol and the salt you measured in steps 1-3. 5. Stir the mixture until the salt is completely dissolved. 6. Wait for the liquids to separate in a sealed container.
7. Weigh a graduated cylinder. →
8. Carefully pour off some of the purified alcohol from the top layer into the graduated cylinder recording thevolume of the alcohol →
9. Reweigh the graduated cylinder.→
10. Calculate the weight of the alcohol (subtract) →
11. Calculate the density of the alcohol →
12. Look up the density of isopropyl alcohol in the table at the end of this packet →
13. The density of water is	1.000 g/mL
14. Calculate the purity of your product using this formula: (Final density =density times percent alcohol + density times percent water)So … → [box 11] [box 13] [box 12] Solve this for %_alc

Do your math computations here:

2 Growing Crystals Using Supersaturation

Every solid that can be dissolved in water has a solubility, which is the largest quantity of the solid that can be dissolved in the water to make a clear solution. When the water starts getting cloudy and you can see solid particles floating around, that means no more solid can dissolve into the water and the solution (water and solid mixture) is saturated. But, the solubility of most solids increases as the mixture is heated, so more of the solid can be dissolved in hot water than in cold water. We are going to use this characteristic of solutions to make crystals in this lab.

When the molecules of the crystal come together, impurities (which are the unwanted products of the chemical reaction) do not fit into the structure, much like the wrong piece of a puzzle does not fit. So, if the crystal forms slowly enough, the impurities will be rejectedbecause they do not fit correctly, and instead, they remain in the solution and float away. But if a solution is cooled too quickly, there is not time to reject the impurities and instead, they become trapped in the crystal structure and the pattern is disturbed. Timing is important.

Procedure:

1. You will need Choose a salt from the front table. Available chemicals including: NaCl (table salt), CuSO₄∙5H₂O (root killer), Na₂B₄O₇∙10H₂O (borax).

2. Measure an amount of water that will fill a beaker to the half way mark, and pour it in the beaker.

3. Heat the water to about 50°C using a thermometer to monitor its temperature.

4. Look-up your salt up in the solubility graph above.

5. Measure out the amount of salt you will need to make a saturated solution. To do this, take the number of grams the graph indicated for the temperature of your water and multiply it by the number of mL of water in your beaker, then divide it by 100 mL, the amount of water the graph used as a basis.

6. Add the salt slowly while your partner stirs it in.

7. Stir patiently, but if the salt does not dissolve completely, heat it to a slightly higher temperature, or add a few drops of water. It is vital that the solution be saturated.

8. Using a pencil and some thread, hang a seed crystal in the solution, or just drop a crystal into your solution as a “seed”. The crystal will grow around the seed crystal and enlarge it.

9. Let the solution cool and crystalize.

10. Remove your crystal and dry it for grading.

11. Categorize it using the picture to the right.

Chemistry Lab

11 Ionic Salts 2: Growing Crystals Using Supersaturation

Why Do Some Solids Dissolve in Water?

	H₂O
NaCl(s)		Na⁺(aq)	+	Cl^-(aq)

Solubility Equilibria

	H₂O
C₁₂H₂₂O₁₁(s)		C₁₂H₂₂O₁₁(aq)

diagram and ionic solids dissociate to give solutions of the positive and negative ions they contain.

	H₂O
NaCl(s)		Na⁺(aq)	+	Cl^-(aq)

When the salt is first added, it dissolves and dissociates rapidly. The conductivity of the solution therefore increases rapidly at first.

	dissolve
NaCl(s)		Na⁺(aq)	+	Cl^-(aq)
	dissociate

			associate
Na⁺(aq)	+	Cl^-(aq)		NaCl(s)
			precipitate

Solubility Rules

A salt is soluble if it dissolves in water to give a solution with a concentration of at least 0.1 moles per liter at room temperature.
A salt is insoluble if the concentration of an aqueous solution is less than 0.001 M at room temperature.
Slightly soluble salts give solutions that fall between these extremes.

Solubility Rules for Ionic Compounds in Water

Soluble Salts

1. The Na⁺, K⁺, and NH₄⁺ ions form soluble salts. Thus, NaCl, KNO₃, (NH₄)₂SO₄, Na₂S, and (NH₄)₂CO₃ are soluble.

2. The nitrate (NO3-) ion forms soluble salts. Thus, Cu(NO₃)₂ and Fe(NO₃)₃ are soluble.

Insoluble Salts

1. Sulfides (S2-) are usually insoluble. Exceptions include Na₂S, K₂S, (NH₄)₂S, MgS, CaS, SrS, and BaS.

2. Oxides (O2-) are usually insoluble. Exceptions include Na₂O, K₂O, SrO, and BaO, which are soluble, and CaO, which is slightly soluble.

3. Hydroxides (OH-) are usually insoluble. Exceptions include NaOH, KOH, Sr(OH)₂, and Ba(OH)₂, which are soluble, and Ca(OH)₂, which is slightly soluble.

4. Chromates (CrO42-) are usually insoluble. Exceptions include Na₂CrO₄, K₂CrO₄, (NH₄)₂CrO₄, and MgCrO₄.

5. Phosphates (PO43-) and carbonates (CO32-) are usually insoluble. Exceptions include salts of the Na⁺, K⁺, and NH₄⁺ ions.

Solubility: How solubility is measured

Solubility of Common inorganic compounds in grams solute per 100 mL of water

Substance	0^oC	10^oC	20^oC	30^oC	40^oC	50^oC
KI, potassium iodide	127.5	136	144	152	160	168
KCl, potassium chloride	27.6	31.0	34.0	37.0	40.0	42.6
NaCl, sodium chloride	35.7	35.8	36.0	36.3	36.6	37.0
NaHCO₃, sodium bicarbonate	6.9	8.15	9.6	11.1	12.7	14.45
NaOH, sodium hydroxide	-----------	-------------	109	119	145	174
MgSO₄• 7 H₂O, epsom salts magnesium sulfate heptahydrate	------------	23.6	26.2	29	31.3	-------------

NaCl(s) <---> Na¹⁺(aq) + Cl^1-(aq)

Reading the Solubility Plot

2 Growing Crystals Using Supersaturation

Every solid that can be dissolved in water has a solubility, which is the largest quantity of the solid that can be dissolved in water to make a clear solution. When the water starts getting cloudy and you can see solid particles floating around, that means no more solid can dissolve into the water and the solution (water and solid mixture) is saturated. But, the solubility of most solids increases as the mixture is heated, so more of the solid can be dissolved in hot water than in cold water. We are going to use this characteristic of solutions to make crystals in this lab.

When the molecules of the crystal come together, impuritiesdo not fit into the structure, much like the wrong piece of a puzzle does not fit. So, if the crystal forms slowly enough, the impurities will be rejectedbecause they do not fit correctly, and instead, they remain in the solution and float away. But if a solution is cooled too quickly, there is not time to reject the impurities and instead, they become trapped in the crystal structure and the pattern is disturbed. Timing is important.

Procedure:

12. You will need to choose a salt from the front table. Available chemical include thing like: NaCl (table salt), CuSO₄∙5H₂O (root killer), Na₂B₄O₇∙10H₂O (borax).

13. Measure an amount of water that will fill a beaker to the half way mark, and pour it in the beaker. (Be sure to record the measurement.)

14. Heat the water to about 50°C using a thermometer to monitor its temperature.

15. Look-up your salt up in the solubility graph above.

16. Measure out the amount of salt you will need to make a saturated solution. To do this, take the number of grams the graph indicated for the temperature of your water and multiply it by the number of mL of water in your beaker, then divide it by 100 mL, (the amount of water the graph used as a basis.)

17. Add the salt slowly while your partner stirs it in.

18. Stir patiently, but if the salt does not dissolve completely, heat it to a slightly higher temperature, or add a few drops of water. It is vital that the solution be saturated.

19. Using a pencil and some thread, hang a seed crystal in the solution, or just drop a crystal into your solution as a “seed”. The crystal will grow around the seed crystal and enlarge it.

20. Let the solution cool and crystalize.

21. Remove your crystal and dry it for grading.

22. Categorize it using the picture to the right.

Writing up your report

Writing lab reports are a vital part of laboratory science, and a large part of you grade. The report must summarize what you hoped to accomplish, how you went about doing it, and the results you observed. Labs vary, but here are some suggestions:

1. Write your nameand the title of the lab at the top of the page.

2. Start with a description andpurpose of the lab.

3. List the materials you used and substitutions you made.

4. Write up the procedure you planned to follow.

5. Make a table out of the measurements you took and observations you made.

6. Record the formulas you used and the mathematics you did on the sheet.

7. List adaptations and changes you had to make as the lab progressed.

8. Write a conclusion, that is, what your results mean and answers to the questions posed by the purpose listed in #2.

9. List inherent problems with your approach, and sources of inaccuracy and error.

10. If possible, suggest an improved approach.

Every student is to submit a lab report. The report is not a group project.

Chemistry Lab

12 Ionic Salts 3: Electroplating With A Metal Ion Salt

Why Do Some Solids Dissolve in Water?

Ionic salts dissolve in water because energy is given off when the polar salt molecules form intermolecular bonds with the polar water molecules. The weak bonds that form between the solute and the solvent compensate for the energy needed to disrupt the structure of both the pure solute (the salt) and the solvent (the water). In the case of CuSO₄ and water, this process works enough to dissolve over 30 grams in a liter of water.

	H₂O
NaCl(s)		Na⁺(aq)	+	Cl^-(aq)

We can generally assume that salts dissociate into their ions when they dissolve in water.Ionic compounds dissolve in water if the energy given off when the ions interact with water molecules compensates for the energy needed to break the ionic bonds in the solid and the energy required to separate the water moleculesso that the ions can be inserted into solution.

	H₂O
NaCl(s)		Na⁺(aq)	+	Cl^-(aq)

When the salt is first added, it dissolves and dissociates rapidly. The conductivity of the solution therefore increases rapidly at first.

	dissolve
NaCl(s)		Na⁺(aq)	+	Cl^-(aq)
	dissociate

Solubility: How solubility is measured

NaCl(s) <---> Na¹⁺(aq) + Cl^1-(aq)

Reading the Solubility Plot

Determining Concentration and Electrical Power

Text Box: Conductivities of electrolytes all diminish with concentration -Steven Lower 2017 Professor Emeritus (Chemistry) at Simon Fraser University First, more salt may not make much difference. The conductivity of a solution does not rise evenly with its concentration. Many substances produce ions that interfere with their own ability to migrate with electrical currents because of interionic attractions. Adding more salt can have little effect.

Secondly, more conductivity will produce faster plating, but also not allow the copper time to adhere. Copper plating is sensitive to the level of the electric current amperage– that is the number of electrons,and its voltage– the amount force behind their movement. These values are in turn affected by resistance – the tendency to block electricity or turn it into heat. Which, in turn, is affected by the solution concentration. The math can be daunting. At this level, keep the electricity and concentrations low and you should do well.

3 Electroplating With A Metal Ion Salt

Electroplating is a form of electrolysis in which the electrodes play a bigger role than just conducting the current. Inelectrolysis, electricity is passed through an ionic substance to produce chemical reactions at the electrodes. Inelectroplating, the electrode itself goes into solution and solidifies again at the other electrode.Using electricity, you can coat the metal of one electrode with the metal of the other. Jewelry and silverware can be silver- or gold-plated, while zinc is often used to coat iron to protect against rust. Professional electroplatersuse specialized chemicals and equipment to ensure a high-quality coat.

Equipment:

A low voltage D.C.power supply. You can use a lab power supply, the peacock, ET or a battery.
Two alligator clip leads or insulated wires
A beaker or glass container
0.1M Copper sulfate solution
Copper electrode (wire, pipe, or strip all work well)
Object to coat (brass key, iron nail, washer, spring or something you own)
Eye protection

Procedure:

Choose a metal object to electroplate. Galvanized nails work poorly. They are coated already with zinc by electroplating or hot-dipping.
Prepare the object for copper-plating by cleaning it with toothpaste, Cleanser, or soap and water. Dry it off on a paper towel. Try not to touch it with your hands. The oil and acids on your hands will affect the process.
Pour some of your your 0.1 M copper sulfate solution into a beaker.
Use one alligator clip to attach the copper electrode to the positive terminal (+) of the power source (this is now the anode)and the other to attach your object to the negative terminal (now called the cathode).The anodeabsorbs electrons, the cathodeprovides them.
Partially suspend the object to be plated in the solution by wrapping the wire lead loosely around a pencil and placing the pencil across the mouth of the beaker. The alligator clip should not touch the solution.
Place the piece of copper into the solution, making sure it doesn't touch the object you want to electroplate. Be sure the solution level is below the alligator clip. An electrical circuit has now formed and current is flowing.
Leave the circuit running for 20-30 minutes, or until you are happy with the amount of copper on your object.

Make any innovations you wish to bring about a better result for more credit in this lab.

Theory:

The copper sulfate solution is an electrolyte. Electrolytes are ionic compounds that dissolve and allow water to conduct electricity from one electrode to the other.Pure water is nonconductive. Small quantities of salts commonly found in tap water are what makes it a danger to people working with electricity. When the current is flowing, oxidation(loss of electrons) happens at the copper anode(+), adding copper ions to the solution. Those ions travel on the electric current to the cathode(-), where reduction (gain of electrons) happens, plating the copper ions onto the object you attached. There were already copper ions present in the copper sulfate solution before you started, but the oxidation reaction at the anode will keep replacing them in the solution as they are plated onto the cathode. That will keep the reaction going.

Writing up your report

11. Write your nameand the title of the lab at the top of the page.

12. Start with a description andpurpose of the lab.

13. List the materials you used and substitutions you made.

14. Write up the procedure you planned to follow.

15. Make a table out of the measurements you took and observations you made.

16. Record the formulas you used and the mathematics you did on the sheet.

17. List adaptations and changes you had to make as the lab progressed.

18. Write a conclusion, that is, what your results mean and answers to the questions posed by the purpose listed in #2.

19. List inherent problems with your approach, and sources of inaccuracy and error.

20. If possible, suggest an improved approach.

Every student is to submit a lab report. The report is not a group project.

Chemistry Lab

13 Building a Battery

Battery Theory[6]

Electricity is the flow of electrons through a conductive path like a wire. This path is called a circuit.

Batteries have three parts, an anode (-)[7], a cathode (+)[8], and the electrolyte[9]. The cathode and anode (the positive and negative sides at either end of a traditional battery) are hooked up to an electrical circuit.

The chemical reactions in the battery causes a buildup of electrons at the anode. This results in an electrical difference between the anode and the cathode. You can think of this difference as an unstable build-up of the electrons. The electrons wants to rearrange themselves to get rid of this difference. But they do this in a certain way. Electrons repel each other and try to go to a place with fewer electrons.

In a battery, the only place to go is to the cathode. But, the electrolyte keeps the electrons from going straight from the anode to the cathode within the battery. When the circuit is closed (a wire connects the cathode and the anode) the electrons will be able to get to the cathode. In the picture above, the electrons go through the wire, lighting the light bulb along the way. This is one way of describing how electrical potential causes electrons to flow through the circuit.

However, these electrochemical processes change the chemicals in anode and cathode to make them stop supplying electrons. So there is a limited amount of power available in a battery.

When you recharge a battery, you change the direction of the flow of electrons using another power source, such as solar panels. The electrochemical processes happen in reverse, and the anode and cathode are restored to their original state and can again provide full power.

Anode and Cathode Ideas

Batteries work because the anode and cathode have a different electronegativity. There is a picture of Tomas Edison’s Nickle Iron battery (NiFe). It has Nickle oxide hydroxide NiO(OH) as a positive plate and iron as a negative plate. It produced 1.4 volts and was known for its durability often lasting for more than 20 years.

You are likely going to use two metals for the sake of simplicity. The table to the left was provided by the University of Texas to help their students decide which metals to use in their battery projects. Recall that two reactions are occurring at the same time. A reduction reaction takes place at the anode. Electrons are taken from metal ions and the metal deposits on the anode itself. The second reaction is an oxidation reaction at the cathode. Metal atoms are ionized and go into solution. Figure out the maximum voltage by subtracting the reduction potential from the oxidation potential.

Some home schoolers make a coke can battery constructed from a piece of thick copper grounding wire, a strip of aluminum from a soda can, and a glass of Coke. The aluminum was cut with a pair of tin snips and then sanded to remove the paint. [10]

This is an idea from a school like ours: To make a bleach battery, fill a beakers a third of the way with water. Add ¼ that amount of bleach. Fold one strip of copper and one strip of aluminum over the side of each cup with as much of the metal in the solution as possible.[11]

One of the first batteries, invented by Alessandro Volta, is the voltaic pile. It is a stack of alternating zinc and copper sheets separated by paper soaked in salt water or vinegar, creating a series of thin battery cells.[12]

dickens water battery The Dickens battery uses the magnesium as its source of electricity, which many you probably already know if you’ve ever used a fire starter, is a very energy dense material. The design is simple enough that pretty much anyone can make it.

You start out with thick, magnesium rods, which you can buy on Ebay. After that, you’ll need to fasten a metal electrode to the rod with a hose clamp. The metal used for this step is never specified, so feel free to try out a few different metals to see what nets you the best results (more on that in a moment).

After that, you wrap the rod in porous foam, and then coil copper wire around the foam. The idea is to allow water to pass through the foam, but to keep the copper from touching the electrode. Doing so won’t cause anything catastrophic, but your battery will stop producing energy.

After it’s all said and done, it should look like this:

From there, you’ll need a small jar to store this contraption, and you’ll have to puncture holes in the lid to allow the positive and negative contacts to push through. Fill the jar with tap water up to the top of the foam, and close the lid with the contacts exposed. You’ll also need to use something like caulk to seal the holes in the lid, thus keeping the water from evaporating. And that’s it! Your magnesium battery is all done.[13]

If you need more voltage, make a Daniell’s cell, invented by John Fredric Daniell. A Daniell’s cell is made up of a copper strip in a copper sulfate solution and a zinc strip in a zinc sulfate solution. A salt bridge connects the two electrolyte solutions.[14]

See the last page of this lab for a better table of anions and cations.

Making the electrolyte for your battery[15]

Fill a beaker half full of water. For an electrolyte solution, distilled water is the best choice. It will minimize the possible contaminants in the solution. Some contaminants could cause a reaction with the electrolyte ions. For example, if you are mixing a solution of NaCl and the water contains low levels of lead, you will get a precipitate coming out of solution. The removal of some of the ions from solution changes the strength of the solution.

Choose an electrolyte that supports the application best. For batteries, you should select an electrolyte that includes an element used in one or both of the half-cells. For example, if one of the half-cell reactions is with copper, a good choice of an electrolyte is CuCO3 or CuCl2. Both of these will support the half-cell by ensuring that there are Cu2+ ions in solution. You should choose a strong acid, a strong base or the salt of one of these. The high dissociation value of these compounds enhances the ability of the electrolyte solution to transport charge.

Text Box: Record the formula for your cathode
Record the formula for your anode
Record the formula for your electrolyte

Include notes on why your design is good, issues you have and inventions. Measure enough strong acid, strong base or salt to generate an electrolyte solution of sufficient strength to support the demands of the electrochemical cell. If the concentration of the electrolyte is too low, it can inhibit the operation of the electrochemical cell. The electrolyte concentration should be in the range of 1M. Strong acid, bases and salts therefore work better than weak acid and bases due to the higher degree of dissociation.

Making Molar Solutions[16]

Text Box: Record the atomic symbols and weights for each element in your electrolyte Molar (M) solutions are based on the number of moles of chemical in 1 liter of solution. A mole consists of 6.02×1023 molecules or atoms. Molecular weight (MW) is the weight of one mole of a chemical. Determine MW using a periodic table by adding the atomic mass of each atom in the chemical formula. Example: For the MW of CaCl2, add the atomic mass of Ca (40.01) to that of two Cl (2 x 35.45) to get 110.91 g/mole. Therefore, a 1M solution of CaCl2 consists of 110.91 g of CaCl2 dissolved in enough water to make one liter of solution.

Once the molecular weight of a chemical is known, the weight of chemical to dissolve in a solution for a molar solution less than 1M is calculated by the formula:

grams of chemical = (molarity of solution in mole/liter) x (MW of chemical in g/mole) x (ml of solution) ÷ 1000 ml/liter

For example, to make 100 ml of 0.1 M CaCl2 solution, use the previous formula to find out how much CaCl2 you need:

grams of CaCl2 = (0.1) x (110.91) x (100) ÷ (1000) = 1.11 g

Now you can make your solution: dissolve 1.11 g of CaCl2 in sufficient water to make 100 ml of solution. The amount of water needed will be slightly less than 100 ml.

Text Box: Multiply the weights by the quantities listed in the formula, then add them. Record the sum.

For example for H2O Hydrogen would be multiplied by two and oxygen by one. Because 2+16=18, you would enter 18 if you used H2O A balance and a volumetric flask are used to make molar solutions. A procedure for making a molar solution with a 100 ml volumetric flask is as follows:

Calculate the weight of chemical needed to make 100ml of solution using the above formula.

Weigh out amount of chemical needed using a balance.

Transfer the weighed out chemical to a clean, dry 100ml volumetric flask.

Fill to line

Slowly add distilled water to the volumetric flask. Wash all the chemical into the bottom of the flask as you do so. Keep adding water until you reach the 100ml mark on the neck of the flask.

Place the stopper in the flask and gently swirl the flask until all the chemical is dissolved.

If you don’t have a volumetric flask you can use a 100ml graduated cylinder instead. Just add the chemical to the graduated cylinder and then add distilled water until you reach the 100ml mark in the side of the cylinder.

Building your battery

You are going to need to combine the anode, cathode and electrolyte in a container like a beaker. Here are some pictures what other folks did:

Related image Salt Bridges[17]

Salt bridges maintains electrical neutrality within the internal circuit, preventing the cell from rapidly running its reaction to equilibrium. If no salt bridge were present, the solution in one half cell would accumulate negative charge and the solution in the other half cell would accumulate positive charge as the reaction proceeded, quickly preventing further reaction, and hence production of electricity.

Glass tube bridges[18]

One type of salt bridge consists of a U-shaped glass tube filled with a relatively inert electrolyte; usually potassium chloride or sodium chloride is used, although the diagram here illustrates the use of a potassium nitrate solution. The electrolyte is so chosen thatit does not react with any of the chemicals used in the cellthe anion and cation have similar conductivity, and hence similar migratory speed.

The electrolyte is often gelified with agar-agar to help prevent the intermixing of fluids which might otherwise occur.The conductivity of a glass tube bridge depends mostly on the concentration of the electrolyte solution. At concentrations below saturation, an increase in concentration increases conductivity. Beyond-saturation electrolyte content and narrow tube diameter may both lower conductivity.

Filter paper bridges[19]

The other type of salt bridge consists of a filter paper, also soaked with a relatively inert electrolyte, usually potassium chloride or sodium chloride because they are chemically inert. No gelification agent is required as the filter paper provides a solid medium for conduction.

Conductivity of this kind of salt bridge depends on a number of factors: the concentration of the electrolyte solution, the texture of the filter paper and the absorbing ability of the filter paper. Generally, smoother texture and higher absorbency equates to higher conductivity.

A porous disk or other porous barrier between the two half-cells may be used instead of a salt bridge; however, they basically serve the same purpose.

Calculating Voltage

Electrode Potential, written E⁰, is the cell voltage. It varies with concentration, surface area, temperature, pressure, lighting, and many other factors, but a mathematically perfect calculation is possible, if the chemistry of both the anode (-) and the cathode (+) are known. Here is the equation:

Procedure Notes:

Notes on chemicals

	Formula	Atomic Symbols	Atomic Weight	Quantity	Quantity Total	Molecular Weight
Cathode
Anode
Electrolyte



Rationale behind choice:

Notes on electrolyte preparation:

Volumetric flask capacity in mL
Molecular Wt. of electrolyte Salt for 1000 mL of 1 molar solution
Divide by 1000 (To get the mass require for 1mL of solution.)
Multiply by capacity in mL of Volumetric Flask
Multiply by the desired molarity
Amount weighed out. (Account for weighing inaccuracy)
Actual molarity of electrolyte

Salt bridges keep half reactions apart

Describe design of salt bridge and theory

Why this design is better:

Materials used:

Text Box: Draw a

picture of

your device

here

Battery effectiveness:

Cathode E⁰		Look up in table
Anode E⁰		Look up in table
Optimal Voltage E⁰_cath –E⁰_an		Subtract
LED test		None / Faint / Bright
Max Measured Voltage		Use a mutimeter
Max Measured Amperage		Use a mutimeter
Using internet, write out chemical reaction:

Conclusions:(Example: did it work? Why not optimal?)

Improvements you would make: (REQUIRED.)

14 Titration of Hydrochloric Acid with Sodium Hydroxide[20]

Cautions: Hydrochloric acid solution is a strong acid. Sodium hydroxide solution is a strong base. Both are harmful to skin and eyes. Affected areas should be washed thoroughly with copious amounts of water.

Purpose: The purpose of this lab is to determine the concentration of a hydrochloric acid solution using acid‐base titration.

Background: Titration is a technique that chemists use to determine the unknown concentration of a known solution (we know what chemical is dissolved, but not how much in a solution). Because we know what the chemical is, we know how it will react with other chemicals and we can use that reaction to determine the concentration of the solution by measuring the formation of product(s). In the case of an unknown concentration of acid, we can use a known concentration of hydroxide base. This type of reaction is a neutralization reaction, where salt and water are products of the reaction:

Acid+ Base à Salt + H₂O

We can use a pH indicator, a chemical that changes color depending on the pH, to show us when the reaction has completely neutralized. This point, where all acid was consumed and there is no excess of base, is called the equivalence point. We can use this equivalence point to determine the initial concentration of acid using a series of calculations. The goal of the titration is to get as close as possible to the equivalence point by careful addition of the base; this will ensure the calculated acid concentration is as close to the true value as possible. You will do three titrations and average the trials.

The terms below will help you understand the terminology used throughout the experiment:

• Titrant—the solution of known concentration is also called the standardized solution. In this lab, the titrant is sodium hydroxide solution.

• Buret—a long, cylindrical piece of glass that can be used to determine small, accurate quantities of a solution. A buret is controlled by a stopcock, a white Teflon piece that can be turned to deliver the solution. The markings on the buret are such thatyou must subtract the initial reading(where the titrant level is initially) from the final reading to determine the volume of base delivered. The buret measures 2 digits after the decimal point accurately.

• Volumetric pipette/pipette bulb—a thin glass tube with only one marking used to measure a very specific volume of liquid. You will use a pipette bulb to draw the liquid into the pipette.

• Text Box: Materials:
• 50‐mL Buret with clamp
• Phenolphthalein indicator
• 125 mL or 250‐mL Erlenmeyer flasks
• Buret funnel
• 250‐mL beaker
• 25‐mL volumetric pipette
• Pipette bulb Phenolphthalein—a pH indicator. In acidic and neutral solutions, the indicator is colorless, but in a basic solution, the color is a vibrant pink. The higher the pH is, the stronger the pink color is. The equivalence point will be when the color is a very faint pink color. Keep your flask with acid and indicator over a white piece of paper to ensure you can see the color change.

Also of importance in titrations are the calculations you need to determine the unknown concentration of the acid. These calculations are outlined below. You may want to refer to your notes from lecture for additional examples.

• Text Box: Molar value for standardized NaOH: Determination of moles of base delivered:After each titration, you will need to determine the number of moles of sodium hydroxide used. First, you will need to know the molarity of the solution (the solution has been previously standardized, meaning it has a very accurate molarity that has been experimentally determined). Write this down when you start the titration. Next, you must determine the volume of the solution delivered to reach the equivalence point. Next, you will find the moles of base used in the titration: Note that the volume of base is in L, not in mL

• Determine number of moles of HCl in flask: If you write the balanced reaction for the neutralization of sodium hydroxide and hydrochloric acid, you will see that the reaction proceeds in a 1:1 fashion. That is, for every hydroxide (OH^‐) ion added, it can neutralize exactly one hydronium (H⁺) ion. This is not always the case for neutralization reactions, and is thus not always the case for acid‐base titrations. The general formula is below, where the determined moles of base from the equation above are multiplied by the stoichiometric[21] ratio found by looking at the balanced equation:

• Determination of acid concentration: Now that you know the number of moles of acid in the flask (at the start of the titration, by the end, there is only water and salt), you can determine its initial concentration. Because you know the initial volume of acid used, you can use the following to determine the concentration:

Procedure:You will do at least three titrations. If you add too much base and the solution is too bright pink, you will need to discard the data and do another run. Also, if your titrations are greater than 1% different from each other, you will need to conduct additional titrations. (4 columns of data are provided for these purposes.) Patience in this lab will prevent you from having to do extra trials.

1. Record the molarity of the sodium hydroxide solution on the data sheet

2. Obtain about 100 mL of the sodium hydroxide solution in a clean beaker. This should be enough for the initial cleaning of your buret and for your first 3 trials.

3. Clean your buret: Add about 5 mL of the base solution from the beaker to the buret (use a funnel to pour). Move the funnel around while adding to ensure the sides of the buret are coated with base. Alternatively, you can remove the buret with the 5 mL of titrant from the buret stand and carefully tilt and rotate to coat all interior surfaces with the titrant. Drain the solution through the stopcock into a waste beaker. Repeat this rinse with a second 5 mL portion of base.

4. Pour more of the sodium hydroxide solution into the buret until it is near the 0.00 mL mark. Open the stopcock to allow several drops to rinse through the tip of the buret. This should eliminate any air bubbles in the buret tip. Record your initial buret reading on the data sheet for trial 1 (the volume does not need to be exactly 0.00 mL).

5. Draw 25.00 mL of the acid solution into the volumetric pipette and transfer this solution into an Erlenmeyer flask. Add 2‐3 drops of phenolphthalein to the acid solution in the flask.

6. Place the flask under the buret and start adding the base solution to the Erlenmeyer flask. Have one lab partner swirl the flask while the other controls the stopcock. When pink starts to develop, add the solution more slowly. At this point you should add one drop at a time followed by swirling until a very light pink color persists for at least 30 seconds. Remember, the lighter the pink the better.

7. Record the final reading of the buret. Wash the contents of the flask down the drain with water.

8. Refill the buret with more sodium hydroxide solution if necessary. Record the new volume under trial 2 on the data sheet. Pipette another sample of acid and add the phenolphthalein as before and titrate as before.

9. Conduct additional titrations until three of them differ by no more than 1.0%.

10. Complete the data sheet and post‐lab questions. Show your work for full credit.DATA SHEET

Name: Lab Partner:____

Text Box: Attach all of your calculations for full credit

Concentration of sodium hydroxide: __________________M

Balanced Chemical Equation of the titration reaction:

	Trial 1	Trial 2	Trial 3	Trial 4
Initial buret volume (mL)
Final buret volume (mL)
Volume of base (mL)
Volume of base (L)
Moles of base (mol)
Acid to Base Mole Ratio
Moles of acid (mol)
Volume of acid (L)
Acid concentration (M)
Average concentration (M)
Percent Difference

POST‐LAB QUESTIONS

Name: Lab Partner:____

1. How would it affect your results if you used a beaker with residual water in it to measure out your standardized sodium hydroxide solution?

______________________________________________________________________________ ______________________________________________________________________________

______________________________________________________________________________

2. How would it affect your results if you used awet Erlenmeyer flask instead of a dry one when transferring your acid solution from the volumetric pipette?

______________________________________________________________________________ ______________________________________________________________________________

______________________________________________________________________________

3. How do you tell if you have exceeded the equivalence point in your titration?

______________________________________________________________________________

4. Vinegar is a solution of acetic acid (CH₃COOH) in water. For quality control purposes, it can be titrated using sodium hydroxide to assure a specific % composition. If 25.00 mL of acetic acid is titrated with 9.08 mL of a standardized 2.293 M sodium hydroxide solution, what is the molarity of the vinegar?

Vinegar molarity: ____________________________

5. For the same data in question 4 above, what is the % composition (wt/wt%)?

Text Box: Attach all of your calculations for full credit

Vinegar % composition: _______________________

6. How will you know when your titration is finished?

______________________________________________________________________________

7. Label the pH scale below with acid, base, and neutral, indicating numbers for each.

8. On the scale above, use an arrow to show where your equivalence point is located.

Chemistry Lab

15 Gas Laws and Molar Mass

DETERMINE MOLAR MASS FROM VAPOR DENSITY [22]

According to the Ideal Gas Law, PV = nRT. Rearranging this equation to put n, the number of moles, onone side gives us: n = (PV)/(RT). For a sample of a gas in a container, R (the ideal gas constant) is known, and the pressure (P), volume (V), and temperature (T) are easy to determine experimentally. With these four values known, determining n, the number of moles of gas in the container, is a simple calculation. Because the number of moles in the sample (n) equals the mass of the sample divided by the molar mass of the substance, the only additional datum we need to calculate the molar mass of the substance is the mass of the sample.

Jean Baptiste André Dumas.jpg In 1826. the French chemist Jean Baptiste Andre Dumas developed a method and an apparatus for determining molar mass from vapor density.

Text Box: REQUIRED EQUIPMENT AND SUPPLIES
1. Goggles, gloves, and protective clothing
2. Balance
3. Barometer (or internet)
4. Thermometer graduated cylinder: 100 mL
5. Erlenmeyer flask. 250 mL or larger
6. Beaker, 250 mL.
7. Ring stand
8. Support clamp (for flask)
9. Warm water
10. Pin or needle
11. Aluminum foil
12. Acetone (- 5 mL) CAUTIONS Although only a small amount is used in this lab, acetone is extremely flammable. Be careful not to expose acetone liquid or vapor to an open flame. Use an exhaust hood, or work outdoors or in a well-ventilated area. Water at 50°C is hot enough to burn you badly. Wear splash goggles, gloves, and protective clothing.

First, accurately determine the mass of an empty Erlenmeyer flask and apiece of aluminum foil used to stopper it. Then introduce a small amount of liquid acetone and immerse the flask in a warm water bath to boil the acetone, converting it to vapor. As the acetone vaporizes, it displaces the airin the flask through a small pinhole in the aluminum foil. When all of the acetone has vaporized, we record the temperature (T), and then cool the flask, causing the vapor to condense to liquid acetone. We reweigh the flask to determine by difference the mass of acetone vapor it contained. We then determine the volume (V) of the flask, and the atmospheric pressure (P). Using those data, we calculate the molar mass of acetone.

PROCEDURE

1. If you have not already done so, put on your splash goggles, gloves, and protective clothing.

2. Loosely crimp a piece of aluminum foilaround the mouth of the flask.

3. Weigh the flask and aluminum foil, and record the mass to 0.01 g on line A of Table 14-5.

4. Transfer about 5 mL of acetone to the flask. The exact amount is not critical.

5. Crimp the aluminum foil tightly around the neck of the flask, covering the mouth, and use the pin or needle to Poke one tiny hole in the center of the foil.

6. Assemble the vapor density apparatus, as shown in the figure above.

7. Immerse the flask in the water bath, with the flask tilted slightly to make the liquid more visible, and gradually heat the flask until the liquid acetone boils. As the acetone boils, acetone vapor displaces the air in the flask.

8. Just as the final drop of acetone boils away, note the temperature of the water bath (which, if you've been heating it very gradually, is the same as the temperature inside the flask). Record the temperaturein kelvins on line B of Table 14-5.

9. Immediately remove the flask from the water bath, and run cold water over it. As the flask cools, the acetone vapor condenses inside the flask.

10. Carefully dry the outside of the flask with a paper towel.

11. Reweigh the flask, aluminum foil, and condensed acetone and record that mass to 0.01 g on line C of Table 14-5.

12. Remove the aluminum foil from the flask, pour out as much of the condensed acetoneas possible, and then replace the flask in the warm water bath. Allow the flask to warm for a minute or two while you complete the following calculations.

13. Calculate the mass of condensed acetone, and enter that value on line D of Table 14-5.

14. Calculate the number of moles of condensed acetone, and enter that value on line E of Table 14-5.

15. Remove the flask from the warm water bath and make sure that no liquid remains in the flask. Use a graduated cylinder to fill the flask with water, noting the total value of water held by the flask when it is full to the brim. Record that volumeas accurately as possible on line F of Table 14-5.

16. Use the barometer (or internet) to determine the atmospheric pressure. If we lived anywhere else, we would have to adjust the pressure reading for altitude differences, but this is flat South Florida. Record the atmospheric pressure on line G of Table 14-5.

17. With the temperature, mass of acetone, volume, and pressure known, you have all the information you need to calculate the molar mass of acetone. Make that calculation, and enter the value on line H of Table 14-5.

Text Box: The units of the universal gas constantR is derived from equation PV=nRT. R stands for Regnault.
Ppressure in atmospheres (atm),
Vvolume in liters (L),
n is moles (mol) of gas,
T temperatureis in Kelvin (K),
R is 0.082057L⋅atm/mol⋅K

Table 14-5
A	Mass of flask + foil
B	Temperature of water K°
C	Mass of flask + foil + condensate
D	Mass of condensed acetate
E	Moles of acetate
F	Volume of flask
G	Atmospheric pressure
H	Molar mass of acetone

Math Corner

Text Box: PV=nRT
So by dividing,
n= PV/RT P=box G

V=box F

n=moles

R=.082

T=box B

Solve for n

Chemistry Lab

16 Build a Blimp Using Gas Laws

Basic Idea

We are going to design and build a hot air balloon and fly it. In this first lab, we will measure the lifting capacity the candles will produce as they heat the air in the blimp and compare it to the mass of the materials used to build the blimp.

Testing

	Materials	Measurement
1	Mass of bag in g
2	Number of candles
3	Mass - one candle in g
4	Enter Line 2 x Line 3 à
5	Number of Straws
6	Mass - one straw in g
7	Enter Line 5 x Line 6 à
8	Mass of 10 cm of tape in g
9	Mass payload in g
	Total lines 1,4,7,8,9à

1 - Weighing: First choose the materials you will use to build your balloon. We are going to use 6 to 12 candles, four straws, a thin garbage bag, and some tape. You choose how many candles to use. The candles are a major contributor to weight, but they also are the power source that heats the air, making your air ship lighter.

2 After you choose and weigh your materials, calculate the volume of your blimp.

Calculating the blimp’s volume is easy. It is a cylinder and the top half of a sphere.

1. Measure the diameter of the filled bag.

2. Measure the height of the filled bag.

3. Measure the height of the bulge at the top of the bag.

4. Use the volume of a cylinder formula to get the volume of the main body of the blimp

5. Text Box: Volume of a cylinder
v=πr^2 h
Volume of a sphere
v=4/3 πr^3
h is the height of the bag
r is half the diameter
π is 3.14 Use the volume of a sphere formula to calculate the volume at the top bulge.

	Materials	Measurement
10	Height of bag in cm
11	Diameter in cm
12	=
13	=
14	Divide #13 by 2
	Total Volume = #12 + #14à

Hot Air Lifting Force[23]

The lifting force from a hot air balloon depends on the density difference between balloon air and surrounding air, and the balloon volume. The lifting force can be calculated as

F_l = V (ρ_c - ρ_h) a_g (1)

where

F_l = lifting force (N, lb_f)

V = balloon volume (m³, ft³)

ρ_c = density cold surrounding air (kg/m³, slugs/ft³)[24]

ρ_c = density hot balloon air (kg/m³, slugs/ft³)

a_g = acceleration of gravity (9.81 m/s², 32.174 ft/s²)

Example - Lifting Force created by a Hot Air Balloon

A hot air balloon with volume 10 m³(353 ft³) is heated to 100 ^oC (212 ^oF). The temperature of the surrounding air is 20 ^oC (68 ^oF). The air density at temperature 100 ^oC is 0.946 kg/m³ (0.00184 slugs/ft³) and the air density at temperature 20 ^oC is 1.205 kg/m³ (0.00234 slugs/ft³).

The lifting force can be calculated as

F_l = (10 m³) [(1.205 kg/m³) - (0.946 kg/m³)] (9.81 m/s²)

= 25.4 N

Weight - or gravity force - can be calculated as

Fg = m a_g (2)

where

Fg = weight - gravity force (N, lb_f)

m = mass (kg, slugs)

Since lifting force of a flying air balloon equals weight (F_l = F_g) - the lifted mass can be expressed by combining (1) and (2) to

m = F_l / a_g

= (25.4 N) / (9.81 m/s²)

= 2.6 kg

The calculation of lifting force can be done in Imperial units as

F_l = (353 ft³) [(0.00234 slugs/ft³) - (0.00184 slugs/ft³)] (32.174 ft/s²)

= 5.7 lb_f

Hot Air Balloon - Specific Lifting Force

Specific lifting force (force per unit air volume) created by an hot air balloon - balloon temperature vs. surrounded air temperature - are indicated in the charts below.

Chemistry Lab

17 Make a Ferromagnetic Liquid

Magnetite, a Simple method! [25]

You need only two readily available chemicals for this lab: Iron SulfateFeSO₄ that you can get from almost any gardening store as Moss killer for lawns(it is also pretty easy to make if you wanted to), and simple Washing Soda Na₂CO₃. If for some bizarre reason you can’t get washing soda, use baking soda that has been heated to a high heat for a while. Heating will convert it into washing soda.

1. Dissolve about 5 gof each in water.It is true that a mole of iron sulfate isabout 10 grams heavier than a mole ofsodium carbonate, but the reaction works best when there is an excess of sodium carbonate and it also keeps the procedure simple.

2. Mix both solutions, stirring well. The mixture will instantly make a horrible gray/green "mud" and thicken up a little too. This is normal; keep mixing. The “mud” is FeCO₃, iron carbonate.

Double substitution reaction: Na₂CO₃ (aq) + FeSO₄ (aq) à FeCO₃ (s) + NaSO₄ (aq)

3. When it`s all mixed, feel free to add more water. Mix it really well.

4. Let it stand for a while. The iron carbonate will start to settle to the bottom leaving a murky liquid on the top.

5. Decant this liquid carefully so as not to lose any iron carbonate.

6. Wash the product by adding water, mixing it and decant the liquid. Do this at least 4 more times. Make sure you wash out as much of the sodium sulfate NaSO₄solution as you can.

7. Filter the iron carbonate, a plain coffee filter is ideal for this, it should catch all the Iron Carbonate that you`ve made and get rid of what should be just water by now. Keeping the iron carbonate in the filter, paper put it somewhere to dry out, a sheet of plastic out in the sun is fine. When it is dry, it will crumble very easily and look just like rust.

8. You now need to heat the filtered powder up to decompose the carbonate. I used a crucible and Bunsen burner: You`ll notice during heating that the Brown rusty carbonate will go Black during heating; this is normal. Keep heating the powder, and keep the lid ON during this this stage. Don`t allow air to get in, if you do, you`ll end up with an impure product.

9. Let the powder cool. It will take on a deep red to black color as shown (it`s a bit more red as I took the lid off to watch so I could give you more data).

Congrats, you have just made a load of Magnetite, Fe₃O₄

a simple test with a magnet:

My result is a little bit on the RED side, but I`ll provide a further picture a bit later of the pure Black stuff.

Please notice the magnet is in a plastic bag, how else would I get it off the magnet if the powder decided to cover it!

As you can see the stuff even sticks to the spatula that isn`t even Magnetic; at least it shouldn`t be.

Bottom of Form

Procedure for Making the Ferromagnetic Liquid out of Magnetite[26]

1. Add 15 ml of ammonia to a beaker.

2. Now, add about 5 ml of the magnetite you created earlier.

3. Heat the mixture to about 70 degrees celcius, [27]

4. Then add about ½ ml of oleic acid.This will produce ammonia gas, so do this in a well-ventilated area! The oleic acid reacts with ammonia to form ammonium oleate. Heat causes the oleate ion to enter the mixture, while the ammonia escapes as a gas. When the oleate ion binds to the magnetite Fe₃O₄ it is reconverted to oleic acid. Oleic acid is a surfactant. It keeps the Fe₃O₄in nanoparticle size by preventing clumping.

5. Now,stir in 10 ml of motor oil.Only the particles coated with oleic acid will dissolve in the oil, and the oil will act as a carrier for your magnetite.

6. Let the mixturecool,

7. then pour off any water left on top of the mixture.

8. To purify the precipitates, [28]

a. dissolve it in petroleum in another beaker, only the particles coated with oleic acid will be dissolved.

b. Put a strong magnet under the beaker to keep the undissolved precipitates at the bottom. Pour the solution to another beaker.

[1] Thompson, Robert Bruice, Illustrated Guide to Home Chemistry Experiments pp.116-119.

[2] Standardized solutions containing a precisely known concentration of an element or a substance. A known weight of solute is dissolved to make a specific volume.

[3] Categorize shape: Linear, bent, trigonal, tetrahedral, trigonal pyramidal, trigonal bipyramidal, square planar, square pyramidal, octiheral

[4]http://www.chem.ucla.edu/~harding/lewisdots.html

[5] From Chapter 7 https://portal.gssd.ca/class/7b7uocm/Pages/Level%201/chapter-7.aspx

[6] Northwestern University: http://www.qrg.northwestern.edu/projects/vss/docs/power/2-how-do-batteries-work.html

[7] The positively charged electrode by which the electrons leave a device

[8] The negatively charged electrode by which electrons enter an electrical device

[9] A liquid or gel that contains ions and can be decomposed by electrolysis. Electrolyte serves as catalyst to make a battery conductive by promoting the movement of ions from the cathode to the anode on charge and in reverse on discharge. Ions are electrically charged atoms that have lost or gained electrons. The electrolyte of a battery consists of soluble salts, acids or other bases in liquid, gelled and dry formats.

[10]http://sciphile.org/lessons/survey-homemade-batteries

[11] Mr. Fleming, at https://www.zizzers.org/site/handlers/filedownload.ashx?moduleinstanceid=4599&dataid=6469&FileName=WPHS Chemistry Battery Papers 2015 - 2^ndDrafts.pdf

[12]https://sciencing.com/possible-materials-could-use-make-battery-11403.html

[13] Joshua Krause, http://readynutrition.com/resources/how-to-make-a-battery-that-lasts-practically-forever_21062015/

[15]https://sciencing.com/make-electrolyte-8668566.html

[16]https://learning-center.homesciencetools.com/article/making-chemical-solutions-science-lesson/

[17]Hogendoorn, Bob (2010). Heinemann Chemistry Enhanced (2). Melbourne, Australia: Pearson Australia. p. 416.

[20]CHEM 1011, Austin Peay State University Department of Chemistry,https://www.coursehero.com/file/14830197/SP12-1011-Titration-of-Hydrochloric-Acid-with-Sodium-Hydroxide-0/

[21]Stoichiometric /ˌstɔɪkɪəˈmɛtrɪk/ adjective (chem) concerned with, involving, or having the exact proportions for a particular chemical reaction: a stoichiometric mixture, http://www.dictionary.com/browse/stoichiometric

[22] Thompson, Robert Bruice, Illustrated Guide to Home Chemistry Experiments pp.264-267.

[23]https://www.engineeringtoolbox.com/hot-air-balloon-lifting-force-d_562.html

[24]Density (at 60°F and 1 atm): 1.208 kg/m3 = 0.00234 slug/ft³ = 0.0754 lb/ft³

[25]Derived from: By YT2095, June 12, 2008 in Experimentshttps://www.scienceforums.net/topic/30907-magnetite-a-simple-method/

[26] From “Liquid of the future” https://fear-of-lightning.wonderhowto.com/how-to/make-ferrofluid-liquid-future-0132750/

[27] “Ferrofluid Synthesis” https://www.linkedin.com/pulse/ferrofluid-synthesis-li-jen-chu

Half Cell Reaction E^o, V
Acidic Solution
F₂(g) + 2e^- → 2 F^-(aq)	+2.866
O₃(g) + 2H⁺(aq) + 2e^-→ O₂(g) + H₂O(l)	+2.075
S₂O₈^2-(aq) + 2e^- → 2SO₄^2-(aq)	+2.01
H₂O₂(aq) + 2H⁺(aq) +2e^- → 2H₂O(l)	+1.763
MnO₄^-(aq) + 8H⁺(aq) + 5e^- → Mn²⁺(aq) + 4H₂O(l)	+1.51
PbO₂(s) + 4H⁺(aq) + 2e^- → Pb²⁺(aq) + 4H₂O(l)	+1.455
Cl₂(g) + 2e^- → 2Cl^-(aq)	+1.358
Cr₂O₇^2-(aq) + 14H⁺(aq) + 6e^- → 2Cr³⁺(aq) + 7H₂O(l)	+1.33
MnO₂(s) + 4H⁺(aq) +2e^- -> Mn²⁺(aq) + 2H₂O(l)	+1.23
O₂(g) + 4H⁺(aq) + 4e^- → 2H₂O(l)	+1.229
2IO₃^-(aq) + 12H⁺(aq) + 10e^- → I₂(s) + 6H₂O(l)	+1.20
Br₂(l) + 2e^- → 2Br^-(aq)	+1.065
NO₃^-(aq) + 4H⁺(aq) + 3e^- → NO(g) + 2 H₂O(l)	+0.956
Ag⁺(aq) + e^- → Ag(s)	+0.800
Fe³⁺(aq) + e^- → Fe²⁺(aq)	+0.771
O₂(g) + 2H⁺(ag) + 2e^- → H₂O₂(aq)	+0.695
I₂(s) + 2e^- → 2I^-(aq)	+0.535
Cu²⁺(aq) + 2e^- → Cu(s)	+0.340
SO₄^2-(aq) + 4H⁺(aq) + 2e^- → 2H₂O(l) + SO₂(g)	+0.17


Sn⁴⁺(aq) + 2e^- → Sn²⁺(aq)	+0.154
S(s) + 2H⁺(aq) + 2e^- → H₂S(g)	+0.14
2H⁺(aq) + 2e^- → H₂(g)	0
Pb²⁺(aq) + 2e^- → Pb	-0.125
Sn²⁺(aq) + 2e^- → Sn(s)	-0.137
Fe²⁺(aq) + 2e^- → Fe(s)	-0.440
Zn²⁺ + 2e^- → Zn(s)	-0.763
Al³⁺(aq) + 3e^- → Al(s)	-1.676
Mg²⁺(aq) + 2e^- → Mg(s)	-2.356
Na⁺(aq) + e^- → Na(s)	-2.713
Ca²⁺(aq) + 2e^- → Ca(s)	-2.84
K⁺(aq) + + e^- → K(s)	-2.924
Li⁺(aq) + e^- → Li(s)	-3.040

Basic Solution
O₃(aq) + H₂O(l) + 2e^- → O₂(g) + 2OH^-(aq)	+1.246
OCl^-(aq) + H₂O(l) + 2e^- → Cl^-(aq) + 2OH^-(aq)	+0.890
O₂(g) + 2H₂O(l) +4e^- → 4OH^-(aq)	+0.401
2H₂O(l) + + 2e^- → H₂(aq) + 2OH^-(aq)	-0.0828

Chemistry Lab

Table of Reactions Obtained with Borax

Name:____________________________ Lab Partner:________________________________

Name:____________________________ Lab Partner:________________________________

Hot Air Lifting Force[23]

Example - Lifting Force created by a Hot Air Balloon

Hot Air Balloon - Specific Lifting Force

[20]CHEM 1011, Austin Peay State University Department of Chemistry,https://www.coursehero.com/file/14830197/SP12-1011-Titration-of-Hydrochloric-Acid-with-Sodium-Hydroxide-0/

[21]Stoichiometric /ˌstɔɪkɪəˈmɛtrɪk/ adjective (chem) concerned with, involving, or having the exact proportions for a particular chemical reaction: a stoichiometric mixture, http://www.dictionary.com/browse/stoichiometric

Name: Lab Partner:____

Name: Lab Partner:____