Hot Ice

Sodium acetate or hot ice is an amazing chemical you can prepare yourself from baking soda and vinegar. You can cool a solution of sodium acetate below its melting point and then cause the liquid to crystallize. The crystallization is an exothermic process, so the resulting ice is hot. Solidification occurs so quickly you can form sculptures as you pour the hot ice.

 

Sodium Acetate or Hot Ice Materials

• 1 liter clear vinegar (weak acetic acid)
• 4 tablespoons baking soda (sodium bicarbonate)

Prepare the Sodium Acetate or Hot Ice

1. In a saucepan or large beaker, add baking soda to the vinegar, a little at a time and stirring between additions. The baking soda and vinegar react to form sodium acetate and carbon dioxide gas. If you don't add the baking soda slowly, you'll essentially get a baking soda and vinegar volcano, which would overflow your container. You've made the sodium acetate, but it is too dilute to be very useful, so you need to remove most of the water. Here is the reaction between the baking soda and vinegar to produce the sodium acetate:
   Na⁺[HCO₃]^– + CH₃–COOH → CH₃–COO^– Na⁺ + H₂O + CO₂
   
2. Boil the solution to concentrate the sodium acetate. You could just remove the solution from heat once you have 100-150 ml of solution remaining, but the easiest way to get good results is to simply boil the solution until a crystal skin or film starts to form on the surface. This took me about an hour on the stove over medium heat. If you use lower heat you are less likely to get yellow or brown liguid, but it will take longer. If discoloration occurs, it's okay.
3. Once you remove the sodium acetate solution from heat, immediately cover it to prevent any further evaporation. I poured my solution into a separate container and covered it with plastic wrap. You should not have any crystals in your solution. If you do have crystals, stir a very small amount of water or vinegar into the solution, just sufficient to dissolve the crystals.
4. Place the covered container of sodium acetate solution in the refrigerator to chill.

Activities Involving Hot Ice

The sodium acetate in the solution in the refrigerator is an example of a supercooled liquid. That is, the sodium acetate exists in liquid form below its usual melting point. You can initiate crystallization by adding a small crystal of sodium acetate or possibly even by touching the surface of the sodium acetate solution with a spoon or finger. The crystallization is an example of an exothermic process. Heat is released as the 'ice' forms. To demonstrate supercooling, crystallization, and heat release you could:

• Drop a crystal into the container of cooled sodium acetate solution. The sodium acetate will crystallize within seconds, working outward from where you added the crystal. The crystal acts as a nucleation site or seed for rapid crystal growth. Although the solution just came out of the refrigerator, if you touch the container you will find it is now warm or hot.
• Pour the solution onto a shallow dish. If the hot ice does not spontaneously begin crystallization, you can touch it with a crystal of sodium acetate (you can usually scrape a small amount of sodium acetate from the side of the container you used earlier). The crystallization will progress from the dish up toward where you are pouring the liquid. You can construct towers of hot ice. The towers will be warm to the touch.
• You can re-melt sodium acetate and re-use it for demonstrations.

Hot Ice Safety

As you would expect, sodium acetate is a safe chemical for use in demonstrations. It is used as a food additive to enhance flavor and is the active chemical in many hot packs. The heat generated by the crystallization of a refrigerated sodium acetate solution should not present a burn hazard.

Hot Ice Help

Answers to common questions about hot ice are available that should help solve any problems you may encounter with this project. There is also a video tutorial showing how to make hot ice.

Black Snakes or Glow Worms

Black snakes, sometimes called glow worms, are small tablets that you light, using a punk or a lighter, that burn to produce long black 'snakes' of ash. They produce some smoke (which had a characteristic, probably toxic odor), but no fire or explosion. The original fireworks used to contain salts of a heavy metal (such as mercury), so while they were marketed for kids to play with, they really weren't that much safer than conventional fireworks, just dangerous in a different way. However, there is a safe way to make black snakes. You can heat baking soda (sodium bicarbonate) with sugar (sucrose) to produce carbon dioxide gas that puffs up black carbon ash.

Soda & Sugar Black Snake Materials

• sand
• alcohol or fuel oil (I didn't have any high-proof alcohol on hand, so I used lighter fluid left over from the handheld fireballs project)
• baking soda
• sugar (I used powdered sugar, but you can grind table sugar in a coffee grinder)

Make Snakes

• Mix 4 parts powdered sugar with 1 part baking soda. (I used 4 tsp sugar and 1 tsp baking soda.)
• Make a mound with the sand. Push a depression into the middle of the sand.
• Pour the alcohol or other fuel into the sand to wet it.
• Pour the sugar and soda mixture into the depression.
• Ignite the mound, using a lighter or match.

At first, you'll get a flame and some small scattered blackened balls. Once the reaction gets going, the carbon dioxide will puff up the carbonate into the continuously extruded 'snake'. Actually, you don't even need the sand. I tried this project using baking soda and sugar in a metal mixing bowl, added the fuel, and lit the mixture. It worked fine. The old firework snakes had a distinct smell. These have a smell too... burnt marshmallows! If you use pure ethanol, sugar, and baking soda, then there is nothing toxic about this project. One caution: Don't add fuel to the burning snake, since you risk igniting the alcohol stream. How Black Snakes Work
The sugar and baking soda snake proceeds according to the following chemical reactions, where sodium bicarbonate breaks down into sodium carbonate, water vapor, and carbon dioxide gas while burning the sugar in oxygen produces water vapor and carbon dioxide gas. The snake is carbonate with black carbon particles:
2 NaHCO₃ -> Na₂CO₃ + H₂O + CO₂
C₂H₅OH + 3 O₂ -> 2 CO₂ + 3 H₂O
These instructions were adapted from a tutorial given on Boing Boing which in turn came from a Russian site. The Russian site goes on to suggest two additional ways to make chemical snakes:
Ammonium Nitrate Black Snake
This works the same way as the sugar and baking soda snake, except you use ammonium nitrate (niter) instead of sugar. Mix one part ammonium nitrate and one part baking soda. This recipe is more like what you would see in commercial black snake fireworks, which are supposedly composed of soda with nitrated naphthalenes and linseed oil. It's another very safe demonstration, though not safe enough to eat, like sugar and baking soda.
Ammonium Dichromate Green Snake
The green snake is a variation on the ammonium dichromate volcano. The volcano is a cool chemistry demonstration (orange sparks, green ash, smoke), but it's a chemistry-lab-only demonstration (not safe for kids at all) because the chromium compound is toxic. The green soda snakes are made from:

• two parts of ammonium nitrate
• one part of powdered sugar
• one part of ammonium dichromate

Mix the ingredients, add a small amount of water, and roll the result into a snake shape (gloves please!). Allow the snake to dry (the tutorial suggests using a hairdryer to speed the process). Light one end of the snake. In this case, an orange snake burns to green ash. It's worth knowing how to do this demonstration if you have ammonium dichromate and ammonium nitrate on hand, otherwise let the Russian photos suffice and play with the sugar and baking soda snakes instead. Another (spectacular) form of black carbon snake results from reacting sugar and sulfuric acid.

Egg in Bottle

The egg in a bottle demonstration is an easy chemistry or physics demonstration you can do at home or in the lab. You set an egg on top of a bottle (as pictured). You change the temperature of the air inside the container either by dropping a piece of burning paper into the bottle or by directly heating/cooling the bottle. Air pushes the egg into the bottle.

 
Egg in a Bottle Materials

• peeled hard-boiled egg (or soft-boiled, if a yolk mess interests you)
• flask or jar with opening slightly smaller than the diameter of the egg
• paper/lighter or very hot water or very cold liquid

In a chemistry lab, this demonstration is most commonly performed using a 250-ml flask and a medium or large egg. If you are trying this demonstration at home, you can use a glass apple juice bottle. I used a Sobe™ soft drink bottle. If you use too large of an egg, it will get sucked into the bottle, but stuck (resulting in a gooey mess if the egg was soft-boiled). I recommend a medium egg for the Sobe™ bottle. An extra-large egg gets wedged in the bottle. Perform the Demonstration

• Method 1: Set a piece of paper on fire and drop it into the bottle. Set the egg on top of the bottle (small side pointed downward). When the flame goes out, the egg will get pushed into the bottle.
• Method 2: Set the egg on the bottle. Run the bottle under very hot tap water. Warmed air will escape around the egg. Set the bottle on the counter. As it cools, the egg will be pushed into the bottle.
• Method 3: Set the egg on the bottle. Immerse the bottle in a very cold liquid. I have heard of this being done using liquid nitrogen, but that sounds dangerous (could shatter the glass). I recommend trying ice water. The egg is pushed in as the air inside the bottle is chilled.

How It WorksIf you just set the egg on the bottle, its diameter is too large for it to slip inside. The pressure of the air inside and outside of the bottle is the same, so the only force that would cause the egg to enter the bottle is gravity. Gravity isn't sufficient to pull the egg inside the bottle.
When you change the temperature of the air inside the bottle, you change the pressure of the air inside the bottle. If you have a constant volume of air and heat it, the pressure of the air increases. If you cool the air, the pressure decreases. If you can lower the pressure inside the bottle enough, the air pressure outside the bottle will push the egg into the container.
It's easy to see how the pressure changes when you chill the bottle, but why is the egg pushed into the bottle when heat is applied? When you drop burning paper into the bottle, the paper will burn until the oxygen is consumed (or the paper is consumed, whichever comes first). Combustion heats the air in the bottle, increasing the air pressure. The heated air pushes the egg out of the way, making it appear to jump on the mouth of the bottle. As the air cools, the egg settles down and seals the mouth of the bottle. Now there is less air in the bottle than when you started, so it exerts less pressure. When the temperature inside and outside the bottle is the same, there is enough positive pressure outside the bottle to push the egg inside.
Heating the bottle produces the same result (and may be easier to do if you can't keep the paper burning long enough to put the egg on the bottle). The bottle and the air are heated. Hot air escapes from the bottle until the pressure both inside and outside the bottle is the same. As the bottle and air inside continue to cool, a pressure gradient builds, so the egg is pushed into the bottle.
How to Get the Egg Out
You can get the egg out by increasing the pressure inside the bottle so that it is higher than the pressure of the air outside of the bottle. Roll the egg around so it is situated with the small end resting in the mouth of the bottle. Tilt the bottle just enough so you can blow air inside the bottle. Roll the egg over the opening before you take your mouth away. Hold the bottle upside down and watch the egg 'fall' out of the bottle. Alternatively, you can apply negative pressure to the bottle by sucking the air out, but then you risk choking on an egg, so that's not a good plan.

Fireballs

Fire is made up of light and heated gases from combustion. You can control the temperature of fire by selecting a fuel that burns with a cool flame. If you pour the fuel onto a substance that won't burn, you can make a fireball that you can hold in your hand or juggle. Here are written instructions for making your own handheld fireballs. There is also a step-by-step video tutorial of this fire project if you would like see what to expect.
 

Materials Needed to Make Fireballs

• 2" x 5" strip of cotton cloth (like from a t-shirt)
• 100% cotton thread
• needle
• naphtha lighter fluid (e.g., Ronsonol™)
• match or lighter

How to Make a Fireball

• Thread the needle with cotton thread.
• Tightly roll the cotton strip into a ball.
• Pierce the ball with the needle and wrap the ball with the thread. End by running the needle through the ball one more time and break off the thread.
• Soak the ball with lighter fluid. Don't soak your hands.
• Don't ignite the the fireball while you are holding it. Set the ball on a fire-proof surface. I used a frying pan from my kitchen.
• If you want to hold the fireball, my recommendation is to pick it up with tongs and carefully/slowly set it on your hand. That way you'll be able to tell if you can take the heat or not. Once you gain some confidence, you can pick the fireball up using your fingers.

Safety & Additional Information

• It's best to use 100% cotton fabric and thread. If the fiber is synthetic (like nylon or polyester) it might burn or melt, with unpleasant consequences.
• The 'trick' to this demonstration is the fuel. It needs to be naphtha or kerosene. I have had good luck with Ronsonol™ and Zippo™ (not the butane stuff... read your ingredient list). Rubbing alcohol (isopropyl alcohol) works, but it burns a little hotter.
• It's pretty hard to blow the fireball out. You either need to blow hard or else suffocate the flame to extinguish it. You can set a saucepan lid over the fireball.
• The fireballs are reusable. Put them out when they run out of fuel or else the cotton will burn (you can tell this is happening when the ball starts to blacken and produce sooty smoke). If you get to the point where the cotton itself is burning, the fireball will be too hot to hold. Ideally you want to extinguish the fireball before it consumes all of its fuel. Simply soak it in more lighter fluid and relight it to reuse it.
• Regarding holding these in your hand or doing tricks with them... the cone of the flame is hot, especially above the ball, however, the fuel burns at a relatively low temperature. The flashpoint of Ronsonol™ brand of naphtha is 6°C or 43° F, with combustion mainly around 400°F. To put that in perspective, touching the fireball is a lot like touching a hot pizza right out of the oven (except without the sticky cheese part).

Fireballs are great fun to make, but like all fire projects, use proper safety precautions and common sense. Don't get burned or set your house or yard on fire. This is a project which requires adult supervision.

Green Fire

It's easy to make brilliant green fire. This cool chemistry project requires only two household chemicals.

 
Green Fire Materials

• Boric Acid
  Medical grade boric acid can be found in the pharmacy sections of some stores for use as a disinfectant. It is a white powder. It's not the same chemical as borax. I used Enoz Roach Away™, which is 99% boric acid, sold with household insecticides.
• Heet™ Gas Line Antifreeze and Water Remover
  Heet™ is sold with automotive chemicals.
• Metal or Stoneware Container
• Lighter

Instructions for Making Green Fire

1. Pour some Heet™ into the container. How much you use will determine how long your fire will burn. I used about a half cup of Heet™ for approximately 10 minutes of fire.
2. Sprinkle some boric acid into the liquid and swirl it around to mix it up. I used 1-2 teaspoons of powder. It won't all dissolve, so don't worry about some powder at the bottom of the container.
3. Set the container on a heat-safe surface and ignite it with a lighter. I have a video of green fire, if you would like to see what to expect.

Green Fire Tips & Safety Information

• Boric acid is a relatively safe household chemical. You can rinse the residue remaining in the container down the drain.
• This is an outdoor project. There isn't a lot of smoke produced, nor is it particularly toxic, but the heat is intense. It will set off your smoke alarm.
• Be sure to set your container on a heat-safe surface. Do not follow my extremely bad example and set it on your glass patio table. Similarly, don't use any container that might shatter dangerously. Use metal or possibly stoneware, not glass, wood, or plastic.
• Heet™ primarily is methanol (methyl alcohol). You could try this project with other types of alcohol. Possibilities include ethanol, such as vodka or Everclear, or isopropyl alcohol (rubbing alcohol). You might also try other common household metal salts for different flame colors.
• For example, I susbstituted rubbing alcohol (isopropyl alcohol) for the Heet™. The result was a fire that alternated from orange to blue to green. It wasn't as spectacular as the Heet™ fire, but it was still pretty cool.
• The green fire could be used as a stunning Halloween decoration in a cauldron or possibly inside a jack-o-lantern.
• Keep the chemicals for this project out of reach of children or pets, since methanol is harmful if swallowed. Read and follow any safety precautions listed on the labels of the specific products you use.

White smoke

React a jar of liquid and an apparently empty jar to make smoke. The white smoke chemistry demonstration is easy to perform and visually appealing.

Difficulty: Easy

Time Required: Minutes

 

Here's How:

1. Pour a small volume of hydrochloric acid into one of the jars. Swirl it around to coat the jar, and pour the excess back into its container. Place a square of cardboard over the jar to cover it.
2. Fill the second jar with ammonia. Cover it with the square of cardboard, which will now be separating the contents of the two containers.
3. Invert the jars, so the ammonia is on top and the apparently empty jar is on the bottom.
4. Hold the jars together and pull the cardboard away. Both jars should immediately fill with a cloud or 'smoke' of tiny ammonium chloride crystals.

Tips:

1. Wear gloves and safety goggles and perform the demonstration in a fume hood. Both ammonia and hydrochloric acid can give nasty chemical burns. As always, observe safe lab procedure.

What You Need:

• Ammonia (NH3)
• Hydrochloric Acid (HCl)
• 2 clean glass jars, both the same size, about 250 ml
• square of cardboard large enough to cover the mouth of the jar

Instant Fire

Potassium chlorate and ordinary table sugar are combined. When a drop of sulfuric acid is added, a reaction is catalyzed which produces heat, an amazing bright/tall purple flame, and a lot of smoke.

Difficulty: Easy

Time Required: minutes

 

Here's How:

1. Mix equal parts potassium chlorate and table sugar (sucrose) in a small glass jar or test tube. Choose a container you don't value, as the demonstration will probably cause it to shatter.
2. Place the mixture in a fume hood and equip lab safety gear (which you should be wearing anyway). To initiate the reaction, carefully add a drop or two of sulfuric acid to the powdered mixture. The mixture will burst into into a tall purple flame, accompanied by heat and a lot of smoke.
3. How it works: potassium chlorate (KClO₃) is a powerful oxidizer, used in matches and fireworks. Sucrose is an easy-to-oxidize energy source. When sulfuric acid is introduced, potassium chlorate decomposes to produce oxygen:2KClO₃(s) + heat —> 2KCl(s) + 3O₂(g)
   The sugar burns in the presence of oxygen. The flame is purple from the heating of the potassium (similar to a flame test).

Tips:

1. Perform this demonstration in a fume hood, as a considerable quantity of smoke will be produced. Alternatively, perform this demonstration outdoors.
2. Granulated table sugar is preferable to powdered sugar which is in turn preferable to reagent grade sucrose. The powdered sugar is capable of smothering the fire, while the granules of the reagent-grade sucrose may be too large to support a good reaction.
3. Follow proper safety precautions. Do not store the potassium chlorate and sugar mixture, as it can react spontaneously. Use care when removing the potassium chlorate from its container, to avoid sparking, which can ignite the container. Wear the usual protective gear when performing this reaction (goggles, lab coat, etc.).
4. The 'Dancing Gummi Bear' is a variation on this demonstration. Here, a small quantity of potassium chlorate is carefully heated in a large test tube, clamped to a ring stand over a flame, until it has melted. A Gummi Bear candy is added to the container, resulting in a vigorous reaction. The bear dances amidst bright purple flames.

What You Need:

• potassium chlorate
• powdered (confectioners) sugar or table sugar (sucrose)
• sulfuric acid
• small glass jar or test tube

Smoke Bomb Materials

The smoke bomb you would purchase from a fireworks store usually is made from potassium chlorate (KClO₃ - oxidizer), sugar (sucrose or dextrin - fuel), sodium bicarbonate (otherwise known as baking soda - to moderate the rate of the reaction and keep it from getting too hot), and a powdered organic dye (for colored smoke). When a commercial smoke bomb is burned, the reaction makes white smoke and the heat evaporates the organic dye. Commercial smoke bombs have small holes through which the smoke and dye are ejected, to create a jet of finely dispersed particles. Crafting this type of smoke bomb is beyond most of us, but you can make an effective smoke bomb quite easily. There are even colorants you can add if you want to make colored smoke. Let's start out with instructions for the easiest/safest type of smoke bomb you can make:
Smoke Bomb Materials

• sugar (sucrose or table sugar)
• potassium nitrate, KNO₃, also known as saltpeter (buy it online or you can find this at some garden supply stores in the fertilizer section, some pharmacies carry it too)
• skillet or pan
• aluminum foil

Once you've gathered your smoke bomb materials, it's easy to make the smoke bomb...

Make Copper Sulfate

Copper sulfate crystals are among the most beautiful crystals you can grow, but you might not have access to a chemistry lab or want to order the copper sulfate from a chemical supply company. That's okay, because you can make copper sulfate yourself using readily-available materials.

Materials for Making Copper Sulfate
There are actually a few different ways you can make copper sulfate yourself. This method relies on a little electrochemistry to get the job done. You will need:

• copper wire - which is high purity copper
• sulfuric acid - H₂SO₄ - battery acid
• water
• 6-volt battery

Make Copper Sulfate

1. Fill a jar or beaker with 5 ml concentrated sulfuric acid and 30 ml of water. If your sulfuric acid solution is already diluted, add less water.
2. Set two copper wires into the solution so that they are not touching each other.
3. Connect the wires to a 6-volt battery.
4. The solution will turn blue as copper sulfate is produced.

When you run electricity through copper electrodes which are separated from each other in a dilute sulfuric acid bath the negative electrode will evolve bubbles of hydrogen gas while the positive electrode will be dissolved into the sulfuric acid and oxidized by the current. Some of the copper from the positive electrode will make its way to the anode where it will be reduced. This cuts into your copper sulfate yield, but you can minimize the loss by taking some care with your set-up. Coil the wire for the positive electrode and set it at the bottom of your beaker or jar. Slip a piece of plastic tubing (e.g., a small length of aquarium hose) over the wire where it extends up from the coil to keep it from reacting with the solution near the anode. (If you had to strip your wire, just leave the insulating coating on the part that runs down into the liquid). Suspend the negative copper electrode (anode) over the cathode coil, leaving a good amount of space. When you connect the battery, you should get bubbles from the anode, but not the cathode. If you get bubbling at both electrodes, try increasing the distance between the electrodes. Most of the copper sulfate will be at the bottom of the container, separated from the anode. Collect Your Copper Sulfate
You can boil the copper sulfate solution to recover your copper sulfate. Because the solution contains sulfuric acid, you won't be able to boil the liquid off completely (and you need to take care not to touch the liquid, which will become concentrated acid). The copper sulfate will precipitate out as a blue powder. Pour off the sulfuric acid and reuse it to make more copper sulfate!
If you would prefer to have copper sulfate crystals, you can grow them directly from the blue solution that you prepared. Just allow the solution to evaporate. Again, use care in recovering your crystals because the solution is very acidic.

Fire Magic Trick

Do you like to play with fire? Try some fire magic tricks. This is a collection science magic tricks that involve flames or fire. You can color fire, hold it in your hand, and appear to bend it to your will.

Burning Money Fire Magic Trick

ICHIRO, Getty Images

Light a bill on fire and watch it not-burn. This fire magic trick works because paper currency isn't merely paper.

Ammonia - How to Prepare Ammonia Gas

These are instructions for preparing ammonia gas (NH₃)from ammonium chloride and calcium hydroxide in water.
Reactants

ammonium chloride (NH₄Cl)
calcium hydroxide [Ca(OH)₂]
Gas Preparation
Gently heat a mixture of ammonium chloride and calcium hydroxide in water. Collect the ammonia from the upward displacement of air in a hood.
Reaction
Ca(OH)₂ + 2NH₄Cl --> 2NH₃ + CaCl₂ + 2H₂O

Calculating Concentration

The concentration of a chemical solution refers to the amount of solute that is dissolved in a solvent. We normally think of a solute as a solid that is added to a solvent (e.g., adding table salt to water), but the solute could just as easily exist in another phase. For example, if we add a small amount of ethanol to water, then the ethanol is the solute and the water is the solvent. If we add a smaller amount of water to a larger amount of ethanol, then the water could be the solute!
Units of Concentration

Once you have identified the solute and solvent in a solution, you are ready to determine its concentration. Concentration may be expressed several different ways, using percent composition by mass, volume percent, mole fraction, molarity, molality, or normality.
1. Percent Composition by Mass (%) This is the mass of the solute divided by the mass of the solution (mass of solute plus mass of solvent), multiplied by 100.
Example:
Determine the percent composition by mass of a 100 g salt solution which contains 20 g salt.
Solution:
20 g NaCl / 100 g solution x 100 = 20% NaCl solution

2. Volume Percent (% v/v)
Volume percent or volume/volume percent most often is used when preparing solutions of liquids. Volume percent is defined as:
v/v % = [(volume of solute)/(volume of solution)] x 100%
Note that volume percent is relative to volume of solution, not volume of solvent. For example, wine is about 12% v/v ethanol. This means there are 12 ml ethanol for every 100 ml of wine. It is important to realize liqud and gas volumes are not necessarily additive. If you mix 12 ml of ethanol and 100 ml of wine, you will get less than 112 ml of solution.
As another example. 70% v/v rubbing alcohol may be prepared by taking 700 ml of isopropyl alcohol and adding sufficient water to obtain 1000 ml of solution (which will not be 300 ml).

 3. Mole Fraction (X)

This is the number of moles of a compound divided by the total number of moles of all chemical species in the solution. Keep in mind, the sum of all mole fractions in a solution always equals 1.
Example:
What are the mole fractions of the components of the solution formed when 92 g glycerol is mixed with 90 g water? (molecular weight water = 18; molecular weight of glycerol = 92)
Solution:
90 g water = 90 g x 1 mol / 18 g = 5 mol water
92 g glycerol = 92 g x 1 mol / 92 g = 1 mol glycerol
total mol = 5 + 1 = 6 mol
x_water = 5 mol / 6 mol = 0.833
x _glycerol = 1 mol / 6 mol = 0.167
It's a good idea to check your math by making sure the mole fractions add up to 1:
x_water + x_glycerol = .833 + 0.167 = 1.000

4. Molarity (M)

Molarity is probably the most commonly used unit of concentration. It is the number of moles of solute per liter of solution (not necessarily the same as the volume of solvent!).
Example:
What is the molarity of a solution made when water is added to 11 g CaCl₂ to make 100 mL of solution?
Solution:
11 g CaCl₂ / (110 g CaCl₂ / mol CaCl₂) = 0.10 mol CaCl₂
100 mL x 1 L / 1000 mL = 0.10 L
molarity = 0.10 mol / 0.10 L
molarity = 1.0 M

5. Molality (m)
Molality is the number of moles of solute per kilogram of solvent. Because the density of water at 25°C is about 1 kilogram per liter, molality is approximately equal to molarity for dilute aqueous solutions at this temperature. This is a useful approximation, but remember that it is only an approximation and doesn't apply when the solution is at a different temperature, isn't dilute, or uses a solvent other than water.
Example:
What is the molality of a solution of 10 g NaOH in 500 g water?
Solution:
10 g NaOH / (4 g NaOH / 1 mol NaOH) = 0.25 mol NaOH
500 g water x 1 kg / 1000 g = 0.50 kg water
molality = 0.25 mol / 0.50 kg
molality = 0.05 M / kg
molality = 0.50 m

6. Normality (N)
Normality is equal to the gram equivalent weight of a solute per liter of solution. A gram equivalent weight or equivalent is a measure of the reactive capcity of a given molecule. Normality is the only concentration unit that is reaction dependent.
Example:
1 M sulfuric acid (H₂SO₄) is 2 N for acid-base reactions because each mole of sulfuric acid provides 2 moles of H⁺ ions. On the other hand, 1 M sulfuric acid is 1 N for sulfate precipitation, since 1 mole of sulfuric acid provides 1 mole of sulfate ions.

Dilutions

You dilute a solution whenever you add solvent to a solution. Adding solvent results in a solution of lower concentration. You can calculate the concentration of a solution following a dilution by applying this equation:
M_iV_i = M_fV_f
where M is molarity, V is volume, and the subscripts i and f refer to the initial and final values.
Example:
How many millilieters of 5.5 M NaOH are needed to prepare 300 mL of 1.2 M NaOH?
Solution:
5.5 M x V₁ = 1.2 M x 0.3 L
V₁ = 1.2 M x 0.3 L / 5.5 M
V₁ = 0.065 L
V₁ = 65 mL
So, to prepare the 1.2 M NaOH solution, you pour 65 mL of 5.5 M NaOH into your container and add water to get 300 mL final volume.

How Carbon Monoxide Detectors Work

Carbon monoxide is an invisible odorless gas that is the leading cause of accidental poisoning deaths in America. Cabon Monoxide Detectors can alert you to dangerous levels of carbon monoxide.

How the First Carbon Monoxide Detectors Worked

Originally, carbon monoxide detectors were simple opto-chemical detectors that indicated the presence of carbon monoxide by exhibiting a color change when carbon monoxide reacted with a chemical on a white pad, producing a brownish or black color. These detectors do not require an external power source to function, but modern designs use audible alarms to confer a higher level of protection:

Biomimetic Carbon Monoxide Sensors

A opto-chemical or gel sensor interacts with synthetic hemoglobin, darkening in color when carbon monoxide is present and lightening in color when carbon monoxide concentrations are low. A light sensor may be used to detect the change in light levels to sound an alarm.

Semiconductor Carbon Monoxide Detectors

An integrated circuit monitors a sensor, tripping the alarm when concentrations of carbon monoxide are high. The sensor is made from thin wires of semiconducting tin dioxide that rest on an insulating ceramic base. Increasing carbon monoxide concentration reduces the electrical resistance of the sensor, causing the alarm to sound.

Electrochemical Carbon Monoxide Detectors

This is an electrochemical cell that is designed to produce current in relation to the amount of carbon monoxide present in the air. Carbon monoxide is oxidized to carbon dioxide at one electrode while oxygen is consumed at the other electrode. Sulfuric acid is the usual electrolyte that separates the electrodes. The current triggers the alarm or can even be used to quantify the amount of carbon monoxide that is present.

Chemical Oxygen Demand (COD)

Basic: 
Oxidation of organic substances with potassium dichromate and silver sulfate excess in boiling sulfuric acid. The amount of potassium dichromate is not reduced during the oxidation reaction is determined by how titrimetrik with standard solution of ferrous ammonium sulfate (FAS) and indicators feroin. Chloride ion concentration above 1 mg / L is closed with mercury sulfate.  
Reaction:
CxHyOz + Cr2O72-→ CO2 + H2O + Cr3+  
Cr2O72-(excess) + Fe2 + → Fe3 + + 2Cr3+ + H2O
Reagents:  
1. Examples  
2. Sulfuric acid 98% 
3. Silver sulfate solution 
4. Mercury sulfate solution  
5. Potassium dichromate solution 
6. Ferrous ammonium sulfate solution
7. Solution of potassium hydrogen phthalate
8. Feroin indicator solution 
Tools:
1. Pumpkin peck 50 ml, 100 ml, and 1000 ml
2. Goblets 250, 500, and 1000 ml
3. 25.50 measuring cup, and 100 ml
4. Watch glass
5. 25 ml Burette + clamp + Standards
6. The temperature of the appliance COD (cod tube, air condensers, equipment destruction, ice water bath, and shelves)
7. Electronic Scales 4 decimal places (Mettler AE 260)
8. Magnetic stir bar magnet and 2.5 cm
 
How it works:
1. In the pumpkin destruction COD included consecutive 20 ml sample or have been diluted, 5 ml of mercury sulfate. 10 ml of 0.1 N potassium dichromate and boiling stones
2. Air condenser mounted above the pumpkin COD then added 40 ml 98% sulfuric acid through the upper air condenser carefully
3. Destruction tool heated to the perfect red dots mark
4. Examples that have been prepared (no. 2) is inserted into the device destruction and didestruksikan for 120 minutes
5. When didestruksi / reflux lasted 10 minutes add 5 ml of silver sulphate through the upper condenser
6. Examples that have been didestruksi cooled in the open air to room temperature
7. Added 25 ml of distilled water through the upper air condenser as a rinse
8. Air condenser removed and cooled in ice water sample
9. Example titrated with 0.1 N ferrous ammonium sulfate with feroin indicator until the end point (color change from blue green to red brown)
10. Blank determination made by the same treatment as samples, but samples of distilled water replaced
11. Done setting the standard solution of potassium hydrogen phthalate with the same treatment as an example, but an example is replaced with a standard solution of potassium hydrogen phthalate
12. Added 10 ml of 0.1 N potassium dichromate solution into a flask COD sample / standard / blank that has been titrated to end point. Then titrated back with a solution of 0.1 N ferrous ammonium sulfate until the endpoint (this step for the determination of FAS solution of normality 0.1 N)

Calculation:
1. The calculation of the concentration of ferrous ammonium sulfate (FAS)
FAS = ml K2Cr2O7 N x N K2Cr2O7
                             ml Fas  
1. COD value calculation COD (mg O2 / L) = (a-b) x c x 8000 x d  
                                                                                     v  


Description:
a = volume of FAS to titrate the blank. ml
b = volume umtuk FAS titration sample.
c = normality of FAS. N
d = dilution factor
v = volume of sample. ml

Titration

Titration there are times when people refer to as the volumetric method, this is due to the volume of solution in the titration measurements play an important role. From taking a particular analyte by the volume until the reading of the volume of titrant used for titration's end affect all the results of the analysis. Therefore, the proper use of equipment in the titration also should not be overlooked.Volumetry method differentiated the types of reactions involved between titrant and analyte are:Acid-Base. There are a lot of acid and alkaline compounds which can be determined by titration. Both strong acid or strong base, the end point titrasipun very easily observed with the use of acid-base indicators such as fenolphtalein (PP), red metal, metallic orange, and others. At the equivalent point obtained, the solution is neutral but with a little addition of titrant to reach the end point of titration is enough to change the color of acid-base indicator. Another way is to use pHmeter. Weak acid and weak base can also be titrated as well as the organic acid is titrated with non-water solvent.Reduction-Oxidation. Substances that are oxidizing agents such as KMnO4, K2CrO4, I2, and substances that are reducing agents such as H2C2O4, Fe2, Sn2 can be determined by this titration method. Redox reactions involved as titrant and analyte react. Some of the redox titration method does not require the indicator to look like a titration end point titration of KMnO4 and H2C2O4 caused KMnO4 itself is colored. Amylum usually used for titrations involving I2.Kompleksometri. Complex formation reaction between EDTA and metal ions underlies this method. EDTA is a widely used type of titrant for titration kompleksometri and react with many metals, reaksinyapun can be controlled by controlling the pH of the solution.Precipitation. Reaction formation of sediment form the basis of this method. Titrant and analyte react to form a precipitate, such as the determination of chloride ion using AgNO3 titrant. Indicators can be used to determine the end point titration for example K2CrO4 for titration using silver nitrate titrant.