Acids vs. Bases

A fairly quick overview of the procedure:

1. Use a pipet bulb to pipet 10mL of HCL solution into a 250mL beaker. Add 50mL of distilled water.
2. Place the beaker on a magnetic stirrer and add a stirring bar.
3. Use a utility clamp to suspend a pH Sensor on a ring stand. Position the pH Sensor in the HCL    Solution and adjust its position so that it is not in the way of the stirring bar.
4. Obtain a 50mL buret and use a utility clamp to attach the buret to the rin stand. Fill the buret with 10mL of NaOH Solution.
5. Prepare the computer for data collection in LoggerPro.
6. Before adding NaOH titrant, click the collect button and moniot the pH for about 5 seconds. Once the pH reading has stabilized, click the keep button. In the edit box, type "0."
7. Begin the titration process and enter the values as you go.
8. Clean Up!

Here is our lovely graph!

Last Reflection Eva!

So it's almost the complete end of school for me! I'm very pleased with this class and my performance in it. It's been really fun and I've felt that I've learned a lot. I completed all the labs, and I did them correctly and safetly the entire year. My last Acid vs. Bases lab is the only thing I wish I could improve. I didn't get to do any analysis on it, but I have NO TIME. So it will just have to stay as it is. Overall, I'm really glad I took this class and I've learned more about Chemistry. Now off to graduation!

Beer's Law Lab, but It Isn't About Beer, Just Concentration, but Not That Concentration Like Focus, It's About Solutions and Their Concentration, but Not Solutions Like Solutions to a Problem. . . .

The main point of this lab was to determine the concentration of an unknown solution. In order to do this we:

1.   Added 30mL of .40 NiSo4 solution into a beaker, and 30mL of water into a separate beaker.

2.   We took four clean test tubes and labeled them 1 through 4. Then, with a pipet, we put 2, 4, 6, and 8 mL of .4 M NiSO4 solution into test tubes 1-4 respectively. We repeated the process with water putting 8, 6, 4, and 2 mL of distilled water into Tubes 1-4 respectively. Then stir!

3.   On a computer, open the LoggerPro file "Exp 11 Colorimeter."

4.   Calibrate the Colorimeter

5.   Fill a cuvette 3/4 full of the solution from test tube 1. Then place it in the Colorimeter and collect the data on LoggerPro.

6.   Repeat these steps for test tubes 2-4. Afterwards, take a sample of the .40 M NiSO4 solution and record its data as well. 

7.   Examine the data by clicking the Linear Regression button, this will then but a line of best fit into the graph.

8.   Once this is finished, find the beaker of unknown solution, or in this case, the three bottles. Test each of these samples in the Colorimeter. With the information you gain from testing each of these samples you should be able to determine the concentration of each solution. 

For Unknown Solution #1, we found an absorbance of .248. Plugging that number into our graph above means that the mystery solution has a concentration of about .1474 mol/L.

For Unknown Solution #2, there was an absorbance of .558. That means that the solution had a concentration of about .3676 mol/L/

Finally, for Unknown Solution #3, there was an recorded absorbance of .430, meaning that the solution's concentration was around .2602 mol/L.

Crystal Lab!

To make a crystal you must:

1. Take a beaker and fill it with distilled water.

2. Use a hot plate to warm up the beaker and begin to saturate the water with Aluminum Potassium Sulfate.

3. Continue to allow the beaker to heat up until all of the Aluminum Potassium Sulfate is dissolved. 

4. Let the solution sit and cool completely.

5. Once this is done there should be one or more seed crystals at the bottom of the beaker. 

6. Take out the seed crystal you want to use and tie it to a stick.

7. Clean out the beaker and then repeat steps 1-4.

8. After the solution is cooled enough, take the seed crystal so that is dangles in the solution. If everything works out well, you crystal will begin to grow and grow.

9. Simply repeat the steps in order to make the crystal grow larger and larger.

THERE YOU GO!

Independent Research Project

Ultralight Metallic Microlattice on Prezi

And Here Are Some
WORKS CITED

• Berglund, Jennifer. "Technology / Materials Science." Engineers â ¨Concoct the World's Lightest Material. Discover Magazine, 29 Mar. 2012. Web. 01 Apr. 2012. <http://discovermagazine.com/2012/mar/31-engineers-concoct-the-worlds-lightest-material>. 

• Edward, Lynch. "Micro-Lattice, The Worldâ  s Lightest Material." Laughing Squid. 28 Nov. 2011. Web. 01 Apr. 2012. <http://laughingsquid.com/micro-lattice-the-worlds-lightest-material/>. 

• Lee, Kevin. "World's Lightest Material Takes Hits Like a Champ." PCWorld. Geektech, 21 Nov. 2011. Web. 01 Apr. 2012. <http://www.pcworld.com/article/244412/worlds_lightest_material_takes_hits_like_a_champ.html>. 

• Liggett, Brit. "Scientists Develop World's Lightest Metal, 100x Lighter than Styrofoam | Inhabitat - Sustainable Design Innovation, Eco Architecture, Green Building." Inhabitat. Inhabitat.com, 20 Nov. 2011. Web. 02 Apr. 2012. <http://inhabitat.com/scientists-develop-worlds-lightest-metal-100x-lighter-than-styrofoam/>. 

• McCormack, Simon. "Lightest Material On Earth Invented (VIDEO)." The Huffington Post. TheHuffingtonPost.com, 18 Nov. 2011. Web. 02 Apr. 2012. <http://www.huffingtonpost.com/2011/11/18/lightest-material-on-earth_n_1101909.html>. 

• "Metallic Microlattice." Wikipedia. Wikimedia Foundation, 04 July 2012. Web. 03 Apr. 2012. <http://en.wikipedia.org/wiki/Metallic_microlattice>. 

• Netburn, Deborah. "Scientists Invent Lightest Material on Earth. What Now?" The Sydney Morning Herald. SMH.com, 18 Nov. 2011. Web. 02 Apr. 2012. <http://www.smh.com.au/technology/sci-tech/scientists-invent-lightest-material-on-earth-what-now-20111118-1nmok.html>. 

• "World's 'lightest Material' Unveiled by US Engineers." BBC News. BBC, 18 Nov. 2011. Web. 29 Mar. 2012. <http://www.bbc.co.uk/news/technology-15788735>.

Mirror Mirror On The Wall, Who's The Smartest of Them All?

I am! So says me in my THIRD QUARTER REFLECTION! This quarter was a bunch of math and calculations and numbers. All of which I am pretty good at, and relatively enjoy. I really did understand finding the limiting and excess reactant, and such things like that. It made sense to me, and I was able to do the math fairly well. I did really well on my tests, and I got some positive feedback on my blogs. I understand percent yield, and I'm able to tell you what my data says. I've been working really hard on keeping up with my work and putting out good stuff. The only trouble I had this quarter was being absent during this class because of other stuff I had to do. Yeah, I'm awesome.

Yet Another Stoichiometry Lab, Only This Time It's Dry

Technically called Stoichiometry of a Double Replacement Reaction, but anyway. Since this is too dangerous/toxic/nasty to do in real life, we were given a piece of paper, and told to do this has a dry lab. So:

Procedure:

1. Clean and dry two beakers
2. Record the number of you chemical: Lab #5
3. Find the mass of lead(II) nitrate: 1.03 grams
4. Find the mass of potassium iodide: 1.33 grams
5. Add about 50mL of DI water to each beaker
6. Stir the beakers until both chemicals have dissolved
7. Combine the contents of both beakers into one beaker
8. Wash any remaining solution into the one beaker
9. Stir the solution
10. Measure the mass of filter paper: 1.39 grams
11. Filter out the precipitate into the filter paper
12. Wash any remaining precipitate into the filter paper 
13. Allow the precipitate and filter paper to dry overnight
14. Find the mass of the precipitate: .95 grams

The balanced equation for the chemical reaction is: 2KI + Pb(NO3)2 -> 2KNO3 + PbI2

The limiting reagent for this experiment was: Lead (II) Nitrate

Theoretically, there should be: 1.43 grams of lead(II) iodide formed from the experiment

In reality, there was actually: .95 grams of lead(II) iodide formed from the experiment

The percent yield of lead(II) from the experiment was: 66%

Here is my work to figure out all of the lovely answers. It's just proof, don't expect me to explain this all again.

Evaporation and Intermolecular Attraction Lab

Because it is impossible to make a punny title out of that.

Materials:

• Computer
• Serial Box Interface ULI
• Data Logger
• Two Probes
• 6 Pieces of Filter Paper
• 2 Small Rubber Bands
• Methanol
• Ethanol
• 1-Propanol
• 1-Butanol
• n-Hexane
• n-Heptane

Procedure:

1. Open "Experiment 9" from the Chemistry with Computers from the experiment files of Logger Pro.
2. Wrap filter paper around the tips of both probes, and secure it with the rubber bands. 
3. Place the both probes into a container holding a sample of liquid.
4. After the probes have been in the liquids for around 30 seconds, begin data collection. Monitor the temperature for around 15 seconds to establish the initial temperature of each liquid.
5. Simultaneously remove the probes from the liquids and tape them with the tips of the probes extended of the edge of a table.
6. When both temperatures have reached minimums, stop data collection.
7. Find the maximum and minimum temperatures, then subtract them to find the change of temperature during evaporation.
8. Remove the rubber bands and discard the used filter paper.
9. Predict, using the data just collected, how the next set of liquids will react.
10. REPEAT!

Heptane (red) and Hexane (green) Test

Ethanol (red) and Methanol (green) Test

Propanol (red) and Butanol (green) Test

Data Analysis:

1. n-Heptane and 1-Butanol had nearly the same molecular weights, but had significantly different changes in temperature due to hydrogen bonds. 1-Butanol contains hydrogen bonds, but n-Heptane does not. This made it easier for n-Heptane to evaporate compared to 1-Butanol.
2. Methanol had the weakest intermolecular forces, while 1-Butanol had the strongest. We can see that methanol had the weakest bonds because of the change of temperature. Methanol changed 18.4 degrees. This shows that it had weak bonds, allowing it to evaporate more and quicker than the other liquids. 1-Butanol had the strongest bonds because its temperature did not change very much at all. It's temperature only went down by about 3.07 degrees. This means that it had stronger intermolecular forces that did not allow it to evaporate as much or as quickly as methanol. 
3. In regards of alkanes, n-Hexane had the weakest intermolecular forces, and n-Heptane had the strongest molecular forces. n-Hexane changed a total of 17 degrees, and n-Heptane changes only 9.41 degrees.
4. 
   

So a Limiting and Excess Reactant Walk Into a Bar. . .

. . . and they ask for a 100% yield of H20! Hahahaha! I'm so funny right?!

Ok, for real now! For the last couple of days we have been working on finding the limiting and execs reactants in a reaction, also known as stoichiometry. We were also introduced to percent yield today, and all of this will be shown now.

Ok so let's do problem 21 of chapter 12. This problem plays out as such:

Photosynthesis reactions in green plants use carbon dioxide and water to produce glucose (C6H12O6) and oxygen. Wite the balanced chemical equation for the reaction. If a plant has 88g carbon dioxide and 64g water available for photosynthesis, determine:

1. The limiting reactant
2. The excess reactant ant the mass in excess
3. The mass of glucose produced.

Alright, so this is what the equations will look like to get the answer:

First the equation needs to be balanced, it will look like 6CO2 + 6H2O = C6H12O6 + 6O2. Now we can start with the actual calculation part.

Let's start with finding what is the limiting and execs reactants. We'll start with carbon dioxide because it comes first. The problem gives you that there is 88grams of CO2. We'll then convert the grams of carbon dioxide to mols. So, there are 44 grams for one mol of CO2 (yes it's labeled wrong in my notes). Now we can find the amount of glucose through the mol ratio between glucose and carbon dioxide. So, for every 6 CO2s there is 1 C6H12O6. The next part is to find how many grams are in one mol of C6H12O6. Putting it all together, there are 180grams of glucose per one mol. Multiply everything together, and then do proper division, and you'll get an answer of 60 grams of Glucose.

If you repeat the same process, but with water instead of carbon dioxide, you'll get 106.6 grams of glucose. This means that the carbon dioxide is the limiting reactant, and water is the execs reactant.
_________________________________________________________________________________
Now to get the percent yield. To get the percent yield we need to use a different problem, in which we are given the actual percent yield. We'll try problem 27 out of the same chapter. It says:

If 14 grams of aluminum hydroxide is present in an  antacid tablet, determine the theoretical yield of aluminum chloride produced when the tablet reacts with stomach acid. If the actual yield of aluminum chloride from this tablet is 22 grams, what is the percent yield?
Al(OH)2 + 3HCl -> AlCl3 + 3H2O

If you do all the math and such you will get a theoretical yield of 23.9 grams of AlCl3. Our actual percent yield is 22 grams of AlCl3. To get the percent yield we need to take the actual yield divided by the theoretical yield.

That would give us: 22/23.9 = .882.
Multiply this by 100 and you get a total of 88.2% yield of Ag.   DONE!

I'm Awesome

So to show you that I know what I'm doing in this class and I'm awesome, here is my test that I aced. It is known as a Chemistry Homework Quiz, and I completed it 14 out of 14. It was over percent composition, empirical and molecular formulas, and hydrations. So yeah, I know my concepts.

And this is what happens when I get distracted while doing my blogs:

An Apple product, within an Apple product, within an Apple product; with the picture being taken with another iPad. It's like a string quartet for Apple.

Silver/Copper Replacement Lab

Which I was gone for, but here it is regardless.

Procedure (Day 1):

1. Obtain 30cm of bare copper wire. Clean the wire, then coil it around a pencil to form a loose spring. Make sure that the other end of the wire reaches the top of your test tube and is uncoiled.
2. Weight the coil with a balance, then place the copper wire in the test tube to check that is is the correct length. Remove the coil and set it aside for later.
3. Weight the weighing dish of silver nitrate and record it's number. 
4. Transfer the contents of silver nitrate to your test tube.
5. Pour distilled water into your test tube until the water is about 2 cm from the top.
6. Cover the top of the test tube with parafilm.
7. Place your thumb on top of the test tube and invert it until all the silver nitrate is disolved. 
8. Weigh the empty weighing dish so that you can determine the mas of silver nitrate that was added to the test tube.
9. Add the copper coil to the test tube.
10. Set the tube aside until the next day or class period.
11. Make sure all data is recorded.

Procedure (Day 2):

1. Weigh a piece of fiter paper for use in separating the silver.
2. Shake the test tube and copper wire to dislodge the silver.
3. Set up a funnel with your filter paper in it.
4. With a waste beaker beneath the funnel, lift the copper wire out of the test tube and hold it over the filter system.
5. Using a water bottle, let D.I. water fun down the wire so that any silver will wash into the filter.
6. Lay the copper wire on a labeled piece of paper to allow it to dry.
7. Pour off the solution in the test tube through the filter paper into the filter, trying to keep the silver precipitate in the test tube.
8. Rinse with distilled water and decant several times into the filter to wash the silver.
9. Finally wash the silver from the test tube onto the filter paper.
10. Allow to dry overnight. 

Procedure (Day 3):

1. Weigh the copper coil and record it's mass.
2. Weigh the silver and filter paper - record the mass.

_________________________________________________________________________________

Mass of silver nitrate = 1.028 grams

Mass of copper coil before reaction = 3.408 grams

Mass of copper coil after reaction = 3.193 grams

Mass of copper reacted = .215 grams

Mass of filter paper & silver = 1.726 grams

Mass of filter paper = 1.420 grams

Mass of silver produced in reaction =.360 grams

Moles of silver produced = .0028 mols

Moles of copper consumed during reaction = .0033 mols

_________________________________________________________________________________

The balanced equation fo this reaction is:  2AgNO3 + Cu  ->  Cu(No3)2 + 2Ag

.65 grams of Ag were formed

.77 grams of Cu became Cu(No3)2

The Baking Soda Lab of Dreams

Materials:

• 20 ml vinegar in a large pipette
• Balance
• 100 ml beaker
• 1 gram of baking soda

Procedure:

1. Get a large plastic pipette filled with vinegar. Measure the mass of the pipette and record the mass.
2. Measure the mass of an empty, clean 00 ml beaker.
3. Transfer about 1 gram of baking soda to the beaker. Record the exact mass of the beaker with the powder.
4. Add vinegar, from the pipette, to the beaker. Swirl the contents and observe the reaction. Continue to add vinegar until no more bubbles are produced. This will take a while so be patient and pay close attention to the reaction.
5. Find the mass of the left-over vinegar in the pipette and record. Subtract the original mass of this pipette to find the mass of the vinegar used in the reaction.

From the picture above, you can see the data we collected throughout the lab. When all was said and done, we calculated that the moles of sodium hydrogen carbonate used was .012 moles. We did this by taking the net mass of baking soda (1.01grams) and dividing it by the mass of NaHCO3 (84grams). 

Next, we calculated the grams of acetic acid added to the beaker for the reaction. First, we were given that there are 5grams of acetic acid per 100grams of vinegar. We then decided to set up a ratio, knowing that we had used 13.34grams of vinegar. If there were 5grams of acetic per 100grams of vinegar, then that was equal to xgrams of vinegar per 13.34grams of vinegar. We cross multiplied to get 100x=66.7, thus resulting in x=.667grams of acetic.

We then calculated the moles of acetic acid used in the reaction. We took the .667grams of acetic acid divided by the actual mass of acetic acid (HC2H3O2). This gave us .011g/mol of acetic acid.

We then found the ratio between the moles of sodium hydrogen carbonate and the moles of acetic acid used in the experiment. Since the .011moles of acetic acid was the smallest number, it served as the 1 in our ratio. We then took the .012 moles of sodium hydrogen carbonate and divided it by the .011 moles of acetic acid, which gave us a ration of 1:1. This, our experimental mole ratio, corresponds to the mole ratio of the balanced equation: NaHCO3 + HC2H3O2 -> CO2 + NaC2H3O2 + H2O, which is a 1:1 ratio.

Mole Classwork

Here is just a lovely picture of some of the work I've been doing in class. I've been working in our Glenco Chemistry book.  Page 221, Chapter Eleven, Problems 42-57 to be exact and if it helps. These have all been over subjects the include: empirical formula, moles, and molecular formula. For example, problem 50 asks:

"What is the empirical formula for a compound that contains 10.89% magnesium, 31.77% chlorine, and 57.34% oxygen?"


To solve this problem I first wrote down everything I knew which looked a lot like this:

Magnesium: 10.89%

Chlorine: 31.77%

Oxygen: 57.43%

Then I found the relative atomic mass in grams of each element (24.3 for Mg, 35.5 for Cl, and 16 for O). I then took the percent of the element and divided it by its weight, getting numbers such as: .448, .895, and 3.584 respectively. 

Now for an empirical formula, you have to find the smallest working ratios between all the elements. Because Mg came up with the smallest number, .448, it will serve as 1 in the ratio. Then take the remaining numbers, .895 and 3.584, and divide them by .448. You will then get about 2 and 8 respectively. This completes your ratio, which is 1:2:8. This means that more every one Magnesium there are two chlorines, and eight oxygens. The complete empirical formula woud then be MCl2O8.

And that's what I've been doing in class. Lots of math. 

  

Hydrates Lab

Procedure:
1. Figure out the molar mass of CuSO4*5H2O
2. Figure out the mass of 5H2O,  then make a prediction about what percent of the entire thing is water.
3. After making this prediction, weigh out 5grams of Copper Sulfate Pentahydrate.
4. Take the Copper Sulfate and place it in a test tube. Use a bunsen burner to heat the test tube until the Copper Sulfate is not longer blue.
5. Wait until the tube is cooled, then weight the Copper Sulfate again.
6. Record the differences of the Copper Sulfate before and after it was heated to find the weight lost. Then calculate exactly how much water was lost.
7. For fun, take remains of the heated Copper Sulfate and pour water on it.

My Group:  
1. My group figured out that the molar mass of CuSo4*5H2O was 250 grams.
2. The mass of 5H2O was 90 grams. Therefore our prediction was that 36% of the Copper Sulfate was water.
3. We completed our procedure and weighted our 5g sample after begin heated. After a date with the bunsen burner, the Copper Sulfate weighed 3.28 grams, being that there was a total weight loss of 1.72 grams.
4. Therefore, we concluded that 34% of our sample of Copper Sulfate was actually water.
5. *SPOILERS* After the Copper Sulfate is heated it turns white. By adding water back to it, it returns to its original blue color!

How Much Water is in Popcorn? Well Let Me Tell You a Story. . .

Procedure:
1. Add enough vegetable oil to cover the bottom of the beaker. Make a tight cover for the beaker for the aluminum foil. Weight the beaker as it is and record the weight.
2. Add 25-30 kernels of popcorn to the beaker and replace the lid. Weight the beaker again, with the popcorn in it.
3. Poke several holes in the aluminum foil and place it on the ring stand over the bunsen burner.
4. When the popcorn is finished popping, remove the heat and carefully take off the aluminum cover.
5. Let the beaker stand until it is cool. Once it has cooled, weigh the beaker with the popped corn and the cover on. Record!

Alright! Our group's results:
Beaker, Oil, & Aluminium: 200g
Beaker, Oil, 30 Kernels, Aluminium: 204.8g
Beaker, Oil, Popped Popcorn, Aluminium after heated: 204.3g

We lost approximately 1% of mass due to getting rid of the water. We believe that we would have lost more mass if we would have followed the directions exactly. We did not remove the aluminum cover when we were supposed to, because popping popcorn we could eat was a lot more important. :D

Mole Lab!

I've been sick! I'm back now, and even though I missed this lab, some of my classmates were kind enough to lend me their data. I've also had this lab explained to me, so allow me to explain it to you!

The class separated into different groups and were given different containers of different elements. It's all different. You have no information on anything about the element other than how many moles it has. So! You take the elements in their containers and weigh them! You now have the number of moles and the weight of every element.

Now you take the weight of grams divided by the number of moles. Once you have gathered this number you compare it to a periodic table. It should match up with just one element of the table, and there you go! You have found your element!

I promise my next lab write-up will be better because I won't be dead or dying.

2nd Quarter Self-Analysis

I did awesome this quarter! I really understood most of the concepts, especially chemical names. I remember a lot of the concepts and tricks from Intro to Chemistry, so I really understood how to name chemicals and elements and all that jazz. Not to mention that we got a lot of help through cheat-sheets, which makes it even easier to name things. There really wasn't a lot of concepts I didn't understand. . . Ok there were no concepts that I didn't understand. Overall, this has been a really good quarter for me, as I have understood everything really. . . So yeah. Merry Christmas!!!!

Chemistry Ornaments: Star Wars for Science

So for our last lab this semester, we made ornaments!

1. First you grab a lovely piece of metal. It can be what ever size you want, small or big.

2. Then you tape the entire piece of metal until it looks like a little mummy.
3. Once your metal is mummified, you can cut a design into the tape. Whatever you cut out will eventually rust, so make sure that any metal showing is what you actually want to rust out.

4. Once your design is cut out, place the piece of metal in a container full of hydrochloric acid. Make sure that the piece of metal is completely submerged, and keep it in the acid for around 15 to 20 minutes.

5. After your piece of metal is ready, use a pair of tongs to pull it out from the acid. Then place the metal in tub full of copper sulfate for another 20 minutes, or until you see your design rusted out.

Your final product should look something like this:

Metal Activity for the Win!

We did a metal lab a few weeks ago but I am barely writing it up right now. :D

Procedure:

1. Place 8 drops of copper (II) nitrate in the first column
2. Place 8 drops of magnesium nitrate in the second column
3. Place 8 drops of zinc nitrate in the third column
4. Place 8 drops of silver nitrate in the fourth column
5. Place a small copper wire in each of the wells in the first row
6. Place a small magnesium ribbon in each well in the second row
7. Place a few zinc grannies in the third row
8. Make sure that the metals are submerged in the liquids
9. After 5 minutes, observe each reaction that occurs on the solid metal and record on a chart
10. Clean up!

If there is a R in the box it means there was a reaction.

If there is a NR in the box it means that there was no reaction.

The 7 Reaction Labs of Christmas

Types of Reactions:

• Synthesis: A + B -> AB        Example: N2 + H2 -> NH3
• Decomposition: AB -> A +B    
• Single Replacement: A + BC -> B + AC
• Double Replacement: AB + CD -> CB + AD
• Combustion: Fuel + O2 -> H2O + CO2

_________________________________________________________________________________

Ok so the Labs went like this:

1. Fill 1/2 of a test tube with hydrochloric acid, place a piece of magnesium metal in the tube, and hold a second test tube over the tube containing the acid and magnesium. Hold the test tube on top until the reaction has stopped. Light a match and place it inside the inside of the top test tube.

• The Mg disappeared, and there was a loud popping noise when the match encountered the gas. 
• 2HCl + Mg -> H2 + MgCl2
• Sing Replacement Reaction

----------------------------------------------------------------------------------------------------------------------------------
     2. Light a Bunsen burner and take a small strip of Mg and hold it in the flames using a tongs.

• There was sparks and a bright flare of light. The Mg was reduced into a white powder. Overall, this was a very fast reaction.
• Mg2 + O2 -> 2(MgO)
• Synthesis Reaction

----------------------------------------------------------------------------------------------------------------------------------

      3. Shine a copper wire using a piece of sand paper, and using tongs, hold the wire in the flames of a   Bunsen burner for several minutes

• When you hold the copper in the fire long enough the flames turn green. The copper eventually began to melt.
• Cu + O2 -Heat> CuO
• Synthesis Reaction

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       4. Add a small amount of ammonium carbonate into a test tube, and using a test tube holder, hear the test tube. Waft the gas coming from the test tube towards your nose. 

• Starts to smell fine and even a little sweet, but the smell soon turns sour and smells like ammonia. The substance evaporated and their was vapor at the top of the test tube. 
• (NH4)2 + CO3 -> 2(NH3) + CO2 + H2O
• Decomposition Reaction

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      5.  Fill a test tube half full of hydrogen peroxide, then add a small amount of manganese dioxide. Using a test tube holder, hold a tube over the first tube. When the reaction is complete, hold the test tube upside down and place a glowing splint into the inverted test tube. 

• The mixture ended up turning black, and the stick relit in when it was entered into the tube. The smoke also settled at the bottom of the tube. 
• 2(H2O2) -MnO2> O2 + 2(H2O)
• Decomposition Reaction

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      6.  Pour a small amount of potassium iodide into a test tube, and then add in a small amount of lead (II) nitrate. 

• The reaction turned yellow.
• 2(KI) + Pb(NO3)2 -> KNO3 + PBI2
• Double Replacement Reaction

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     7. Place a small amount of copper (II) carbonate into a test tube. Heat the test tube while having a second inverted test tube over it in order to collect any gas given off. Light a wood splint and place it into the test tube that was on the top. 

• The calcium in the tube burned up. The smoke from the reaction collected at the bottom of the tube and the top of the tube seemed to be wet. 
• CuCO3 -Heat> CO2 + CuO
• Decomposition