file:///Users/cmathews/Pictures/iPhoto%20Library/Originals/2010/Pictures_2/DSCN1404_0016.jpg

Final Thoughts

So guys, we have had our ups and downs. I have enjoyed having you in class (besides the paper balls and laser pointers). For your final post, due today, I would like to know what part of this class you liked best or learned the most about chemistry.

I have counted up blogs and will post it today so you know where you stand.

Good luck and I hope to see you next year or the year after in AP Chemistry.

Dr. B

5/17/10

Hey guys I though I would just remind everyone that it is seriously important to do well on your exams this semester as well as the chemistry exam. The second semester of sophomore year and the first semester of junior year are what most colleges tend to look at the most. If you don't have very challenging exams one day (like religion or ethics) it is proabably a good idea to start studying for harder exams like math and chemistry on that night. Good luck to all.

5-15-10

I am guessing that I still need to get some more blog points so I am putting one up.

We have a test over chapters 14 and 15 on Monday.

That plus the exam is all the chemistry we have left.

5-14-10

Hello fellow students. We had quite a fiasco yesterday with paper balls and laser pointers. Let's try to do better today.

Ok, so blog-wise, we have a test on Monday over chapter 14 and chapter 15. And our exam is on thursday. So it looks like we are going to have to hit it hard this weekend, like Ben and those balls.

good luck, we REALLY need it.

5/12/10

Okay guys, we had a quiz today over acids and bases. Is it just me or did Dr. B say yesterday that it was going to be on stoichiometry? But anyways we have a test on Monday that is going to be over both chapters 14 and 15. She wants us to have both calculations and conceptual questions. I suggest that you start studying now for the test because it's an exam day and we have shorter class periods.

5/11/10: Exam Review

Hey guys! Guess how smart Jim is! I'll give you a hint: he forgot to bring home his Chemistry notebook two days in a row. But fear not! I've recreated an outline of the stuff we need to know for the exam with my book. This probably covers more than we talked about in class, so feel free to correct me in the comments.

• The exam will go from chapter 7 to chapter 15. Here are most of the important terms and concepts from each chapter, as found in the end of chapter summaries.
  
• Ch. 7: Formula mass, percentage composition, empirical formula
• Ch. 8: Chemical equations, precipitate, types of reactions, electrolysis, activity series
• Ch. 9: Stoichiometry, limiting reactant, excess reactant, theoretical yield, actual yield, percentage yield
• Ch. 10: Kinetic-molecular theory, ideal gas, real gas, fluid, surface tension, capillary action, crystalline solids, amorphous solids, phase, equilibrium, molar enthalpy of vaporization/fusion, triple point, critical point, critical temperature, critical pressure
• Ch. 11: Pressure, mm Hg, pascal, atm, Dalton's Law of Partial Pressure, Boyle's law, Charles' law, Gay-Lussac's law, combined gas law, Gay-Lussac's law of combining volumes of gases, Avogadro's law, ideal gas law, ideal gas constant, Graham's law of effusion
• Ch. 12: Solubility rules, solvent, solute, solition, suspension, colloid, electrolytes, saturation, miscible/immiscible, Henry's law, effervescence, enthalpy of solution, concentration, molarity, molality
• Ch. 13: Net ionic equation, spectator ions, ionization, strong/weak electrolytes, colligative properties, molal freezing-point constant, freezing-point depression, molal boiling-point constant, boiling-point elevation, osmosis, osmotic pressure
• Ch. 14: Binary acid, oxyacid, Arrhenius acid/base, Bronsted-Lowry acid/base, monoprotic acid, polyprotic acid, Lewis acid/base, conjugate acid/base, neutralization
• Ch. 15: pH, pOH, titration, equivalence point, end point, standard solution, primary standard
  

Also, I would like to remind everybody that you should eat chocolate before all of your exams.
http://www.slashfood.com/2006/05/29/chocolate-stimulates-the-brain/

5/11/10

Hey Guys, there was no blog last night so I decided to make one. The notebooks are due tomorrow and she is counting up our blogs. I hope everyone has enough!

5/7/10

Once again guys, I need blogs so I guess I will keep on posting until we have an assigned blogger. On Friday we retook the quiz and I think most people did much better. Thank you Dr. B. We took about 15 min. before the quiz to review the material on it and it really helped me. After the quiz we just chilled out for the rest of class. Our test should be coming up fairly soon so start studying.

5/6/10

Ok I don't know if there was a scheduled blogger for tonight but I need posts so I am going to go ahead and give one.

I (and many others) failed the quiz today so Dr. B. is letting us retake it tomorrow. You need to know how to calculate molarity and how to use acids (H3O) and bases (OH) in calculations. You also need to know how to apply pH and pOH in calculations. In addition, you need to be able to tell weak acids apart from strong acids and weak bases apart from strong bases. Good Luck!

5.5.10.

Hey guys, just thought I'd go ahead and get a blog up here.

Well, today in class Dr. B showed us how to do all of the needed calculations.

In conclusion, lab is due tomorrow as is the homework from the text book; remember to get it all done.

*NOTE* - on your lab, be sure to show calculations, sig figs, and label your work for clearer understanding; you can never use too much partial credit. Put the pre-lab and post-lab questions INTO your Lab Notebook as well.

Well, have a good night guys.

5/4/10

Wasn't there supposed to be a blog last night? Well, yesterday was the last day for lab. Everyone remember to do your post lab questions!

5/3/10

Tomorrow is LAST DAY for lab. Remember you are supposed to have 3 titrations with KHP and 3 titrations with Vinegar.

4/30/10

On Friday, we continued our lab. Our goal was to complete 3 titrations. Did everyone meet the goal?

THE LAB...

All right, everyone. I'm pretty sure everyone is pretty confused on what we are supposed to do for THE LAB, so I'm going to try and put out there what I think we are supposed to do. Scully, I'm counting on you especially to call me out for being wrong on anything.

Ok, I'll start off with some definitons:

1. titration-adding NaOH to KHP until it turns a "the lightest shade of pink"
2. KHP-potassium hydrogen phthalate; it's the stuff that look like big salt
3. NaOH-sodium hydroxide; it's what goes in the buret, and is located in a keg-thing by the microwaves
4. phenolphthalein-[fee-nawl-thal-een] stuff that is in X-lax (not the lacrosse possition!) it's what makes the titration, your mixture of NaOH and KHP turn pink

Next off, some general starter. I'll leave some stuff some other guys to comment though

1. --get 2, TWO, samples of KHP; keep it as close as you can to 0.50 g, but it's better to go lower
2. --dissolve KHP in a beaker of ionized water and 2, TWO, drops of pheolthinggummy; swirl it around until it is completely dissolved
3. --while one guy is doing step two, the other can use the tiny beaker Dr. B gave us, fill it with NaOH from the keg-thing, and pour it into the buret.
4. initial NaOH volume reading=0.something(measure from zero); record this before you titrate;
5. stick the beaker of dissolved KHP under the buret and open the valve. THIS IS YOU TITRATING
6. When you're titrating, let in 5 or so mL of NaOH, then swirl, let in another 5 and swirl, let in another 5 and swirl, AND THEN start adding it in drops sssslllooooowwlllllyyyyyyyy (you add it slowly because the point where it is "the faintest shade of pink" arrives really fast)
7. after you titrate, make your Final NaOH volume reading
8. Final NaOH volume reading= somewhere in the teens(measure from zero)

Now all you do is repeat the steps, except you don't have to necessarily add more NaOH. You can just go from where the last one was. You just have to take that into account in your volume readings.

Ok, I really hoped that helped some of you. If it didn't, well I went to St. Ann, you can't expect me to be that good.

AND IF YOU THINK I'M WRONG, FREAKIN' TELL ME. INFALLIBLE IS NOT ME. that guy's German, and is probably just waking up right now to celebrate some Jesus in Rome. If he isn't somewhere else in the world, but I digress.

goodnight!

4-28-10

Hey, everyone. I'm going to do a quick blog tonight because i want to go to bed early, and, well, we didn't have that much to do.

We worked on those worksheets while Scully and Dr. B were at their chemistry luncheon.

Everyone, make sure you have your pre-lab finished for tomorrow. We start the lab, and Dr. B said we will be working on it for a couple of days. I think the lab is kind of hard to understand, I mean it keeps throwing around words like titration and penolphthalein (I still have no idea how to pronounce that). Soooo, I read a couple of times, and it sunk in. I just hope that Dr. B runs through it a little tomorrow.

allllllllllllllllllllllllllll right, comment time!

4-27-10

hey wassup everybody here's what we did in class today and feel free to comment on anything I left out.

Indicators & pH meters

1. Transition interval

• the pit range over which an indicator changes color
• Indicators that change color at pH lower than 7 are stronger acids than other types of indicators
• Ionizes more than the others
• Indicators that undergo transition in higher pH range are weaker acids

2. pH meter

• determines the pH of a solution by measuring the voltage between the two electrodes that are placed in the solution
• voltage change as water ion concentration in the solution changes
• measures pH more precisely than indicators

3. Titration

• controlled addition and measurement of the amount of a solution of known concentration required to react completely with a measured amount of a solution of unknown concentration

Remember to finish up your pre-lab if you did

have a good night

4/26/10

Ch. 14 contd.

-Any species that can react as either an acid or base is described as amphoteric. A good exmple of this is water

-Water can act as a base
-Water can also act as an acid

-The covalently bonded OH group in an acid is referred to as a hydroxyl group.
-Molecular compounds containing OH groups can be acidic or amphoteric.
-The behavior of oxygen atoms bonded to the atom connected to the OH group.

-In aqueas solutions, neutralization is the reaction of hydronium ions and OH ions to form water molecules.

-A salt is an ionic compound composed of a cation from a base and an anion from an acid.

-NO, NO2, CO2, SO2, and SO3 gasses from industrial processes can dissolve in atmospheric water to produce acidic solutions.

-Very acidic rain is known as acid rain. Acid rain is not as rare as people might think. It is actually fairly common but it is normally not found in large concentrations

END OF CH 14. (Followed by DR. B's famous smile that we have grown to hate when ch's end...)

Ch. 15

-In the self-ionization of water, two water molecules produce a hydronium ion and a hydroxide ion by transfer of a proton.

-Solutions in which H3O=OH, its a neutral solution
-Solutions in which H3O>OH, its acidic
-Solutions in which H3O
-Strong acids and bases are considered completely ionized or dissociated in weak aquous solutions.

-The pH of a solution is defined as the negative of the common logorithm of the Hydrogen ion concentration.



There is a link on the HW website with the pre-lab Q's that are due tomorrow and Dr. B has asked us to bring in our Lab Books to class tomorrow. Have a good night.


-

Was there no Blog for this weekend? If there wasn't, well Friday all we did was watch a movie.
That is all.

Notes for Thursday, April 22, 2010

So, everybody has your tests in hand. Keep in mind that they will be graded out of 33, not 35, and there may be a chance for extra credit.

• An acid reacting with a base will produce a salt (not NaCl, although that is a type of salt) and water.
• The stronger an acid, the weaker its conjugate base.
• The stronger a base, the weaker its conjugate acid.
• Proton transfer reactions favor the production of the weaker acid and the weaker base.
• There are some conjugate acid-base pairs that we will need to know. There's a table, but since I don't have my book on hand, feel free to comment with the page in the book where it can be found.

Well, that was pleasantly brief. Good night everybody!

4/21/10

Today we had a test. It was over chapter 12 and 13. How did everyone do?

Notes for Monday, April 19, 2010

Test Wednesday, everybody! If you can't do molarity problems, molality problems, or problems on colligative properties, then start practicing. Feel free to comment with practice problems.

• A monoprotic acid is an acid that can donate only one proton.
• A polyprotic acid can donate more than one proton per molecule.
• A diprotic acid can donate to protons per molecules, and a triprotic acid can donate three protons per molecule.
• A Lewis acid is an atom, ion, or molecule that accepts one electron pair to form a covalent bon. This is the broadest of the three acid definitions. A bare proton is a Lewis acid. This is the only definition that is not based on hydrogen. A silver ion can be a Lewis acid, for instance.

Here's a table for remembering the different definitions of acid.

-	Acid	Base
Arrhenius	H, H3O producer	OH producer
Bronsted-Lowry	proton donor	proton acceptor
Lewis	electron-pair acceptor	electron-pair donor

• The species that remains after a Bronsted-Lowry acid has given up its proton is the conjugate base of that acidd.
• Bronsted-Lowry acid-base reactions involve 2 acid-base pairs, known as conjugate acid-base pairs.

And no, I don't have a clue why Blogger put a giant gap before the table.

4/16/10

Today, Dr. Bautista was not here. We watched a movie on Mars. It was very interesting and showed the time and dedication of the teams operating the rovers. Very soon we will know more about Mars.

Chemistry Notes for Thursday, 4/15

In class we discussed how Dr. B checked and graded the R constant labs. I however, do not have those notes, and leave it to you guys to fill that part in.
CLASS NOTES:
- Arrhenius Acids are molecular compounds with ionizable hydrogen atoms
-Their water solutions are known as aqueous acids
-All aqueous acids are elctrolytes
- A strong acid ionizes completely in aqueous solutions.
-Strong acids are strong electrolytes
Ex: HClO4 HCl, HNO3

-A weak acid relaes few ions in aqueous solutions
- Hydronium ions, anions, and dissolved acid molecules in aqueous solutions
Ex: HCN
-organic acids (____COOH molecules), such as acetic acid are weak acids

-Most bases are inoic compounds containing metal cations and the OH- anion
-Ammonia, NH3, is molecular
-Ammonia produces OH- ions when it reacs with water molecules
-The strength of a base depends on the extent in which it dissociates in a solution
-Strong bases --> Strong electrolytes
Amount of H3O+'s to OH-'s in types of solutions
Acidic
H3O+ > 10^-7 Moles >OH-
Nuetral
H3O+ = 10^-7 Moles = OH-
Basic
H3O+ < 10^-7 Moles< OH-

Bronsted Lowry Acids and Bases
- A Bronsted Lowry-Acid is a molecule or ion that is a proton donor
-HCl acts as a Bronsted-Lowry acid when it reacts with ammonia
-Water can act as a Bronsted-Lowry Acid
-A Bronsted-Lowry Base is a molecule or ion that Acts as a proton acceptor
-Ammonia accepts a proton from HCl, and is therefore and Bronsted-Lowry Base
- the OH- ion produced in solutions by acids is a Bronsted-Lowry Base

- In a Bronsted-Lowry Acid-Base reaction, protons are transferred from one reactant(the acid) to another (the base)

The blog by chris mathews

today we talked about acids, but not before we went over some very challenging chemistry questions. Good news of the day is that kyle "Half Caff McCaffery" (if i spelled it wrong please correct me) was moved and now gets to sit with us in the front

back to the notes

first acids are aqueous solutions that usually have a sour taste, they can change the color of an acid-base indicator, react with active metals to release H2 gases, react with bases to produce water and salts, and finally conduct electricity.

binary acid- acid that contains only two elements, one of which has to be hydrogen and the other and electronegative element.

examples are HF HCl HBr

the name starts with -hydro for the hydrogen element

then to the root of the second element comes into play

finally all of this is followed by an "ic"

ex. HF is know as hydrofluoric acid notice the suffix and prefix usage

next we have the oxyacid-is an acid that is a compound of oxygen, hydrogen, and usually a nonmetal. these names also follow a pattern

names of anion are based on the names of the acid. next if the anion ends in ate it is an "ic" ending but if the anion ends with ite its ending is "us"

now we have sulfuric acid

apparantly its one of the most commonly produced industrial chemicals in the world, but dont clean the container holding it, or it will corrode!!!

nitric acid, phosphoric acid had no definitions...

hydrochloric acid- concentrated solutions of this acid are commonly called muriatic acid

acetic acid-pure acetic acid is clear, colorless, and pungent smelling liquid that is also known as

glacial acetic acid

i do believe this is the point in class where colby maybry made the brilliant oration on the different levels of acid strength, all of which are related to disassociation.

bases

they taste bitter, change with acid base test, dilute solutions feel slippery, react with acids to produce water and salt, and conduct electricity.

this is all i got for the day. Thank you and have a wonderful evening

ps quiz tomorrow over ch 13 calculations, this should be a good way to bring those averages up guys and remember that the test will be next thursday.

-Electrolytes & Colligative Properties cont'd.

-the actual values of the colligative properties for all strong electrolytes are almost always what would be expected based on the number of particles they produce in solution

-the differences are caused by attractive forces between dissociated ions in aqueuous solution

-according to Debye & Huckel, a cluster of hydrated ions acts as a single ion rather than as individual ions, causing effective total concentration to be less

-Ions of higher charge have lower effective concentrationss than ions with smaller charge

Multiple Choice

1. Acetic acid is a weak electrolyte because it

D. ionizes only slightly in aqueuous solution

2. Which of the following solutions would contain the highest concentration of hydronium ions?

A. 0.10 M HCl

3. Which of the following is the best representation of the precipitation reaction that occurs when aqueuous solutions of sodium carbonate and calcium chloride are mixed?

C. CO3(2-) (aq) + Ca(2-) --> CaCO3(s)

4. Which of the following is not a colligative property?

A. molality

5. Solution A contains 0.1 mol of sucrose, C12H22O11, dissolved in 5oo g of water. Solution B contains 0.1 mol of sodium chloride, NaCl, in 500 g of water. Which of the following is true?

C. Solution A would freeze at a higher temperature than Solution B would.

colligative properties of solutions

colligative properties

• Properties that depend on the concentration of solute particles but not on their identity

the boliing point of a solution differ from those of the pure solvent

nonvolatile substance

• substance that has little tendency to become a gas under existing conditions

nonelectrolyte solutions of the same molality have the same conecntration of partilces.

freezing-point depression

• the freezing-point depressino of a 1 m solution of any nonelectrolyte solute in water is found by experiment to be 1.86 degrees Celsius lower than the freezing point of water.

molal freezing-point constant

• the freezing point depresino of the solvent in a 1-molal solutino of a nonvolatile, noneletrolyte solute

freezing point depression

• the difference between the freezing points of the pure solvent and a solution of a nonelectrolyte in that solvent, and it is directly proportional to the molal concentration of the solution

Ksub f is expressed as degrees Celsius/m

each solvent has its own characteristic molal freezing point constant

boiling point elevation

• the boiling point of a liquid is the temperature at which the vapor pressure of the liquid is equal to the prevailing atmospheric pressure
• a change in the vapor pressure of the liquid will cause a corresponding change in the boiling point

molal boiling point constant

• the boiling point elevation of the solvent in a 1-molal solution of a nonvolatile, nonelectrolyte solute.
• the boiling point elevation of a 1-molal solution of any nonelectrolyte solute in water has been found by experiment to be 0.51 degrees C

boiling point elevation

• the difference between the boiling points of the pure solvent and a noneletrolyte solution of that solvent, and is directly proportional to the molal concentration of the solution
• change in temperature = (constant of temperature) (m)

osmotic pressure

• the external pressure that must be applied to stop osmosis

osmosis

• the movement of solvent through a semipremeable membrane from the side of lower solute concentration to the side of higher concentration

semipermeable membrane

• membrane that allows the passage of some particles while blocking the passage of others

Weekend Blog

Strong electrolyte - compound in an aqueous solution that conduct electricity well/many ions/strong acids contain strong electrolytes

Weak electrolyte - poor conductors of electricity/ very few dissolved ions/ no ionic compounds

Colligative properties - depend on concentration of solution

Vapor-Pressure Lowering

Freezing-Point Depression

Boiling-Point Elevation

Osmotic Pressure

Nonvolatile substance- substance with little tendency to become gas

In nonvolatile substances Boiling-Point rises Freezing-Point lowers

the blog

nobody blogged so i decided to do it anyways.

today we took the pop quiz that everyone failed including top ten chem student in shelby county michael scully (congrats on the award we know the quiz was a fluke). also the answers to all the homework problems were put on the board and everyone copied them down. if you have any issues use the email and be ready for any more disastrous quizzes in the future.

4.7.10.

Sorry the blog took so long to get up, the rain and everything messed with my computer and it kept crashing on me. But here you go.

The seperation of ions that occurs when an ionic compound dissolves is dissociation.
see book for example equations.
review book pg 436 example problem A.

Although no ionic compound is completely insoluble, compounds of very low solubility can be considered insoluble for most practical purposes.

GENERAL SOLUBILITY GUIDELINES -
1. Sodium, potassium, and ammonium compounds are soluble in water.
2. Nitrates, acetates, and chlorates are soluble.
3. Most chlorides are soluble, except those of silver, mercury (I), and lead. Lead(II) chloride is soluble in hot water.
4. Most sulfates are soluble, except those of barium, strontium, lead, calcium, and mercury.
5. Most carbonates, phosphates, and silicates are insoluble, except those of sodium, potassium, and ammonium.
6. Most sulfides are insoluble, except those of calcium, strontium, sodium, potassium, and ammonium.

NOTE:
To decide whether a precipitate can form, you must know the solubilities of those two compounds.

Net Ionic Equation -
includes only those compounds and ions that undergo a chemical change in a reaction in an aqueous solution.
see book pg 439 for example equation.

Spectator Ions -
ions that do not take part in a chemical reaction and are found in solution both before and after the reaction.
see book pg 439 for example equation.

review example problem B, pg 440.

ionization -
ions are formed from solute molecules by the action of the solvent in a process called ionization.

When a molecular compound dissolves and ionizes in a polar solvent, ions are formed where none existed in the undissolved compound.

Hydrogen chlorie, HCl, is a molecular compound that ionizes in aqueous solutions. It contains a highly polar bond. The attraction between a polar HCl molecule and the polar water molecules is strong enough to break the HCl bond, forming hydrogen ions and chloride ions.
see book pg 441 for example equation.

The Hydronium Ion -
some compounds ionize in an aqueous solution to release H+. The H+ ion attracts other molecules or ions so strongly that it does not normally exist alone.
see book for example equation pg 441.

The H3O+ ion is known as the hydronium ion.

Substances that yield ions and conduct an electric current in solution are electrolytes.
Substances that do not yield ions and do not conduct an electric current in solutions are non-electrolytes.

Strong Electrolyte-
is any compound whose dilute aqueous soltions conduct electricity well; this is due to the presence of all or almost all of the dissolved compound in the form of ions.

Weak Electrolyte -
is any compound whose dilute aqueous solutions conduct electricity poorly; this is due to the presence of a small amount of the dissolved compound in the form of ions.

STRONG AND WEAK ELECTROLYTES DIFFER IN **THE DEGREE OF IONIZATION OR DISSOCIATION.**

Have a good night guys.
Colby & Co.

Tuesday, April 6

Today in Dr. B's class we handed in our homework whch was #7-15 (additional problems). Dr. B then handed back the test on chapter 10 & 11, the test is out of thirty (27 on the scantron, 3 written answers). After we looked over the test Dr.B assigned the rest of the additional problems for homework. Thats pretty much it.

Blog?

Was there no blog for tonight, because I need posts and we're running shorter and shorter on time..

Friday, March 26th

Molarity and Molality

(M) Molarity- number of mols of a solute dissolved in one liter solution

(m) Molality- number of mols of a solute dissolved in one kilogram solvent
One Molal Solution- one mol solute dissolved in one kilogram solvent





also the lab will be collected on Monday

March 25th, 2010

Here are the R values for everyone--

• Justin and Griff--0.07563
• Joe and Ben-- 0.07718
• Taylor and Nick-- 0.0727
• Henry and Evan-- 0.06379
• Matt Orians and Kyle-- 0.06697
• Charlie and Chris-- 0.0814
• Ryan and Jim-- 0.07667
• Donnie and Colby-- 0.07833
• Alex and Jacob-- 0.06356
• Will Long and Matt Johnson-- 0.07997
• Riley and Patrick Blose-- 0.07383
• Matt Farrel and Andrew Parmenter-- 0.06745
• Patrick "The Sweenenator" Sweeney and Peter-- 0.07885

Notes:

• The effect of temperature on the solubility of solids in liquids is more difficult to predict
• increasing the temperature increases the solubility of solids
• equivalent temperature increase can result in a large increase in solubility for some solvents and only a slight change for others
• in some cases, solubility of a solid decreases with an increase in temperature
• the formation of a solution is accompanied by an energy change.
• if you dissolve some KI in water you will find that the outside of the container feels cold to the touch
• if you dissolve some NaOH in water, the outside of the container feels hot
• the formation of a solid-liquid solution can apparently either absorb or release energy as heat
• before dissolving begins, solvent molecules are held together by intermolecular forces.
• in the solute, molecules are held together by intermolecular forces
• energy is required to separate solute molecules and solvent molecules from their neighbors
• a solute particle that is surrounded by solvent molecules is said to solvated
• the net amount of energy absorbed as eat by the solution when a specific amount of solute dissolves in a solvent is the enthalpy of solution
• the enthalpy of solution is negative (energy released) when the sum of attractions from Steps 1 and 2 is less than Step 3
• the enthalpy of solution is positive (energy absorbed) when the sum of attractions from steps 1 and 2 is less than Step 3
• the concentration of a solution is a measure of the amount of solute in a given amount of solvent or solution
• "dilute" just means that there is a relatively small amount of solute in a solvent
• note that these terms are unrelated to the degree to which a solution is saturated. a saturated solution of a substance that is not very soluble might be very dilute
• Molarity is the number of moles per solute in one liter of a solution
• to relate the molarity of a solution to the mass of solute present, you must know the molar mass of the solute. for example, a "one molar" solution of NaOH contains one mole of NaOH in every liter of solution
• the symbol for molarity is M, and the concentration of one molar solution of sodium hydroxide is written as 1 M NaOH
• one mole of NaOH has a mass of 40.0 g. If this quantity of NaOH is dissolved in enough water to make exactly 1.00 L of solution, the solution is a 1 M solution

molarity (M) = amount of solute (mol)/volume of solution (L)

• note that a 1 M solution is not made by adding 1 mol of solute to 1 L of solvent. in such a case, the final total volume of the solution might not be 1 L.
• the resulting solution is carefully diluted with more solvent to bring the total volume to 1 L
• see pages 420 and 421 for the answers and explanations of Sample Problems A, B, and C
• the lab is due tomorrow. dont forget to find the class avg and % error of the class average, as well as writing your conclusion

Notes from Wednesday, March 24, 2010

So, we started the day off with some demonstrations that taught us a few things. For instance, the volume of a mixture of two liquids is not necessarily the sum of the volume of its components, supersaturated solutions form crystals when disturbed by the addition of more solute, and water and methanol can be made immiscible by adding a salt (in our case, KCO3). Why? Comment with your theories.

Here's the other things from today.

• Pressure increases the solubility of gas in liquid. Increased pressure causes more gas particles to dissolve in the liquid, and decreased pressure allows more dissolved gas to escape from the liquid.
• Henry's law: The solubility of a gas in a liquid is directly proportional to the partial pressure of that gas on the surface of the liquid.
• The rapid escape of gas from a liquid in which it is dissolved is known as effervesence.
• Unlike a solid, increasing the temperature usually decreases gas solubility.

Since we have a short blog today, I leave you with another video, about the awesomeness of effervesence. Have a good night!

Solubility

Solubility
-every combination of solute and solvent has a limit to the amount of solute that can be dissolved
-depends on nature of solute, solvent, and temperature
-solute molecules leave the solid surface and move around randomly
-when max solubility is reached, the amount of solid dissoluting is the same as the amount crystallizing
-solubility equilibrium- dissolution and crystallization of a solute occur at the same rate
Saturated vs. Unsaturated
-saturated-contains the maximum amount of dissolved solute
-unsaturated-contains any amount less than the maximum amount of dissolved solute
-supersaturated-upon cooling, the solute does not settle out of the solution, creating a solution with more than the maximum amount of solute dissolved at certain conditions
-"Like Dissolves Like"
-dissolution happens between two like substances
-type of bondin
-polarity of molecules
-intermolecular forces
-solutions with water involved as the solvent are hydrated
-liquids that do not dissolve in each other are immiscible
-liquids that dissolve freely in each other are miscible

Chapter 12 Section 2

Section 2 The Solution Process

Factors Affecting the Rate of Dissolution

• Because the dissolution process occurs at the surface of the solute, it can be sped up if the surface area of the solute is increased


• Stirring or shaking helps to disperse solute particles and increase contact between the solute surface. This speeds up the dissolving process


• At higher temperatures, collisions between solvent molecules and solvent are more frequent and are of higher energy than at lower temperature. This helps to disperse solvent molecules among the solvent molecules, and speeds up the dissolving process

Solubility

• If you add spoonful after spoonful of sugar to tea, eventually no more sugar will dissolve
• This illustrates the fact that for every combination of solvent with a solid solute at a given temperature there is a limit to the amount of solid that can be dissolved
• The point at which this limit is reached for any solvent-solute combination is difficult to predict precisely and depends on the nature of the solute, the nature of the solvent and the temperature.
• When a solute is first added to a solvent-solute molecules leave the solid surface and more about at random in the solvent
• As more solute is added, more collisions occur between dissolved solute particles. Some of the solute molecules return to the crystal
• When maximum solubility is reached molecules are returning to the solid form at the same rate at which they are going into the solution
• Solution equilibrium is the physical state in which the opposing processes of dissolution and crystallization of a solute occur at which the same rates

Saturated vs. Unsaturated Solutions

• A solution that contains the maximum amount of dissolved solute is described as a saturated solution
• If more solute is added to a saturated solution it falls to the bottom of that container and doesn't dissolve
• A solution that contains less solute than a saturated solution under the existing conditions is called unsaturated solution
• When a saturated solution is cooled the excess solute usually comes out of solution, leaving the solution at the lower temperature
• But sometimes the excess solute doesn't separate a supersaturated solution is produced which is a solution that contains more dissolved solution that contains a more dissolved solute than a saturated solution under the same conditions
• The solubility of a sub solution is the amount of that sub solution required to form a saturated solution with a specific amount of solvent at a specific temperature
• Solubility's vary widely and must be determined experimentally
• They can be found in chemical hand/books and are usually given as gimmes of solute per 100g of solvent at a given temperature

Solute-Solvent Interactions

• Solubility varies greatly with the type of compound involved
• "Like" dissolves like it is a rough but useful rule for predicting whether one substance will dissolve in another
• What makes substances similar depends on type of bonding, polarity and non polarity of molecules, and intermolecular forces between solute and solvent

Dissolving Ionic Compounds in Aqueous Solutions

• The polarity of water molecules plays an important role in the formation of solutions of ionic compounds in water
• The slightly charged parts of water molecules attract ions in the ionic compounds and surround them to keep them separated from other ions in the solution
• This solution process with water as the solvent is referred to as hydration

Non polar Solvents

• Ionic compounds are generally not soluble in non polar solvents such as carbon tetrachloride and toluene
• The non polar solvent molecules do not attract the ions of the crystal strongly enough to overcome the forces holding the crystal together
• Ionic and non polar substances differ widely in bonding, polarity, and intermolecular forces

Liquid Solutes and Solvents

• Liquids that are not soluble in each other are immiscible
• Liquids that dissolve freely in one another in any proportion are said to be miscible

Test Study Material

Guys, these videos are awesome. I couldn't help but to share them:

http://khanacademy.org/

for these chapter, click on the ones about ideal gases and states of matter. I'm sure the next chapters will have videos, too.

KEEP BLOGGIN'!

acid

today in chemistry, we did the lab where we found the partial pressure of the gas created by the magnesium and and acid. Don't forget to finish your labs over the next couple of days and STUDY FOR THE TEST ON THURSDAY.

Chapter 12 Solutions

Section 1 Types of Mixtures

Solutions

• You know from experience that sugar dissolves in water. Sugar is described as "soluble in water". By soluble we mean capable of being dissolved.
• When sugar dissolves, all its molecules become uniformly disturbed among the water molecules. The solid sugar is no longer visible
• Such a mixture is called a solution. A solution is a homogeneous mixture of two or more substances in a single phase.
• The dissolving medium in a solution is called the solvent and the substances dissolved in a solution is called the solute
• Solutions may exist as gases, liquids, or solids. Therefore many possible solute-solvent combinations between gases, liquids, or solids.
• example: Alloys are solid solutions in which the atoms of two or more metals are uniformly mixed
• Solution = Solvent + Solute

Suspensions

• If the particles in a solvent are so large that they settle out unless the mixture is constantly stirred or agitated, the mixture is called a suspension.
• example: A jar of muddy water consists of soil particles suspended in water. The soil particles will eventually all collect or the bottom of he jar because the soil particles are denser than the solvent
• Particles over 1000mm in diameter are 1000 times as large as atoms molecules, or ions suspensions

Colloids

• Particles that are intermediate in sizes between those in solutions and suspensions from mixtures known are as colloidal dispersions or simply colloids
• The particles in a colloid are small enough to be suspended throughout the solvent by the constant movement of the surrounding molecules
• Colloidal particles make up the dispersed phase, and water is the dispersing medium

Tyndall Effect

• Many colloids look similar to solutions because their particles can not be seen
• The Tyndall Effect is a property that can be used to distinguish between a solution and a colloid

Solutes: Electrolytes Versus Nonelectrolytes

• A substance that dissolves in water to give a solution that conducts electric currents is called an electrolyte
• Any soluble ionic compound such as sodium chloride, is an electrolyte
• The positive and negative ions separate from each other in solution and are free to move, making it possible for an electric current to pass through the solution
• A substance that dissolves in water to give a solution that doesn't conduct electric current is called a nonelectrolyte
• Sugar is an example of a nonelectrolyte
• Neutral solute molecules do not contain mobile charged particles, so a solution of a nonelectrolyte cannot conduct electric currents

Movie Day, No Test

Today, our class was scheduled to take the ch. 10&11 test. Instead Dr. Bautista put in a movie , which i am getting ready to summarize, because we were very good for her. The test was rescheduled for after spring break, though, along with the gas lab. Here are some of the key points stated in the movie:

- In the beginning of the film, a group of scientists discussed some of the volcano eruptions and history of the eruptions at Montserrat, an island in the Lesser Antilles.

- The dangerous part of this volcano is its deadly flow of pyroclast. Pyroclast will burn, crush, and suffocate anything in its way.

- The ash from this volcano prevents any outdoor activity for civilians

- 18000 meteorites hit the earth every year

- Most of this is space junk made from molten rock of planet earth.

- Many pacific islands, like Hawaii, are made from these earth shaping meteorites and volcano lava.

- Studies show that under the plates we live on, there are remains of ancient volcanic rock.

- In this volcanic area, there lives more types of bacteria then people living on earth

- Sedimentary Rock is made from rock being in water, and age of earth.

- Stromatolides are rocks found especially in regions of Australia, that are around 3,000 million years old.

- Volcanic lava will destroy all vegetation in its way

- Africa, Asia, and South America once were bonded togethor as one big supercontinent. Similar rock structures from each of these continents proves this theory.

- At the very bottom of the volcano type ocean floor there are some of the most unique life forms in the entire earth.

- Iceland is the only place where the midatlantic ridge is evident above water.

If I missed anything then comment it

3/1/10

TEST TOMORROW!!

The Chapter 10 and 11 test is tomorrow! Be sure to study, especially calculations! Also tomorrow, an outline of chapter 11 is due. Today we worked on our outlines during class, so here's all the blog notes from this chapter compiled into one to help with your outline:

Please comment with study tips or online quizzes etc.

ch 11

the pressure formula is p=f/a such that p-pressure, f=force, and a is area

always remember that area is squared

the SI unit for area is N which means newton. it will increase the speed of one kilogram mass by one meter per second that that force is applied.

pressure is a force per unit area, therefore pressure of a 500 N on a floor with an Area of 325 cm^2 is:

500 n / 325 cm^2 = 1.5 N/Cm^2

*the greater the force -> greater pressure

*smaller the area-> greater pressure KNOW THOSE TWO THINGS

introduced by Evangelista Toricelli who was constantly picked on by his parents, the ultimate teaser being his girlish name....

water pumps can raise about 34 feet

thought that it must be dependent on weight and weight of air

reasoned that since mercury was 14 times less dense than water, it would be 1/14 of 34 feet

tested it and it in rose 30 in.

pressures can be also measured in units of atmospheres. Because the average pressure is 760 mm of Hg or one torr named after evangelista.

In pascals pressure is exerted by one N on one square meter

Daltons law

the measure of full pressure in a gas is the sum of the measure of partial pressures

- To determine the pressure of a gas inside a collection bottle, you would use this equation, which is an instance of Dalton's Law of Partial Pressure:

- P(atmosphere) = P(gas) + P(water)

- If you raise the bottle until the water levels inside and outside the bottle are the same, the total pressure outside and inside will be the same.

- Reading the atmospheric pressure on a barometer and looking up the value of P(water) at the temperature of the experiment in a table (p. 859 in our book), you can calculate the P(gas).

Sample Problem B

- Oxygen gas from the decomposition of potassium chlorate, KClO3, was collected by water displacement. The barometric pressure and the temperature during the experiment were 731.0 torr and 20.0 degrees C respectively. What was the partial pressure of the oxygen collected?

G: Total Pressure = P(atmosphere) = 731.0 torr

P(water) = 17.5 torr (vapor pressure of water at 20.0 C from table A)

P(atmosphere) = P(oxygen) + P(water) ; solve for P(oxygen)

P(oxygen) = P(atm) - P(water)

substitute: P(oxygen) = 731.0 torr - 17.5 torr = 713.5 torr

- Boyle's Law

- Robert Boyle discovered that doubling the pressure on a sample of gas at constant temperature reduced its volume by 1/2

- explained by the Kinetic-Molecular Theory (Dr. B said to make sure and know this!)

- the pressure of a gas is caused by moving molecules hitting the container walls

- if the volume of the container is decreased, more collisions will occur and the pressure will increase

- if the volume of the container is increased, less collisions will occur and the pressure will decrease

- Boyle's Law states that the volume of a fixed mass of gas varies inversely with the pressure at constant temperature.

- Formula: PV=K (P=pressure, V=volume, K= constant)

-the inverse would be a straight line (V=K/P)

- Because of the transitive property of equality, since two different quantities are equal to the same thing (volume x pressure = K), it can be concluded that two separate sets of conditions are equal to each other (P1V1 = P2V2)

Charles Law (cont)

Charles Law: Volume-Temperature between volume and temperature was discovered by the French scientist Jacques Charles in 1787.

Charles found that the volume changes by 1/273 of the original volume for each Celsius degree, at a constant pressure and at an initial temperature of 0 degrees C.

The temperature 273 is absolute zero and is given a value of zero in the Kelvin temperature scale. The relationship between the two temperature scales is K=273.15 +degrees C.

Charles Law states that the volume of a fixed mass of gas at a constant pressure varies directly with the Kelvin temperature.

Gas volume and Kelvin temperature are directly proportional to each other at constant pressure.

Mathematically, Charles Law can be expressed as: V=KT or V/T=K where V is the volume, T is the Kelvin temperature, and K is a constant. the ratio V/T for any set of volume temperature values always equals the same K.

The equation reflects the fact that volume and temperature are directly proportional to each other at constant pressure.

The form of Charles Law that can be applied directly to most volume- temperature gas problems is: V1/T1 = V2/T2.

V1 and T1 represent initial conditions, and V2 and T2 represent another set of conditions.

Given three of the four values, V1, T1, V2, and T2, you can use this equation to calculate the 4th value for a system at constant pressure.

Gay-Lussacs Law: Pressure Temperature Relationship

At a constant volume, the pressure of a gas increases with increasing temperature.

Gas pressure is the result of collisions of molecules with container walls.

The energyu and frequency of collisions depend on the average kinetic energy of the molecues.

Pressure is directly proportional to Kelvin temperature.

Gay Lussacs Law: The pressure of a fixed mass of gas at constant volume varies directly with the Kelvin temperature.

This law is named after Joseph Gay-Lussac, who discovered it in 1802.

Mathematically, Gay Lussacs Law can be expressed as P=KT or P/T=K where P is pressure, T is the Kelvin temperaure, and K is a constant. The ratio P/T for any set of volume-temperature values always equals the same K.

• Boyle's Law, Charles' Law, and Gay-Lussac's Law can be combined into a single equation that can be used for situations in which temperature, pressure, and volume all vary at the same time.
• This is the combined gas law, PV/T=k, or P1V1/T1=P2V2/T2.
• Each gas law can be derived from the combined gas law when the proper variable is kept constant.
• Sample Problem F can be found in your book.
• In the early 1800s, French chemist Joseph Gay-Lussac observed that 2L hydrogen can react with 1 L oxygen to form 2L water vapor.
• This reaction shows a simple 2:1:2 ratio in the volumes of reactants and products. This same ratio applies to any volume proportions.
• Gay-Lussas's law of combining volumes of gases (that's a mouthful) states that at constant temperature and pressure, the volumes of gaseous reactants and products can be expressed as ratios of small whole numbers.
• 1811: Avogadro explained Gay-Lussac's law of combining volumes of gases without violating Dalton's idea of indivisible atoms.
• Avogadro reasoned that, instead of always being in monoatomic form, when they combine to form products, gas molecules can contain more than one atom.
• Avogadro's law: equal volumes of different gases contain the same number of molecules, at given pressure and temperature. Also, gas volume is directly proportional to the amount of gas at a given temperature or pressure. V=kn.
• Dalton had guessed that the formula for water was HO, but Avogadro's reasoning established that water must contain twice as many hydrogen atoms as oxygen atoms because of the volume ratios in which the gases combine.
• Ergo, Avogadro's idea of diatomic gases was consistent with all other knowledge and laws.

• You can use the volume ratios as conversion factors in gas stoichiometry problems as you would mole ratios.
• Ideal Gas Law
• You have learned about equations describing the relationships between 2 or 3 of the 4 variables - Pressure, Volume, Temperature and number of moles - needed to describe a sample at a time.
• All of the laws you have learned thus far can be combined into a single equation, the IDEAL GAS LAW: the mathematical relationship among pressure, volume, temperature, and number of moles of a gas.
• R is a constant
• PV=nRT
• In the equation representing the Ideal gas law, R = idea gas constant
• Its value depends on the units chosen for pressure, volume, and temperature in the rest of the equation.
• Measured values of P, V, Temp., and n for a gas at near-ideal conditions can be used to calculate R
• R = 0.082058

At the begining of the lab Dr. B will give us the barometric pressure, but not the units we need, we will have to convert it as part of the lab.
During the lab there will be no chewing on any substances. HCl will be used during the lab and it will burn your skin.
......................................................Missed some info...................................................
Add HCl the the tube then hold the tube with HCl at an angle when squeezing the water slowly into the tube, try not to let the water and HCl mix because it will cause the reaction to go slower. When putting the capper and magnesium in the tube put it in close to the top and not far in the tube. Then you put your finger over the top of the tube and flip it over and into the beaker with water. It should start to bubble and gas will start to form at the what used to be the bottom of the tube, now the top.
NOTES - CH11 sec 3 cont.
The Ideal Gas Law cont.
The Ideal Gas Constant cont.

• The calculated value of R is usually rounded to 0.0821(L x atm)(md x k)
• Dr. B wants Rto equal 0.08206
• use this value in ideal gas law calculations when the volume is in liters, the preasure is in atmospheres, and the temp is in kelvins
• The ideal gas law canbe applied to determine the existing conditions of a gas sample when three of the four values; P,T,V, and n; are known
• be sure to match the units of the known quantities and the units of R
• Sample problem 1: what is the pressure in atmospheres exerted by a 0.500mol smple of nitrogen gas in a 10.0 L container at 298K?........P=nRT/V..........P=(0.500mol)(0.08206L x atm)(298K)/10.0L 122atm

Graham's Law of Efusion

• Rates of effusion and diffusion depend on the relative velocities of gas molecules. the velocity of a gas varies inversely with the square root of its molar mass.
• recall that the average kinetic energy of the molecules in any gas depends on the temperature and equalys(1/2)mv^2
• for 2 different gases, A and B, at the same temperature, the following relationship is true 1/2 MaVa^2= 1/2 MbVb^2
• from the equation relationg the kinetic energy of two different gases at the same conditions, one can derive an equation relating the rates of effusion of two gases with their molecular mass. Rate of effusion of A/rate of effusion of B= square root of Mb/ square root of Ma
• this equation is known as Graham's law of effusion which states that the rates of effusion of gasses at the same temperature and pressure are inversely proportional to the square roots of their molar mass

Friday, 2/26/10

In class we went over the answers to the homework #12 which are:

1. a. 105kPa b. 5.0 mL c. 42.4kPa d. 6.78 x 10^-3 dm^3 e. 1.24 atm f. 1.5 m^3
2. 8.0 m^3
3. 2.5 x 10^-2 atm
4. 8.01 x 10^-2 dm^3
5. a. 234 K b. 1.2 dm^3 c. -269.17 C d. 8.10 x 10^-2 L e. 487 cm^3 f. 68.2 m^3
6. 1.45 cm^3
7. -40.2 C

This weekend's homework is the worksheet she handed out.

2-25-10

The one question we did today in class:

(rate of effusion H)/(rate of effusion O)= Squareroot(O's molar mass)/Squareroot(H's molar mass)

AKA: H has a higher rate of effusion, because O is bigger.

Wednesday of Next week is when the test will be

The Homework is Pg. 390: 7-11, 18-26

GETCHU SUM BLOG POINTS

where is the blog??

Is anyone else wondering where the blog might be, because i thought Dr.B said she would have something we would comment on up in seventh period?

Pre-lab stuff + parts of Sec 3 and 4 notes

At the begining of the lab Dr. B will give us the barometric pressure, but not the units we need, we will have to convert it as part of the lab.
During the lab there will be no chewing on any substances. HCl will be used during the lab and it will burn your skin.
......................................................Missed some info...................................................
Add HCl the the tube then hold the tube with HCl at an angle when squeezing the water slowly into the tube, try not to let the water and HCl mix because it will cause the reaction to go slower. When putting the capper and magnesium in the tube put it in close to the top and not far in the tube. Then you put your finger over the top of the tube and flip it over and into the beaker with water. It should start to bubble and gas will start to form at the what used to be the bottom of the tube, now the top.
NOTES - CH11 sec 3 cont.
The Ideal Gas Law cont.
The Ideal Gas Constant cont.

• The calculated value of R is usually rounded to 0.0821(L x atm)(md x k)
• Dr. B wants Rto equal 0.08206
• use this value in ideal gas law calculations when the volume is in liters, the preasure is in atmospheres, and the temp is in kelvins
• The ideal gas law canbe applied to determine the existing conditions of a gas sample when three of the four values; P,T,V, and n; are known
• be sure to match the units of the known quantities and the units of R
• Sample problem 1: what is the pressure in atmospheres exerted by a 0.500mol smple of nitrogen gas in a 10.0 L container at 298K?........P=nRT/V..........P=(0.500mol)(0.08206L x atm)(298K)/10.0L 122atm

Graham's Law of Efusion

• Rates of effusion and diffusion depend on the relative velocities of gas molecules. the velocity of a gas varies inversely with the square root of its molar mass.
• recall that the average kinetic energy of the molecules in any gas depends on the temperature and equalys(1/2)mv^2
• for 2 different gases, A and B, at the same temperature, the following relationship is true 1/2 MaVa^2= 1/2 MbVb^2
• from the equation relationg the kinetic energy of two different gases at the same conditions, one can derive an equation relating the rates of effusion of two gases with their molecular mass. Rate of effusion of A/rate of effusion of B= square root of Mb/ square root of Ma
• this equation is known as Graham's law of effusion which states that the rates of effusion of gasses at the same temperature and pressure are inversely proportional to the square roots of their molar mass

Boom

• You can use the volume ratios as conversion factors in gas stoichiometry problems as you would mole ratios.
• Ideal Gas Law
• You have learned about equations describing the relationships between 2 or 3 of the 4 variables - Pressure, Volume, Temperature and number of moles - needed to describe a sample at a time.
• All of the laws you have learned thus far can be combined into a single equation, the IDEAL GAS LAW: the mathematical relationship among pressure, volume, temperature, and number of moles of a gas.
• R is a constant
• PV=nRT
• In the equation representing the Ideal gas law, R = idea gas constant
• Its value depends on the units chosen for pressure, volume, and temperature in the rest of the equation.
• Measured values of P, V, Temp., and n for a gas at near-ideal conditions can be used to calculate R
• R = 0.082058
• P.S. I left the floor open for comments containing some of the notes, I purposely left out some bullet points.

Notes from Friday, Feb. 19

Homework for this weekend: Read ahead in the chapter. Print out the ideal Gas Law Constant Lab and do the pre-lab for Monday.

• Boyle's Law, Charles' Law, and Gay-Lussac's Law can be combined into a single equation that can be used for situations in which temperature, pressure, and volume all vary at the same time.
• This is the combined gas law, PV/T=k, or P1V1/T1=P2V2/T2.
• Each gas law can be derived from the combined gas law when the proper variable is kept constant.
• Sample Problem F can be found in your book.
• In the early 1800s, French chemist Joseph Gay-Lussac observed that 2L hydrogen can react with 1 L oxygen to form 2L water vapor.
• This reaction shows a simple 2:1:2 ratio in the volumes of reactants and products. This same ratio applies to any volume proportions.
• Gay-Lussas's law of combining volumes of gases (that's a mouthful) states that at constant temperature and pressure, the volumes of gaseous reactants and products can be expressed as ratios of small whole numbers.
• 1811: Avogadro explained Gay-Lussac's law of combining volumes of gases without violating Dalton's idea of indivisible atoms.
• Avogadro reasoned that, instead of always being in monoatomic form, when they combine to form products, gas molecules can contain more than one atom.
• Avogadro's law: equal volumes of different gases contain the same number of molecules, at given pressure and temperature. Also, gas volume is directly proportional to the amount of gas at a given temperature or pressure. V=kn.
• Dalton had guessed that the formula for water was HO, but Avogadro's reasoning established that water must contain twice as many hydrogen atoms as oxygen atoms because of the volume ratios in which the gases combine.
• Ergo, Avogadro's idea of diatomic gases was consistent with all other knowledge and laws.
  

2/18.10

Today in class we went over Tuesday nights homework and took notes. Also, there is a quiz tomorrow on Chapter 10 the section is unknown so know all of them. For those that qualify: The make up test will be on the 24th after school.

Tuesday night's homework answers.

19. a. 2264 J/g

b. 333.5 J/g

20. 84.3 J/g

21. a. 0.277 mol

b. 15.5 kJ/mol

22. a. 21.8 mol

b. 15.5 kJ/mol

NOTES

Charles Law (cont)

Charles Law: Volume-Temperature between volume and temperature was discovered by the French scientist Jacques Charles in 1787.

Charles found that the volume changes by 1/273 of the original volume for each Celsius degree, at a constant pressure and at an initial temperature of 0 degrees C.

The temperature 273 is absolute zero and is given a value of zero in the Kelvin temperature scale. The relationship between the two temperature scales is K=273.15 +degrees C.

Charles Law states that the volume of a fixed mass of gas at a constant pressure varies directly with the Kelvin temperature.

Gas volume and Kelvin temperature are directly proportional to each other at constant pressure.

Mathematically, Charles Law can be expressed as: V=KT or V/T=K where V is the volume, T is the Kelvin temperature, and K is a constant. the ratio V/T for any set of volume temperature values always equals the same K.

The equation reflects the fact that volume and temperature are directly proportional to each other at constant pressure.

The form of Charles Law that can be applied directly to most volume- temperature gas problems is: V1/T1 = V2/T2.

V1 and T1 represent initial conditions, and V2 and T2 represent another set of conditions.

Given three of the four values, V1, T1, V2, and T2, you can use this equation to calculate the 4th value for a system at constant pressure.

Gay-Lussacs Law: Pressure Temperature Relationship

At a constant volume, the pressure of a gas increases with increasing temperature.

Gas pressure is the result of collisions of molecules with container walls.

The energyu and frequency of collisions depend on the average kinetic energy of the molecues.

Pressure is directly proportional to Kelvin temperature.

Gay Lussacs Law: The pressure of a fixed mass of gas at constant volume varies directly with the Kelvin temperature.

This law is named after Joseph Gay-Lussac, who discovered it in 1802.

Mathematically, Gay Lussacs Law can be expressed as P=KT or P/T=K where P is pressure, T is the Kelvin temperaure, and K is a constant. The ratio P/T for any set of volume-temperature values always equals the same K.

2-17-10 THE 100TH BLOG OF THE YEAR!!!

This blog is a momentous occasion, for it is the 100th blog of our sophomore year of honor's chemistry. Instead of holding a fiesta at my house to celebrate this incredible event, I've gathered an extensive list of truly fascinating facts about the number 100:

• On average, 100 people choke to death on ball-point pens every year (how sad).
• Ben Franklin is on the face of the $100 bill.
• In Greece, India and Israel, 100 is the police telephone number.
• Nicodemus brought 100 pounds of myrrh & aloes to embalm Jesus after his crucifixion.

and here's the 100th blog.......

- To determine the pressure of a gas inside a collection bottle, you would use this equation, which is an instance of Dalton's Law of Partial Pressure:

- P(atmosphere) = P(gas) + P(water)

- If you raise the bottle until the water levels inside and outside the bottle are the same, the total pressure outside and inside will be the same.

- Reading the atmospheric pressure on a barometer and looking up the value of P(water) at the temperature of the experiment in a table (p. 859 in our book), you can calculate the P(gas).

Sample Problem B

- Oxygen gas from the decomposition of potassium chlorate, KClO3, was collected by water displacement. The barometric pressure and the temperature during the experiment were 731.0 torr and 20.0 degrees C respectively. What was the partial pressure of the oxygen collected?

G: Total Pressure = P(atmosphere) = 731.0 torr

P(water) = 17.5 torr (vapor pressure of water at 20.0 C from table A)

P(atmosphere) = P(oxygen) + P(water) ; solve for P(oxygen)

P(oxygen) = P(atm) - P(water)

substitute: P(oxygen) = 731.0 torr - 17.5 torr = 713.5 torr

- Boyle's Law

- Robert Boyle discovered that doubling the pressure on a sample of gas at constant temperature reduced its volume by 1/2

- explained by the Kinetic-Molecular Theory (Dr. B said to make sure and know this!)

- the pressure of a gas is caused by moving molecules hitting the container walls

- if the volume of the container is decreased, more collisions will occur and the pressure will increase

- if the volume of the container is increased, less collisions will occur and the pressure will decrease

- Boyle's Law states that the volume of a fixed mass of gas varies inversely with the pressure at constant temperature.

- Formula: PV=K (P=pressure, V=volume, K= constant)

-the inverse would be a straight line (V=K/P)

- Because of the transitive property of equality, since two different quantities are equal to the same thing (volume x pressure = K), it can be concluded that two separate sets of conditions are equal to each other (P1V1 = P2V2)

That's the blog, have a wonderful night.