Thursday, December 9, 2010

Today we started class off with reviewing Tammy's blog. If you were absent, her blog would be a good review of what we did that day (packet pages 1-3). In her blog, she incorporated the equation: 1 atm=14.7 BSI=760 mmHg=760 torr.

Next, we reviewed the equation: P*V=n*R*T which would used for today's entire lesson. With this equation, the P stands for pressure, V stands for volume, n stands for the number of moles, R stands for the gas law content, and the T stands for temperature.

From that formula, there are other equations that are formed when P, V, or T is constant. The first equation is: V=K*T (assuming that the pressure is constant).This equation is direct which means that as the volume of a sample of the gas increases, the Kelvin temperature of the gas increases also. An example of this would be (#3a, packet page 5): As the Kelvin temperature is increased by a factor of 2, the volume of the gas sample is increased by a factor of 2. When a sample of gas with volume V1 and temperature T1 are changed to a new volume V2 or a new temperature T2, you form a 2-point equation. The theory relating volume and temperature for a two-point equation is called Charle's Law. This equation would be: V1/T1=V2/T2. On packet page 6, number 6 is an example of how to apply the Another equation that can be formed using Charle's Law. For number 6a, you would take the initial temp and the inital volume and plug that into v1 and t1 (5L/300K). Next, you would plug in 600K for t2 and "X" for v2. After solving the proportion (multiplying both sides by 600), the answer would become 10L for the final volume.

Another equation derived off of the forumla PV=nRT is: PV=K (assuming that temperature is constant). This equation is inversely proportional meaning that when the pressure of a sample of gas is increased, the volume of the sample decreases. An example of this would be (#10): As the pressure is increased by a factor of 2, the volume of the gas sample is decreased by a factor of two (or 1/2 of the original value). If a sample of gas with a volume of V1 and a pressure of P1 is changed to a new volume V2 and a new pressure P2, a two-point equation is formed. This theory is called Boyle's Law. The Boyle's Law equation is: P1V1=P2V2. An example of Boyle's law is number 13a on packet page 7. In this problem, you would multiply the p1(1.0 atm) and v1(5.0L). After you get the answer to that, you divide that2.0 atm(p2) to get the final volume(V2 or "X").

The last equation that is derrived off of the formula is: P=KT. This is equation is direct meaning that as the pressure of the gas increases, so does the Kelvin temperature of the gas. If a sample of gas with a pressure P1 and a temperature of T1 is changed to a new pressure P2 and a new temperature T2, you would use a two-point equation. The theory of this equation is called John's Law. The equation is: P1/T1=P2/T2. An example of John's Law is on packet page 8 and number 20a. You would plug in 300K for T1 and .528atm for P1, and 600K for T2 and "x" for P2. After that, you would solve the porportion (multiplying both sides by 600k to get 1.056 atm for the final pressure.

After going over packet pages 5-8, we went to the science computer lab to work on our web assigns. That concluded our day! HW: web assign due Friday

Wednesday, December 8

We started class today by doing page 1 in the unit packet.
The following statements are false:

• Gases are massless matter. Gases are made up of atoms, which have mass.
• Gases will condense when heated. Gases expand when heated.
• Gases are colorless. I2(g) is purple, NO2(g) is brown.
• Gases push on the walls of their container, but ONLY in a downwards direction due to the influence of gravity. Consider a balloon; gas is pushed in ALL directions on the wall
• The volume of a gas is inversely related to the temperature of a gas. They are directly related. As temperature increases, volume increases.
• The particles of a sample of gas are relatively motionless. The particles move very quickly.
• The entire column of our atmosphere applies pressure on objects on Earth; this pressure is so small that it is of little consequence and influence.

We then did page 2, #4. The units for pressure are atm, psi, mm Hg, torr, and Pa. The units for volume are L, m3, and mL. The units for temperature are celcius and Kelvin. The answers are:

a) P

b)V

c)V

d)P

e)P

f)T

g)V

h)P

i)P

j)T

We then did #5 on page 2. To do this, we had to use the equation °C+273=K

For a, we were given 100°C to convert to K. The answer was 373 K since 100+273=373.
For b, we were given 400 K to convert to °C. The answer was 127°C since 400-273=127.
For c, d, and e we had to use converting factors.
1 atm=14.7 psi=760 mm Hg
The answers were:
c) 1.03 atm
d) 3.0 atm
e) 684 mm Hg (torr)

We then did page 3, #8 and 9. To determine the pressure, you take the difference in height and add or subtract from 760 mm Hg. For the second manometer in #8 you had becuase P has more pressing power than Hg. For that one, P=700 mm Hg + 760 mm Hg = 1460 mm Hg. In the third one, the gas outside has more pressing power, so you subtract the difference in height. P=760 mm Hg - 560 mm Hg=200 mm Hg.

For #9 we did something similar to measure pressure. For A, we had to convert 75 cm (the differnence in heights) to 750 m Hg. For B, we subtracted the two heights, 82 cm- 25 cm, to get 57 cm which we then converted to 570 mm Hg. For C, we also subtracted the two heights, 125 cm-50 cm, to get 75 cm which we converted to 750 mm Hg.

We then wrote down the answers to the Chapter 5.2-3 reading sheet on page 27 about gas laws. The answers are:

1. a, b

2. b, c, d

3. a

4. b

5. d

6. c

7. a, ,b c, d

8. a

9. d

10. c

11. d

12. b, c, a

13. a

14. b, c

We ended class with some interesting demos. In one, Mr. H put a vlown up balloon in liquid nitrogen. That caused the volume to decrease because the the nitrogen is very cold, so the temperature had decreased. It was so cold that the air turned into liquid air. He also froze a water bottle by putting it in the liquid nitrogen. Another demo was to take the shrunken balloon and quickly put it in an empty bottle. As the temperature increases, the volume of the balloon increases causing there to be a blown up balloon inside a water bottle. At the very end of class, Mr. H poured the liquid nitrogen all over the floor where we watched it seem to disappear as it turned into a gas.

Today's chemistry class had us begin a new unit. We took the unit 5 test friday, so we started unit 6: Gases today. The first thing we did was meet in the chemistry lab to do a chemthink activity.

This link will take you to the chemthink website: http://chemthink.com/chemthink.html

*your login is your school i.d. and password

The Chemthink was on the Behavior of Gases. We started out with the tutorial which was very imformative.

*The first thing we learned was the effect temperature has on gases. As the temperature increases, gas particles begin to move faster, and when the temperature decreases, the particles move slower.

*The next thing we learned was the larger a particle is, the slower it moves. Take helium(He), neon(Ne), and argon(Ar). Argon, the largest particle, moves the slowest, Neon, the second largest, moves medium speed, and Helium, the smallest, moves the fastest.

As we continued through the chemthink, we learned that Pressure=Force/Area. Pressure also relates to Temperature, number of atoms, and volume.

This picture shows the relationship of how pressure varies directly with changes in temperature.

Formula: P/T=constant

This picture shows how pressure varies directly with the number of gas particles.

Formula: P/n=constant

This picture shows how pressure varies inversely with the volume of a gas sample

Formula: P*V=constant

This infomation is an overview of what the Chemthink was, and what the basics of the Gases unit are. After the Chemistry lab, we returned to the room and went over the previous nights webassign homework. The answers for page 27 in your unit packet are:

1. D

2. C

3.B

4.D

5.A

6.A

7.B

8.B

9.860

10. A

11.E

12.C

13.C

Notes from Class Discussion:

Volume (V)- Liters

=Amount of 3D space occupied by an object

*Further information can be found on page 7 in your book

Temperature (T)- Kelvin

=measure of the average amount of energy of motion (kinetic energy) of particles

*further information can be found on pages 7-8 in your book

Number of Moles (n)

1 mole=6.022 x 10^23

*further information can be found on page 55 in your book

Pressure (P)

=Force/unit area

=For gases, pressure results fromparticle collisions with container walls.

*further information can be found on page 104 in your book

Class then ended with Mr.H saying that we would finish going over problems 11-13 tomorrow, and that class the next day would have plenty of demonstrations regarding the new unit!

Thursday, December 2

Today we started off class by reviewing the 3.4 Reading Sheet. After looking over that, Mr. Henderson passed out calculators and told us how to use them more efficiently. There is a button (sto) that can store a numbers in the calculator. For example, if you wanted to store a molar mass because you would be using it in the problem a lot, you would plug that number under sto 1. So when you used that number, you would not have to type it in; you would just have to the recall 1 button. He also said that we should either be comfortable with his calculators or bring your own to use. If you bring your own calculator it cannot be an inspire.

We then went over what would the format of the test would like like.

-14 multiple choice questions

-3 pages of short answers (stoichiometry problems)

-1 question on the lab

To study for these types of problems there are review questions on the website and moodle with answers. A great way to study would be to use your packet. Everyone should have almost a completely filled out packet with correct answers. You could rework a problem that is difficult for you and check your answer with your packet. To prepare for the test in class, we completed review questions 17-19 which simulated what the questions about the lab would look like.

Furthermore, Mr. Henderson lectured us about percent yield in detail. To find the percent yield of a substance you need the theoretical yield (mass of the expected product) and the actual yield (mass of the measured product). You discover the theoretical yield by using the normal stoichiometry conversion factors, and you discover the actual yield by massing what your results are in the lab. The equation for the percent yield is:

Percent yield = Actual Yield * 100

Theoretical Yield

Lastly, we attempted problems involving this percent yield concept on a worksheet Mr. Henderson passed out. The toughest problem was number 5; the answer to it was 709.75 grams. Today was a very productive class period as we got in needed review time for the test. Finally, I would like to wish everyone good luck!

Wednesday, December 1

One-half the class was at the "Aids Day Workshop". Those who were present received a quick review of percent yield problems. We discussed the percent yield equation (see graphic below). We discussed how the yield in a given reaction refers to the amount of product that is produced. When conducting an experiment, the actual amount of yield is often less than the theoretical amount of yield. Spillage and other means of losing materials is the cause of the actual yield being less than the theoretical yield. The theoretical yield of product (i.e., the mass of product that we expect to obtain in a reaction) is related to the mass of reactants via the usual stoichiometry calculations.

We then took time to work on a handout on percent yield. Access Handout on course website.

The answers to the handout are shown below. For the first two columns of the table, think stoichiometry.

We will do more tomorrow on the topic of percent yield and review for the Friday exam. A review sheet for the exam should be up by the end of day on Wednesday.