Thursday, November 1, 2007
Thurs-Day 1
Bio- we reviewed our unit on Mitosis, Cytokinesis, DNA, and Types of Asexual Reproduction.
We also practiced TEST-TAKING skills, which are to be employed as you practice for tomorrow's test and as you take tomorrow's test so that you MAXIMIZE your score and take pride in your effort; you will also have the satisfaction of knowing that you could not try any harder when you properly employ all of the techniques together: as soon as the test BEGINS, no sooner, drawing out things that you KNOW will be tested, jotting down any information that you might forget, UNDERLINING/circling key terms in the question and ASSOCIATING those terms with your predicted answer, which you then write down.
In short answer questions, make SURE that your answer includes the KEY TERMS from the questions but does not MERELY repeat those terms; your answers will have to use those terms in a detailed description or explanation. Remember, DRAWINGS help explanations and labeled drawings are a large part of any good explanation.
The answers to the DNA tutorial handout will be posted on Blackboard this afternoon. Check your work.
Nobody came to extra help this morning. That means that everybody is very confident with the material in this unit or that you are not getting help that is available when you require it.
Chem 7- we reviewed our Graham's Law of diffusion/effusion explanation and then went on to a very important mathematical concept in Chemistry: the "MOLE" is a term used to indicate Avogadro's number of particles = 6.02 x 10^23 particles.
One thing that has been experimentally determined and related to Avogadro's Law is that ONE MOLE of any ("ideal" - follows the gas laws and tenets of K-M Theory) gas at STP takes up 22.4 L of space. So if you have, 22.4 L of oxygen or ozone or nitrogen or carbon monoxide or carbon dioxide gas at STP, there ARE 6.02 x 10^23 molecules of any of those gases in that 22.4 L volume container.
Using our reference tables, we are fortunate that we can just look at the "atomic mass" of ANY element and know that the number seen is the MASS of that element IN GRAMS for one mole of atoms of that element, e.g. one mole of Argon atoms has a mass of 39.948 grams, whereas one mole of Neon atoms has a mass of 20.179 grams. Thus each Neon atom weighs about half as much as each Argon atom.
From this, we began to do some gas density calculations ONLY for gases at STP (1 mol of any gas at STP has a volume of 22.4 L)!!!
Chem 8/9- we delved into Graham's (like the crackers) Law, which state the relationship between the relative masses of gas molecules and their respective rates of diffusion and effusion. Graham found that, at the same temperature and pressure, the heavier the gas, the slower its rate of effusion/diffusion and the lighter the gas, the faster its rate of effusion/diffusion.
We explained this phenomenon by looking at the Kelvin Temperature - average kinetic energy connection of Kinetic Molecular Theory. At the same temperature, any two samples of gas molecules will have the SAME average kinetic energy BUT, there are TWO factors that contribute to kinetic energy: velocity AND mass. So, if two molecules have the same KE, the heavier one MUST be going slower (thus diffusing slower) and the lighter one must be going faster (thus diffusing faster). This accounts for Graham's "empirical" (EXPERIMENTALLY found, not just theoretically predicted!) discovery.
We saw how to calculate the relative masses of molecules by using the "atomic mass" numbers listed in the Periodic Table. For example, N= 14.3 atomic mass units and H= 1.0 atomic mass units so, ammonia, NH3 has an atomic mass of 17.3 (one N and three H's). Carbon dioxide, CO2, has an atomic mass of 44.0 atomic mass units
( one C = 12.0 amu and two O's = 2 x 16.0 amu) = 44.0 amu ; so, under the same conditions (T and P), ammonia will diffuse/effuse faster than carbon dioxide!
We did a lab showing the effect of varying temperature on the volume of a gas that is at constant pressure and a constant number of molecules or "moles". We saw that, as the little sample of air in a capillary tube is cooled at constant external pressure, the volume of that sample of air decreases. We will discuss our data from that Charles's Law lab on Monday.
The lab write-up for the "crush the can" (Gay-Lussac's Law) lab is due Monday, after we discuss some of the questions from that lab, tomorrow.
We also practiced TEST-TAKING skills, which are to be employed as you practice for tomorrow's test and as you take tomorrow's test so that you MAXIMIZE your score and take pride in your effort; you will also have the satisfaction of knowing that you could not try any harder when you properly employ all of the techniques together: as soon as the test BEGINS, no sooner, drawing out things that you KNOW will be tested, jotting down any information that you might forget, UNDERLINING/circling key terms in the question and ASSOCIATING those terms with your predicted answer, which you then write down.
In short answer questions, make SURE that your answer includes the KEY TERMS from the questions but does not MERELY repeat those terms; your answers will have to use those terms in a detailed description or explanation. Remember, DRAWINGS help explanations and labeled drawings are a large part of any good explanation.
The answers to the DNA tutorial handout will be posted on Blackboard this afternoon. Check your work.
Nobody came to extra help this morning. That means that everybody is very confident with the material in this unit or that you are not getting help that is available when you require it.
Chem 7- we reviewed our Graham's Law of diffusion/effusion explanation and then went on to a very important mathematical concept in Chemistry: the "MOLE" is a term used to indicate Avogadro's number of particles = 6.02 x 10^23 particles.
One thing that has been experimentally determined and related to Avogadro's Law is that ONE MOLE of any ("ideal" - follows the gas laws and tenets of K-M Theory) gas at STP takes up 22.4 L of space. So if you have, 22.4 L of oxygen or ozone or nitrogen or carbon monoxide or carbon dioxide gas at STP, there ARE 6.02 x 10^23 molecules of any of those gases in that 22.4 L volume container.
Using our reference tables, we are fortunate that we can just look at the "atomic mass" of ANY element and know that the number seen is the MASS of that element IN GRAMS for one mole of atoms of that element, e.g. one mole of Argon atoms has a mass of 39.948 grams, whereas one mole of Neon atoms has a mass of 20.179 grams. Thus each Neon atom weighs about half as much as each Argon atom.
From this, we began to do some gas density calculations ONLY for gases at STP (1 mol of any gas at STP has a volume of 22.4 L)!!!
Chem 8/9- we delved into Graham's (like the crackers) Law, which state the relationship between the relative masses of gas molecules and their respective rates of diffusion and effusion. Graham found that, at the same temperature and pressure, the heavier the gas, the slower its rate of effusion/diffusion and the lighter the gas, the faster its rate of effusion/diffusion.
We explained this phenomenon by looking at the Kelvin Temperature - average kinetic energy connection of Kinetic Molecular Theory. At the same temperature, any two samples of gas molecules will have the SAME average kinetic energy BUT, there are TWO factors that contribute to kinetic energy: velocity AND mass. So, if two molecules have the same KE, the heavier one MUST be going slower (thus diffusing slower) and the lighter one must be going faster (thus diffusing faster). This accounts for Graham's "empirical" (EXPERIMENTALLY found, not just theoretically predicted!) discovery.
We saw how to calculate the relative masses of molecules by using the "atomic mass" numbers listed in the Periodic Table. For example, N= 14.3 atomic mass units and H= 1.0 atomic mass units so, ammonia, NH3 has an atomic mass of 17.3 (one N and three H's). Carbon dioxide, CO2, has an atomic mass of 44.0 atomic mass units
( one C = 12.0 amu and two O's = 2 x 16.0 amu) = 44.0 amu ; so, under the same conditions (T and P), ammonia will diffuse/effuse faster than carbon dioxide!
We did a lab showing the effect of varying temperature on the volume of a gas that is at constant pressure and a constant number of molecules or "moles". We saw that, as the little sample of air in a capillary tube is cooled at constant external pressure, the volume of that sample of air decreases. We will discuss our data from that Charles's Law lab on Monday.
The lab write-up for the "crush the can" (Gay-Lussac's Law) lab is due Monday, after we discuss some of the questions from that lab, tomorrow.
# posted by C @ 11:33 AM 
