Friday, September 12, 2008
Fri-Day 1
AP Chem- Good news! Blackboard is now available for our course. I will catch up on posting previous class notes as well as putting up extra solved problems, worksheets, links, powerpoints, and videos! Check your Blackboard account and you should see the AP Chemistry course listed.
I will post most of the files in the "Course Documents" section.
We did a few more Law of Multiple Proportion problems. Some of you had questions about what the mathematical ratios of the element masses meant. Going back to our first example involving the two "oxides of carbon", which are really CO and CO2, when you put the mass ratio of O to C in CO2 OVER the ratio of O to C in CO, you are really calculating the relative masses from OXYGEN (because O is in BOTH numerators) in the two compounds for a given mass of C. Thus, the answer is 2 to 1 as you can see if you took one mole of CO2 and one mole of CO, you'd get 12 g of C from each compound BUT you'd get 32 g of O from CO2 and 16 g of O from CO, which is a 2 to 1 ratio. You'd get the same answer no matter what size samples of CO2 and CO that you used, as we demonstrated by example in class.
If you flipped the ratio to C to O for both compounds, with CO2 in the numerator and CO in the denominator, you'd get an answer of 0.5 or 1 to 2 because, for a given mass of oxygen in both compounds, there is only half as much carbon in CO2 as there is in CO. For proof, just take one mole of CO2 and TWO moles of CO (so that you can get the SAME 32 g of O extracted from each sample): decompose the samples and you get 12 g of carbon and 32 g of oxygen from the one mole of CO2; you get 24 g of carbon from the 32 g of oxygen in CO. So you get 12 g of C from CO2 to 24 g of C from CO, which is in a 1 to 2 ratio.
Either way, it shows that atoms combine in simple whole number ratios and that several compounds of the same elements will always have INTEGER/whole number mass ratios of their respective elements.
We also discussed Avogadro's famous insight that equal volumes of any gases must contain the same number of molecules. This law, when combined with the law of multiple proportions can be used to deduce the molecular formula of some compounds. Avogadro's law also can help us determine the stoichiometry of a gaseous chemical reaction.
We worked a little with glass tubing today to see how certain parts of glassware can be made by softening/melting and molding glass.
Bio 6/7- CHECK OUT Blackboard this weekend; I'll put up your HW answers as well as other extra help files. I will have extra help in Room 308 on MONDAY morning (8AM) so come prepared with any questions that gave you trouble.
we discussed some of the tools of the biologist including the RULES for making and recording measurements. The UNIT of measurement and the final "guessed" digit tells other scientist the precision and relative cost and type of instrument used in making the measurement. Remember, even if your measurement seems "exact", it is NOT; you must ALWAYS guess (even if that guess is a 0) to ONE more decimal place than the lowest decimal place that is PHYSICALLY drawn, etched, or somehow marked on the measuring device (unless the device has a digital display).
We also saw how to phrase the RELATIONSHIP between two variables, how to properly stain a microscope specimen on a slide; we also practiced a variety of ways to state the same hypothesis.
We then observed a stained amoeba specimen and practiced our microscopy.
Bio 8- CHECK OUT Blackboard this weekend; I'll put up your HW answers as well as other extra help files. I will have extra help in Room 308 on MONDAY morning (8AM) so come prepared with any questions that gave you trouble.
we discussed some of the tools of the biologist including the RULES for making and recording measurements. The UNIT of measurement and the final "guessed" digit tells other scientist the precision and relative cost and type of instrument used in making the measurement. Remember, even if your measurement seems "exact", it is NOT; you must ALWAYS guess (even if that guess is a 0) to ONE more decimal place than the lowest decimal place that is PHYSICALLY drawn, etched, or somehow marked on the measuring device (unless the device has a digital display).
We also saw how to phrase the RELATIONSHIP between two variables, how to properly stain a microscope specimen on a slide; we also practiced a variety of ways to state the same hypothesis.
I will post most of the files in the "Course Documents" section.
We did a few more Law of Multiple Proportion problems. Some of you had questions about what the mathematical ratios of the element masses meant. Going back to our first example involving the two "oxides of carbon", which are really CO and CO2, when you put the mass ratio of O to C in CO2 OVER the ratio of O to C in CO, you are really calculating the relative masses from OXYGEN (because O is in BOTH numerators) in the two compounds for a given mass of C. Thus, the answer is 2 to 1 as you can see if you took one mole of CO2 and one mole of CO, you'd get 12 g of C from each compound BUT you'd get 32 g of O from CO2 and 16 g of O from CO, which is a 2 to 1 ratio. You'd get the same answer no matter what size samples of CO2 and CO that you used, as we demonstrated by example in class.
If you flipped the ratio to C to O for both compounds, with CO2 in the numerator and CO in the denominator, you'd get an answer of 0.5 or 1 to 2 because, for a given mass of oxygen in both compounds, there is only half as much carbon in CO2 as there is in CO. For proof, just take one mole of CO2 and TWO moles of CO (so that you can get the SAME 32 g of O extracted from each sample): decompose the samples and you get 12 g of carbon and 32 g of oxygen from the one mole of CO2; you get 24 g of carbon from the 32 g of oxygen in CO. So you get 12 g of C from CO2 to 24 g of C from CO, which is in a 1 to 2 ratio.
Either way, it shows that atoms combine in simple whole number ratios and that several compounds of the same elements will always have INTEGER/whole number mass ratios of their respective elements.
We also discussed Avogadro's famous insight that equal volumes of any gases must contain the same number of molecules. This law, when combined with the law of multiple proportions can be used to deduce the molecular formula of some compounds. Avogadro's law also can help us determine the stoichiometry of a gaseous chemical reaction.
We worked a little with glass tubing today to see how certain parts of glassware can be made by softening/melting and molding glass.
Bio 6/7- CHECK OUT Blackboard this weekend; I'll put up your HW answers as well as other extra help files. I will have extra help in Room 308 on MONDAY morning (8AM) so come prepared with any questions that gave you trouble.
we discussed some of the tools of the biologist including the RULES for making and recording measurements. The UNIT of measurement and the final "guessed" digit tells other scientist the precision and relative cost and type of instrument used in making the measurement. Remember, even if your measurement seems "exact", it is NOT; you must ALWAYS guess (even if that guess is a 0) to ONE more decimal place than the lowest decimal place that is PHYSICALLY drawn, etched, or somehow marked on the measuring device (unless the device has a digital display).
We also saw how to phrase the RELATIONSHIP between two variables, how to properly stain a microscope specimen on a slide; we also practiced a variety of ways to state the same hypothesis.
We then observed a stained amoeba specimen and practiced our microscopy.
Bio 8- CHECK OUT Blackboard this weekend; I'll put up your HW answers as well as other extra help files. I will have extra help in Room 308 on MONDAY morning (8AM) so come prepared with any questions that gave you trouble.
we discussed some of the tools of the biologist including the RULES for making and recording measurements. The UNIT of measurement and the final "guessed" digit tells other scientist the precision and relative cost and type of instrument used in making the measurement. Remember, even if your measurement seems "exact", it is NOT; you must ALWAYS guess (even if that guess is a 0) to ONE more decimal place than the lowest decimal place that is PHYSICALLY drawn, etched, or somehow marked on the measuring device (unless the device has a digital display).
We also saw how to phrase the RELATIONSHIP between two variables, how to properly stain a microscope specimen on a slide; we also practiced a variety of ways to state the same hypothesis.
# posted by C @ 1:41 PM 
