1. Bellwork Friday, 10/9 Find your notes on isotopes. At the top of pg. 2: Fill in the following… Finish this and we’ll check it.

2. 6 protons, 8 neutrons, 6 electrons 6 protons, 5 neutrons, 6 electrons How many protons, neutrons, and electrons? Draw good nuclide symbols.

3. Let’s check… The only stable isotope of naturally occurring fluorine has a mass number of 19. How many p+, n 0 and e- are in an atom of F? Write the isotope symbol. 3

4. 4 F199

5. Word Wall Isotope Atomic number Mass number Nuclide

6. Chapter 3 Calculating average atomic mass

7. 7 In chemistry there are many different concepts of mass… We need to talk about one of them. Atomic mass

8. Please don’t panic… I like to rewrite my notes from my slides. I think it helps you later when you study. (Y’all are all studying every night, right??) Also… there are different terms for things and it depends on the textbook, the professor… and it gets confusing when there are truly different types of mass… I’ll do my best to keep it clear for you. But promise not to panic on me.

9. Promise?!

10. Isotopes have mass numbers Isotopes are atoms of the same element (the same number of p+ and e-) but with different masses due to having different numbers of n 0. Right? 10

11. What’s the difference between atomic mass (on the periodic table) and mass number? Today we get to work it out! Question of the day:

12. We know that the mass of an atom is very, very, very, very small. The mass of an atom is incredibly small and it doesn’t make sense to use units like grams or even picograms to measure it… It’s easier to measure it in atomic mass units (amu) 12

13. Scientists like a standard… The chosen isotope is carbon-12. It was assigned a mass of exactly 12 atomic mass units. (This is handy because it has 6 p+ and 6 n 0.) Sooooo… An amu is defined as one-twelfth the mass of a carbon-12 atom. (our textbook uses just plain “u”) 13

14. Scientists use a reference isotope for these units. This means amu is a relative measurement. (relative to carbon-12).

15. Average atomic mass This is the number you see on the periodic table. 15

16. Average atomic mass. The average atomic mass of an element is the weighted average of the masses of its isotopes on this scale… Remember relative abundance? Think about “weighted” like your grades: 40% formative and 60% summative. 16

17. 17 The average atomic mass of an element depends of both the mass and relative abundance of each of the element’s isotopes. Most elements occur in nature as a specific mixture of isotopes. They don’t occur equally… some are more abundant than others.

18. The average atomic mass is the weighted average of all of the naturally occurring isotopes of the element. Isotope Percent Abundance Atomic Mass, amu 12 C 98.90%12.00000 13 C 1.10%13.00335

19. 19 Elements rarely occur as only one isotope. They exist as mixtures of different isotopes of various masses. Using average atomic mass means the less common isotopes are accounted for.

20. 20 Don’t confuse mass number with average atomic mass. Mass number is the mass of one particular atom (isotope). Average atomic mass is the average mass of a group of atoms of the same element that takes into consideration all the isotopes.

21. Go back and answer the question at the beginning of your notes. Make sure you both understand this concept.

22. How do we calculate the average atomic mass? The average atomic mass of an element depends of both the mass and relative abundance (percent abundance) of each of the element’s isotopes. Let’s look at a problem. 22

23. Copper consists of 69.15% copper-63, which has an atomic mass of 62.929 601 amu, and 30.85% copper- 65, which has an atomic mass of 64.927 794 amu. Calculate the average atomic mass for Cu.

24. Breaking down the calculation: Step 1: multiply the atomic mass of each isotope by its relative abundance. Step 2: add the results. 0.6915 x 62.929 amu + 0.3085 x 64.927 794 amu = 43.52 amu + 20.03 amu = 63.55 amu 24

25. Here’s another one! Oxygen has three naturally occurring isotopes in the following proportions: O-16 99.762 % (15.99491 amu); O-17 0.038000% (16.99913 amu); O-18 0.20000% (17.99916 amu). What is the average atomic mass of oxygen?

26. Did you get it? 1).99762 x 15.99491 = 15.957 2).00038000 x 16.99914 = 0.0064597 3).0020000 x 17.999 =.035998 15.999 amu 26

27. More problems for practice! Textbook, pg. 117 Look at the sample problem then work problems #1 and 2.

29. Bellwork 10/13 Calculate the atomic mass of bromine. The two isotopes of bromine have atomic masses and relative abundance of 78.92 amu (50.69%) and 80.92 amu (49.31%).

30. Bullseye comparison Isotope Atom Atomic mass Electron Mass number Nucleus Atomic number Periodic table Proton neutron

31. Activity Time! The atomic mass of candium