1. Atomic Mass The Atomic Mass of Candium Lab

2. Atomic Mass This is the weighted average mass of the atoms in nature of that element A weighted avg mass reflects both the mass & the relative abundance of the isotopes

3. Atoms of the same element with a different number of neutrons

4. Different isotopes have different properties Name followed by mass number to identify Boron – 10 and Boron - 11: which is more abundant if the mass of B is 10.81 amu?

5. To calculate atomic mass based on relative abundance: The number of stable isotopes of the element The mass of each isotope The natural % abundance of each isotope

6. EX Problem: Element J has 2 natural isotopes. The isotope with a mass of 10.012 amu has a relative abundance of 19.91%. The isotope with a mass of 11.009 amu has a relative abundance of 80.09%. Find the atomic mass of the element.

7. Knowns: Isotope 1 Mass = 10.012 amu RA = 19.91% or 0.1991 Isotope 2 Mass = 11.009 amu RA = 80.09% or 0.8009

8. The mass each isotope contributes to the element’s atomic mass can be calculated by multiplying the isotope’s mass by its RA. The atomic mass of the element is the sum of these individual contributions.

9. Contribution isotope 1 = 10.012 x 0.1991 = 1.993 amu Contribution isotope 2 = 11.009 x 0.8009 = 8.817 amu Add these values = 10.810 amu The calculated value is closer to the more abundant isotope so our answer makes sense.

10. Now you try: Copper has naturally occuring isotopes with mass numbers 63 & 65. The RA and atomic masses are 69.2% for mass = 62.93 amu & 30.8% for mass 64.93 amu. Calculate the avg atomic mass of copper.

11. Did you get 63.56 amu? Remember sig figs.. This answer is closer to the more abundant isotope’s mass so it makes sense. This is what you will be finding in lab tomorrow. Let’s look at it now.