Tuesday, September 23, 2014 Objective: Students will calculate the average atomic mass of common isotopes of an element. Warm-Up What do the following.

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Presentation transcript:

Tuesday, September 23, 2014 Objective: Students will calculate the average atomic mass of common isotopes of an element. Warm-Up What do the following element descriptions actually mean? (think about how many subatomic particles they have, how they are different or similar, etc.) Silicon-28 Agenda Average Atomic Mass Notes and Practice

Average Atomic Mass

Average Atomic Mass All elements exist as a collection of isotopes. Isotope’s Natural Abundance: how commonly an isotope is found in nature. Since all elements exist as a collection of isotopes, we need to find the average atomic mass of the element in order to use it for calculations. The average atomic mass of an element is the sum of the masses of its isotopes each multiplied by its natural abundance. Average Atomic Mass is a weighted average!

Calculating Average Atomic Mass Change the percent abundance to a decimal. Multiply the decimal amount by the mass number. Repeat steps 1 and 2 for all isotopes. Add the answers together to get the average atomic mass. The unit is amu, atomic mass unit.

Brain Break: Dice Intervals 1: jumping jacks 2: 30 sec wall sit 3: 10 squats 4: Ski jumps 5: run in place 30 seconds 6: 10 wall push-ups

Example Nitrogen-14 99.63% Nitrogen-15 0.37%

Example Lithium-6 7.5% Lithium-7 92.5%

Example Magnesium-24 78.99% Magnesium-25 10.00% Magnesium-26 11.01%