Atomic Mass and Intro to the Mole

1. Atomic Mass and Intro to the Mole

2. Class Opener • How do isotopes of the same element differ from each other?

3. Objectives • To define atomic mass and to demonstrate how the average atomic mass of an element is determined • To introduce the concept of the mole

4. Atomic Mass • Decimal number listed on the periodic table. • Weighted average of all the naturally occurring isotopes of that element. • No atom has this exact mass.

5. You will notice that the average atomic mass of an element is often closest to the atomic mass of the most abundant element.

6. Study Check • Chlorine has two naturally occurring isotopes, chlorine-37 and chlorine-35. Which isotope is more abundant? How do you know?

7. Calculating Average Atomic Mass • The average atomic mass of an element depends on both the mass and the relative abundance of each of the element’s isotopes. Naturally occurring copper consists of 69.15% copper-63, and 30.85% copper-65.

8. To calculate the atomic mass, take the mass times the decimal form of the percent abundance of each. Then you add the numbers together to get the number that appears on the periodic table.

9. Atomic Mass for Copper Naturally occurring copper consists of 69.15% copper-63, and 30.85% copper-65 0.6915 x 62.9296= 43.52 0.3085 x 64.9278= 20.03 63.55 (This is what appears on the periodic table)

10. Study Check • Naturally occurring carbon consists of 98.93% carbon-12 with an atomic mass of 12.00, and 1.07% carbon -13 with an atomic mass of 13.00. Calculate the atomic mass.

11. The Mole An Introduction to Chemistry’s Favorite Number

12. Quantities in Chemistry • Imagine trying to count all the grains of sand in the castle • Easier way: • Count all the grains of sand in 1 gram of sand • Weigh all the sand and convert using your previous measurement

13. Quantities in Chemistry • Another easy way: • Count all the grains of sand in 1 liter of sand • Measure the volume of the sand and convert using your previous measurement

14. How Does This Relate to Chemistry? • Chemicals react in fixed ratios at the atomic level • In order to predict how reactions will occur, chemists need to know how many atoms or molecules they have

15. How Does This Relate to Chemistry? • Imagine you want to burn the spoonful of sugar on the right • Every sugar molecule reacts with a fixed number of oxygen molecules • In order to know how much carbon dioxide and water will be produced, you need to know how many molecules of sugar you start with

16. SI unit for amount of substance is called mole. • A mole measures the number of particles within a substance. • A mole refers to a specific number of particles. (Counting Unit) • Particles can be atoms, ionic compounds, or molecules

17. 1 mole = 6.02 x 1023 particles • 6.02 x 1023 is also known as Avogadro’s Number 1 mole aluminum = 6.02 x 1023 Al atoms 1 mole copper = 6.02 x 1023Cu atoms 1 mole lead = 6.02 x 1023 Pb atoms

18. Although 1 mole always contains the same number of particles, the mass of one mole varies depending on the substance. • Molar Mass – mass of one mole of a substance. • Mass of one mole of an element is equal to its ATOMIC MASS expressed in grams. • 1 mole of aluminum = 26.98 grams • 1 mole of copper = 63.55 grams • 1 mole of lead = 207.2 grams

19. Relating Mass to Moles • Molecular mass is the sum of the atomic masses of the component elements • H2O = 2 hydrogen, 1 oxygen = 2(1) + 1(16) = 18 • CH4 = 1 carbon, 4 hydrogen = 1(12) + 4(1) = 16 • NaCl = 1 sodium, 1 chlorine = 1(23) + 1(35) = 58 • 1 mole of molecules will have a weight that equals the molecular mass • 1 mol H2O = 18 g • 1 mol CH4 = 16 g • 1 mol NaCl = 58 g

20. Representative Particles & Moles