1. Section 6.1 Atoms and Moles 1.Students will be able to describe the concept of average mass. 2.Students will be able to demonstrate how counting can be done by weighing. 3.Students will be able to define atomic mass and use it to convert between mass and number of atoms. 4.Students will be able to define a mole and Avogadro’s number. 5.Students will be able to complete calculations converting between mass, moles, and number of atoms. Objectives

2. Section 6.1 Atoms and Moles Class Warmup

3. Section 6.1 Atoms and Moles Counting by Weighing How many pennies are in the jar? Count or Weigh? To determine by weighing what information is necessary? 1.Mass of one penny 2.Total mass of the jar of pennies 3.Mass of the jar

4. Section 6.1 Atoms and Moles Counting by Weighing Is there a flaw in this method? What if all of the pennies do not have the same mass? Find the avg. mass. How is this relevant to counting the number of atoms? Atoms of the same element can have different masses (isotopes)

5. Section 6.1 Atoms and Moles Counting by Weighing Let’s say you wanted to place 1000 pennies into a container. You can do this without actually counting them if you have a balance. If you know the avg. mass of a penny, you can determine the total mass of 1000 pennies.

6. Section 6.1 Atoms and Moles B. Atomic Masses: Counting Atoms by Weighing Atoms have very tiny masses so scientists made a unit to avoid using very small numbers. 1 atomic mass unit (amu) = 1.66  10 -24 g Scientists defined the atomic mass unit: 1 Carbon-12 atom = 12 amu Since carbon-12 has a sum of 12 protons and neutrons and the protons and neutrons have very similar masses: 1 proton ~ 1 amu 1 neutron ~ 1 amu

7. Section 6.1 Atoms and Moles B. Atomic Masses: Counting Atoms by Weighing The mass of an atom is approximately equal to its mass number 65 Cu ~ 65 amu 35 Cl ~ 35 amu For carbon 12, the atomic mass = 12 amu (exactly) For all other isotopes, the exact mass is slightly different from the mass # 35 Cl = 34.969 amu 37 Cl = 36.966 amu

8. Section 6.1 Atoms and Moles B. Atomic Masses: Counting Atoms by Weighing The atomic mass listed on the periodic table is the weighted average of all naturally occurring isotopes of that element

9. Section 6.1 Atoms and Moles C. The Mole Objectives: 1.Define a mole 2.Use the mole to calculate atoms, molecules, and mass.

10. Section 6.1 Atoms and Moles C. The Mole The Mole A counting unit used to count atoms, molecules, formula units, electrons, etc. 1 mole = 6.02 x 10 23 (Avogadro’s Number) How was the mole determined? 1 mole equals the number of carbon-12 atoms is exactly 12 grams of carbon 12.

11. Section 6.1 Atoms and Moles C. The Mole Because of the way atomic mass units and the mole are defined, there is a relationship between grams and amu So… 1 atom of sodium (Na) = 22.98977 amu 1 mole of sodium (Na) = 22.98977 grams

12. Section 6.1 Atoms and Moles C. The Mole One mole of anything contains 6.022 x 10 23 units of that substance. –Avogadro’s number is 6.022 x 10 23. A sample of an element with a mass equal to that element’s average atomic mass (expressed in g) contains one mole of atoms.

13. Section 6.1 Atoms and Moles C. The Mole