Unit 7: Chemical Quantities

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Presentation transcript:

Unit 7: Chemical Quantities Moles <--> Particles

After today, you will be able to… Using dimensional analysis, calculate how many atoms, molecules, or formula units are in one mole of that substance Use correct significant figures and units in these calculations Identify key scientists in the development of the mole

Substances can be measured in different ways…. The Mole can be related to each of these types of measurements. Count – 5 apples, dozen eggs Mass – 2lbs of bananas, 16oz box of cereal Volume – gallon of milk, pint of cream

A Brief History on the Mole… Avogadro: (1811) An Italian scientist who studied the behavior of gases. Theorized:“The volume of a gas at a specific temperature and pressure contains equal numbers of atoms or molecules regardless of the nature of the gas.”

Loschmidt (1865): Estimated the average diameter of the molecules in air and was able to calculate the number of particles in a given volume of gas. Millikan (1910): Measured the charge on an electron. From the charge on a mole of electrons, he divided the two and obtained Avogadro’s number.

Perrin (1926): Earned the. Nobel Prize for computing Perrin (1926): Earned the Nobel Prize for computing Avogadro’s number using many different methods and named this constant in honor of Avogadro. Used oxygen as a standard and proposed “Avogadro’s number is the number of molecules in exactly 32-grams of oxygen.”

The standard was later changed to the carbon-12 isotope. The presently accepted definition of the mole is: “The amount of any substance that contains as many elementary entities as there are in 12 grams of pure carbon-12.”

The mole (mol) as a unit in chemistry serves as a bridge between the atomic and macroscopic worlds. In Latin, mole means “huge pile.”

1 mole = 6.02x1023 atoms, molecules, or formula units = 602, , , , , , , 000 000 000 000 000 000 000 = six-hundred and two sex-tillion

This equality can be used as a conversion factor to convert particles into moles, or moles into particles.

Example: How many atoms are in a 1.22 mole sample of sodium? K: 1.22 mol Na K stands for what we know! U stands for what we are solving for! U: ? atoms Na 1.22 mol Na 6.02x1023 atoms Na x = 1 1 mol Na 7.34x1023 atoms Na

Example: How many moles are equal to 4.79x1024 molecules of CO? K: 4.79x1024 molec. CO U: ? mol CO 4.79x1024 molec. CO 1 mol CO . = x 1 6.02x1023 molec. CO 7.96 mol CO