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Math All the math

The Molar Mass Molar Mass – the mass of one mole of any substance, reported in grams (gram atomic mass)

The Molar Mass of a Compound or Molecule Molar mass – sum of the molar masses of the elements in the compound Example: CH4 The molar mass of CH4=16.05 g/mole Element # mol element/mol compound Molar Mass Mass contribution C 1 12.01 g/mol 12.01 g H 4 1.01 g/mol 4.04 g

For example, the molar mass for the element, Al, is 27. 01 g For example, the molar mass for the element, Al, is 27.01 g. This means that 1 mole of Al atoms has a mass of 27.01g. In other words, 6.02 x 1023 Al atoms has a mass of 27.01 g.

Percent Composition Relative amounts of each element in a compound Can be used as a conversion factor

Practice What is the percent by mass of nitrogen in NH3? MM of NH3: 17.04 g/mol Mass contribution of N: 14.01 g/mol (14.01g/mol)÷(17.04g/mol) * 100 = 82.22%

The mole We use it to count, just like we use dozen or pair. If we have: 1 dozen roses= 12 roses If we have one mole of anything: We have 6.02x1023 of that thing Avogadro’s number

Empirical Formulas A chemical formula showing the lowest whole # ratio of elements in a compound Can be calculated from percent composition Example: Find the empirical formula for a compound that is composed of 62.1% C, 13.8% H, and 24.1% N.

Empirical Formula Contd To determine formula: Convert the masses of each element into moles. Divide each mole value by the smallest mole value to get mole ratio Write empirical formula using symbols, and use the whole-number ratio as subscript. Non-whole numbers – multiply by 2 or 3 to make them whole

If percents are given, assume that the % = the number of grams. C 62.1g H 13.8g N 24.1 g

Molecular Formulas will be identical to or some multiple of the empirical formula. Show the actual number of each kind of atom present Example: C2H6=Molecular Formula CH3=Empirical Formula Example: CH4=Molecular Formulas CH4=Empirical Formula

Obtaining Molecular Formula from Empirical Formula + Molar Masses Steps *Find empirical formula *Calculate molar mass for empirical formula *Divide MM molecular/MM empirical *Multiply this factor by each of subscripts in empirical formula to give molecular formula

Dimensional Analysis Three steps: List the knowns and the unknown. Solve for the unknown. Select the appropriate unit equality and make it into a conversion factor, with the unknown’s unit in the numerator and the known’s unit in the denominator. Multiply the known quantity by the conversion factor. Make sure to eliminate all but the unknown’s unit. Evaluate - Does the result make sense? Chapter 10

Converting Moles to Particles and vice versa. We use moles to count particles like: atoms, molecules, formula units, or ions. One mole of moles (the animal) would have the equivalent mass to 60x the Earth’s oceans.

Mole to particles How many atoms are there in 1 mole of sodium? 6.02x1023 atoms How many formula units are there in 1 mole of KOH? 6.02x1023 formula units How many molecules are there in one mole of water? 6.02x1023 molecules

Particles to mole Example: How many moles of water molecules are in 1.21 x 1024 water molecules?

Mole to Particles Example: How many Cu atoms are in 2.54 moles of Cu atoms?

Mass to mole Calculations If you are given a certain mass of a substance, you can find out how many moles you have by using the molar mass. Example: How many moles of carbon atoms are found in 24.02 g of carbon? 24.02g x (1mole) = 2.000 moles C (12.01g)

The Molar Volume of a Gas One mole of any gas at STP occupies a volume of 22.4L. S.T.P.-Standard Temperature and Pressure. The values are 0 ◦Celsius and 1 atmosphere of pressure. Example: What is the volume of 2.55 moles of hydrogen gas @STP?

The mole is the heart of the matter molar mass 1.00 mol 6.02x1023particles 1.00 mol MOLE Particles Mass 1.00 mol molar mass 1.00 mol 6.02x1023particles 1 mol 22.4 L 22.4 L 1 mol Volume (gas at STP)

Stoichiometry The calculation of quantities in chemical reactions A balanced chemical equation becomes a recipe that we can use

What do equations tell us? 2Na + Cl2  2NaCl In the above equation, the coefficients tell you: The relative number of particles The relative number of moles If and only if all substances are gaseous, the relative number of liters @STP

Mole to Mole Conversions How many moles of chlorine gas are needed to react with excess sodium to produce 5.00 moles of sodium chloride?

Molarity A unit of concentration Units are mole solute/liter solution Conversion factor! 

Practice If you dissolve 5 moles of NaCl in enough water to produce 1 L of solution, what is the concentration?

Gram to Gram Conversions How many grams of water can be produced when 64.00 grams of oxygen react with excess hydrogen?

Volume to Volume Conversions 2H2 (g) + O2(g)  2H2O (g) Example: How many liters of oxygen @ S.T.P. are needed to completely react with 6.0 liters of hydrogen to produce water? Notice, what is different about this reaction from the one we’ve previously been working with?

Percent Yield Measurement of efficiency of a reaction

Practice A factory worker in a paint company follows a batch slip to prepare a 1500 gallon batch of paint. However, when the batch goes to the filling department only 1489 gallons are filled. What is the % yield?

Limiting reactant/reagent As in the cookie recipe, if you run out of eggs (and can’t get more), but have the rest of the recipe, you can’t make any more cookies. In chemistry we call the reactant that runs out first, the limiting reactant/reagent. All other reactants that we have in excess we call excess reagents.

Practice 2H2 + O2  2H2O If 40.00g of hydrogen react with 25.00 g of oxygen, what mass of water will be produced?