1.  Unit 4: Counting Particles Too Small to See Miss Kelley Chemistry Spring 2015

2. Bell ringer 1/5  How could you find the number of Styrofoam packaging peanuts of a bag full of Styrofoam packaging peanuts?  Suppose that you had found that 100 peanuts had a mass of 5.5 g  This was the challenge that early chemists faced when they wanted to get some sense of how many atoms or molecules were present in a sample they were investigating.

3.  The pressure at a fixed mass of gas at a constant volume varies directly with the temperature  P/T=k Lussac’s Law

4. 1.Avogadro's Hypothesis  "equal volumes of gases at the same temperature and pressure contain the same number of molecules regardless of their chemical nature and physical properties”  V/n= k  V is the volume of the gas.  n is the amount of substance of the gas.  k is a proportionality constant.  https://www.youtube.com/watch?v=9FU0C569jI4 https://www.youtube.com/watch?v=9FU0C569jI4  The hypothesis allows us to determine the number of molecules that will react with one another based on volumes & enables us to demine the relative masses of elements and compounds.

5. 2. Bell ringer 1/12  What is relative mass?  Is there a connection between the concet of relative mass and the masses of the elements on the periodic table?

6. 2. Relative mass  Relative mass- proportion of the mass of a substance A in comparison to the mass of another substance B  A:B A/B or B/A  Relative atomic mass (symbol: A r ) is a dimensionless physical quantity, the ratio of the average mass of atoms of an element (from a single given sample or source) to 1 ⁄ 12 of the mass of an atom of carbon-12 (known as the atomic mass unit or amu). The relative atomic mass is a statistical term, referring to an abundance-weighted figure involving measurement of many atoms. As in all related terms, the word "relative" refers to making the figure relative to carbon-12, so that the final figure is dimensionless.carbon-12  The term relative atomic mass is exactly equivalent to atomic weight

7. Avagadro’s number  This number (Avogadro's number) is 6.022 X 10 23. It is the number of molecules of any gas present in a volume of 22.41 L and is the same for the lightest gas (hydrogen) as for a heavy gas such as carbon dioxide or bromine. It is also known as a Mole.

8. Mole  Even the smallest measurable mass of matter contains 6 sextillion atoms, so chemists use a unit of amount called the mole (abbreviated mol).  By definition, one mole is the number of atoms in 12 g of carbon-12. This number, called Avogadro's number, has been measured to be approximately 6.022 x 10 23 (to 4 s.f).

9. Perspective  Chemists use the mole in the same way that grocers use the dozen for groups of 12 and stationers use the ream for groups of 500.  By grouping numbers together, we get a smaller number to use in practical situations, 2 gross of paper clips for example, instead of 288 paper clips.  MOLE WS

11. 3 Molar Mass  The mole is just a number; it can be used for atoms, molecules, ions, electrons, or anything else we wish to refer to.  Molar mass allows chemists to FIND Composition of Compounds. Molar mass is used as a conversion factor to relate the amount of a substance to its mass.  amu= g/mol  Molar mass is generally calculated with two places after the decimal point  EX. Because we know the formula of water is H 2 O, for example, then we can say one mole of water molecules contains one mole of oxygen atoms and two moles of hydrogen atoms.

12. 4 &5 Molar Mass  The molar mass of glucose, C 6 H 12 O 6, is:  How many moles in 500. g?  EMPRICAL LAB  In general, the molar mass of any molecular compound is the mass in grams numerically equivalent to the sum of the atomic masses of the atoms in the molecular formula.  1 amu= 1g/mol

13. Extra Credit +5 -7points Empirical Lab Day 2

14. Bell ringer 1/13 Question  Ibuprofen C 13 H 18 O 2 is the active ingredient of many painkillers. If a bottle has a mass of 33 g, and there are 100 pills in a bottle, how many moles are in 1 pill? Solution

15. 6. Number of particles  The number of atoms in a sample of an element can be counted by weighing, in the same way that banks count pennies.  Ex: Determine how many moles are in a sample of silicon that has a mass of 5.23 g. Then determine the number of atoms

16. Mole city

17. Day 2 empirical lab  How could we drive off the hygroscopic liquid left in the beaker?  Strongly heat, but not strongly enough to produce smoke  After heating allow the sample to cool then find the mass.  Repeat the above step until the mass is within.02g

18. Day 2 empirical lab calculations  Calculate questions 1-5  Show all work  Lab write up due Thursday conclusions  Answer questions 1-4  Use complete sentences

19. Bell ringer 1/20 7. Number of molecules/atoms in a compound  First-century Roman doctors believed that urine whitened teeth and also kept them firmly in place. As gross as that sounds, it must have worked because it was used as an active ingredient in toothpaste and mouthwash well into the 18th century. Would you believe it’s still used today? Thankfully,not in its original form! Modern dentists recognized that it was the ammonia that cleaned the teeth, and they still use that. The formula for ammonia is NH3. How many moles are in 0.75 g of ammonia? How many molecules? How many atoms of N? of H? Atoms  Moles Molecules

20. 8. Percent composition  Percent Composition - The percent composition of a component in a compound is the percent of the total mass of the compound(all atoms) that is due to that component (1 type of atom, anion, or cation).  Percent composition- Mass ratio given by a %, showing how much of 1type of atom is in a compound  Ex: Calculate the % composition of Nitrogen in ammonia ( from previous example)

21. 9. Empirical Formula  Empirical Formula- simplest whole number ratio of a compound, where the subscripts of the atoms represent the relative mass  Based on empirical evidence (number stuff) or data  Relative number of moles  Ex: Suppose we have 100g sample where 73% is Mercury & 27% is chloride. What is the empirical formula of this compound?

22. 10. Molecular formula  Molecular formula- whole number ratio showing how many atoms from each element are present in a compound  based on their prefixes or charge … see unit 3 for naming rules  Empirical formula * the ratio of Divide the molar mass of the compound by the empirical formula’s mass of the compound  Emp. Form * MM/EM  Ex: Suppose we have a sample of 40.00% C, 6.72% H, 53.29% O. Which has a molarmass of 180 g/mol.  What is it’s Empirical formula?  molecular formula?