1. Things you need to review and may need help with when we return in the fall. You will find notes and examples within this PowerPoint as well as useful video links.

2. History and Experiments Dalton’s Atomic Theory – two parts are incorrect all atoms are not identical atoms can be divided JJ Thomson – neutral atom, p+ = e- discovered electrons with cathode tube experiment Lord Kelvin – Plum Pudding model Protons and electrons are mixed together in equal quantities Millikan – oil drop exp quantitative charge and mass of e- Rutherford - gold foil exp small dense positive nucleus atom mostly empty space Bohr – with Rutherford planetary model, only works for H

3. Atoms

4. More about atoms Are there any other particles in the atom? – Scientists have found the subatomic particles can be subdivided into smaller parts, for example quarks. Quarks are the elementary particles that protons and neutrons are made up of. (1968) What is an atomic number? Where do we find it? How can we use it? – Atomic # = # of protons – Found on the periodic table – Used to identify an element, an element only has one atomic number In an atom, what is the relationship between protons and electrons? – An atom is neutral, so that means the protons=the electrons We never change or manipulate protons, which would mean we are changing the element and we aren’t nuclear chemists.

5. Isotopes (discovered by Chadwick) Definition – atoms of an element with the same number of protons, but different numbers of neutrons Impact on atom - changes the mass number and the mass of the atom Discredited part of Dalton’s atomic theory Name of an isotope – name of element – mass number, example carbon – 12 Symbol of an Isotope X = element symbol AA = mass number = (# of p + ) + (# of n 0 ) ZZ = atomic number = # of p+

6. Calculating Average Atomic Mass Average atomic mass – the weighted average of the mass of all the isotopes that exist for an element, – given the unit amu, – found on the periodic table AAM =  [(isotope mass) x (decimal abundance of isotope)] For each isotope you multiply its mass by its decimal abundance. Then you add these products together. Example #1 – Calculate Average Atomic Mass Naturally occurring chlorine is 75.53% 35 Cl, which has an atomic mass of 34.969 amu, and 24.47% 37 Cl, which has a mass of 36.966 amu. Calculate the average atomic mass of chlorine.

7. Mass Spectrometry and Isotopes https://youtu.be/tOGM2gOHKPc This is an excellent video that using a simple analogy to introduce the concept and then applies that knowledge to the actual working of a mass spectrometer. https://youtu.be/NuIH9-6Fm6U Excellent video on mass spectrometry. Gives great insight into how data is collected.

8. @ STP, 1mol = 22.4L Use Avogadro’s # Use molar mass MOLES Used to convert gases only! Liters Grams Particles

9. Percent Composition Steps for Calculation of Percent Composition ①Calculate the molar mass of the compound. ②Divide the mass of the element in 1 mol of compound by the mass of 1 mol of the compound. ③Multiply by 100. ④Check the sum of the mass percents; they should equal 100. Examples #1 and #2: CCl 4 NaOH

10. Calculating Empirical Formula Steps for Calculating Empirical Formula ①Obtain the mass of each element, generally given, but may involve a subtraction step. ②Convert grams to moles. ③Divide the number of moles of each element by the smallest number of moles to convert the smallest number to one. If all calculated numbers are whole numbers, these are the subscripts in the empirical formula, OTHERWISE go to step four. ④Multiply all the numbers from step three by the smallest whole number that will convert all of them to whole numbers. Example #1:Calculate the empirical formula of a compound using a sample that contains 4.151 g Al and 3.692 g of O.

11. Calculating Molecular Formula Calculation of Molecular Formulas To calculate the molecular formula, three things are needed: ①the empirical formula ②the empirical formula molar mass ③the molecular molar mass (always given) To find the molecular formula, the empirical formula molar mass if compared to the molecular molar mass. Steps for Calculating Molecular Formulas ①Calculate the empirical formula, if necessary. ②Calculate the empirical formula molar mass. ③Divide the molecular molar mass by the empirical formula molar mass. ④Multiply the empirical formula be the result of step 3. Example #1: Calculate the molecular formula of a compound that has a molecular molar mass of 283.88 g and an empirical formula of P 2 O 5.

12. Naming and Formula Writing Please see the notes on the wiki page.