1. Chemistry Chapter Three – Atoms: The Building Blocks of Matter South Lake High School Science Department Ms. Sanders

2. Foundation of Chemistry Law of Conservation of Mass – mass is neither created nor destroyed during ordinary chemical reactions or physical changes Law of Definite Proportions – chemical compounds contain the same elements in exactly the same proportions by mass regardless of sample size or source of compound Law of Multiple Proportions – if two or more different compounds are composed of the same two elements, then the ratio of the elements is always a ratio of small whole numbers

3. Foundation  Law of Conservation of Mass

4. Foundation  Law of Multiple Proportions

5. Atomic Theory  Dalton’s Atomic Theory 1.All matter is composed of atoms 2.Atoms of the same element are identical whereas atoms of different elements differ 3.Atoms cannot be subdivided, created, or destroyed 4.Atoms of different elements combine in simple whole-number ratios to form chemical compounds 5.In chemical reactions, atoms are combined, separated, or rearranged  Modern Atomic Theory  Atoms are divisible into smaller particles (quarks, mesons)  One element can have atoms with different masses (isotopes)

6. Atom Structure  Atom – smallest particle of an element that retains the chemical properties of that element  2 Regions 1.Nucleus – make up the mass of the atom  Positive protons  Neutral neutrons 2.Area surrounding Nucleus  Negative Electrons  Subatomic particles = protons, neutrons, and electrons

7. Discoveries  Discovery of electron was a result of electricity and matter investigations  Cathode – ray tubes (cathode - negative, anode – positive) particles that compose cathode rays are negatively charged (electrons)  J.J. Thomson – concluded all cathode rays are composed of identical negatively charged particles – electrons

8. Discoveries  Cathode Ray

9. Discoveries  Robert Millikan measured the charge of an electron, led to ability to calculate the mass of an electron  Ernest Rutherford (associates Geiger and Marsden) bombarded a thin piece of gold foil with fast-moving alpha particles, some particles deflected therefore they concluded that there is a small space with a positive electric charge within atoms – atom mostly empty space with a small, positive center

10. Discoveries  Gold Foil Experiment

11. Discoveries  Gold Foil Experiment

12. Counting Atoms  Atomic number – number of protons of each atom of that element  Found on the periodic table  Atomic # = # Protons  Atomic # = # Electrons  # Electrons = # Protons  Mass number – total number of protons and neutrons that make up the nucleus of an isotope  Found on the periodic table  Mass # = # Protons + # Neutrons  Comes from the nucleus  To find the number of Neutrons: subtract atomic # from mass #  # Neutrons = Mass # - Atomic #  # Neutrons = Mass # - # Protons

13. Counting Atoms  Isotopes – atoms of the same element that have different masses, different number of neutrons { mass number – atomic number = number of neutrons}  2 Methods for designating Isotoptes 1.Hyphen Notation - mass number is written with a hyphen after the name of the element – i.e. Hydrogen – 3 (tritium) 2.Nuclear Symbol – shown the composition of the nucleus – i.e. 92 235 U (the mass number is the top number, the atomic number is the bottom number)  Nuclide – term for a specific isotope of an element  Isotopes of Carbon:  Carbon-12: has 6 neutrons (12 – 6 = 6)  Carbon-13: has 7 neutrons (13 – 6 = 7)

14. Counting Atoms  Isotopes of Carbon

15. Molar Conversions  Atomic mass – 1 amu (atomic mass unit), exactly 1/12 the mass of a carbon-12 atom  Average atomic mass – weighted average of the atomic masses of the naturally occurring isotopes of an element, in grams  Mole (mol) – amount of a substance that contains as many particles as there are atoms in exactly 12 g of carbon-12  Molar mass – mass of one mole of pure substance [Avogadro’s number = 6.022 x 10 23 ], in grams/mole (used as a conversion factor)  Formula mass – mass of a pure substance, in grams

16. Molar Conversions  How to calculate molar mass:  Get the mass from the periodic table  Multiply the mass by the number of atoms  Add up the total mass  Example:  Molar Mass of NO ₂  N 1x14 = 14  O 2x16 = +32  46 g/mole

17. Molar Conversions Dimensional Analysis:  Given x Unknown = Unknown  Page 85:  1) What is the mass in grams of 2.25 moles of Fe? (need to get molar mass from periodic table.)

18. Molar Conversions Dimensional Analysis:  Given x Unknown = Unknown  Page 85:  1) What is the mass in grams of 2.25 moles of Fe? (need to get molar mass from periodic table.)  2.25 mole Fe x 56 g Fe = 126 g Fe 1 mole Fe

19. Molar Conversions  Continue page 85:  1) How many moles of calcium, Ca, are in 5 g of calcium? (need to get the molar mass from the periodic table.)

20. Molar Conversions  Continue page 85:  1) How many moles of calcium, Ca, are in 5 g of calcium? (need to get the molar mass from the periodic table.)  5 g Ca x 1 mole Ca = 1.125 mole Ca 40 g Ca

21. Molar Conversions  Avogadros Number  Avogadro’s number = 6.022 x 10 23  6.022 x 10 23 molecules or atoms or particles = 2 mole  Page 86:  1) How many moles of lead, Pb, are in 1.50 x 10¹² atoms of lead? (need to use Avogadros number to convert)

22. Molar Conversions  Continue page 86:  1) How many moles of lead, Pb, are in 1.50 x 10¹² atoms of lead? (need to use Avogadros number to convert)  1.50 x 10 ¹² atoms Pb x 1 mole Pb = 2.49 x 10ˉ¹² moles Pb 6.02 x 10²³ atoms Pb