1. Atoms, Molecules, and Ions Chapter II A Quick Survey

2. Fundamental Chemical Laws These laws led to the discovery that all matter is composed of atoms (1) Law of Conservation of Mass Frenchman Antoine Lavosier Matter can’t be created or destroyed

3. Fundamental Chemical Laws cont… (2) Law of Definite Proportion (Constant Composition) Frenchman Joseph Proust A compound always contains the same percentage by mass of each element in it. For example all samples of CO 2 will always be 1 part C (12.01 g C) to 2 parts O (32 g O)

4. Fundamental Chemical Laws cont… (3) Law of Multiple Proportions Englishman John Dalton When two elements will form more than one compound those two elements will always combine in whole-number ratios For example: 1 gram of oxygen will combine with 1.750, 0.875, or 0.4375 grams of nitrogen

5. Dalton’s Atomic Theory Elements are made of tiny particles called atoms All atoms of a given element are identical, atoms of different elements are different Chemical compounds form when atoms of different elements combine with one another A chemical reaction merely involves the rearrangement of atoms, hence mass is conserved

6. Atomic Theory History The electron Discovered by J.J. Thomson during experiments with cathode ray tubes Cathode Ray tube The same results were obtained regardless of various metals being used as the electrodes. The same amount of deflection Since all atoms are neutral he proposed there must be an equal amount of positive charge in atoms But he did not discover the source of positive charge Plum-pudding model of the atom

7. Atomic Theory History cont… The source of positive charge was discovered by Thomson’s student Ernest Rutherford. Rutherford is best remembered for his Gold – Foil Experiment Demonstration of the Experiment

8. Atomic Theory History cont… Millikan’s Oil Drop Experiment allowed for the calculation of the charge of an individual electron The Oil Drop Experiment By varying the voltage in the two plates Millikan could determine the amount of charge needed to counterbalance the force of gravity. Using this information with the known charge to mass ratio of the electron, he could calculate the mass of an electron. 9.11 x 10 -31 kg

9. The Modern View of Atomic Structure The atom contains a tiny nucleus approximately 10 -13 cm in diameter with electrons circling it at a distance of about 10 -8 cm from the center Protons and neutrons make up the nucleus and are held together by the strong nuclear force

10. The Modern View of Atomic Structure Mass of subatomic particles Charge of subatomic particles While all atoms are composed of the same components it is the arrangements of the electrons that give atoms their different chemical properties

11. The Modern View of Atomic Structure Writing the symbol of the elements and counting the number of protons, neutrons, and electrons The same atom with a different number of neutrons and hence mass are known as isotopes

12. Molecules and Ions Atoms are held together in compounds by chemical bonds Chemical bonds fall into 2 major categories (there are others, but for now we will focus on only these two) Covalent bonds – sharing of electrons between atoms resulting in molecules Ionic bonds – resulting from a transfer of electrons from one atom to another. The resulting ions formed are attracted to one another forming an individual unit cell (not a molecule) The compounds formed from ionic bonds are known as salts

13. The periodic table Review of organization Groups Periods Metals Nonmetals Metalloids

14. Nomenclature of inorganic compounds Simple binary ionic compounds Type 1 – when the cation only forms one charge Cations keep their names Anions have the end of their name changed it ide Examples:

15. Nomenclature of inorganic compounds Simple binary ionic compounds Type II – when the cation forms more than one charge Cation must be identified by following it with a Roman numeral to represent the charge on the cation Anion ending is still just change to ide Examples

16. Nomenclature of inorganic compounds Ionic compounds that contain polyatomic ions The same rules for ionic compounds already introduced still apply The polyatomic ions name is never changed Examples

17. Nomenclature of inorganic compounds Covalent compounds Use prefixes Mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca The first element only gets a prefix if there is more than one The second element always gets a prefix Examples

18. Simple Organic Compounds Hydrocarbons Nothing but carbon and hydrogen Methane, ethane, propane, pentane, hexane, heptane, octane, nonane decane Formula is C n H (2n + 2) All of these fit into a category known as alkanes

19. Simple Organic Compounds Alkenes Hydrocarbon with a double bond C n H 2n Examples Alkynes Hydrocarbon with a triple bond C n H (2n – 2) Examples

20. Hydrocarbon Derivatives Alcohols Hydrocarbons with one of the H’s of the end removed and replaced with OH The OH is called a hydroxyl group and is not to be confused with the hydroxide ion that has a 1 - charge on it. Same name just change the ending to ol Examples

21. Hydrocarbon Derivatives Carboxyllic Acids Aldehydes Ketones Ethers Amines Esters

22. Naming Acids Non- Oxygen Acids Give a prefix hydro and a suffix ic followed by the word acid Examples HCl, HCN, HBr, HI, HF

23. Oxyacids Contain a polyatomic ion If the polyatomic ion end in ate change the ending to ic Examples H 2 SO 4, H 3 PO 4, HC 2 H 3 O 2, HClO 4 If the polyatomic ion ends in ite change the ending to ous Examples H 2 SO 3, H 3 PO 3, HNO 2, HBrO 2

24. Practice 68 – 71 p. 77 End of chapter 2