I. Oxidation Numbers II. Nomenclature III. The Mole

Transcription

1 I. Oxidation Numbers II. Nomenclature III. The Mole 1

2 I. Oxidation Numbers 2

3 A positive or negative whole number assigned to an element in a molecule or ion on the basis of a set of formal rules; to some degree it reflects the positive or negative character of an atom 1. Keep track of electrons 2. Tell if electrons gained, lost, or unequally shared 3. Allows us to predict formulas of chemical compound Not to be confused with oxidation, a process in which a substance loses one or more electrons. 3

4 1. Electronegativity Higher EN = negative oxidation number Lower EN = positive oxidation number 4

5 Remember always refer back to: Electronegativity Electrostatic forces Bonding Characteristics There will be exceptions to the following rules 5

6 1. Oxidation # of free atoms, pure elements, and polyatomic elements is ZERO Both atoms have equal EN, no transfer or shift of electrons 6

7 2. Oxidation # of monatomic ion is equal to the charge of the ion If an atom loses 2e-, the other atom(s) must gain those 2e-. Electrons DO NOT just float around. 7

8 3. The sum of all oxidation numbers in a compound must be ZERO. Compounds are not electrically charged! 8

9 4. Alkali metals always have a +1 oxidation # when not free Hydrogen is not an alkali metal although it is in group 1 9

10 5. Alkaline Earth Metals always have a +2 oxidation number. They form +2 ions when they bond. 10

11 6a. Certain elements have the same oxidation # in almost all their compounds. Halogens have oxidation number -1 when bond to metals Halogen with higher EN than other bonded nonmental is assigned the negative number 11

12 6b. Hydrogen assigned a +1 oxidation # in most compounds BUT.. Hydrogen + metal = metallic hydrides Hydrogen has a -1 oxidation number Hydrogen more EN than any other metal 12

13 6c. Oxygen has an -2 oxidation number in most compounds Oxygen is bonded to highly EN elements does not have -2 oxidation number Oxygen is VERY EN and pulls e- from most other elements Exception: Perioxide ion O 2 2- where O has -1 O.N. 13

14 7. Oxidation number of all atoms in a polyatomic ion add up to the charge of the ion 14

15 Follow rules in order If rules contradict each other, the rule listed first should be followed Write an algebraic equation to solve for unknown oxidation numbers in compounds Ionic Compounds Criss-Cross method Use charge of one ion as the subscript for the other Simplify ratios of atoms Exception: Peroxide ion O 2-2 NaO example: Na 2 O 2, not 15

16 II. Nomenclature 16

17 IUPAC developed a systematic way to name compounds. Names reveal the composition and qualities of certain substances Indicate the types of bonds and intermolecular attractions Covalent Compounds Binary Ionic Compounds 17

18 Composed of ONLY nonmetals Two word names Use Greek Prefix System Least EN element first, then more EN element Ending of last element is changed to -ide. Number Prefix 1 Mono* 2 Di 3 Tri 4 Tetra 5 Penta 6 Hexa 7 Hepta 8 Octa 9 Nona 10 deca *omit mono- for first atom 18

19 1. Antimony tribromide 2. Hexaboron silicide 3. Chlorine dioxide 4. Iodide pentafluoride 5. P 4 S 5 6. Si 2 Br 6 7. CH 4 8. NF 3 SbBr 3 B 6 Si ClO 2 IF 5 Tetraphorsphorus pentasulfide Disilicon hexabromide Methane Nitrogen trifluoride 19

20 Compounds formed between metals and nonmetals! Ionic compounds DO NOT use Greek Prefix System Named according to the two elements or polyatomic ions Positive ions use the same name as their parent atoms (ex. sodium atoms form sodium ions). Named FIRST. Negative ions have an ide ending. Named SECOND. Polyatomic ions made up of 2+ types of atoms with covalent bonds! They act like a single unit. 20

21 Oxyanions: anions composed of oxygen and one other element/ polyatomic ion (table p. 174). 2 Forms of oxyanion: More oxygens = name ends in -ate Less oxygens = name ends in -ite 3+ Forms of oxyanion: Most oxygens = name per- -ate Least oxygens = name hypo- -ite 21

22 Polyatomic compounds: compounds that contain polyatomic ions (see green handout) Same as binary compounds, but: Name the cation first Name the anion second Replace -ide ending with polyatomic ion name If 2 polyatomic ions, use polyatomic ion names Do not use Greek prefix system 22

23 1. MgO 2. K 2 S 3. Na 2 SO 4 4. Ba(ClO 3 ) 2 5. NH 4 Cl 6. K 2 Cr 2 O 7 7. CaSO 4 8. Zn 3 (PO 4 ) 2 Magnesium oxide Potassium sulfide Sodium sulfate Barium chlorate Ammonium chloride Potassium dichromate Calcium sulfate Zinc phosphate 23

24 Stock System (Roman Numeral System): If an element can have more than 1 oxidation state, a roman numeral is placed in parenthesis after the element s name. Transition metals Common Name Use suffix at the end of the first element (metal) Smaller oxidation number -ous Larger oxidation number -ic 24

25 1. Hg 2 I 2 2. CuBr 3. FeCl 2 4. Co 2 (C 2 O 4 ) 3 5. SnO 6. SnO 2 7. PbSO 4 8. Pb(SO 4 ) 2 Mercury (I) iodide Copper (I) bromide Iron (II) chloride Cobalt (III) oxalate Tin (II) oxide Tin (IV) oxide Lead (II) sulfate Lead (IV) sulfate 25

26 Hydrates are compounds that have water molecules in their crystalline structure. Hold water = water of hydration Formulas written followed by a dot and number of water molecules. Compounds name + greek prefix + hydrate Ex. Na 2 CO 3 7H 2 O sodium carbonate heptahydrate Anhydrates = compounds with NO water in their structures 26

27 Binary acids form when binary compounds dissolve in water hydro - acid HCl: hydrogen chloride hydrochloric acid HBr: hydrogen bromide hydrobromic acid H 2 S: hydrogen sulfide hydrosulfuric acid 27

28 Ternary acids contain three elements, generally containing a polyatomic ion, or combination of hydrogen, oxygen, and a nonmetal. Rule1: Addition of 1 oxygen to the acid: Per -ic acid. Rule 2: Subtraction of 1 oxygen from the acid: -ous acid. Rule 3: Subtraction of 2 oxygens from the acid: hypo -ous acid. 28

29 III. The Mole 29

30 It is a counting number (like a dozen) Used to count really, really small things Avagadro s Number (N A ) 1 MOLE = 6.022x10 23 units (4 sig figs) A mole is a amount!!!!! That is why it is used to count very small things, like atoms and molecules 30

31 1 mole of hockey pucks would equal the mass of the moon! 1 mole of basketballs would fill a bag the size of the earth! 1 mole of pennies would cover the Earth ¼ mile deep! 1 mole of sand would fill all the Great Lakes 10 times! 1 mole of popcorn kernels would cover the United States 9 miles deep! 31

32 Remember! 1 Mole of ANYTHING is equal to 23 items x

33 We can use the concept of the mole to solve for problems like: How many copper atoms are in one penny? What is the mass of a single atom? What is the mass of moles of helium atoms? How many atoms are in 33mg of gold? 33

34 Because a mole is so large, measurements aren t counted, they are weighed. Molar mass is the mass (grams) of 1 mole (N A ) of particles of an element or compound 34

35 The same number! Different units! Look at the periodic table Scientists chose Avagadro s # (N A ) to: relate atomic mass units to the larger, more practical unit of GRAMS. Represent the # of particles in a mole so that the atomic mass of an element and mass of a mole of the element have the same numeric value, just different units! (Hydrogen experiment) 35

36 Carbon Aluminum Zinc g/mol g/mol g/mol 36

37 Sodium bicarbonate NaHCO (16.00) = g/mol Use atomic mass from periodic table 37

38 Examples: H 2 O NaCl 2(1.01) = g/mol = g/mol 38

39 P.184 in textbook flow chart Mass (g) Moles Number of units (atoms, etc..) 39

40 How many molecules or atoms are in a certain amount of a substance? How many grams are there in a mole of a substance? How many moles are there in??? grams of a substance? What is the percent composition of a substance? (how much do each of the different types of atoms weigh in the compound?) 40

41 Use dimensional analysis to SOLVE problems Sample Problems: P. 184 P. 185 P. 186 P. 187 P. 189 P. 190 P. 192 P

42 Example: How many moles are in 22 grams of copper metal? In all problems like this, you need to go through four steps to find a solution. 42

43 Step 1: Figure out how many parts in your calculation you will have by using this diagram 43

44 Step 2: Make a T-chart, and put whatever information the problem gave you in the top left. After that, put the units of whatever you were given in the bottom right of the T, and the units of what you want to find in the top right. 44

45 Step 3: Put the conversion factors into the T- chart in front of the units on the right. 45

46 Step 4: Cancel out the units from the top left and bottom right, then find the answer by multiplying all the stuff on the top together and dividing it by the stuff on the bottom. Pau! 46

47 Continue by adding another section in the T-chart repeat steps and there you go. 47

48 Molarity is the amount of a substance dissolved in one liter of solution. Molarity (M) = moles/ liters of solution 48

49 Chemical compounds contain two or more atoms chemically combined to behave as one unit. Masses of compound units can be found by adding the masses of the atoms contained in them. Formula unit = a single unit of a compound NaCl = one unit of sodium chloride 49

50 The mass of a mole of a substance: Gram-atomic mass = mass of a mole of atoms Gram-molecular mass = mass of a mole of molecules Gram-formula mass= mass of a mole of formula units in an ionic compound All have the units g/mol 50

51 Formulas tell us the proportions of atoms in a molecule or ionic compound We can relate # of atoms # of moles Example: gram-molecular mass of NH 3? 1 mole of nitrogen = (1) = g/mol 3 moles of hydrogen = (3) = g/mol = g/mol 51

52 Example: gram-formula mass of Al 2 (SO 4 ) 3? Each formula unit contains two Al, three S, and 12 O. A mole of Al 2 (SO 4 ) 3 consists of 2 moles of Al atoms, 3 moles of S atoms, and 12 moles of O atoms. Find the molar mass g/mol Al 2 (SO 4 ) 3 52

53 Structural Formulas Show the types of atoms Exact composition of each molecule Arrangement of chemical bonds 53

54 Molecular Formula TRUE FORMULA Shows the types of atoms Exact composition of ATOMS in each molecule Does not show shape, location of bonds, or bond type H 2 O = water C 2 H 4 = ethene Cl 2 = chlorine 54

55 Empirical Formulas Tell what elements are present in simple ratios Used for both ionic compounds and molecules Careful when writing empirical formulas for molecules May be the actual molecular composition OR May only show the simplest ratio of atoms in the molecule H 2 O = water Empirical formula = H 2 0 C 2 H 4 = ethene Empirical formula = CH 2 Cl 2 = chlorine Empirical formula = Cl 55

56 Empirical Formulas REDUCE SUBSCRIPTS Example: C 2 H 6 CH 3 1. FIND MASS (OR %) OF EACH ELEMENT 2. Find moles of each element 3. Divide moles by the smallest # to find subscripts 4. When necessary, multiply subscripts by 2,3, or 4 to get whole # s 56

57 Find the empirical formula for a sample of 25.9% N and 74.1% O. 57

58 Empirical formula: N 1 O 2.5 Need to make subscripts whole numbers x2 N 2 O 5 58

59 Mole ratio in the EF mass ratio Example: Water formula H moles hydrogen for every 1 mole of oxygen Express in mass 1 mole of water contains 2.016g of hydrogen atoms and 16.00g of oxygen atoms Convert moles to mass (grams) 2 mole H (1.008g/1 mole H) = 2.016g H 1 mole O (16.00g/ 1 mole O) = 16.00g O Total mass of water = 2.016g g = 18.02g Now find % composition.. 59

60 Molecular Formula: 1. Find the empirical formula 2. Find the empirical formula mass 3. Divide the molecular mass by the empirical mass 4. Multiply each subscript by the answer from step 3 60

61 The empirical formula for ethene is CH2. Find the molecular formula is the molecular mass is 28.1 g/mol. Empirical mass = g/mol (28.1 g/mol)/ (14.03 g/mol) = 2.00 (CH 2 ) 2 C 2 H 4 61

62 Percent composition = the mass composition of a compound All other formulas describe # of atoms %Comp. describes masses of atoms Which atoms make up the most mass in a compound or molecule? 62

63 Percent = per hundred General setup: Part/whole x 100% Example: Lab analysis of 30.00g Al 2 (SO 4 ) g Al (4.731/30.00) x 100% = 15.77% Al 8.433g S (8.433/30.00) x 100% = 28.11% S g O (16.836/ 30.00) x 100% = 56.12% O 63

64 Sample Problem p