REVIEW. Copyright (c) 2011 by Michael A. Janusa, PhD. All rights reserved.

Transcription

1 REVIEW 1 Copyright (c) 2011 by Michael A. Janusa, PhD. All rights reserved.

2 Measurement and Significant Figures To indicate the precision of a measured number (or result of calculations on measured numbers), we often use the concept of significant figures. Significant figures are those digits in a measured number (or result of the calculation with a measured number) that include all certain digits plus a final one having some uncertainty (first digit basically guessing). 2 ml 1.5 ml 1 ml cm cm cm 2

3 Rules for Significant Figures: All nonzero digits are significant. i.e SF SF Zeros between significant figures are significant. i.e , SF 20,006 5 SF Zeros preceding the first nonzero digit are not significant. i.e SF SF Zeros to the right of the decimal after a nonzero digit are significant. i.e SF SF SF SF Zeros at the end of a nondecimal number may or may not be significant. (Use scientific notation.) i.e , 2, or 3 SF 3 SF 3

4 Scientific notation is the representation of a number in the form A. x 10 n, where A is a number (sign digits only) with a single nonzero digit to the left of the decimal point and n is an integer or whole number x SF 9.0 x SF 9.00 x SF 300,000,000 (write with 3 SF) 3.00 x SF x SF x SF 4.21 x SF 6.39 x x Note: exp or EE represents x 10 4

5 Number of significant figures refers to the number of digits reported for the value of a measured or calculated quantity, indicating the precision of the value. [Basically means if all quantities have X sign fig can t report final answer with more than X sign figs: measurement or calculation dictates sign figs.] When multiplying and dividing measured quantities, give as many significant figures as the least found in the measurements used. 2.1 x 3.52 = = 7.4 Which gets us to rounding: left most digit to be dropped 5 or greater add 1 to last digit to be retained, less than five leave alone Multiple step calculation - Guard digit: When adding or subtracting measured quantities, give the same number of decimals as the least found in the measurements used (3 sign) (5 sign) (4 sign) arithmetic rules if combined ( ), x /, + - 5

6 = = x x 10-1 = = 3.58 x x 5.8 = = = or 14.9 gives x ( ) = x (0.37) = = = = = =

7 Measurement and Significant Figures (cont d) An exact number is a number that arises when you count items or when you define a unit (conversion 12 in = 1 ft). For example, when you say you have nine coins in a bottle, you mean exactly nine ( infinite). When you say there are twelve inches in a foot, you mean exactly twelve. Note that exact numbers have no effect on significant figures in a calculation. HW 1 7

8 The Periodic Table Metals, Nonmetals, and Metalloids generally, left of staircase metals, staircase metalloids, right of staircase nonmetals. This is important for determining bond type, using proper terminology, and making decisions. metals nonmetals 8

9 Chemical Formulas; Molecular and Ionic Substances involves ionic bond transfer electrons between atoms attraction between charged particles typically metal/nonmetal or polyatomic ions Molecular substances A molecule is a definite group of atoms that are chemically bonded together through sharing of electrons (covalent bonding, generally nonmetalnonmetal including H). A molecular substance is a substance that is composed of molecules, all of which are alike. A molecular formula gives the exact number of atoms of elements in a molecule (i.e. C 2 H 6 O). Structural formulas show how the atoms are bonded to one another in a molecule. i.e. ethanol, CH 3 CH 2 OH involves covalent bond share electrons between atoms typically nonmetal/nonmetal Na + Cl - C : C 9

10 Ionic substances Although many substances are molecular, others are composed of ions (charged particles, transfer of electrons, ionic bonding, generally metal-nonmetal). An ion is an electrically charged particle obtained from an atom or chemically bonded group of atoms by adding or removing electrons. Sodium chloride is a substance made up of ions. Na + 1e - Cl - 10

11 Chemical Formulas; Molecular and Ionic Substances Ionic substances The formula of an ionic compound is written by giving the smallest possible whole-number ratio of different ions in the substance. The formula unit of the substance is the group of atoms or ions explicitly symbolized by its formula. Covalent bond (share e - ) Ionic bond (transfer e - / attraction charged particles C : O nm nm Molecules m nm and charged ions Formula unit Na + Cl - Molecular substance Molecular formula Ionic substance formula 11

12 Ionic substances When an atom gains extra electrons, it becomes a negatively charged ion, called an anion (more electrons than protons). An atom that loses electrons becomes a positively charged ion, called a cation (more protons than electrons). An ionic compound is a compound composed of cations and anions. ionic or molecular; formula unit or molecule; ionic or covalent bonds involved? NaCl CaBr 2 Na 2 SO 4 CO 2 ionic substance; formula unit; ionic bond ionic substance; formula unit; ionic bonds ionic substance; formula unit; ionic and covalent bonds molecular substance; molecule; covalent bonds 12

13 Ions in Aqueous Solution Many ionic compounds (ionic bond/m-nm) dissociate into independent ions when dissolved in water NaCl (s) Na + (aq) + Cl - (aq) Soluble salt Soluble ionic compounds dissociate 100% - referred to as strong electrolytes breaks into charged particles 13

14 Ions in Aqueous Solution Most molecular (covalent bond/nm-nm) compounds dissolve but do not dissociate into ions, exception acids. C 6 H 12 O 6 (s) C 6 H 12 O 6 (aq) These compounds are referred to as nonelectrolytes; no charged particles; soluble but no ions formed. How would the sodium sulfate dissolve? Na 2 SO 4 (s) 2Na + (aq) + SO 4 2- (aq) 14

15 Chemical Substances; Formulas and Names Ionic compounds Most ionic compounds contain metal and nonmetal atoms; for example, NaCl. You name an ionic compound by giving the name of the cation followed by the name of the anion. Sodium chloride, NaCl Calcium Iodide, CaI 2 Potassium Bromide, KBr A monatomic ion is an ion formed from a single atom. 15

16 How get charge for ions? Rules for predicting charges on monatomic ions Most of the main group metals form cations with the charge equal to their group number. The charge on a monatomic anion for a nonmetal equals the group number minus 8. Most transition elements form more than one ion, each with a different charge (exceptions Cd 2+, Zn 2+, Ag + ). Other important elements with variable charge Pb 4+, Pb 2+ Sn 4+, Sn 2+ As 5+, As 3+ Sb 5+, Sb varies

17 Rules for naming monatomic ions Monatomic cations are named after the element. For example, Al 3+ is called the aluminum ion. If there is more than one cation of an element (charge), a Roman numeral in parentheses denoting the charge on the ion is used. This often occurs with transition elements. Na + sodium ion Fe 2+ iron (II) ion Ca 2+ calcium ion Fe 3+ iron (III) ion Older name: higher ox state (charge) ic, / lower, -ous Fe 3+ ferric ion Fe 2+ ferrous ion Cu 2+ cupric ion Cu + cuprous ion Hg 2+ mercuric ion Hg 2 2+ mercurous ion The names of the monatomic anions use the stem name of the element followed by the suffix ide. For example, Br - is called the bromide ion. Br bromine 17

18 The formula of an ionic compound is written by giving the smallest possible whole-number ratio of different ions in the substance. Sodium chloride Na + Cl - NaCl Iron (III) sulfate Fe 3+ SO 4 2- Fe 2 (SO 4 ) 3 Chromium (III) oxide Cr 3+ O 2- Cr 2 O 3 Calcium nitrate Ca 2+ NO 3 - Ca(NO 3 ) 2 Sodium phosphate Na + PO 4 3- Na 3 PO 4 Strontium oxide Sr 2+ O 2- SrO 18

19 Naming Binary Compounds NaF - sodium fluoride LiCl - lithium chloride MgO - magnesium oxide MnBr 2 - manganese (II) bromide Co 2 O 3 - cobalt (III) oxide CuCl 2 - copper (II) chloride or cupric chloride 19

20 Chemical Substances; Formulas and Names Polyatomic ions A polyatomic ion is an ion consisting of two or more atoms chemically bonded together and carrying a net electric charge. Table in book lists some common polyatomic ions. Most are oxo anions consists of oxygen with another element (central element). NO 3 - nitrate SO 4 2- sulfate NO - 2 nitrite SO 2-3 sulfite Most groups ate, -ite differ by O Mn, Br, Cl, I per- -ate, -ate, -ite, hypo- -ite 20

21 Ions You Should Know Polyatomic ions NH Ammonium OH - - Hydroxide CN - - Cyanide SO Sulfate SO Sulfite ClO perchlorate ClO chlorate ClO chlorite ClO - - hypochlorite Hg mercury (I) or mecurous S 2 O thiosulfate SCN - - thiocyanate CNO - - cyanate MnO permanganate O Peroxide PO Phosphate PO Phosphite CO Carbonate HCO Bicarbonate or Hydrogen Carbonate N azide NO nitrate NO nitrite C 2 H 3 O acetate Cr 2 O dichromate CrO chromate C 2 O oxalate HSO bisulfate or hydrogen sulfate H 2 PO dihydrogen phosphate 21

22 SnSO 4 Na 2 SO 3 Ca(ClO) 2 Ba(OH) 2 KClO 4 Cr 2 (SO 4 ) 3 tin (II) sulfate or stannous sulfate sodium sulfite calcium hypochlorite barium hydroxide potassium perchlorate chromium (III) sulfate Mg 3 N 2 Fe 3 (PO 4 ) 2 magnesium nitride iron (II) phosphate or ferrous phosphate Ti(NO 3 ) 4 titanium (IV) nitrate 22

23 Chemical Substances; Formulas and Names molecular compounds Binary compounds composed of two nonmetals are usually molecular and are named using a prefix system (name same as ionic except must indicate how many atoms are present using mono, di, tri, etc.). No charges involved with molecular compounds but we typically put more metallic compound first. NF 3 F 3 N 23

24 Chemical Substances; Formulas and Names Binary molecular compounds The name of the compound has the elements in the order given in the formula. You name the first element using the exact element name. Name the second element by writing the stem name of the element with the suffix ide. If there is more than one atom of any given element, you add a prefix (di, tri, tetra, penta, hexa, hepta, octa, etc.) 24

25 Binary molecular compounds N 2 O 3 dinitrogen trioxide SF 4 sulfur tetrafluoride ClO 2 chlorine dioxide SF 6 sulfur hexafluoride Cl 2 O 7 dichlorine heptoxide HCl (g) hydrogen chloride Name this compound but think about bonding: MgCl 2 magnesium chloride; ionic bond no prefix Older names: water - H 2 O, ammonia NH 3, hydrogen sulfide H 2 S, nitric oxide NO, hydrazine N 2 H 4 25

26 Acids Chemical Substances; Formulas and Names Acids are traditionally defined as compounds with a potential H + as the cation. Binary acids consist of a hydrogen ion and any single anion. For example, HCl (aq) is hydrochloric acid. An oxoacid is an acid containing hydrogen, oxygen, and another element. An example is HNO 3, nitric acid. 26

27 oxoacids Anion prefix/suffix per- -ate ion -ate ion -ite ion hypo- -ite ion NO - 3 nitrate ion NO - 2 nitrite ion ClO - 4 perchlorate ion acid prefix/suffic per- -ic acid -ic acid -ous acid hypo- -ous acid HNO 3 nitric acid HNO 2 nitrous acid HClO 4 perchloric acid SO 4 2- sulfate ion H 2 SO 4 sulfuric acid PO 4 3- phosphate ion H 3 PO 4 phosphoric acid HW 2 27

28 Molecular Weight and Formula Weight, Molar Mass The molecular weight of a substance is the sum of the atomic weights of all the atoms in a molecule of the substance. For, example, a molecule of H 2 O contains 2 hydrogen atoms (at 1.01 amu each) and 1 oxygen atom (16.00 amu), giving a molecular weight of amu. Molecular wt mass one molecule or do for 1 mole of substance called molar mass: g H 2 O/mol H 2 O 28

29 Working with Solutions Molar Concentration When we dissolve a substance in a liquid, we call the substance the solute (being dissolved) and the liquid the solvent (doing the dissolving). The general term concentration refers to the quantity of solute in a standard quantity of solution. There are many concentration terms but we will concentrate on one. 29

30 Working with Solutions Molar Concentration Molar concentration, or molarity (M), is defined as the moles of solute dissolved in one liter (cubic decimeter) of solution. Molarity (M) moles liters of of solute solution solute + solvent volume 30

31 Working with Solutions Molar Concentration The molarity of a solution and its volume are inversely proportional. Therefore, adding water makes the solution less concentrated. Most of time will be using a stock solution and diluting to new concentration. Basically using C c V c C d V d So, as water is added, increasing the final volume, V f, the final molarity, M f, decreases. Thing to realize here is that M x V = mols: want new concentration of substance take mols and divide by total volume 31

32 Mixture example A solution is prepared by mixing 12.9 ml of M HCl and 56.7 ml of Na M HCl, 2 SO then 4 add ml of water. Assuming the liquid volumes are additive, calculate the molarity of HCl in the resulting solution. mol x ml = L mmol = mol tot ml L mol mmols of HCl (12.9 ml) x (0.245 HCl) (56.7 ml ) L 3.161mmol HCl mmol HCl mmol HCl x (0.847 mol L HCl) mmol mmol HCl M HCl 12.9 ml 56.7 ml ml mmol HCl M HCl ml M HCl HW 3 32

33 Solubility Rules for Ionic Compounds (Dissociates 100%) 1.) All compounds containing alkali metal cations and the ammonium ion are soluble. 2.) All compounds containing NO 3 -, ClO4 -, ClO3 -, and C2 H 3 O 2 - anions are soluble. 3.) All chlorides, bromides, and iodides are soluble except those containing Ag +, Pb 2+, or Hg ) All sulfates are soluble except those containing Hg 2 2+, Pb 2+, Ba 2+, Sr 2+, or Ca 2+. Ag 2 SO 4 is slightly soluble. 5.) All hydroxides are insoluble except compounds of the alkali metals and Ca 2+, Sr 2+, and Ba 2+ are slightly soluble. 6.) All other compounds containing PO 4 3-, S 2-, CO 3 2-, CrO 4 2-, SO 3 2- and most other anions are insoluble except those that also contain alkali metals or NH 4+. Generally, compound dissolves > 0.10 M - soluble (aq) < 0.01 M - insoluble (s) in between - slightly soluble Hg 2 Cl 2 (s) KI (aq) Pb(NO 3 ) 2 (aq) insoluble soluble soluble (this class we will assume slightly soluble as soluble) 33

34 Strong Acids (Ionizes 100%) HCl, HBr, HI, HClO 4, HNO 3, H 2 SO 4 Strong Bases (Dissociates 100%) NaOH, KOH, LiOH, Ba(OH) 2, Ca(OH) 2, Sr(OH) 2 34

35 Ions in Aqueous Solution Molecular and Ionic Equations A molecular/formula unit equation is one in which the reactants and products are written as if they were molecules/formula units, even though they may actually exist in solution as ions. Calcium hydroxide + sodium carbonate M.E. Ca(OH) 2 (aq) strong base + Na 2 CO 3 CaCO NaOH (aq) (s) (aq) soluble salt insoluble salt strong base s solid l liquid aq aqueous (acid/bases and soluble salts dissolve in water) g gases 35

36 An total ionic equation, however, represents strong electrolytes as separate independent ions. This is a more accurate representation of the way electrolytes behave in solution. M.E. A complete ionic equation is a chemical equation in which strong electrolytes (such as soluble ionic compounds, strong acids/bases) are written as separate ions in solution. (note: g, l, insoluble salts (s), weak acid/bases do not break up into ions) Ca(OH) 2 (aq) + Total ionic Ions in Aqueous Solution Molecular and Ionic Equations Na 2 CO 3 (aq) CaCO 3 (s) + 2 NaOH (aq) strong base soluble salt insoluble salt strong base Ca 2+ (aq) + 2OH - (aq) + 2Na + (aq) + CO 3 2- (aq) CaCO 3 (s) + 2Na + (aq) + 2OH - (aq) 36

37 Net ionic equations. A net ionic equation is a chemical equation from which the spectator ions have been removed. A spectator ion is an ion in an ionic equation that M.E. does not take part in the reaction. Ca(OH) 2 (aq) + Na 2 CO 3 (aq) CaCO 3 (s) + 2 NaOH (aq) Total Ionic Ca 2+ (aq) + 2OH - (aq) + 2Na + (aq) + CO 3 2- (aq) CaCO 3 (s) + 2Na + (aq) + 2OH - (aq) Net Ca 2+ (aq) + CO 3 2- (aq) CaCO 3 (s) 37

38 Types of Chemical Reactions Oxidation-Reduction Reactions (Redox rxn) Oxidation-reduction reactions involve the transfer of electrons from one species to another. Oxidation is defined as the loss of electrons. Reduction is defined as the gain of electrons. Oxidation and reduction always occur simultaneously. 38

39 27.1 Reduction and Oxidation Redox reactions transfer of e - reduction oxidation reactions Reduction gain of e - / gain of H / lost of O Fe e - Fe 2+ (lower ox state) note: must balance atoms and charges 39

40 Oxidation - loss of e - / loss of H / gain of O Fe 2+ Fe e - (higher ox state) H 2 O + BrO - 3 BrO H + + 2e - (Br oxidized: charge 5+ 7+) Br + 3(-2) = -1 Br = = +5 Br + 4(-2) = -1 Br = = +7 2H + + 2e - H 2 (H reduced: charge 1+ 0) Oxidizing agent is species that undergoes reduction. Reducing agent is species that undergoes oxidation. Note: need both for reaction to happen; can t have something being reduced unless something else is being oxidized. 40

41 27.3 Balancing Redox Reactions - Must know charges (oxidation numbers) of species including polyatomic ions - Must know strong/weak acids and bases - Must know the solubility rules Oxidation Numbers hypothetical charge assigned to the atom in order to track electrons; determined by rules. 41

42 Rules to balance redox 1.) Convert to net ionic form if equation is originally in molecular form (eliminate spectator ions). 2.) Write half reactions. 3.) Balance atoms using H + / OH - / H 2 O as needed: acidic: H + / H 2 O put water on side that needs O or H (comes from solvent) basic: OH - / H 2 O put water on side that needs H but if there is no H involved then put OH - on the side that needs the O in a 2:1 ratio 2OH - / H 2 O balance O with OH, double OH, add 1/2 water to other side. 4.) Balance charges for half rxn using e -. 5.) Balance transfer/accept number of electron in whole reaction. 6.) Convert equation back to molecular form if necessary (re-apply spectator ions). 42

43 Zn (s) + AgNO 3(aq) Zn(NO 3 ) 2(aq) + Ag (s) Total ionic: Zn(s) + Ag + (aq) + NO 3- (aq) Zn 2+ (aq) + 2NO 3- (aq) + Ag(s) Net ionic: Zn(s) + Ag + (aq) Zn 2+ (aq) + Ag(s) 43

44 Net: Zn (s) + Ag + (aq) Zn 2+ (aq) + Ag (s) Oxidation: Zn(s) Zn 2+ (aq) + 2e - Reduction: [ ] 2 1e - + Ag + (aq) Ag(s) Balanced net: Zn(s) + 2 Ag + (aq) Zn 2+ (aq) + 2 Ag(s) Balanced eq: Zn (s) + 2 AgNO 3(aq) Zn(NO 3 ) 2(aq) + 2 Ag (s) 44

45 Net: H + MnO 4 - (aq) + Fe 2+ (aq) Mn 2+ (aq) + Fe 3+ (aq) Ox: [ Fe 2+ (aq) Fe 3+ (aq) + 1e - ] 5 Red: 5e H + (aq) + MnO 4- (aq) Mn 2+ (aq) + 4 H 2 O(l) Balanced net: 8 H + (aq) + MnO 4- (aq) + 5 Fe 2+ (aq) Mn 2+ (aq) + 5 Fe 3+ (aq) + 4 H 2 O(l) 45

46 KMnO 4(aq) + NaNO 2(aq) + HCl (aq) NaNO 3(aq) + MnCl 2(aq) + KCl (aq) + H 2 O (l) Net: MnO - 4 (aq) + NO - Mn 2+ 2 (aq) + H + (aq) NO - 3 (aq) + (aq) + H 2 O(l) Ox: [ H 2 O(l) + NO 2- (aq) NO 3- (aq) + 2 H + (aq) + 2 e - ] 5 Red: [ 5 e H + (aq) + MnO 4- (aq) Mn 2+ (aq) + 4 H 2 O(l) ] 2 Balanced net: 2 MnO 4- (aq) + 5 NO 2- (aq) + 16 H + (aq) + 5 H 2 O(l) 2Mn 2+ (aq) + 8 H 2 O(l) + 5 NO 3- (aq) +10 H + (aq) 2 MnO 4- (aq) + 5 NO 2- (aq) + 6 H + (aq) 2Mn 2+ (aq) + 3 H 2 O(l) + 5 NO 3- (aq) Balanced eq: 2 KMnO 4 (aq) + 5 NaNO 2 (aq) + 6 HCl(aq) 2MnCl 2 (aq) + 3 H 2 O(l) + 5 NaNO 3 (aq) + 2 KCl 46

47 Net: OH - CrI 3 (s) + Cl 2 (g) CrO 4 2- (aq) + IO 4 - (aq) + Cl - (aq) Ox: [ 32 OH - (aq) + CrI 3 (s) CrO 2-4 (aq) + 3 IO 4- (aq) + 16 H 2 O(l) + 27 e - ] 2 Red: [ 2 e - + Cl 2 (g) 2 Cl - (aq) ] 27 Balanced net: 64 OH - (aq) + 2 CrI 3 (s) + 27 Cl 2 (g) 2 CrO 4 2- (aq) + 6 IO 4- (aq) + 54 Cl - (aq) + 32 H 2 O(l) HW 4 47