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REVIEW

REVIEW. 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.

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REVIEW

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  1. REVIEW

  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).

  3. Rules for Significant Figures: • All nonzero digits are significant. i.e. 111 1286 • Zeros between significant figures are significant. i.e. 1001 20,006 • Zeros preceding the first nonzero digit are not significant. i.e. 0.0002 0.00206 • Zeros to the right of the decimal after a nonzero digit are significant. i.e. 0.00300 9.00 9.10 90.0 • Zeros at the end of a nondecimal number may or may not be significant. (Use scientific notation.) i.e. 900 900.

  4. Scientific notation – is the representation of a number in the form A. x 10n, 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. 900 300,000,000 0.0000301 843.4 0.00421 6.39 x 10-4 3.275 x 102 Note: exp or EE represents “x 10”

  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 = 7.392 = 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 – 1.2143 -- 1.21 • Multiple step calculation - Guard digit: 1.214 • When adding or subtracting measured quantities, give the same number of decimals as the least found in the measurements used. • 84.2 (3 sign) • +22.321 (5 sign) • 106.521 • 106.5 (4 sign) arithmetic rules if combined ( ), x / , + -

  6. 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 (9.00000…. - 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. 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.

  8. Chemical Formulas; Molecular and Ionic Substances • A molecule is a definite group of atoms that are chemically bonded together through sharing of electrons (covalent bonding, generally nonmetal-nonmetal including H). • Molecular substances • 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. C2H6O). • Structural formulas show how the atoms are bonded to one another in a molecule. • i.e. ethanol, CH3CH2OH

  9. Although many substances are molecular, others are composed of ions (charged particles, transfer of electrons, ionic bonding, generally metal-nonmetal). • Ionic substances • 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.

  10. 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.

  11. When an atom gains extra electrons, it becomes a negatively charged ion, called an anion (more electrons than protons). • Ionic substances • 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. • NaCl • CaBr2 • Na2SO4 • CO2

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

  13. Ions in Aqueous Solution • Most molecular (covalent bond/nm-nm) compounds dissolve but do not dissociate into ions, exception acids. C6H12O6 (s)  C6H12O6 (aq) These compounds are referred to as nonelectrolytes; no charged particles; soluble but no ions formed.

  14. 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, CaI2 • Potassium Bromide, KBr • A monatomic ion is an ion formed from a single atom. • Table in book lists some common monatomic ions of the main group elements.

  15. Most of the main group metals form cations with the charge equal to their group number. How get charge for ions? Rules for predicting charges on monatomic ions • 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 Cd2+, Zn2+, Ag+?). • Other important elements with variable charge • Pb4+, Pb2+ Sn4+, Sn2+ As5+, As3+ Sb5+, Sb3+

  16. Monatomic cations are named after the element. For example, Al3+ 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 Ca2+ calcium ion • Fe2+ iron (II) ion Fe3+ iron (III) ion • Older name: higher ox state (charge) – ic, / lower, -ous • Fe3+ ferric ion Fe2+ ferrous ion Cu2+ cupric • Cu+ cuprous ion Hg2+ mercuric ion Hg22+ 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 • Rules for naming monatomic ions

  17. 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- Iron (III) sulfate Fe3+ SO42- Chromium (III) oxide Cr3+ O2- Calcium nitrate Ca2+ NO3- Sodium phosphate Na+ PO43- Strontium oxide Sr2+ O2-

  18. Naming Binary Compounds • NaF - - Lithium chloride • MgO - • MnBr2 - - Cobalt (III) oxide - Copper (II) chloride

  19. Chemical Substances; Formulas and Names • 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). • Polyatomic ions Most groups –ate, -ite differ by O Mn, Br, Cl, I per- -ate, -ate, -ite, hypo- -ite

  20. O22- - Peroxide PO43- - Phosphate PO33- - Phosphite CO32- - Carbonate HCO3- - Bicarbonate or Hydrogen Carbonate N3- - azide NO3- - nitrate NO2- - nitrite C2H3O2- - acetate Cr2O72- - dichromate CrO42- - chromate C2O42- - oxalate HSO4- - bisulfate or hydrogen sulfate H2PO4- - dihydrogen phosphate Ions You Should Know Polyatomic ions • NH4+ - Ammonium • OH- - Hydroxide • CN- - Cyanide • SO42- - Sulfate • SO32- - Sulfite • ClO4- - perchlorate • ClO3- - chlorate • ClO2- - chlorite • ClO- - hypochlorite • Hg22+ - mercury (I) or mecurous • S2O32- - thiosulfate • SCN- - thiocyanate • CNO- - cyanate • MnO4- - permanganate

  21. SnSO4 sodium sulfite Ca(ClO)2 barium hydroxide potassium perchlorate Cr2(SO4)3 magnesium nitride Fe3(PO4)2 titanium (IV) nitrate

  22. 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.

  23. 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. Table in book lists the Greek prefixes used.

  24. N2O3 • SF4 • chlorine dioxide • sulfur hexafluoride • Cl2O7 • HCl (g) • Older names: water - H2O, ammonia – NH3, • hydrogen sulfide – H2S, nitric oxide – NO, hydrazine – N2H4 • Binary molecular compounds

  25. Chemical Substances; Formulas and Names • Acids • 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 is hydrochloric acid. • An oxoacid is an acid containing hydrogen, oxygen, and another element. An example is HNO3, nitric acid. • Table in book lists some oxoanions and their oxoacids.

  26. oxoacids Anion prefix/suffixacid prefix/suffic per- -ate ion per- -ic acid -ate ion -ic acid -ite ion -ous acid hypo- -ite ion hypo- -ous acid NO3- nitrate ion HNO3 nitric acid NO2- nitrite ion HNO2 nitrous acid ClO4- perchlorate ion HClO4 perchloric acid HW 2

  27. 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 H2O contains 2 hydrogen atoms (at 1.01 amu each) and 1 oxygen atom (16.00 amu), giving a molecular weight of 18.02 amu. • Molecular wt – mass one molecule or • do for 1 mole of substance called molar mass: 18.02 g H2O/mol H2O

  28. Working with SolutionsMolar 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. Working with SolutionsMolar Concentration • Molar concentration, or molarity (M), is defined as the moles of solute dissolved in one liter (cubic decimeter) of solution.

  30. Working with SolutionsMolar 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 • So, as water is added, increasing the final volume,Vf, the final molarity, Mf, decreases. Thing to realize here is that M x V = mols: want new concentration of substance take mols and divide by total volume

  31. Mixture example • A solution is prepared by mixing 12.9 mL of 0.245 M HCl and 56.7 mL of 0.847 M HCl, then add 630.4 mL of water. Assuming the liquid volumes are additive, calculate the molarity of HCl in the resulting solution. HW 3

  32. Solubility Rules for Ionic Compounds (Dissociates 100%) 1.) All compounds containing alkali metal cations and the ammonium ion are soluble. 2.) All compounds containing NO3-, ClO4-, ClO3-, and C2H3O2- anions are soluble. 3.) All chlorides, bromides, and iodides are soluble except those containing Ag+, Pb2+, or Hg22+. 4.) All sulfates are soluble except those containing Hg22+, Pb2+, Ba2+, Sr2+, or Ca2+. Ag2SO4 is slightly soluble. 5.) All hydroxides are insoluble except compounds of the alkali metals and Ca2+, Sr2+, and Ba2+ are slightly soluble. 6.) All other compounds containing PO43-, S2-, CO32-, CrO42-, SO32- and most other anions are insoluble except those that also contain alkali metals or NH4+. Generally, compound dissolves > 0.10 M - soluble (aq) < 0.01 M - insoluble (s) in between - slightly soluble (this class we will assume slightly soluble as soluble)

  33. Strong Acids (Ionizes 100%) HCl, HBr, HI, HClO4, HNO3, H2SO4 Strong Bases (Dissociates 100%) NaOH, KOH, LiOH, Ba(OH)2, Ca(OH)2, Sr(OH)2

  34. Ions in Aqueous SolutionMolecular 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 F.U.

  35. Ions in Aqueous SolutionMolecular and Ionic Equations • 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. • 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) F.U. Ca(OH)2 (aq) + Na2CO3 (aq)  CaCO3 (s) + 2 NaOH (aq) Total ionic

  36. A net ionic equation is a chemical equation from which the spectator ions have been removed. • A spectator ion is an ion in an ionicequationthat does not take part in the reaction. Net ionic equations. F.U. Ca(OH)2 (aq) + Na2CO3 (aq)  CaCO3 (s) + 2 NaOH (aq) Total Ionic Ca2+(aq) + 2OH-(aq) + 2Na+(aq) + CO32- (aq)  CaCO3(s) + 2Na+(aq) + 2OH-(aq) Net

  37. 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. Redox reactions – transfer of e- • reduction – oxidation reactions • Reduction – gain of e- / gain of H / lost of O • Fe3++ 1e- Fe2+ (lower ox state) • note: must balance atoms and charges

  39. Oxidation - loss of e- / loss of H / gain of O Fe2+ Fe3++ 1e- (higher ox state) H2O + BrO3-  BrO4- + 2H++ 2e- (Br oxidized: charge 5+  7+) 2H++ 2e-  H2 (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. Balancing Redox equations • - 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. 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- / H2O as needed: • acidic: H+ / H2O put water on side that needs O or H (comes from solvent) • basic: OH- / H2O 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- / H2O 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. FU: Zn(s) + AgNO3(aq) Zn(NO3)2(aq) + Ag(s) • Total ionic: • Net ionic:

  43. Net: Zn(s) + Ag+(aq) Zn2+(aq) + Ag(s) • Ox: • Red: • Balanced net: • Balanced FU:

  44. H+ • Net: MnO4-(aq) + Fe2+(aq) Mn2+(aq) + Fe3+(aq) • Ox: • Red: • Balanced net:

  45. FU: • KMnO4(aq) + NaNO2(aq) + HCl(aq) NaNO3(aq) + MnCl2(aq) + KCl(aq) + H2O(l) • Net: • Ox: • Red: • Balanced net: • Balanced FU:

  46. FU:OH- • CrI3 (s) + Cl2 (g) CrO42-(aq) + IO4-(aq) + Cl-(aq) • Ox: • Red: • Balanced net: HW 4

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