Understanding Acids, Bases, and Salts in Chemistry
Dive into the world of acids, bases, and salts with definitions, examples, and explanations of Brønsted-Lowry acids, autoprotolysis, pH scale, strong acids, and weak acids. Explore the fundamentals of chemical reactions and their behavior in aqueous solutions.
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Presentation Transcript
Announcements Quiz 3 Today Today s Lecture Chapter 6 Acids, Bases and Salts Chapter 7 Titrations Overview and Definitions
Acids, Bases and Salts Definitions of Acids and Bases - Lewis Acids/Bases (defined before, most general category) - Br nsted-Lowry Acids/Bases: acid = proton donor base = proton acceptor (must have free electron pair so also is a Lewis base) - definitions are relative
Brnsted-Lowry Acids - examples HCO2H(aq) + H2O(l) acid base conjugate conjugate HCO2- + H3O+ base acid CH3NH2(aq) + H2O(l) base acid conjugate conjugate CH3NH3++ OH- acid base H2SO4+ CH3CO2H(l) acid base HSO4-+ CH3CO2H2+ conjugate conjugate base acid
Brnsted-Lowry Acids Note: for most acids, the reaction with water is simplified: Example: HNO2(nitrous acid) HNO2 H++ NO2-
Autoprotolysis and the pH Scale Autoprotolysis refers to proton transfer in protic solvents like water: H2O(l) H++ OH- K = Kw= [H+][OH-] = 1.0 x 10-14(T = 25 C) In pure water [H+] = [OH-] = Kw0.5= 1.0 x 10-7M pH = -log[H+] = 7.0 Acidic is pH < 7; basic is pH > 7
Strong Acids Strong acids completely dissociate in water (except at very high concentrations) HX(aq) H++ X-(no HX(aq) exists) Ka> 1 Major strong acids: HCl, HNO3, H2SO4 Note: For H2SO4, 1stdissociation is that of a strong acid, but 2nddissociation is that of a weak acid (Ka~ 0.01)
Weak Acids Partially dissociate in water Most have H that can dissociate HX(aq) H++ X- Example: HNO2 Degree of dissociation given by Kavalue Ka= [H+][NO2-]/[HNO2] Metal cations can be acids through the reaction: Mn++ H2O(l) MOH(n-1)++ H+ (although for +1 and some +2 metals the above reactions favor reactants so strongly the metals can be considered neutral ) (HX(aq) exists) H++ NO2-
Ionic Compounds in Water First step should be dissociation to respective ions: example: NaCl(s) Na++ Cl- In subsequent steps, determine how anion/cation react: - anions usually only react as bases - cations may react as acids - see if ions are recognizable conjugate acids or bases - polyprotic acids are somewhat different
Ionic Compounds in Water Conjugate bases of weak acids are basic. NO2-+ H2O(l) HNO2(aq) + OH- Conjugate bases of weaker weak acids are stronger bases. Kb= Kw/Ka CN-is a stronger base than NO2-because Ka(HCN) = 6.2 x 10-10and Ka(HNO2) = 7.1 x 10-3
Acidity of Ionic Compounds Determine if the ionic compounds are acidic or basic in the following examples: 1. NaCl 2. NH4Cl 3. NaCH3CO2 4. Fe(NO3)3 5. NH4CN
Polyprotic Acids Examples: H2SO4, H3PO4, H2C2O4 H3PO4has 3 Kavalues (Ka1, Ka2, Ka3) for 3 reactions losing H+: 1) H3PO4 H2PO4-+ H+ 2) H2PO4- HPO42-+ H+ 3) HPO42- PO43-+ H+ Release more than 1 H+per molecule Ka1 Ka2 Ka3
Chapter 7 - Titrations Introduction Overview A. Chapter 7 covers general titrations (quantitation, practical aspects, types of titrations, shape of precipitation titration curve not covering calculations due to time) B. Chapter 11 covers titration curves for acid- base titrations - covered later C. Other Chapters (12, 16) cover other types of titrations not covered
Titrations Definitions Titrant: Reagent solution added out of buret (concentration usually known) Analyte solution: Solution containing analyte Equivalence Point: point where ratio of moles of titrant to moles of analyte is equal to the stoichiometric ratio titrant analyte solution for: Al3+ + 3C2O42- Al(C2O4)33- n(Al3+)/n(C2O42-) = 1/3 at equivalence pt.
Titrations Practical Requirements The equilibrium constant must be large Size of K value depends on desired precision and concentration of analyte Typically K ~ 106 is marginal, K > 1010 is better The reaction must be fast It must be possible to observe the equivalence point observed equivalence point = end point
Titrations Detection of Endpoints An endpoint is defined as the point in the titration when the equivalence point is observed Ways to detect endpoints: Use of colored reactants example: MnO4- + H2C2O4 (aq) Mn2+ + CO2 (g) PINK Clear Clear Use of indicators An indicator changes color in response to the change in a reactant s concentration Use of simple instruments Must respond quickly, but typical equipment is low cost
Titrations Detection of Endpoints Simple instruments electrodes (typically respond to log of ion concentrations) spectroscopic measurements (measurement of absorption of light) Can improve titration precision vs. using indicators Titration Error = Difference between end point and equivalence point = systematic error Note: It is possible to have large errors or uncertainties in detection of reagent conc. by various methods without having great titration errors to meter
Titrations Other Definitions Standardization vs. Analyte Titrations To accurately determine an analyte s concentration, the titrant concentration must be well known This can be done by preparing a primary standard (high purity standard) Alternatively, the titrant concentration can be determined in a standardization titration (e.g. vs. a known standard) Rationale: many solutions can not be prepared accurately from available standards Example: determination of [H2O2] by titration with MnO4- neither compound is very stable so no primary standard instead, [MnO4-] determined by titration with H2C2O4 in standardization titration then, H2O2 titrated using standardized MnO4-
