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    Precusors to acid Rain ENV 110

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    1 : 1-35
    2 : Acids and Bases! Lab #2 “Class, you may take these notes now! It is late, so you may also look for them on blackboard. Just slides 1-33 Rob!!
    3 : Equilibria in Aqueous Solutions 1 Ionization of Water and the pH Scale 2 Strong and Weak Acids 3 Strong and Weak Bases 4 Solutions of Strong Acids and Bases: Neutralization and Titration 5 Titration and Titration Curves 6 Equilibria with Weak Acids and Bases 7 Indicators
    4 : Equilibria in Aqueous Solutions At Standard Temperature and Pressure (STP), about 25 °C (298 K), Water self ionizes! Kw = [H3O+][OH-] = 1.0×10-14.
    5 : The pH scale .
    6 : The pH scale .
    7 : Ion Selective electrodes .
    8 : Equation for pH
    9 : What is the pH? pH is minus the log of the H+ activity in an aqueous solution where aH is the (dimensionless) activity of hydrogen ions.
    10 : What is the pH? What is the pH of a 0.010 M HCl solution? Since HCl is a strong acid, the hydronium ion concentration will be equal to the HCl concentration: [H3O+] = 0.010 M
    11 : What is the pH? The pH can be found by taking the negative log of the hydronium ion concentration: pH = -log[H3O+] = -log(0.010) = 2.00
    12 : The self-ionization constant The resulting equilibrium constant is called the ionization constant, dissociation constant, or self-ionization constant, or ion product of water and is symbolized by Kw. Kw = Ka [H2O] = Keq [H2O]2 = [H3O+] [OH-] where [H3O+] = molarity of hydrogen or hydronium ion, and [OH-] = molarity of hydroxide ion.
    13 : Let’s Get to the point: No “Song and Dance”: This is the story I was told. Strong acid - Strong base Weak acid - Strong base Strong acid - Weak base Weak acid - Weak base
    14 : Characteristics of Acids Acids have a sour taste Acids turn litmus from blue to red Acids react with bases
    15 : What types of acids? Two types of acids Arrhenius Brønsted-Lowry
    16 : The Brains! Svante Arrhenius(release H+ or OH-) Brønsted-Lowry (conjagate base pairs) G. N. Lewis (Metal Ions) And, Ralph Pearson proposed an advanced qualitative concept known as “Hard Soft Acid Base principle “
    17 : Arrhenius acids and bases Arrhenius Acids: Substances that when placed in water, will dissociate to produce H+ ions. Arrhenius Bases: Substances that when placed in water will dissociate to yield OH- ions.
    18 : Arrhenius acids and bases Arrhenius does not consider acid – base pairing
    19 : Arrhenius Acid Definition An acid is a substance that generates hydronium ions, H3O+ (often described as H+), when added to water. An acidic solution is a solution with a significant concentration of H3O+ ions.
    20 : Arrhenius receiving the Nobel Prize in 1903 .
    21 : Bronstead Lowry systems Introduction to acid-base pair systems The second definition of acids and bases!! Takes into account the concept of a buffer solution
    22 : Bronstead Lowry systems HC2H3O2(aq) + H2O(l) ? C2H3O2–(aq) + H3O+(aq)
    23 : conjugate acid-base pairs acid base conjugate acid conjugate base CO32–(aq) + HC2H3O2(aq) ? C2H3O2–(aq) + HCO3– conjugate acid-base pairs acid base conjugate acid conjugate base H3PO4(aq) + OCl –(aq) ? H2PO4–(aq) + HOCl(aq) conjugate acid-base pairs Bronstead Lowry systems
    24 : Monoprotic and Polyprotic Acids If each molecule of an acid can donate one hydrogen ion, the acid is called a monoprotic acid   If each molecule can donate two of more hydrogen ions, the acid is a polyprotic acid
    25 : Arrehnius acids Hydrochloric acid, HCl(aq) Sulfuric acid, H2SO4 Phosphoric acid, H3PO4
    26 : Diprotic and triprotic acids Sulfuric acid, H2SO4 A diprotic acid that has two acidic hydrogens   Phosphoric acid, H3PO4   A triprotic acid that has three acidic hydrogens
    27 : Sulfuric acid, H2SO4
    28 : Phosphoric acid, H3PO4
    29 : Strong Acid-Strong Base Titrations Here is an example of a titration curve, produced when a strong base is added to a strong acid. This curve shows how pH varies as 0.100 M NaOH is added to 50.0 mL of 0.100 M HCl.
    30 : Strong Acid-Strong Base Titrations “The Curve” Acid + base -> Salt + Water
    31 : 1M HCl with 1M NaOH
    32 : Strong Acid-Strong Base Titrations The equivalence point of the titration is the point at which exactly enough titrant has been added to react with all of the substance being titrated with no titrant left over.
    33 : Strong Acid-Strong Base Titrations At the equivalence point, the number of moles of titrant added so far corresponds exactly to the number of moles of substance being titrated according to the reaction stoichiometry. Moles of Acid = Moles of Base
    34 :
    35 :
    36 : The Calculation The original number of moles of H+ in the solution is: 50.00 x 10-3 L x 0.100 M HCl = 5.00 x 10-3 moles The number of moles of OH- added is : 49.00 x 10-3 L x 0.100 M OH- = 4.90 x 10-3 moles Thus there remains: (5.00 x 10-3) - (4.90 x 10-3) = 1.0 x 10-4 moles H+ (aq) The total volume of solution is 0.04900 L + 0.05000 L = 0.09900 L [H+] = {1.0 x 10-4 moles / 0.09900 L } = 1.0 x 10-3 M pH = 3.00
    37 : The Calculation What is the pH when 49.00 mL of 0.100 M NaOH solution have been added to 50.00 mL of 0.100 M HCl solution? Because it is a strong acid-base reaction, the reaction simplifies to: H+ (aq) + OH- (aq) -> H2O (l)
    38 : Weak acid -Strong base
    39 : Weak acid -Strong base Please immediately compare the two curves! Notice the “Half equivalence point” in the weak acid –strong base case
    40 : Weak acid -Strong base
    41 :
    42 : 1 REGISTERED LATE FOR CLASS 2. HANDED IN LAB THAT WAS 3 WEEKS OVER DUE 3.
    43 : Strong Acid-Strong Base Titrations (In an acid-base titration, there is a 1:1 acid:base stoichiometry, so the equivalence point is the point where the moles of titrant added equals the moles of substance initially in the solution being titrated.) Notice that the pH increases slowly at first, then rapidly as it nears the equivalence point
    44 : List Strong acids and bases Strong acids: Perchloric acid, HClO4 sulfuric acid, H2SO4 hydrochloric acid, HCl Nitric acid, HNO3
    45 : List Strong acids and bases
    46 : 1M HCl with 1M NaOH
    47 : Ion Selective electrodes .
    48 :
    49 : Weak Acid-Strong Base Titrations* .
    50 : The Calculation What is the pH when 30.0 mL of 0.100 M NaOH have been added to 50.0 mL of 0.100 M acetic acid? STEP 1: Stochiometric calculation: The original number of moles of HC2H3O2 in the solution is : 50.0 x 10-3 L x 0.100 M = 5.00 x 10-3 moles HC2H3O2 Similarly, there are 3.00 x 10-3 moles of OH- due to the NaOH solution. The reaction goes to completion: OH- (aq) + HC2H3O2 (aq) C2H3O2- (aq) + H2O (l) OH- HC2H3O2 C2H3O2- INITIAL 3.00 x 10-3 mol 5.00 x 10-3 mol 0 CHANGE -3.00 x 10-3 mol -3.00 x 10-3 mol +3.00 x 10-3 mol FINAL 0 2.00 x 10-3 mol 3.00 x 10-3 mol The total volume is 80.0 mL. We now calculate the resulting molarities : [HC2H3O2] = { 2.00 x 10-3 mol HC2H3O2 / 0.0800 L } = 0.0250 M [C2H3O2-] = { 3.00 x 10-3 mol C2H3O2- } / 0.0800 L } = 0.0375 M STEP 2: Equilibrium calculation, using simplification: Ka = { [H+][C2H3O2-] / [HC2H3O2] } = 1.8 x 10-5 [H+] = { KA [HC2H3O2] / [C2H3O2-] } = { (1.8 x 10-5)(0.0250) / (0.0375) } = 1.2 x 10-5 M pH = -log(1.2 x 10-5) = 4.92
    51 : Weak acid -Strong base  
    52 : Weak base -Strong Acid .
    53 : The Calculation! Here, 0.100 M HCl is being added to 50.0 mL of 0.100 M ammonia solution. As in the weak acid-strong base titration, there are three major differences between this curve (in blue) and a strong base-strong acid one (in black): (Note that the strong base-strong acid titration curve is identical to the strong acid-strong base titration, but flipped vertically.) 1. The weak-acid solution has a lower initial pH. 2. The pH drops more rapidly at the start, but less rapidly near the equivalence point. 3. The pH at the equivalence point does not equal 7.00. POINT OF EMPHASIS : The equivalence point for a weak base-strong acid titration has a pH < 7.00.
    54 :
    55 : Weak Acid-Strong Base Titrations* EXAMPLE: What is the pH when 30.0 mL of 0.100 M NaOH have been added to 50.0 mL of 0.100 M acetic acid? STEP 1: Stochiometric calculation: The original number of moles of HC2H3O2 in the solution is : 50.0 x 10-3 L x 0.100 M = 5.00 x 10-3 moles HC2H3O2 Similarly, there are 3.00 x 10-3 moles of OH- due to the NaOH solution. The reaction goes to completion: OH- (aq) + HC2H3O2 (aq) C2H3O2- (aq) + H2O (l) OH- HC2H3O2 C2H3O2- INITIAL 3.00 x 10-3 mol 5.00 x 10-3 mol 0 CHANGE -3.00 x 10-3 mol -3.00 x 10-3 mol +3.00 x 10-3 mol FINAL 0 2.00 x 10-3 mol 3.00 x 10-3 mol The total volume is 80.0 mL. We now calculate the resulting molarities : [HC2H3O2] = { 2.00 x 10-3 mol HC2H3O2 / 0.0800 L } = 0.0250 M [C2H3O2-] = { 3.00 x 10-3 mol C2H3O2- } / 0.0800 L } = 0.0375 M STEP 2: Equilibrium calculation, using simplification: Ka = { [H+][C2H3O2-] / [HC2H3O2] } = 1.8 x 10-5 [H+] = { KA [HC2H3O2] / [C2H3O2-] } = { (1.8 x 10-5)(0.0250) / (0.0375) } = 1.2 x 10-5 M pH = -log(1.2 x 10-5) = 4.92
    56 : 3.00 x 10-3 mol 5.00 x 10-3 mol 0 CHANGE -3.00 x 10-3 mol -3.00 x 10-3 mol +3.00 x 10-3 mol FINAL 0 2.00 x 10-3 mol 3.00 x 10-3 mol The total volume is 80.0 mL.
    57 : Lewis definition Further information: Lewis acids and bases The hydrogen requirement of Arrhenius and Brønsted-Lowry was removed by the Lewis definition of acid-base reactions, devised by Gilbert N. Lewis in 1923[12], in the same year as Brønsted-Lowry, but it was not elaborated by him until 1938[2]. Instead of defining acid-base reactions in terms of protons or other bonded substances, the Lewis definition defines a base (referred to as a Lewis base) to be a compound that can donate an electron pair, and an acid (a Lewis acid) to be a compound that can receive this electron pair.[13]. In this system, an acid does not exchange atoms with a base, but combines with it. For example, consider this classical aqueous acid-base reaction: HCl (aq) + NaOH (aq) ? H2O (l) + NaCl (aq) The Lewis definition does not regard this reaction as the formation of salt and water or the transfer of H+ from HCl to OH-. Instead, it regards the acid to be the H+ ion itself, and the base to be the OH- ion, which has an unshared electron pair. Therefore, the acid-base reaction here, according to the Lewis definition, is the donation of the electron pair from OH- to the H+ ion. This forms a covalent bond between H+ and OH-, thus producing water (H2O). By treating acid-base reactions in terms of electron pairs instead of specific substances, the Lewis definition can be applied to reactions that do not fall under other definitions of acid-base reactions. For example, a silver cation behaves as an acid with respect to ammonia, which behaves as a base, in the following reaction: Ag+ + 2 :NH3 ? [H3N:Ag:NH3]+ The result of this reaction is the formation of an ammonia-silver adduct. In reactions between Lewis acids and bases, there is the formation of an adduct[13] when the highest occupied molecular orbital (HOMO) of a molecule, such as NH3 with available lone electron pair(s) donates lone pairs of electrons to the electron-deficient molecule's lowest unoccupied molecular orbital (LUMO) through a co-ordinate covalent bond; in such a reaction, the HOMO-interacting molecule acts as a base, and the LUMO-interacting molecule acts as an acid.[13] In highly-polar molecules, such as boron trifluoride (BF3),[13] the most electronegative element pulls electrons towards its own orbitals, providing a more positive charge on the less-electronegative element and a difference in its electronic structure due to the axial or equatorial orbiting positions of its electrons, causing repulsive effects from lone pair-bonding pair (Lp-Bp) interactions between bonded atoms in excess of those already provided by bonding pair-bonding pair (Bp-Bp) interactions.[13] Adducts involving metal ions are referred to as co-ordination compounds
    58 : Equilibria in Aqueous Solutions Almost all the reactions that a chemist is concerned with take place in solution rather than in gaseous or solid phases. Most of these reactions occur in aqueous solution, where water is the solvent. There are good reasons for this preference for liquid media
    59 : Equilibria in Aqueous Solutions Molecules must come into contact to react, and the rates of migration of atoms or molecules within crystals usually are too slow to be useful. In contrast, molecules of gases are mobile, but gas volumes are inconveniently large, and many substances cannot be brought into the gas phase without decomposing.
    60 : Equilibria in Aqueous Solutions Moreover, water molecules dissociate to a small extent into H+ and OH- ions, a property that is important in acid-base reactions. We will be concerned equilibria in aqueous solution, especially those involving acids and bases.
    61 : Equilibria in Aqueous Solutions The oxygen atom draws the electrons of the 0 - H bonds toward itself, acquiring a slight negative charge and leaving small positive charges on the two hydrogen atoms.
    62 : Equilibria in Aqueous Solutions Water therefore can interact with other polar molecules. Moreover, water molecules dissociate to a small extent into H+ and OH- ions, a property that is important in acid-base reactions. This chapter is concerned with reactions and equilibria in aqueous solution, especially those involving acids and bases.
    63 : Equilibria in Aqueous Solutions If reactants and products in a chemical reaction are in solution, the form of the equilibrium-constant expression is the same as for gas reactions, but the logical units of concentration are moles per liter of solution (units of molarity). aA + bB cC + dD
    64 : Equilibria in Aqueous Solutions Some reactions in aqueous solution involve water as a participant. A well-studied example is the hydrolysis ("splitting by water") of the ethyl acetate molecule to yield acetic acid and ethyl alcohol (ethanol):  
    65 : Bronstead Lowry systems Introduction to acid-base pair systems
    66 : Bronstead Lowry systems HC2H3O2(aq) + H2O(l) ? C2H3O2–(aq) + H3O+(aq)
    67 : Bronstead Lowry systems Acid conjugate base HC2H3O2(aq) + H2O(l) ? C2H3O2–(aq) + H3O+(aq) Base conjugate acid pair
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    69 : conjugate acid-base pairs acid base conjugate acid conjugate base CO32–(aq) + HC2H3O2(aq) ? C2H3O2–(aq) + HCO3–(aq) conjugate acid-base pairs acid base conjugate acid conjugate base H3PO4(aq) + OCl –(aq) ? H2PO4–(aq) + HOCl(aq) conjugate acid-base pairs
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    74 : conjugate acid-base pairs acid base conjugate acid conjugate base CO32–(aq) + HC2H3O2(aq) ? C2H3O2–(aq) + HCO3– conjugate acid-base pairs acid base conjugate acid conjugate base H3PO4(aq) + OCl –(aq) ? H2PO4–(aq) + HOCl(aq) conjugate acid-base pairs Bronstead Lowry systems
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