Acid-base+reactions,+Concepts+of+Arrhenius,+Bronsted-Lowry+and+Lewis,+Coordination+complexes,+Amphoterism

__Acid-Base Reactions__ The reactions between acids and bases are called //** neutralization reactions **//. They are often double replacement reactions. In most cases, a neutralization reaction can be described as the mixing of an acid with a base to form a salt and water:

HBr(//aq) +// KOH(//aq)// à KBr(//aq)// + H2O(//l)// (acid) + (base) à (salt) + (water) Note that most acid-base reactions take place in aqueous solution. The symbol (//aq//) is often ommited for clarity.

__Acid and Base Strengths__ **Strong Acids:** Strong acids are acids that dissociate completely into ions when dissolved in water. The six most important strong acids are:


 * ** Strong Acids **

||
 * Hydrochloric acid || HCl à  || H+ + Cl- ||
 * Hydrobromic acid || HBr à  || H+ + Br- ||
 * Hydroiodic acid || HI à  || H+ + I- ||
 * Perchloric acid || HClO4 à  || H+ + ClO4- ||
 * Nitric acid || HNO3 à  || H+ + NO3- ||
 * Sulfuric acid || H2SO4 à  || H+ + HSO4- ||

Here's a way to memorize them all: (To the tune of "Three Blind Mice")

Three blind mice, 'Cause acids aren't nice, Wear your goggles, be careful now, 'Cause the strongest acids will you hurt you (ow!) And now these mice are not allowed Around acids For what they did

Three binary Strong acids There's hydrochloric, HCl And hydroiodic, as well And hydrobromic is quite swell Three binary Strong acids

Four Oxy- Acids, see There's nitric, that's HNO3 And chloric, HClO3 Sulfuric, H2SO4 And wait 'cause there is still one more! Perchloric, HClO4 Yes, that's four Oxyacids galore

Now you know Which acids blow Stay away 'cause they will hurt Wear an apron on your shirt Remember, always stay alert! When dealing with These acids

- Negeen Farsio

All other acids are weak acids. This means that when they dissolve in water, they do not completely ionize. Only a small percentage of the molecules in weak acids dissolve into ions. The strength of an acid is inversely proportional to the strength of the bond between the hydrogen and the remainder of the molecule. Strong acids have very weak bonds, and weak acids have stronger bonds. Acids also increase in strength from left to right of a period in the periodic table, like electronegativity. For example, PH3 is weaker than H2S, which is weaker than HBr. Acids also become stronger from the top of a group to the bottom, like atom size. All mineral acids that are not strong acids are oxoacids, which are acids that contain hydrogen, oxygen and another element. The strength of an oxoacid always depends on the number of oxygen atoms per each hydrogen atom and on the electronegativity of the central atom. Oxoacids that have the same central atom and the same number of hydrogen atoms increase in strength as the number of oxygen atoms increases. As the electronegativity of the central atom increases, the O-H bond weakens which results in a stronger acid. Telluric acid < selenic acid < sulfuric acid Hypochlorous acid < chlorous acid < chloric acid < perchloric acid
 * Weak Acids ** **:**

All metal hydroxides are strong bases. However, most metal hydroxides are only slightly soluble. Only the hydroxides of group IA metals, calcium, strontium and barium are soluble enough to dissociate in water.
 * Strong Bases: **

All hydroxides besides those mentioned above are weak bases. Bases related to ammonia are weak bases.
 * Weak Bases: **

__ Acid-Base Theories __


 * ** Acid-Base Theories ** ||
 * Theory Name || Brief Description ||
 * Arrhenius || An acid adds hydrogen ions to a solution, and a base adds hydroxide ions to a solution. ||
 * Bronsted-Lowry || An acid is a proton donor, and a base is a proton acceptor. ||
 * Lewis || An acid is an electron pair acceptor, and a base is an electron pair donor. ||

**Arrhenius Theory:** An acid is considered to be any substance that increases the concentration of hydrogen ions in an aqueous solution. When an acid is dissolved in water, hydrogen ions are released. Because these ions are unlikely to exist in aqueous solution, they are hydrated or bound to water molecules forming H3O + HCl(//g//) + H2O(//l//) à H3O+(//aq//) + Cl- (//aq//) The Arrhenius theory considers bases to be substances that increase the hydroxide ion, OH-, concentration when dissolved in water.

KOH(//s) //– H2O à  K+ (//aq) +// OH- (//aq)// Acids are proton donors and bases are proton acceptors, for example, Ammonia, NH3, is a base since it accepts protons from water molecules in a reaction to form Ammonium, NH4+ ** Lewis Theory: ** Acids are substances that accept electron pairs from other atoms, ions, or molecules and bases are substances that donate electron pairs in forming chemical bonds. The Lewis Theory is used maily to explain the formation of substances called //complexes//.
 * Bronsted-Lowry Theory: **

__ Conjugate Acid-Base Pairs __ HC2H3O2 + OH- ↔ C2H3O2- + H2O HC2H3O2 is an acid that reacts with the base OH- in the forward reaction. The products, however, are also acids and bases. C2H3O2- is a base that can accept an H+ from water, while water is an acid since it donates a proton in the reverse reaction. HC2H3O2 and C2H3O2- are called a conjugate acid-base pair, same goes for OH- and H2O, which are another acid-base pair.

Note that conjugate acid-base pairs always have formulas that differ by only one H+ : Conjugate acid ↔ Conjugate base + H+ || Strong conjugate base || || Weak conjugate base || || Strong conjugate acid || || Weak conjugate acid ||
 * Weak acid
 * Strong acid
 * Weak base
 * Strong base

__ Amphiprotic (Amophetric) Substances __ Amphiprotic and amophetric are two words that describe the same phenomenon in which a substance can act as both a conjugate base and a conjugate acid.Amphiprotic salts are anions that must have at least one proton to act as bases. The anions of partially neutralized polyprotic acids are always amphiprotic. Water, the hydrogen carbonate ion, the dihydrogen phosphate ion, and the monohydrogen phosphate ion can all accept a proton and act as a base, they can also donate a proton and act as an acid.

__Coordination Complexes__  Co 2+  ( // aq // ) + 4Cl -  ( // aq // )↔ CoCl 4  ( // aq // ) The ions or molecules that bind to transition-metal ions to form these complexes are called **ligands** (from Latin, "to tie or bind"). The number of ligands bound to the transition metal ion is called the ** coordination number .** Some main group elements like aluminum, tin, and lead form complexes such as the AlF63-, SnCl42- and PbI42- ions. isothiocyanato (bonded through nitrogen) || Name the ligand first, then the cation. Ex:  tetraamminecopper(II) ion: Cu(NH3)42+ <span style="font-size: 10pt; color: #000000; font-family: Arial, Helvetica, sans-serif;"> diamminesilver(I) ion: Ag(NH3)2+. tetrahydroxyzinc(II) ion: Zn(OH)4 2- Acid-base reactions may change NH3 into NH4+ (or vice versa) which will alter its ability to act as a ligand. Visually, a precipitate may go back into solution as a complex ion is formed. For example, Cu2+ + a little NH4OH will form the light blue precipitate, Cu(OH)2. With excess ammonia, the complex, Cu(NH3)42+, forms. Keywords such as "excess" and "concentrated" of any solution may indicate complex ions. AgNO3 + HCl forms the white precipitate, AgCl. With excess, concentrated HCl, the complex ion, AgCl2-, forms and the solution clears. Coordination numbers generally range between 2 and 12, with 4 (tetracoordinate) and 6 (hexacoordinate) being the most common.<span style="font-size: 32pt; font-family: 'Times New Roman'; mso-hansi-font-family: 'Times New Roman'; mso-ascii-font-family: 'Times New Roman';">
 * Coordination compounds ** consist of a metal cation or neutral atom to which neutral or negatively charged ligands have bonded ex. the FeCl4- ion and CrCl3 6 NH3. They are also known as **complex ions** or **coordination complexes** because they are Lewis acid-base complexes.
 * <span style="font-size: 10pt; color: #666699; font-family: Arial; text-align: center; mso-bidi-language: AR-AE; msobidilanguage: AR-AE;">Common Ligands ||
 * Ligands || Names used in the ion ||
 * H2O || aqua ||
 * NH3 || ammine ||
 * OH- || hydroxy ||
 * Cl- || chloro ||
 * Br- || bromo ||
 * CN- || cyano ||
 * SCN- || thiocyanato (bonded through sulphur)
 * <span style="font-size: 10pt; color: #003366; font-family: Arial; text-align: center; mso-bidi-language: AR-AE; msobidilanguage: AR-AE;">Some coordination complexes ||
 * molecular formula || ligand || lewis acid || donor atom || coordination number ||
 * Ag(NH3)2+ || NH3 || Ag+ || N || 2 ||
 * [Zn(CN)4]2- || CN- || Zn2+ || C || 4 ||
 * [Ni(CN)4]2- || CN- || Ni2+ || C || 4 ||
 * [PtCl6] 2- || Cl- || Pt4+ || Cl || 6 ||
 * [Ni(NH3)6]2+ || NH3 || Ni2+ || N || 6 ||

__ The pH Scale __ The pH scale is a way of expressing the strength of acids and bases. If the pH is under 7, the substance is acidic. If it is over 7, the substance is basic. A substance with a pH of 7 is neutral, like water. To calculate the pH we use <span style="font-size: 10pt; color: #de7385; font-family: Arial, Helvetica, sans-serif; msobidilanguage: AR-AEmsoBidiLanguage;">pH = - log [H+] <span style="font-size: 10pt; color: #000000; font-family: Arial, Helvetica, sans-serif;">The pH is equal to the negative log of the concentration of H+ Example: If [H+] = 1 x 10-10 pH = - log 1 x 10-10 pH = - (- 10) pH = 10 Example: If [H+] = 1.8 x 10-5 pH = - log 1.8 x 10-5 pH = - (- 4.74) pH = 4.74 Another method to measure the acidity of basicity of a substance is the pOH scale. Finding the pOH is helpful when trying to find the pH of a base, which does not have a [H+]. pOH = - log [OH- ] <span style="font-size: 10pt; color: #000000; font-family: Arial, Helvetica, sans-serif;">We can also use the pOH to find the pH: <span style="font-size: 10pt; color: #690c1c; font-family: Arial, Helvetica, sans-serif; msobidilanguage: AR-AE;">pH + pOH = 14 Another helpful equation is <span style="font-size: 10pt; color: #bf4055; font-family: Arial, Helvetica, sans-serif; msobidilanguage: AR-AEmsoBidiLanguage;">Kw = [H+] [OH-] <span style="font-size: 10pt; color: #bf4055; font-family: Arial, Helvetica, sans-serif; msobidilanguage: AR-AEmsoBidiLanguage;">Where Kw = 1x10-14 Example: What is the pH of a 0.0010 M NaOH solution? [OH-] = 0.0010 (or 1.0 X 10-3 M) pOH = - log 0.0010 pOH = 3 pH = 14 – 3 = 11 OR Kw = [H+] [OH-] 1x10-14 = [H+] 0.0010 [H+] = 1.0 x 10-11 M pH = - log (1.0 x 10-11) = 11.00 Equilibrium constant for weak acids Weak acids have Ka < 1 Equilibrium constant for weak bases Weak base has Kb < 1 Relation between Ka, Kb, [H+] and pH When dealing with equilibria involving a weak base or acid is when you need to use Ka or Kb, Example: You have 0.010 M NH3. Calculate the pH. NH3 + H2O <span style="font-size: 10pt; font-family: Wingdings; mso-char-type: symbol; mso-symbol-font-family: Wingdings; msoasciifontfamily: Arial; msohansifontfamily: Arial; msochartype: symbol; msosymbolfontfamily: Wingdings; msobidilanguage: AR-AE; msobidifontfamily: Arial;">à  NH4+ + OH- Kb = 1.8 x 10-5 (will be given) Step 1. Define equilibrium concentrations using the RICE method || <span style="font-size: 10pt; font-family: Arial; text-align: center; mso-bidi-language: AR-AE; msobidilanguage: AR-AE;">[NH3] || <span style="font-size: 10pt; font-family: Arial; text-align: center; mso-bidi-language: AR-AE; msobidilanguage: AR-AE;">[NH4+] || <span style="font-size: 10pt; font-family: Arial; text-align: center; mso-bidi-language: AR-AE; msobidilanguage: AR-AE;">[OH-] || || By plugging all the numbers into the Kb expression, we can find the value of x which is = <span style="font-size: 10pt; font-family: Arial; text-align: center; mso-bidi-language: AR-AE; msobidilanguage: AR-AE;">[OH-], then it becomes easy to find the pH [OH-] = 4.2 x 10-4 M pOH = - log [OH-] = 3.37 pH + pOH = 14, pH = 10.63
 * <span style="font-size: 10pt; font-family: Arial; text-align: center; mso-bidi-language: AR-AE; msobidilanguage: AR-AE;">Reaction
 * <span style="font-size: 10pt; font-family: Arial; text-align: center; mso-bidi-language: AR-AE; msobidilanguage: AR-AE;">Initial || <span style="font-size: 10pt; font-family: Arial; text-align: center; mso-bidi-language: AR-AE; msobidilanguage: AR-AE;">0.010 || <span style="font-size: 10pt; font-family: Arial; text-align: center; mso-bidi-language: AR-AE; msobidilanguage: AR-AE;">0 || <span style="font-size: 10pt; font-family: Arial; text-align: center; mso-bidi-language: AR-AE; msobidilanguage: AR-AE;">0 ||
 * <span style="font-size: 10pt; font-family: Arial; text-align: center; mso-bidi-language: AR-AE; msobidilanguage: AR-AE;">Change || <span style="font-size: 10pt; font-family: Arial; text-align: center; mso-bidi-language: AR-AE; msobidilanguage: AR-AE;">-x || <span style="font-size: 10pt; font-family: Arial; text-align: center; mso-bidi-language: AR-AE; msobidilanguage: AR-AE;">+x || <span style="font-size: 10pt; font-family: Arial; text-align: center; mso-bidi-language: AR-AE; msobidilanguage: AR-AE;">+x ||
 * <span style="font-size: 10pt; font-family: Arial; text-align: center; mso-bidi-language: AR-AE; msobidilanguage: AR-AE;">Equilibrium || <span style="font-size: 10pt; font-family: Arial; text-align: center; mso-bidi-language: AR-AE; msobidilanguage: AR-AE;">0.010 - x || <span style="font-size: 10pt; font-family: Arial; text-align: center; mso-bidi-language: AR-AE; msobidilanguage: AR-AE;">x || <span style="font-size: 10pt; font-family: Arial; text-align: center; mso-bidi-language: AR-AE; msobidilanguage: AR-AE;">x

__Practice__ The following list applies to questions i & ii only: 1) Cu(NH3)42+ 2) KOH 3) HCO3- 4) CO32- 5) SO32-

i) Which of the above is the product of a lewis acid reacting with a lewis base? a) 1&2 b) 2 c) 1&3 d) 4 e) 4&5 ii) Which is the strongest base? a) 1 b) 2 c) 3 d) 4 e) 5

iii) Which of the following cannot be either a lewis acid or a lewis base? a) CH4 b) Cu2+ c) CO d) Fe3+ e)NH3

iv) In the complex ion Cu(NH3)42+ the NH3 is called a) a cation b) a ligand c) a lewis acid d) an anion e) a conjugate acid

v) Rank the strengths of phosphoric acid, sulfuric acid and perchloric acid. Explain your reasoning.

Previous AP Exam Question (1982): A buffer solution contains 0.40 mole of formic acid, HCOOH, and 0.60 mole of sodium formate, HCOONa, in 1.00 liter of solution. The ionization constant, Ka, of formic acid is 1.8 x 10^-4.

(a) Calculate the pH of this solution.

(b) If 100. milliliters of this buffer solution is diluted to a volume of 1.00 liter with pure water, the pH does not change. Discuss why the pH remains constant on dilution.

(c) A 5.00-milliliter sample of 1.00-molar HCl is added to 100. milliliters of the original buffer solution. Calculate the [H3O+] of the resulting solution.

(d) A 800-milliliter sample of 2.00-molar formic acid is mixed with 200. milliliters of 4.80-molar NaOH. Calculate the [H3O+] of the resulting solution.

**Answers and Explainations:** i) C, only these two species actually exist in solution KOH might be considered as a lewis acid (K+) and a lewis base (OH-). However, KOH really does not exist to any extent in solution ii) D, the copper-ammonia complex has negligible basicity. The carbonate ion will be more basic than the bicarbonate ion. Also, since sulfurous acid is stronger than carbonic acid, their conjugate bases have their strengths reversed and the carbonate ion will be a stronger base than the sulfite ion. iii) A, CH4 lacks the capacity either to accept or to donate electron pairs. iv) B, In this ion, NH3, is the ligand, which is another namer for a lewis base. v) The strengths are ranked in order as perchloric acid > sulfuric acid > phosphoric acid. Stronger acids have weaker bonds between the hydrogen and oxygen in these formulas. The strengths of these acids depend on the electronegativity of the central atom, the more electronegative it is, the more electron density is withdrawn from the O-H bond, weakening it and making a stronger acid. The strength of the acid also depends on the number of oxygen atoms not bonded to hydrogen. These oxygen atoms can attract electrongs from the O-H bond, causing the acid to become stronger. We see that HClO4 has three of these oxygen atoms, sulfuric acid has two, and phosphoric acid has only one. Finally, the stability of the anion, due to delocalizing electrongs over the oxygen not bonded to H, means that the acid is stronger.

__References__ __<span style="font-size: 10pt; font-family: Calibri; mso-bidi-language: AR-AE; mso-bidi-font-family: Arial;">Barron's AP chemistry 2008 book. 4th edition. __ __<span style="font-size: 10pt; color: windowtext; font-family: Calibri; msobidilanguage: AR-AE; msobidifontfamily: Arial;">[] __ __<span style="font-size: 10pt; color: windowtext; font-family: Calibri; msobidilanguage: AR-AE; msobidifontfamily: Arial;">[] __ __ [] [] __