Acids+and+Bases

Acids and Bases (Chapter 16, pages: 674-714)

1. General Concepts: a. Understand what is meant when we read or hear about acids and bases When refering to an acid or base, they are usually dissolved in a liquid, making it an aqueous solution. Characteristics of an acid, are a sour taste, formation of gas bubles when in contact with metal, produces H+ ions, accepts electrons, and has a pH less than 7.0. Acids turn blue litmus paper red. Characteristics of a base are, an alkali, a bitter taste, forms OH-, donates electrons, turns pH paper blue, and has a pH higher than 7.0. Bases turn red litmus paper blue. Acids can exist either as a pure substance or in solution, which is more common. A base is an aqueous substance. A strong acid means that all the molecules of the substance donated its protons. The reaction went forward, entirely to completion. Consequently, the amount of protons (hydronium ions) in solution would be concentrated. A weak acid means that very few of the molecules of the substance donated its proton. The equilibrium for a reaction involving a weak acid is largely to the left and the equilibrium constant would be small. Consequently, because the amount of protons (hydronium ions) is small in solution at equilibrium, the solution is diluted. A strong acid dissociates completely in water to produce a conjugate weak base and hydronium (H+). A weak acid does not completely dissociate; it only partially ionizes in solution.
 * b. Know the difference between strong and weak acids, and the difference concentrated and diluted **

c. Have a frame of reference – the pH scale- with which we can categorize solutions of acids and bases The pH scale is numbers 0-14. Substances lower than 7.0 are acids, and substances higher than 7.0 are bases. Here is an example of a pH scale, and substances at each pH number. d. Know where many common substances fit within the pH scale: Fluid: pH: Stomach Acid 1.5 Lemon Juice 2.0 Vinegar: 3.0 Grapefruit Juice 3.2 Orange Juice 3.5 Urine 4.8-7.5 Rainwater 5.5 Saliva 6.4-6.9 Milk 6.5 Pure Water 7.0 Blood 7.35-7.45 Tears 7.4 Milk of magnesia 10.6 Household ammonia 11.5
 * Table 16.4 pg 678:**

Here are some common household materials and there pH's. Gives you a good idea and understanding.


 * 1.0 battery acid (sulfuric acid)
 * 1.8-2.0 limes
 * 2.2-2.4 lemon juice
 * 2.2 vinegar (acetic acid)
 * 2.46 pepsi
 * 2.8-3.4 fruit jellies
 * 2.9-3.3 apple juice, cola
 * 3.0-3.5 strawberries
 * 3.7 orange juice
 * 4.0-4.5 tomatoes
 * 5.6 unpolluted rain
 * 5.8-6.4 peas
 * 6.0-6.5 corn
 * 6.1-6.4 butter
 * 6.4 cow's milk
 * 6.5-7.5 human saliva
 * 6.5-7.0 maple syrup
 * 7.0 distilled water
 * 7.3-7.5 human blood
 * 7.6-8.0 egg whites
 * 8.3 baking soda
 * 9.2 borax
 * 10.5 milk of magnesia
 * 11.0 laundry ammonia
 * 11.7 chlorine
 * 12.0 lime water
 * 13.0 lye
 * 14.0

The pH is equal to the negative log of the concentration of protons in solution, pH = -log[H+]. This equation can also be written as 10^(-pH)=[H+]. So, if you know the concentration of the protons (hydronium ions) in solution, then you can calculate the pH. Similarly, if you know the pH of the solution, you can calculate the concentration of protons (hydronium ions).
 * e. Be able to do a variety of calculations relating to the –log (negative log) function; commonly denoted as a small case ‘p’ as in pH, pOH, pKa, etc. **

The pOH is equal to the negative log of the concentration of hydroxide ions in solution, pOH = -log[OH-]. This equation can also be written as 10^(-pOH)=[OH-]. So if you know the concentration of hydroxide ions in solution, then you can calculate the pOH.

In addition, you can also calculate the pOH of the solution if you know the pH of the solution. In this case, the pOH of the solution would be equal to the quantity, pOH=14.00-pH. The reason you can do this is because the ion-product constant of water, symbolized by Kw, is equal to 1.0*10^(-14). And the equilibrium constant expression for the Kw is the following expression, Kw=[H+][OH-]. This formula indicates that the concentrations between the protons (hydronium ions) and hydroxide ions at equilibrium is going to be constant, always equal to the value of Kw, 1.0*10^(-14). If you take the log of both sides of this equation you get log(Kw)=log[H+]+log[OH-]. Then if you multiply both sides of the equation by -1, the equation will be written, -log(Kw)= -log[H+] + (-log[OH-]). Next, we substitute. the negative log of Kw, 1.0*10^(-14) is 14, and the negative log of protons (hydronium ions) and hydroxide ions is pH and pOH respectively. Consequently, we come up with the formula 14 = pH + pOH. If we know either the pH or the pOH of the solution, we can calculate the other.
 * (Make sure you use log and not ln, especially remember this if using a TI-89 Calculator, it can be a big headache.)**

A) A 0.000235 M solution of H+ //pH = -log(H+) = -log(0.000235) = 3.63//

A) A 0.000235 M solution of OH- //pOH = -log(0.000235) = 3.63 => pH + pOH = 14 => pH = 14 – 3.63 => pH = 10.37//

2. Be able to define the following concepts: a. Acid - having a pH less than 7 and having a larger amount of H+ ions than OH- ions. Acids tend to have a sour taste, turn blue litmus paper red, and react with certain bases and metals to form salts. Acids can act as a proton donor and can accept a pair of electrons to form a covalent bond. When acids come in contact with metals, bubbles are formed. An acid is a substance that increases H+concentration when added to water. Any specie that donates hydrogen ions. b. Base - having a pH higher than 7 and having more OH- ions than H+. Bases have a bitter taste to them. A base is a substance that increases OH- concentration when added to water. Bases have the ability to accept proton donations from acids. The solution is basic when [H3O+] < [OH-]. Any specie that accepts hydrogen ion. Any specie that produced hydroxil (OH) ions in an aqueous solution. c. Amphiprotic - means its capable for acting as an acid or a base. example: Al2O3; It can act as a base with hydrochloric acid to form salt and water and acts as a acid with sodium hydroxide.

"An //amphiprotic// substance is one which can both donate hydrogen ions (protons) and also accept them."

d. Autoprotolysis- A proton (hydron) transfer reaction between two identical molecules (usually a solvent), one acting as a Brønsted acid and the other as a Brønsted base. For example: 2 H2O --> H3O+ + OH- "The autoionization of water (or similar compounds) in which a proton (hydrogen ion) is transferred to form a cation and an anion" e. pH scale and pOH scale pH scale a scale used to measure acidity. pH= -log[H+] **pH:** the pH of a solution is defined as the negative base-10 logarithm of the hydronium ion concentration and is measured in mol/L. (page 677) The pH scale is comprised of H+ cations and OH- anions. The greater the H+ cations, the greater the acidity and lesser the pH value. The greater the OH- anions, the greater the basicity and greater the pH value. The scale doesn't stop at 14.0 but keeps on going. Most of the bases used in general chemistry are only comprised of the ones distinguishd on the 14.0 scale. When the pH value is >7 the ratio of OH- anions is greater than the H+ cations. When the pH value is <7 the ratio of H+ cations is greater than the OH- anions. A pH value of 7 doesn't necessarily define a solution as being neither acidic or basic. Instead the ration of H+ cations and OH- anions are in equilibrium. This gives a neutral pH value for the solution. (Compliments of Dave Wilson)--

pOH scale: a scale used to measure basicity. pOH= -log[OH-]. The same is true for the opposite: pH = 14 - pOH f. Polyprotic acids an acid with more than one ionizable proton. An example of a polyprotic acid would be phosphoric acid.(H3PO4). Polyprotic acids can donate or release more than one proton (H+). Because of this, they have multiple Ka's, with the first Ka being the strongest. " Species capable of losing more than one hydrogen ion in acid base reactions. The hydrogen ions are lost in sequence, not at the same time. " Because the later Ka's are so much smaller than the original they do not significantly affect the final PH of the solution. Simply use the first KA to determine H+ concentration as the amount of H+ donated by successive ionizations will rounded as insignificant digits. Example: H3A(aq) + H2O(l) H3O+(aq) + H2A−(aq) //K//a1 H2A−(aq) + H2O(l) H3O+(aq) + HA2−(aq) //K//a2 HA2−(aq) + H2O(l) H3O+(aq) + A3−(aq) //K//a3
 * Note: ** adding the pH value and pOH value will equal 14. This means that if you have the pH of a substance and you want the pOH: pOH = 14 - pH

g. salts

An ionic compound made up of the cation from a base and the anion from an acid, in other words, a product. When dealing with acids and bases, if it is a salt, it is a strong base. Hydroxides are ALWAYS salts - (a metal combined with OH- ) Ex: AH + BOH -> H2O + AB (AB being the salt formed by the cation A and the anion B)

You can tell if a substance is an Arrhenius Acid/Base by determining whether the concentration of hydroxide or protons (hydronium) increased when the substance was dumped in water. So, an Arrhenius Acid would be any substance that caused the solution to increase in protons (hydronium ions) when the substance was put in water. Similarly, an Arrhenius base would be any substance that increased the concentration of hydroxide ions in water.
 * 3. Be able to use the three definitions of acids and bases to predict if a molecule is an acid or a base (Arrhenius, Bronsted-Lowry, and Lewis). **

A Bronsted-Lowry Acid/Base is characterized by acceptance or donation of a proton. So, you can tell if a substance is a Bronsted-Lowry Acid if the substance donates a proton and forms a conjugate base. Related, you can tell if a substance is a Bronsted-Lowry Base if the substance accepts a proton and in the process forms a conjugate acid.

A Lewis acid accepts electons, and a Lewis base donates electrons.

4. Recognize general chemical formulas for acids and bases Most acids have the general formula HA, where A- is an anion and most bases have the form BOH, where B+ is an appropriate cation. n aqueous solutions acids increase the hydrogen ion (H+) concentration. On the other hand bases increase the hydroxide ion (OH-) concentration. In more simple terms, acids generally begin with Hydrogen and bases usually end with hydroxide.

5. List the common strong acids and strong bases. Acids: Hydrochloric acid: HCl Hydrobromic acid: HBr Hydroiodic acdi: HI sulfuric acid: H 2 SO 4 Nitric acid: HNO 3 Perchloric acid: HClO 4 chloric acid (HClO 3 ) and nitric acid ( HNO3 )


 * __HF is NOT a strong acid! __**

**__H3PO4 is ALSO NOT a strong Acid!__**

Bases: Lithium hydroxide: LiOH Sodium hydroxide: NaOH Potassium hydroxide: KOH Rubidium hydroxide: RbOH Cesium Hydroxide : CsOH <span style="font-family: Verdana,sans-serif; font-size: 12px; line-height: normal;">Calcium hydroxide: Ca(OH) 2 <span style="font-family: Verdana,sans-serif; font-size: 12px; line-height: normal;">Strontium hydroxide: Sr(OH) 2 <span style="font-family: Verdana,sans-serif; font-size: 12px; line-height: normal;">Barium hydroxide: Ba(OH)2

6. Describe the difference between a strong acid at low [ ] and a weak acid at high [ ] The strength of the acid is dependent on its ability to disassociate its attached H+. The strength of the acid is represented by the Ka of the reaction, and dictates how strong the acid is. Ka being a constant, is effected by the concentration of the reactants. However, the lowest concentrations of a strong acid with always be more acidic than the highest concentrations of a weak acid. This is due to the fact that the concentrations do not compensate for what the Ka represents. The concentrations of the two will not make a large enough difference in the Ka to equate the two separate reactions.

7. Be able to predict the relative strength of an acid based on: a. Acid dissociation constant (ka) A large Ka value indicates a stronger acid, whereas a small Ka value indicates a weaker acid. b. Resonance structures of the conjugate base The more resonance structures the molecule has the stronger acid is With Increasing resonance structure in the conjugated product the acidity increases c. Electronegativity of the atoms- acidity increases with increasing electronegativity of the atom bonded to the hydrogen If the central atom is an electronegative element, or is in a high oxidation state, it will attract electrons, causing the O-H bond to be more polar. This makes it easier for the hydrogen to be lost as H+, making the acid stronger.

8. Be able to correlate the following values a. Ka and Kb- every strong conjugate base has an ionization constant Kb. Every strong conjugate acid has an ionization constant Ka. Ka multiplied by Kb equals Kw = 1.0 x 10^-14. Strong acids have a very high Ka value in which we then state Ka is nonexistent b. pKa and pKb- ﻿These two constants sum to the value of 10. The difference between the pKa and pKb values and the Ka and Kb values are the fact that a "p" relates the number to the logarithmic scale. log is equivalent to 10^x. This factor can be used to convert the value between the two different forms of the constants. c. pH and pO- pH plus pOH equals 14.00 pH + pOH = 14 so if you have two of the three values you can find the other one.

d. [H+] and [OH-] or [H3O+] and [OH-] These are the ions which dissociate from an acid or base. H+ and H3O+ dissociate from acids and OH^- dissociates from bases. In the case of an acid base neutralization, the H+ and the OH- ions bond together to form H2O (water) along with the production of a salt. 9. Be able to predict pH values based on a given strength for an acid or base If you have a strong acid the pH will be from one to 3.5. If you have a strong base the pH will be from 10.5 to 14. If you have a weak acid the pH will be from 3.5 to 7. If you have a weak base the pH will be from 7 to 10.5. The closer the pH is to 0, the stronger the acid. The closer the pH is to 14, the stronger the base. The closer the pH is to 7, the more neutral the solution is.

10. Be able to predict the chemical structure of conjugate pairs A conjugate base is the product of an acid after reacting in equilibrium. A conjugate acid is the product of a base after reacting in equilibrium.

11. Be able to calculate the pH of a solution for: a. Strong acids 1 b. Strong bases 14 c. Weak acids( weak acid ionized less than 100%) 6 d. Weak bases 8 The closer the pH is to 7 the weaker the acid or the base. So the farthest the acid or base is to 7 the stronger it is.

12. Be able to predict the effects on pH after the addition of a salt to the solution Adding a salt could make a solution more acidic, more basic, or even not cause a change in pH. The effects of the the change in pH will be noticable in the chemical structure of the salt and how it dissociates into H+ (AKA H3O+) or OH-