Video 11.1 Acids and Bases.

Video 11.1 Acids and Bases

OBJECTIVES By the end of the video you should be able to…
Identify and explain electrolytes. Identify acids, bases, and salts with their properties using the Arrhenius Theory. Describe and write acid/base reactions.

What is an electrolyte? A substance that dissolves in water and conduct electricity. Acids and bases are electrolytes.

Arrhenius Acid: Substance that, when dissolved in water, increases the concentration of hydrogen H+(or hydronium H3O+) ions. Acids have a sour taste and can burn your skin. Acids react vigorously with metals to make H2 pH is less than 7 On table K

Arrhenius Base: Substance that, when dissolved in water, increases the concentration of hydroxide ions. Bases have a bitter taste and are corrosive. pH>7 On Table L

Arrhenius If an acid is added to a base, it results in neutralization, where all properties of the acid and base are lost. The products are a salt and water. NaOH + HCl  H2O + NaCl Salt is another name for solid ionic compounds containing elements other than H+ and OH-. Salts are electrolytes with high melting points and boiling points.

Label & Name the acid, base & salt:
H2SO4 + LiOH  H2O + Li2SO4 KOH + HNO3  KNO3 + H2O Acid Base Salt Base Acid Salt

Identify the salt produced:
NaOH + HF  Ca(OH)2 + H2SO4  H2O + NaF H2O + CaSO4 Make sure you check your compounds with the criss cross rule for ions!

Acids react with most metals
One hydrogen will be replaced by a more active metal. (Unreactive metals include Cu, Ag, and Au.) 2HCl + Mg  MgCl2 + H2(g) 2HCl + Zn  ZnCl2 + H2

OBJECTIVES Now you should be able to…
Identify and explain electrolytes. Identify acids, bases, and salts with their properties using the Arrhenius Theory. Describe and write acid/base reactions.

Video 11.2 Alternate Acid Base Theory

Contrast concentration and strength of acids and bases. Identify Acids and Bases using the “Alternate” or Bronsted-Lowry Theory. Identify conjugate acids and bases.

Acid and Base Strength Strong acids and bases are completely dissociated in water to make a lot of H+or OH-. These create normal neut6realization reactions. Weak acids and bases only dissociate partially in water to make a small amount of H+or OH-. These do not create normal neutralization reactions. How is strength different from concentration?

Strength versus Concentration
Strength refers to the amount of ions a substance makes when it breaks down. Concentration refers to the amount of the substance initially, before it breaks down. This is usually measured in molarity (mol/L).

Memorize Strong Electrolytes:
ACIDS BASES HCl HBr HI HClO4 H2SO4 H3PO4 HNO3 LiOH NaOH KOH RbOH CsOH Ca(OH)2 Ba(OH)2

Strength versus Concentration
For each decided if they are weak or strong, and dilute or concentrated.

Acid: Proton donor Base: Proton acceptor Brønsted–Lowry : BAAD
Protons refer to hydrogen ions. H+ H

What Happens When an Acid Dissolves in Water?
Water acts as a Brønsted–Lowry base and removes a proton (H+) from the acid. As a result, the conjugate base of the acid and a hydronium ion are formed.

Conjugate Acids and Bases:
Reactions between acids and bases always yield their conjugate bases and acids.

If it can act as either an acid or a base it is amphiprotic.
HCO3− HSO4− H2O These all have hydrogen atoms to donate and a negative charge that would attract H+

Find the conjugate acid:
NH3 HS- HSO4- H2PO4- Bases gain H+ to become a conjugate acid! NH4+ H2S H2SO4 H3PO4

Find the conjugate base:
ClO4- HClO4 HCl H2SO4 H3PO4 Acids lose H+ to become a conjugate base! Cl- HSO4- H2PO4-

Contrast concentration and strength of acids and bases. Identify Acids and Bases using the “Alternate” or Bronsted-Lowry Theory. Identify conjugate acids and bases.

Video 11.3 pH

Find the pH, pOH, [H3O+], [OH-] of any solution. Compare the acidity or alkalinity of solutions based on pH values.

Ion-Product Constant Write the equilibrium expression for water’s ionization H2O(l)  H3O+ + OH- Kc = [H3O+] [OH−] This special equilibrium constant is the ion- product constant for water, Kw. At 25°C, Kw = 1.0  10−14 Now calculate the [H3O+] and [OH-].

Ion-Product Constant In pure water, Kw = [H3O+] [OH−] = 1.0  10−14
Because in pure water [H3O+]=[OH−], [H3O+] = √(1.0  10−14) = 1.0  10−7 What is the pH of water? 7 How??? -log (H3O+)= -log (1.0x10-7) = 7 (it’s the exponent!)

Other “p” Scales The “p” in pH tells us to take the negative log of the quantity (in this case, hydrogen ions). pH = -log [H3O+] pOH= -log [OH−] pH + pOH = 14 [H3O+][OH-] = 1.0x10-14

pH and pOH If the [H3O+] = 1x10-3 the pH is
9 If the [OH-] = 1x10-4 the pOH is 4 If the [OH-] = 1x the pOH is 12 If the pH is 3 the pOH is 11 If the pOH is 12 the pH is 2

How Do We Measure pH? For less accurate measurements, one can use
Litmus paper “Red” paper turns blue above ~pH = 8 “Blue” paper turns red below ~pH = 5 An indicator…

How Do We Measure pH? For more accurate measurements, one uses a pH meter, which measures the voltage in the solution.

Earth’s Natural Litmus
Hydrangeas are blue when the acidity of the soil is between and red if the acidity is between A mix of colors can be seen between

pH is a logarithmic scale
If the pH goes up one, the concentration of hydrogen ions decreases by 10. The solution is more basic. If the pH goes down one, the concentration of hydrogen ions increases by 10. The solution is more acidic.

Examples If the pH changes from 3 to 4, how much more basic is the solution? 10 times If the pH = 3 and the hydrogen concentration increases by 100 times, what is the new pH? 1 If the pH = 8 and the hydrogen ion concentration decreases by 10,000 times, what is the new pH? 12

Strong Acids You will recall that the seven strong acids are HCl, HBr, HI, HNO3, H2SO4, H3PO4, and HClO4. These are, by definition, strong electrolytes and exist totally as ions in aqueous solution. For the monoprotic strong acids, [H3O+] = [acid]. To calculate the pH for strong acids use pH = -log [H3O+]

pH of Acids If the concentration of HA is 2.5x10-3M calculate the pH.
-log(2.5x10-3) = 2.60 If the concentration of HA is 3.5x10-8M calculate the pH. -log (3.5x10-8) = 7.46

Strong Bases Strong bases are the soluble hydroxides, which are the alkali metal and heavier alkaline earth metal hydroxides (Ca+2 and Ba2+). Again, these substances dissociate completely in aqueous solution. And [Base] = [OH-] PH = 14-(-log[OH-])

pH of Bases If the concentration of MOH is 1.5x10-4M calculate the pH.
14-[-log(1.5x10-4)] = 10.18 If the concentration of MOH is 2.2x10-13 M calculate the pH. 14-[-log (2.2x10-13)] = 1.34

Think What’s another way to calculate the pH of 2.2x10-13 M MOH?
1.0x10-14/2.2x10-3 = 0.045 -log(0.045) = 1.34

Find the pH, pOH, [H3O+], [OH-] of any solution. Compare the acidity or alkalinity of solutions based on pH values.

Video 11.4 Indicators

Identify possible indicators and use them to determine acidity or alkalinity (pH) of solutions using table M.

Table M

Indicators If NaOH is tested with methyl orange, what color will it be? At what pH will bromocrescol green turn yellow? What type of solution will turn bromothymol blue, yellow? At what pH will both bromothymol and thymol blue be yellow? yellow 3.8 acidic 7.6-8

Examples Why won’t methyl orange be good at determining the difference between an acid and a base? Which indicator is the best to test the difference between a strong and weak base? What color change will be seen if NaOH is added to HCl with methyl orange? What color change will be seen if nitric acid is added to lithium hydroxide using bromocrescol green? Bases and acids can both be yellow Thymol blue Red to yellow Blue to yellow

Identify possible indicators and use them to determine acidity or alkalinity (pH) of solutions using table M.

Video 11.5 pH of Weak Acids and Bases

Write mass action expressions (K) for acids and bases. Identify strong Acids and Bases using the equilibrium constant. Calculate the pH of weak acids and bases.

Weak Acids A−(aq) + H3O+(aq) HA(aq) + H2O(l) [H3O+] [A−] Ka = [HA]
For a generalized acid dissociation, the equilibrium expression would be This equilibrium constant is called the acid- dissociation constant, Ka. HA(aq) + H2O(l) A−(aq) + H3O+(aq) [H3O+] [A−] [HA] Ka =

Dissociation Constants
The greater the value of Ka, the stronger the acid. The strongest acids don’t report Ka values.

Weak Acids Recall that strong acids completely dissociate, or break down into ions. Therefore the pH is found by taking the –log of the concentration of the acid. Weak acids do not fully break down. Only a very small portion dissociate into ions. Therefore, we must calculate the number of ions that are present after equilibrium is established and then take the –log of the ions.

Calculating pH from Ka [H3O+] [C2H3O2-] [HC2H3O2] Ka =
Calculate the pH of a 0.30 M solution of acetic acid, HC2H3O2, at 25°C. HC2H3O2(aq) + H2O(l) H3O+(aq) + C2H3O2−(aq) Ka for acetic acid at 25°C is 1.8  10−5. First, write the Ka expression: [H3O+] [C2H3O2-] [HC2H3O2] Ka =

Calculating pH from Ka (x)2 1.8  10−5 = (0.30)
[H3O+] [C2H3O2-] [HC2H3O2] Ka = The acid will barely dissociate, so put the initial (given) acid concentration in as the equilibrium acid concentration. Solve for the ions. (1.8  10−5) (0.30) = x2 2.3  10−3 = x Now we know the ions, we can solve like a strong acid: pH = −log [H3O+] = −log (2.3  10−3) = 2.64

Example 1 (x)2 6.8  10−4 = (0.5) [H+] = 5.8  10−3 M
Calculate the pH of 0.50M HF. (Ka=6.8x10-4) [H3O+] [F−] [HF] Ka = (x)2 (0.5) 6.8  10−4 = (6.8  10−4) (0.5) = x2 3.4  10−5 = x2 5.8  10−3 = x [H+] = 5.8  10−3 M pH = −log (5.8  10−3)= 2.23

Example 2 (x)2 (0.9) 6.8  10−4 = (6.8  10−4) (0.9) = x2
Calculate the pH of 0.90M HF. (Ka=6.8x10-4) [H3O+] [F−] [HF] Ka = (x)2 (0.9) 6.8  10−4 = (6.8  10−4) (0.9) = x2 6.1  10−4 = x2 2.4  10−2 = x [H+] = 2.4  10−2 M pH = −log (2.4  10−2)= 1.61 Notice increasing the concentration decreased the pH

Example 3 (x)2 4.5  10−4 = (0.36) (4.5  10−4) (0.36) = x2
Calculate the pH of 0.36M HNO3. (Ka=4.5x10-4) (x)2 (0.36) 4.5  10−4 = [H3O+] [NO3−] [HNO3] Ka = (4.5  10−4) (0.36) = x2 1.6  10−4 = x2 1.3  10−2 = x [H+] = 1.3  10−2 M pH = −log (1.3  10−2)= 1.90

Weak Bases [HB] [OH−] [B−] Kb =
Bases react with water to produce hydroxide ion. [HB] [OH−] [B−] Kb =

Weak Bases Kb can be used to find [OH−] and, through it, pH of weak bases.

[NH4+] [OH−] [NH3] Kb = (x)2 (0.15) 1.8  10−5 = NH3(aq) + H2O(l)
What is the pH of a 0.15 M solution of NH3? NH3(aq) + H2O(l) NH4+(aq) + OH−(aq) [NH4+] [OH−] [NH3] Kb = (x)2 (0.15) 1.8  10−5 = (1.8  10−5) (0.15) = x2 2.7  10−6 = x2 1.6  10−3 = x [OH−] = 1.6  10−3 M pOH = −log (1.6  10−3)= 2.80 pH = − = 11.20

(Remember [H+][OH-] = 1.0x10-14
Ka and Kb Ka  Kb = Kw = 1.0x10-14 (Remember [H+][OH-] = 1.0x10-14

Write mass action expressions (K) for acids and bases. Identify strong Acids and Bases using the equilibrium constant. Calculate the pH of weak acids and bases.

Video 11.6 Acid Base Strength/Properties

Identify factors that define the strength of an acid or bases. Determine if a salt is acidic, basic, or neutral.

Factors Affecting Acid Strength
The more polar the bond between H and it’s neighboring element (or the weaker the H-X bond/IMF) the more acidic the compound (it will lose H+ better because the other bond is more attracted).

In oxyacids, in which an OH is bonded to another atom, Y, the more electronegative Y is, the more acidic the acid.

For a series of oxy-acids, acidity increases with the number of oxygen atoms. The increased polarity and strength of the O-X bonds makes the O-H bonds will be weaker and lose H+ easier.

pH of Salts The cation of a salt comes from the base in a reaction. The anion of the salt comes from the acid. NaOH + HCl  NaCl + H2O SA + SB  Neutral Salt WA + SB  Basic Salt SA+WB Acidic Salt WA + WB depends on Ka and Kb values. Any weak particles will react with the water (hydrolysis)

Identify factors that define the strength of an acid or bases. Determine if a salt is acidic, basic, or neutral.

Video 11.7 Titrations

Define primary standard, standard, standardization, titration, equivalence point, and end point. Calculate the unknown concentration at equivalence point, after a titration. Read a burette.

Neutralization ACID + BASE  SALT + WATER
Salt is another name for any ionic substance. They can conduct electricity. HCl + NaOH  NaCl + H2O acid base salt

Titrations Titrations are used to find the concentration of a solution. Usually the concentration of the base is unknown, so a known acid is added to the base until the solution neutralizes. The endpoint, or the point where the solution is neutral and the titration is over, is marked by a faint pink color in the solution (due to phenolphthalein).

Titration A known concentration of base (or acid) is slowly added to a solution of an unknown acid (or base) to determine it’s concentration. This is also known as standardization if a primary standard (solid, where the moles are constant) is used to determine the molarity of a substance.

Titration A pH meter or indicators are used to determine when the solution has reached the equivalence point, or neutralization point, at which the moles of acid equals that of base. The endpoint is when the indicator actually changes color. You want the endpoint close to the equivalence point.

Titrations To calculate the molarity of the unknown use the formula on Table T: MAVA=MBVB

Examples MAVA=MBVB What is the molarity of NaOH if 100.mL of
3.00M HCl is titrated with 200.mL of NaOH? MAVA=MBVB (3.00)(100) = x(200) x = 1.50M What is the molarity of 100.0mL HCl if it is neutralized by 250mL of 2.0M NaOH? (x)(100) = (2.0)(250) x = 5.0M

Exceptions The titration formula only works if the acid and the base have equal numbers of H+ and OH-. If not, the acid side must be multiplied by the number of hydrogen ions and the base side must be multiplied by the number of hydroxide ions in the formulas. What is the molarity of a solution of Ca(OH)2 if 750mL of it is titrated with 250mL of 3.5M H3PO4? 3MAVA = 2MBVB 3(3.5)(250) = 2(x)(750) x = 1.75M

How much acid is in each burette?
0.6mL 15.4mL How much acid was released? = 14.8mL released from the burette.

Define primary standard, standard, standardization, titration, equivalence point, and end point. Calculate the unknown concentration at equivalence point, after a titration. Read a burette.

Video 11.8 Titration Curves

Identify the standard solution in a titration curve Identify the equivalence point and strength of the acid and base in a titration curve. Estimate the Ka/Kb using the half titration point.

Titration of Strong Acid with Strong Base
From the start of the titration to near the equivalence point, the pH goes up slowly. Just before and after the equivalence point, the pH increases rapidly.

Titration of Strong Acid with Strong Base
At the equivalence point, moles acid = moles base, and the solution contains only water and the salt. The pH equals 7.

The Common-Ion Effect HC2H3O2 + H2O(l) H3O+ + C2H3O2−
If acetate ion is added to the solution, LeChâtelier says the equilibrium will shift to the _______. A weak acid ionizes______when a common ion is added to the solution. Left Less

Buffers Buffers are solutions of a weak conjugate acid-base pair.
They are resistant to pH changes, even when strong acid or base is added. Buffers are used to maintain pH balances. Effectiveness depends on volume used.

Buffers If a small amount of hydroxide is added to an equimolar solution of HF in NaF, for example, the HF reacts with the OH− to make F− and water.

Buffers If acid is added, the F− reacts to form HF and water.

Which are buffer systems?
KH2PO4/H3PO4 NaClO3/HClO3 CH3COOH/CH3COONa Buffers must contain weak acids or bases and their conjugates.

Review Buffers A buffer resists changes in pH since it contains both acidic and basic species. If OH- is added, it reacts with the _____ part of the buffer. If H+ is added, it reacts with the _____ part of the buffer. The acidic and basic species in the buffer cannot consume each other in a _________ reaction, or it won’t work. This is why the buffer must have ___ acids and ____ bases. acidic basic neutralization weak weak

Titration of Weak Acid with Strong Base
Now, the conjugate base of the acid affects the pH when it is formed. It creates a buffer! The pH at the equivalence point will be >7. The weak acid’s titration curve starts at a higher pH and has a slight bump.

Titration of Weak Base with Strong Acid
Again, the conjugate acid affects the pH at the beginning of the reaction due to the formation of a buffer. The pH at the equivalence point in these titrations is < 7. The pH starts very high with a base and should have a slight bump.

The Half Titration Point
The half titration point of a weak acid can be found by halving the volume of NaOH needed to obtain neutralization. At this point the amount of original acid equals the amount of conjugate base.

Ka= [H+][A-] [HA] Ka=[H+] If ½ Eq = 5, [H+] = 1x10-5 so, Ka=1x10-5

Just remember, at half titration: pH=pKa

Identify the standard solution in a titration curve Identify the equivalence point and strength of the acid and base in a titration curve. Estimate the Ka/Kb using the half titration point.

Video 11.1 Acids and Bases.

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Video 11.1 Acids and Bases.

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