CHEM 14A Chapter Notes - Chapter J.2: Deprotonation, Protonation, Acid Strength

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turquoisemoose666

6 Dec 2020

School

UCLA

Department

Chemistry and Biochemistry

Course

CHEM 14A

Professor

Dr.Laurence Lavelle

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Document Summary

Electrolytes are classified as strong or weak according to the extent to which they are present as ions in solution. Completely deprotonated means that almost every acid molecule or ion has lost its acidic hydrogen atom by transferring it as a hydrogen ion to a solvent molecule. Completely protonated means that almost every base molecule or ion has acquired a proton. Incompletely deprotonated means that only a fraction of the acid molecules or ions have lost acidic hydrogen atoms as protons. Incompletely protonated means that only a tiny fraction of the base molecules or ions have acquired protons. Most common strong bases are oxide ions and hydroxide ions. Alkali metal and alkaline earth metal oxides and hydroxides. To summarize, strong acids (the acids listed in table j. 1) are completely deprotonated in solution; weak acids (most other acids) are not. Strong bases (the metal oxides and hydroxides listed in table j. 1) are completely protonated in solution.

Related Questions

Part A A fuel cell differs from a battery because the current is being generated from a reduction-oxidation (redox) reaction in which the reactant is consumed. A classic example of a fuel cell is the phosphoric acid fuel cell (PAFC), which have been used in stationary power generators, buses, and even submarines. Fuel cells require an electrolyte (similar to batteries) to carry electrically charged species between the electrodes, and a PAFC utilizes phosphoric acid (H3PO4) as its electrolyte. Phosphoric acid has three acidic protons that can dissociate in water, as highlighted in red here: H3PO4. However, it is considered a weak acid because some of the acid will remain in molecular form when dissolved in water. Write the complete ionic equation that depicts the dissociation of the first proton for the weak acid phosphoric acid (H3PO4). Make certain you include all charges for any ions produced. Express you answer as a chemical equation including phases. Titrations as analytical neutralization reactions Neutralization reactions can occur between any combination of weak/strong acids and bases. When we want to determine the concentration of a base in solution, we can apply the concept of neutralization by reacting the base with a known amount of an acid in the presence of an indicator (a compound that changes color depending on the pH). This method of determining the concentration of a base is called acid-base titration, and it can also be used to determine the concentration of an acid by titrating it with a known amount of base. A common acid to use for titrations is potassium hydrogen phthalate (written as KHP or KHPh; do not confuse the P for phosphorus). The reason KHP is common is because it is solid at room temperature, and we can accurately mass the amount that is added to solution. As a soluble salt, KHP completely dissociates to form a potassium ion (K+) and an acidic hydrogen phthalate ion (HPâ), which is a weak acid that has only one acidic hydrogen. This ionization is depicted below. KHP(s) â K+(aq)+HPâ(aq) Although ionization (the formation of ions) occurs as a result of acid dissociation for most weak acids, KHP ionizes before acid dissociation occurs. In other words, HPâ is a weak acid ion. Part C If you wanted to determine how much KOH (potassium hydroxide; a strong base) was dissolved in water, you could titrate it with a known amount of KHP (potassium hydrogen phthalate; a weak acid) until a color change indicates that the base is completely neutralized. The molecular equation below depicts this neutralization reaction. KHP(aq)+KOH(aq)âK2P(aq)+H2O(l) Write the net ionic equation that depicts the neutralization of the base by potassium hydrogen phthalate. Instead of using the entire chemical formula of phthalate, use HPâ to represent the acidic hydrogen phthalate ion and P2â to represent the phthalate ion. Express your answer as a net ionic equation including phases.

QUESTION 1

Which of the following substances is most acidic?

		Cowâs milk - pH 6.6
		Apple juice - pH 3.0
		Tomato juice - pH 4.5
		Distilled water - pH 7.0

2 points

QUESTION 2

A buffer is a substance that converts:

Alkaline solutions to neutral solutions.

Acidic solutions to alkaline solutions.

Strong bases or acids to weak bases or acids.

Acidic solutions to neutral solutions.

QUESTION 3

How does a solution of pH 7 compare to a solution of pH 10?

		pH 7 has more OH^- ions
		pH 7 is more basic
		pH 7 has a higher pH
		pH 7 has more H⁺ ions

2 points

QUESTION 4

According to the Arrhenius Theory of acids and bases, a base is:

A substance that decrease the amount of hydroxide ions present.

A substance that increases the amount of hydroxide ions present when it is dissolved in water.

A substance that increase the amount of hydrogen ions present.

A substance that increases the amount of hydrogen ions present when it is dissolved in water.

QUESTION 5

Which of the following are true regarding acid-base titrations?

		They are used to determine the concentration of an unknown acid or base.
		The acid or the base must be a standard solution with a known concentration
		The standard solution is titrated into the chemical with the unknown concentration.
		All of the above are true regarding acid-base titrations.

2 points

QUESTION 6

Which of the following is not an example of a neutralization reaction?

HCl + NaOH NaCl +H₂O

2 NaOH + H₂CO₃ Na₂CO₃ +2 H₂O

H₂SO₄ + 2 NH₄OH (NH₄)₂SO₄ + 2 H₂O

2 AgNO₃ + Zn 2Ag +Zn(NO₃)₂

QUESTION 7

Which acid base theory defines an acid as a proton donor and a base as a proton acceptor?

		Arrhenius Theory
		Bronsted-Lowery Theory
		Conjugate Pairs Theory
		Lewis Theory

2 points

QUESTION 8

The Arrhenius Theory states that acid-base reactions create a solution composed of a water and a salt. This is because mixing an acidic and basic solution together yields:

A neutral solution

A solution with a pH at or near 7

Both A and B are correct

QUESTION 9

Which of the following correctly defines a standard solution?

		A solution which always has the same molecular geometry.
		A solution in which the concentration is precisely known.
		A solution in which the volume is precisely known.
		A solution which always has different levels of solutes and solvents.

2 points

QUESTION 10

Acids naturally present in food are generally safe to eat because they are usually:

		Concentrated
		Dilute
		Weak
		Strong

None of the above are correct

I NEED HELP ANSWERING QUESTION 4 AND 5!!

Experimental Determination of pKa of buffers In this lab you. the student, will titrate several different buffer solutions in order to determine their pKa. Recall from class that (or from the water chemistry chapter of your textbook) that the pKa of a weak acid/base is equal to the pH of the solution when it population of molecules are 50% protonated/deprotonated. In this lab you will be given a butter solution (typically at low pH) and will titrate with 1 M potassium hydroxide (KOH) while measuring the pH during each stage of the process. You will then need to generate from your data a titration plot and then to extrapolate from there the pKa of the weak acid/base. General Procedure Tentative buffers to be tested: water (N = 1), 0.4 M Acetic Acid (N = 3), 0.4 M Glycine (N = 3), [N is the number if replicates] As in any lab put on your personal protective equipment (PPE): lab coal, safety glasses, and gloves Setup burette with 1 M KOH. Be careful not to spill. In a small beaker (100 mL or smaller) pour in 40 mL of buffer to be tested. Titrate until pH passes 12. Be consistent with your dispensing of KOH so your replicate trials line up appropriately. Tabulate and graph data (use formatting that is consistent with Lab 1 and include standard deviation in tables and graphs) Questions and concepts to address in the lab report Describe the buffering capacity of water based on your titration. How do they compare to Acetic Acid and Glycine? For Acetic Acid and Glycine Buffers: For titrations with N = 3. how can you account significant differences between data points? What were the pKas of your buffers based on your titration? How do they compare to those reported in the literature? How do you account for difference in values? Draw the three possible ionic forms of each buffer. Identify which of these predominates at pH 1, pH 6, and pH 11. What are the chemical equations that show how each buffers solutions at either pKa by consuming amounts of added hydrogen or hydroxide ions.

CHEM 14A Chapter Notes - Chapter J.2: Deprotonation, Protonation, Acid Strength

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