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1. The document provides 7 chemistry problems involving calculations of pH, concentrations of products and reactants from acid-base reactions involving strong acids and bases, and identification of types of acids. 2. Problems 1-3 involve calculating the pH of strong acid solutions given molar concentrations. 3. Problems 4-6 require using given volumes, molarities, and extent of ionization to calculate concentrations of products and reactants. 4. Problems 7-9 ask about common weak bases, relationships between types of acids, and identification of more acidic salt solutions.
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![GENERAL CHEMISTRY-II (1412)
S.I. # 24
Calculate the pH of each of the following strong acid solutions:
1. 0.00835 M HNO3
2. 0.525 M HClO4
3. 5.00 mL of 1 M HCl diluted to 0.500 L
4. a mixture formed by adding 50.0 mL of 0.020 M HCl to 150 mL of
0.010 M HI.
5. A 0.100 M solution of brmoacetic acid (BrCH2COOH) is 13.2% ionized.
Calculate [H+] , [BrCH2COO-], and [BrCH2COOH].](https://image.slidesharecdn.com/chem-ii-si-24-1272476882-phpapp02/85/24-1-320.jpg)
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This document provides examples of calculating the pH of various strong acid solutions, including nitric acid, perchloric acid, hydrochloric acid, a mixture of hydrochloric acid and hydroiodic acid, and a solution of bromoacetic acid that is partially ionized. It also asks questions about weak acids, Lewis acids, and which compounds form more acidic solutions based on cation properties.
#19 key
1. 18.9 grams of K2Cr2O7 are present in 50.0 mL of 0.360 M K2Cr2O7 solution.
2. The molarity of the (NH4)2SO4 solution formed from dissolving 4.28 grams of (NH4)2SO4 in 300 mL of water is 0.108 M.
3. 58.7 mL of 0.240M CuSO4 solution contains 2.25 grams of CuSO4 solute.
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- Strong acids and bases completely ionize in solution, allowing their concentrations to determine H3O+ or OH- concentrations.
- Weak acids and bases only partially ionize in solution, with their extent of dissociation determined by their acid or base ionization constant (Ka or Kb).
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Acids, bases + neutralization
neutralization (or neutralisation, see spelling differences) is a chemical reaction in which an acid and a base react to form a salt. Water is frequently, but not necessarily, produced as well. Neutralizations with Arrhenius acids and bases always produce water where acid–alkali reactions produce water and a metal salt.
Acid and bases differentiated assignment p h_answers
Student A's statement that a strong acid has a low pH and a weak acid has a high pH is correct only if comparing acids of the same concentration and basicity. Student B's statement that a dibasic acid has a lower pH than a monobasic acid is only correct if the acids are equally strong and concentrated. The pH of a 0.01 M NaOH solution is calculated to be 12, and the pH of a 0.1 M HCl solution before titration with 0.1 M NaOH is calculated to be 1. After adding 5 cm3 and 24 cm3 of NaOH, the pH is calculated to be 1.18 and 2.69 respectively. The pH after 26 cm3 of
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This document provides examples of calculating the pH of various strong acid solutions, including nitric acid, perchloric acid, hydrochloric acid, a mixture of hydrochloric acid and hydroiodic acid, and a solution of bromoacetic acid that is partially ionized. It also asks questions about weak acids, Lewis acids, and which compounds form more acidic solutions based on cation properties.
#19 key
1. 18.9 grams of K2Cr2O7 are present in 50.0 mL of 0.360 M K2Cr2O7 solution.
2. The molarity of the (NH4)2SO4 solution formed from dissolving 4.28 grams of (NH4)2SO4 in 300 mL of water is 0.108 M.
3. 58.7 mL of 0.240M CuSO4 solution contains 2.25 grams of CuSO4 solute.
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- Weak acids and bases only partially ionize in solution, with their extent of dissociation determined by their acid or base ionization constant (Ka or Kb).
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Student A's statement that a strong acid has a low pH and a weak acid has a high pH is correct only if comparing acids of the same concentration and basicity. Student B's statement that a dibasic acid has a lower pH than a monobasic acid is only correct if the acids are equally strong and concentrated. The pH of a 0.01 M NaOH solution is calculated to be 12, and the pH of a 0.1 M HCl solution before titration with 0.1 M NaOH is calculated to be 1. After adding 5 cm3 and 24 cm3 of NaOH, the pH is calculated to be 1.18 and 2.69 respectively. The pH after 26 cm3 of
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- pH is determined by the negative log of the hydrogen ion (H+) concentration in a solution, while pOH is determined by the negative log of the hydroxide ion (OH-) concentration.
- The pH of water is 7 because at 25°C, the H+ concentration is 1.0 x 10-7 and the log of that value is 7. However, pH varies with temperature as the self-ionization constant of water (Kw) changes with temperature.
- pOH can be calculated from pH using the relationship that pH + pOH must equal 14 for any aqueous solution.
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2) The corresponding pOH values are given as the complement to the pH values.
3) An explanation is provided that even basic solutions contain H+ ions and even acidic solutions contain OH- ions due to autoionization of water.
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2) Common substances span the pH scale from 0 to 14, with examples like stomach acid at pH 1 and bleach/drain cleaners at pH 12-13.
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This document provides calculations to determine the pH at different points in a titration between 0.3 M HC2H3O2(aq) and 0.3 M NaOH(aq). It also provides similar calculations for a titration between 0.1 M NH3(aq) and 0.1 M HCl(aq). ICE tables and equilibrium expressions are used along with mole calculations to determine pH before, at the equivalence point, and after the equivalence point for both strong-weak acid base titrations. Concentration expressions and additional ICE tables are set up as needed when the reaction reaches the equivalence point and changes direction.
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This document discusses water equilibrium and acid-base chemistry. It defines the water ion-product constant Kw as equal to the hydrogen ion concentration times the hydroxide ion concentration. For pure water at 25°C, Kw is 1.0 x 10-14, so the hydrogen and hydroxide ion concentrations are both 1.0 x 10-7 M. The document then provides examples of using Kw to calculate ion concentrations in acidic, basic, and neutral solutions and determine if they are acidic, basic, or neutral based on the relative hydrogen and hydroxide ion concentrations.
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The document describes an acid-base titration experiment involving titrating a weak acid with a strong base. It provides 3 key points:
1) The document outlines the experimental procedure, reagents, and purpose which is to determine the concentration of an unknown acid through titration with a strong base such as NaOH.
2) It discusses acid-base titration curves and indicators, and how indicators work through acid-base equilibria. Buffer solutions are also introduced.
3) Several examples of calculations are shown for strong acid-strong base and strong acid-weak base titrations, including determining concentrations and pH at various points in the titration. Step-by-step workings are provided.
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The document discusses making assumptions to simplify equilibrium calculations when the equilibrium constant (Keq) is extremely small or large. It provides an example of calculating equilibrium concentrations for the decomposition of CO2 at 2000°C assuming negligible decomposition due to the small Keq value of 6.40 x 10-7. The assumptions allow obtaining approximate equilibrium concentrations without solving the full cubic equation.
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The document provides information about acid-base chemistry including:
1) Calculating the concentration of H+ ions and determining if solutions are acidic or basic based on the relative concentrations of H+ and OH- ions.
2) Defining pH as the negative logarithm of the H+ ion concentration and calculating pH values from given [H+] values.
3) Relating pH and pOH through the equation pH + pOH = 14.
4) Listing the 7 strong acids and their formulas as well as the 7 strong bases and their formulas.
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- pH is determined by the negative log of the hydrogen ion (H+) concentration in a solution, while pOH is determined by the negative log of the hydroxide ion (OH-) concentration.
- The pH of water is 7 because at 25°C, the H+ concentration is 1.0 x 10-7 and the log of that value is 7. However, pH varies with temperature as the self-ionization constant of water (Kw) changes with temperature.
- pOH can be calculated from pH using the relationship that pH + pOH must equal 14 for any aqueous solution.
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1. Acids donate hydrogen ions in water, forming hydronium ions, and have characteristics like turning litmus paper red and tasting sour. Bases form hydroxide ions in water and have opposite characteristics.
2. The pH scale measures how acidic or basic a substance is based on its hydronium ion concentration, with values from 0-14. More acidic substances have fewer hydronium ions and lower pH values closer to 0.
3. A substance is neutral when it has equal concentrations of hydronium and hydroxide ions, resulting in a pH of 7.
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This document provides a chemistry problem set on calculating pH and pOH values of various acid and base solutions. It includes:
1) Calculations of the pH of solutions such as 4.5 x 10-3 M HBr and dilutions of 6.0 M HCl.
2) The corresponding pOH values are given as the complement to the pH values.
3) An explanation is provided that even basic solutions contain H+ ions and even acidic solutions contain OH- ions due to autoionization of water.
4) A challenging calculation of pH for a very dilute HBr solution of 1.5 x 10-10 M, which intuitively should not be basic but the actual pH is around 7
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1) Water can act as both an acid and a base, donating or accepting protons in chemical reactions.
2) Water undergoes autoionization, reacting with itself to form hydronium and hydroxide ions. The equilibrium constant for this reaction is known as Kw.
3) In pure water at 25°C, Kw equals 1.0 × 10-14, meaning the concentrations of hydronium and hydroxide ions are both 1.0 × 10-7.
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1. When ions dissolve in water, they are surrounded by water molecules (hydrated).
2. Precipitation will occur when Na2CO3 and AgNO3 or FeSO4 and Pb(NO3)2 are mixed, but not NaNO3 and NiSO4.
3. Acids donate H+ ions, bases accept H+ ions. A monoprotic acid has one ionizable H, a diprotic acid has two. Strong acids fully ionize, weak acids only partially ionize.
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The document provides examples for calculating percentage composition by mass of elements in compounds. It gives the formula for finding percentage composition and worked examples calculating the percentage of elements like nitrogen and hydrogen in ammonia, hydrogen and oxygen in hydrogen peroxide, and water in copper(II) sulfate and iron(III) chloride hexahydrate.
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This document discusses solubility equilibria and the formation of precipitates. It defines key terms like solubility product constant (Ksp), explains how to calculate Ksp values from molar solubility and vice versa, and shows examples of using Ksp to determine whether precipitates will form when solutions are mixed. The key points are that solubility is dependent on equilibrium, saturated solutions have concentrations where Ksp = Q, and precipitates form when mixing produces Q > Ksp (supersaturation).
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3. The solutions are labeled as: KCl neutral, NaNO2 basic, (CH3)2NH2Cl acidic.
Chapter 16 Powerpoint - student version.ppt
This document discusses acid-base equilibria and the key concepts of acids, bases, pH and pOH. It begins with a review of the Arrhenius and Brønsted-Lowry definitions of acids and bases. Key points covered include: water can act as both an acid and a base; conjugate acid-base pairs are formed in acid-base reactions; and the pH scale relates hydrogen ion concentration to acidity. Strong acids and bases are defined as completely dissociating in water.
Acids-Bases-pH-PPT-2018.ppt.organic-chemistry
This document discusses the chemistry of acids and bases. It defines acids as substances that produce hydrogen ions (H+) or hydronium ions (H3O+) in water, and bases as substances that produce hydroxide ions (OH-). It provides examples of common acids and bases and discusses their properties. Key topics covered include the Arrhenius, Brønsted-Lowry, and Lewis definitions of acids and bases; acid-base reactions; pH and pOH scales; strong and weak acids and bases; and acid-base equilibria.
Chapter 16.1 and 2 : Acid-Base Titrations and pH
The document discusses acid-base titrations and pH. It defines pH as the negative logarithm of hydronium ion concentration and explains how to calculate pH from [H3O+] and vice versa. It also describes how to perform and calculate molarity from a titration experiment by determining moles of acid/base reacted using volume and molarity of the titrant. Sample titration problems are worked through demonstrating these concepts.
1. acid base
In the pharmaceutical inorganic chemistry as per PCI for B-pharm FY students easy to understand the topic is Acid and base
Acid Base -Part 2-.pptx
This document discusses different types of hydrolysis reactions and provides examples of acid-base titrations. It describes salt, acid, and base hydrolysis reactions. For acid-base titrations, it explains that a strong acid titrated with a strong base will have a high starting pH declining to 7 at the equivalence point then continuing to a low pH. It also discusses using pH meters and acid-base indicators to determine the equivalence point.
Strong weak acids and bases
Acids are divided into two categories based on the ease with which they can donate protons to the solvent: i) strong acids and ii) weak acids
Strong acids are acids that completely dissociate in water. The reaction of an acid with its solvent (typically H2O) is called an acid dissociation reaction.
Weak acids are acids that dissociate partially in water. The extent of dissociation is given by the equilibrium constant.
Note:
A measure of the relative strength of an acid is: i) the equilibrium constant ka of the dissociation reaction of the acid in water (depends on temperature) ii) the degree of dissociation α of the acid in water (depends on the concentration of the acid an on temperature).
Acids and bases p pt
This document defines acids and bases according to the Arrhenius model and discusses their properties in aqueous solutions. It also explains acid-base titrations for determining concentrations and defines pH. Buffers are formed from a weak acid and its conjugate base or weak base and conjugate acid and do not experience large pH changes upon addition of acids or bases.
Similar to #24 (20)
Csir chemistry acid base equilibrium question paper
Csir chemistry acid base equilibrium question paper
Lect w8 152 - ka and kb calculations_summary exercises_alg
Lect w8 152 - ka and kb calculations_summary exercises_alg
More from jessieo387_1412
#27
This document contains questions about oxidation-reduction reactions and voltaic cells. It asks the student to identify oxidizing and reducing agents, write half reactions, identify anodes and cathodes, and determine the direction of electron and ion flow. It also contains questions about standard cell potentials, units of electrical potential, and a half reaction for a hydrogen electrode.
#26 Key
1. The document provides steps to convert between solubility in grams per liter, molar solubility in mol/L, molar concentrations of ions, and Ksp.
2. It gives the definitions of a reducing agent, oxidizing agent, and half-reaction. A reducing agent causes reduction, an oxidizing agent causes oxidation. Half-reactions involve electrons.
3. An anode is where oxidation occurs and a cathode is where reduction occurs in a voltaic cell. Electrons flow from the anode through the external circuit to the cathode.
#26
The document discusses solubility and solubility products, oxidation-reduction reactions, and voltaic cells. It defines key concepts like reducing agents, oxidizing agents, and half-reactions. It also discusses standard reduction potentials, how to determine if a reaction involves oxidation-reduction, and how to balance half-reactions to indicate oxidation or reduction. Sample problems are provided to calculate equilibrium constants from solubility products and oxidation states in reactions.
#25 Key
1. The common ion effect decreases the ionization of a weak electrolyte when a strong electrolyte with a common ion is added to the solution.
2. Buffer capacity refers to how much acid or base can be neutralized before the pH begins to change significantly.
3. A buffer solution contains a weak acid or base and its salt. It has the ability to resist changes in pH.
4. The pH range of a buffer depends on the acid's Ka value and the relative concentrations of the acid and base in the buffer solution. It is the range over which the buffer acts effectively.
#23 Key
The document discusses chemical reactions and calculations involving acids, bases, and their dissociation in solution. It begins by explaining that the percent dissociation of a weak acid decreases as its concentration increases. It then shows examples of calculating percent dissociation and acid dissociation constants (Ka) for hydrocyanic acid solutions. Similarly, it demonstrates calculating percent dissociation and pKa for ammonia solutions. Finally, it asks and answers questions about whether and why certain acid-base reactions will or will not occur when the solutions are mixed, such as sodium chloride dissolving in water.
#23
The document contains a chemistry practice exam with 10 multiple-choice and short-answer questions testing concepts such as how the percent dissociation of a weak acid changes with concentration, calculating percent dissociation and acid dissociation constants given various parameters, determining pH and pOH values of solutions, and identifying chemical reactions between acids and bases in aqueous solutions.
#20
The document provides a chemistry problem set involving calculations of hydrogen ion (H+) and hydroxide ion (OH-) concentrations in solution, definitions and calculations of pH and pOH, and properties of strong acids, weak acids, and weak bases. It asks students to calculate pH and concentrations of H+ given OH- concentrations and vice versa. It also asks students to list strong acids and bases, properties of strong and weak acids, what weak acids and bases have in common, and why HCl is a strong acid but weak base.
#19 Key
This document discusses Bronsted-Lowry acid-base chemistry. It defines Bronsted-Lowry acids as proton donors and bases as proton acceptors. It explains that conjugate acids are formed by adding a proton to a base and conjugate bases are formed by removing a proton from an acid. Several examples of conjugate acid-base pairs are given. The document also states that the stronger the acid, the weaker its conjugate base, and the stronger the base, the weaker its conjugate acid. It describes how the position of equilibrium favors transfer of a proton to the stronger base. Finally, it provides the autoionization reaction of water and defines the ion product constant, Kw, for water.
#19
This document covers Bronsted-Lowry acids and bases, conjugate acids and bases, and acid-base equilibrium. It defines Bronsted-Lowry acids as proton donors and bases as proton acceptors. Conjugate acids are formed when a base gains a proton, and conjugate bases are formed when an acid loses a proton. Several examples of conjugate acid-base pairs are given. The document also discusses how the strength of an acid or base determines the position of acid-base equilibrium and the autoionization process and ion product constant for water.
#18 Key
The document contains 6 chemistry problems involving gas phase chemical equilibria calculations:
1) Calculating the equilibrium partial pressure of N2O4 and the equilibrium constant Kp for the reaction N2O4 ⇌ 2NO2.
2) Calculating the equilibrium concentrations of H2, Br2, and HBr and the equilibrium constant Kc for the reaction H2 + Br2 ⇌ 2HBr.
3) Calculating the equilibrium constant Kc for the reaction SO2 + Cl2 ⇌ SO2Cl2 given the value of Kp.
4) Calculating the value of Kp for the reverse reaction of 2NO2 ⇌
#18
This document contains 6 chemistry problems involving gas phase chemical equilibria. Problem 1 involves calculating the equilibrium partial pressure and equilibrium constant for the reaction N2O4(g) ⇌ 2NO2(g). Problem 2 involves calculating equilibrium concentrations and the equilibrium constant for the reaction H2(g) + Br2(g) ⇌ 2HBr(g). Problem 3 asks to calculate the equilibrium constant for the reaction SO2(g) + Cl2(g) ⇌ SO2Cl2(g) given the value of Kp at a specific temperature.
#17 Key
1. The document contains 7 chemistry problems involving calculation of equilibrium constants (Keq and Kc) for various chemical reactions.
2. Problem 3 calculates Keq as 1.84x10-2 for the decomposition of HI gas into H2 and I2 gases at 425C, given the partial pressures of reactants and products at equilibrium.
3. Problem 7 involves a multiple step chemical equilibrium problem, calculating the initial and equilibrium partial pressures and concentrations of gases in a reaction between CO2, H2 and H2O, and determining the equilibrium constant Kp.
#17
This document contains 7 questions regarding chemical equilibrium calculations. The questions involve writing equilibrium constants, calculating equilibrium constants based on given partial pressures or concentrations at equilibrium, and determining unknown values at equilibrium given some known values. Specifically, it asks the student to:
1) Write the equilibrium constant expression for a dissolution reaction.
2) Write the equilibrium constant expression for a complex ion formation reaction.
3) Calculate the equilibrium constant for a decomposition reaction given partial pressures at equilibrium.
4) Calculate the equilibrium constant for a decomposition reaction given partial pressures at equilibrium.
5) Calculate the equilibrium constant for a reaction given partial pressures in a container at equilibrium.
6) Calculate unknown concentrations and the equilibrium constant for a reaction given initial
SI #16 Key
The document contains a chemistry exam with multiple choice questions about reaction rates and stoichiometry. Question 1 asks about the relationship between the rates of disappearance of H2 and appearance of HF in the reaction H2 + F2 → 2HF. Question 2 asks about the relationship between the average rate of appearance of product B and disappearance of reactant A in the reaction 3A → 2B. Question 3 asks about the order of reactants and overall order of the reaction for the hydrolysis of t-butyliodide.
SI #16
The document contains a chemistry practice problem set with 10 multiple choice questions about reaction rates and rate laws. The questions cover topics such as how the rates of reactants and products are related in a chemical reaction, how to write rate laws from reaction equations, determining reaction orders from rate equations, and calculating rate constants. It also includes a nomenclature list of 38 chemical species.
Practice Test 3
1. The document contains a practice test with multiple choice questions about chemical kinetics and equilibrium.
2. Questions cover topics like reaction rates, rate laws, reaction orders, equilibrium constants, Le Chatelier's principle, and the Haber process.
3. There is also a bonus section with chemical nomenclature questions matching names to formulas.
Si #15 Key
The document defines key concepts related to reaction rates in chemistry:
1) Reaction rate is defined as the change in concentration of a reactant or product with time.
2) The rate formula relates the rate of a reaction to the change in concentration of a reactant or product over time.
3) The concentration of reactants is directly proportional to the rate of reaction according to the rate formula.
4) A catalyst increases the reaction rate by participating in the reaction without being consumed and allowing collisions between molecules to result in reaction at lower energies.
Si #15
The document discusses reaction rates and kinetics. It defines reaction rate and states that the concentration of reactants is proportional to the rate of reaction. It also defines the rate constant and rate law formula. A catalyst affects the reaction rate by lowering the activation energy, increasing the frequency of collisions, and increasing the probability of successful collisions. Enzymes are biological catalysts and the collision theory states that for a reaction to occur, reactant particles must collide with adequate energy.
SI #13 Key
This document contains information about several chemistry concepts and example calculations:
1) It provides calculations to determine the amount of N2O5 remaining and reaction time for its decomposition reaction at 70°C.
2) It discusses how reaction rates change with changes in reactant concentrations for a reaction following a rate law.
3) It gives the rate law and calculates rates for the decomposition of N2O5 in carbon tetrachloride.
4) It provides the rate law and calculates rates for a reaction that is 1st order in H2 and 2nd order in NO.
5) It states that the law of mass action expresses the relationship between reactant and product concentrations at equilibrium.
Si #12 Key
1. Chemical kinetics is the study of time vs. the rate of chemical change. The three factors that affect chemical reactions are concentration, temperature, and catalysts.
2. Instantaneous rate is the slope of a line drawn tangent to the concentration-time curve at a specific time. Beer's Law states that absorption of electromagnetic radiation by a substance is directly proportional to its concentration.
3. A first-order reaction has a rate proportional to the concentration of a single reactant raised to the first power. The rate is equal to k[A] and a plot of ln[A] versus time gives a straight line with a slope of -k.
More from jessieo387_1412 (20)
#24
- 1. GENERAL CHEMISTRY-II (1412) S.I. # 24 Calculate the pH of each of the following strong acid solutions: 1. 0.00835 M HNO3 2. 0.525 M HClO4 3. 5.00 mL of 1 M HCl diluted to 0.500 L 4. a mixture formed by adding 50.0 mL of 0.020 M HCl to 150 mL of 0.010 M HI. 5. A 0.100 M solution of brmoacetic acid (BrCH2COOH) is 13.2% ionized. Calculate [H+] , [BrCH2COO-], and [BrCH2COOH].
- 2. 6. How many moles of HF (Ka = 6.8 x 10-4) must be present in 0.200 L to form a solution with a pH of 3.25? 7. What are tow kinds of molecules or ions that commonly function as weak bases? 8. If a substance is a Lewis acid, is it necessarily a Bronsted-Lowery acid? Is it necessarily an Arrhenius acid? Explain. 9. Which member of each pair produces the more acidic aqueous solution: a. ZnBr2 or CdCl2 b. CuCl or Cu(NO3)2 c. Ca(NO3)2 or NiBr2