06 microscale titration_fa06
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The document discusses acid-base titration, including the analyte, standard, neutralization reaction, equivalence point, and procedure. It explains that an acid-base titration involves titrating an acid with a base until the moles of H+ equal the moles of OH-. It also discusses using a pH sensor and indicator to determine the equivalence point, and using stoichiometry to calculate unknown concentrations from experimental data.
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Acids, bases + neutralization
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This document provides information about acids and bases including:
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1. The document discusses acids and bases according to the Arrhenius and Brønsted-Lowry definitions. It describes the key characteristics of acids and bases and provides examples of strong and weak acids and bases.
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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.
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The document discusses acid-base titrations, which are procedures used to determine the concentration of an acid or base by titrating with the other substance until reaching the equivalence point where moles of H3O+ and OH- are equal, signaling the end point where an indicator changes color and all the acid or base has been neutralized, forming water and a salt. It also examines the pH ranges that different acid-base indicators change color, types of acid-base reactions, and examples of strong vs. weak acid-base titrations.
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This document provides information about acids and bases including:
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2) Types of acids and bases as strong or weak.
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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.
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The document discusses acid-base titrations, which are procedures used to determine the concentration of an acid or base by titrating with the other substance until reaching the equivalence point where moles of H3O+ and OH- are equal, signaling the end point where an indicator changes color and all the acid or base has been neutralized, forming water and a salt. It also examines the pH ranges that different acid-base indicators change color, types of acid-base reactions, and examples of strong vs. weak acid-base titrations.
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* HCl is a strong acid and will titrate first
* Its equivalence point was at 35.00 mL of NaOH
* NaOH concentration is 0.100 M
* Moles of NaOH used = Volume x Concentration
= 0.03500 L x 0.100 mol/L = 0.003500 mol
* Moles of HCl = Moles of NaOH used = 0.003500 mol
* H3PO4 is a weak acid and will titrate second
* Its equivalence point was at 50.00 mL of NaOH
* Additional NaOH used = 50.00 mL - 35.00 mL = 15.00 mL
* Moles of additional Na
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1) Acid-base indicators change color at a specific pH range near the equivalence point of an acid-base titration. This allows the endpoint to be visually identified.
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pH, pH measurement, and buffer solutions can be summarized as follows:
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2) pH can be measured using pH strips, pH indicators, or a pH meter. Common pH indicators change color depending on the pH range and can indicate if a solution is acidic or basic.
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The document discusses the pH scale, which measures how acidic or basic a substance is. It explains that the pH scale ranges from 0 to 14, with 7 being neutral. Values below 7 are acidic and values above 7 are basic. Each whole number on the pH scale represents a substance that is ten times more acidic or alkaline than the next higher value. The document also provides background on the history of the pH scale and defines acids and bases, noting that acids donate protons while bases receive protons in dissociation reactions.
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This document discusses acid-base theories and reactions. It begins with a brief history of acids and bases, noting they were originally defined based on their sour taste. It then focuses on three main acid-base theories:
1) Arrhenius theory defines acids as producing H+ ions in water and bases as producing OH- ions.
2) Brønsted-Lowry theory provides a more general definition, with acids as proton donors and bases as proton acceptors that form conjugate acid-base pairs.
3) Lewis theory defines acids and bases based on their ability to accept or donate electron pairs in chemical bonds.
The document goes on to discuss the properties of acids and bases, ion
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An acid-base titration is used to determine the concentration of an acid or base. The equivalence point is reached when there are equal moles of H3O+ and OH- ions, resulting in a neutral pH of 7. The titration endpoint is indicated by a color change in the indicator, which should change sharply at the equivalence point. Different types of acid-base reactions result in different pH values at the equivalence point, determined by whether the acid and base are strong or weak.
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1. The document describes acid-base titrations involving strong acids and bases as well as weak acids and bases. Titrations result in titration curves that can be divided into regions based on the relative amounts of acid and base present.
2. Key regions include before equivalence, at equivalence, and after equivalence. The pH at equivalence depends on whether the acid or base is strong or weak. Indicators are used to detect the equivalence point based on their color change near the desired pH.
3. Calculations are provided to determine pH in each region based on acid or base strength, volumes reacted, and equilibrium constants. Assumptions made in the calculations must be checked for validity.
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The document discusses pH titration curves for different acid-base reactions. It explains that the equivalence point occurs when reactants are mixed in exact proportions according to the balanced chemical equation. The end point is seen by a color change in the indicator. Titration curves show a steep pH change near the equivalence point. Curves are provided for strong acid-strong base, strong acid-weak base, weak acid-strong base, and weak acid-weak base reactions. More complex curves are discussed for reactions producing multiple products.
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Neutralization titrations
* HCl is a strong acid and will titrate first
* Its equivalence point was at 35.00 mL of NaOH
* NaOH concentration is 0.100 M
* Moles of NaOH used = Volume x Concentration
= 0.03500 L x 0.100 mol/L = 0.003500 mol
* Moles of HCl = Moles of NaOH used = 0.003500 mol
* H3PO4 is a weak acid and will titrate second
* Its equivalence point was at 50.00 mL of NaOH
* Additional NaOH used = 50.00 mL - 35.00 mL = 15.00 mL
* Moles of additional Na
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This document discusses acid-base titration, including definitions of acids and bases, strong vs weak acids and bases, and the technique of titration. It explains that a titration reaches the equivalence point when the moles of acid equals the moles of base, producing a neutral solution. This can be shown using the equation MAVA = MBVB, where M is molarity, V is volume, and A and B indicate acid and base. An example titration calculation is provided.
Theory of Acid-base Indicators and Acid-base Titration Curves
1) Acid-base indicators change color at a specific pH range near the equivalence point of an acid-base titration. This allows the endpoint to be visually identified.
2) The pH curve for a strong acid-strong base titration shows a sharp change in pH at the equivalence point of 7. A weak acid-strong base titration has a more gradual pH change before and after the equivalence point, which is above 7 due to salt hydrolysis.
3) The suitable indicator depends on the pH changes around the endpoint. It must change color in the steep "vertical" portion of the curve to accurately identify the endpoint.
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This document discusses acid-base titration and theories of acids and bases. It introduces acid-base titration as involving the neutralization reaction between an acid and base. Three theories of acids and bases are described: Arrhenius acid-base theory defines acids as producing H+ ions and bases as producing OH- ions. Bronsted-Lowry proton theory defines acids as proton donors and bases as proton acceptors. Lewis electron theory defines acids as electron pair acceptors and bases as electron pair donors. The document also discusses acid-base indicators and their use in determining the endpoint of a titration, as well as Ostwald and resonance theories of indicators.
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This document discusses acids and bases, including the ionization of water and the pH scale. It defines pH as the negative logarithm of the hydronium ion concentration and explains how pH and pOH are related. Examples are provided to demonstrate how to calculate the hydroxide ion concentration from the hydronium ion concentration or vice versa using the pH scale.
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This document discusses acid-base theories and equilibria in water. It describes the Arrhenius, Bronsted-Lowry, and Lewis theories of acids and bases. The autoionization of water and definitions of pH and pOH are explained. Calculations of pH for strong/weak acids and bases, salts, and buffer solutions using the ionization constant and Henderson-Hasselbalch equation are presented with examples. The role of buffers in resisting pH change upon addition of acids or bases is also summarized.
The PH Scale
The document discusses pH and the pH scale. It defines pH as the negative logarithm of the molar concentration of hydrogen ions and explains that pH values below 7 indicate increasing acidity while values above 7 indicate increasing basicity. It provides examples of calculating pH from given hydrogen ion concentrations and vice versa. It also discusses buffer solutions and how they resist changes in pH when small amounts of acid or base are added.
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pH, pH measurement, and buffer solutions can be summarized as follows:
1) pH is a measure of how acidic or basic a solution is, measured on a scale from 0-14. It is determined by the concentration of hydrogen ions [H+] in the solution.
2) pH can be measured using pH strips, pH indicators, or a pH meter. Common pH indicators change color depending on the pH range and can indicate if a solution is acidic or basic.
3) Buffer solutions help maintain pH when small amounts of acid or base are added. They use a weak acid or base combined with its salt. The Henderson-Hasselbalch equation can be used to calculate the pH of buffer solutions
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Lecture 1 - PH and Litmus paper
In chemistry, pH (potential of hydrogen) is a numeric scale used to specify the acidity or basicity of an aqueous solution. It is approximately the negative of the base 10 logarithm of the molar concentration, measured in units of moles per liter, of hydrogen ions.
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This document discusses salt hydrolysis, which is the reaction of salt ions with water that can result in an acidic, basic, or neutral solution. Salts formed from a weak acid and strong base produce a basic solution, as the conjugate base of the weak acid will accept protons from water. Salts formed from a weak base and strong acid produce an acidic solution, as the conjugate acid of the weak base will donate protons to water. Salts formed from a strong acid and strong base produce a neutral solution. Examples of sodium ethanoate and ammonium chloride are provided to illustrate basic and acidic salt hydrolysis.
Ph scale
Water undergoes self-ionization in which a small percentage of water molecules dissociate into hydronium (H3O+) and hydroxide (OH-) ions. The concentration of these ions is extremely small and the equilibrium lies very much in the forward direction. The self-ionization of water can be represented by the equilibrium constant Kw, which is equal to the product of the hydronium and hydroxide ion concentrations. Kw is temperature dependent and decreases with increasing temperature. The pH scale was developed to quantify the concentration of hydronium ions in solution and thus indicate whether a solution is acidic, basic, or neutral. pH is defined as the negative logarithm of the hydronium ion concentration
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The document discusses the pH scale, which measures how acidic or basic a substance is. It explains that the pH scale ranges from 0 to 14, with 7 being neutral. Values below 7 are acidic and values above 7 are basic. Each whole number on the pH scale represents a substance that is ten times more acidic or alkaline than the next higher value. The document also provides background on the history of the pH scale and defines acids and bases, noting that acids donate protons while bases receive protons in dissociation reactions.
Ainun dan winda
This document discusses acid-base theories and reactions. It begins with a brief history of acids and bases, noting they were originally defined based on their sour taste. It then focuses on three main acid-base theories:
1) Arrhenius theory defines acids as producing H+ ions in water and bases as producing OH- ions.
2) Brønsted-Lowry theory provides a more general definition, with acids as proton donors and bases as proton acceptors that form conjugate acid-base pairs.
3) Lewis theory defines acids and bases based on their ability to accept or donate electron pairs in chemical bonds.
The document goes on to discuss the properties of acids and bases, ion
Acid base titration
An acid-base titration is used to determine the concentration of an acid or base. The equivalence point is reached when there are equal moles of H3O+ and OH- ions, resulting in a neutral pH of 7. The titration endpoint is indicated by a color change in the indicator, which should change sharply at the equivalence point. Different types of acid-base reactions result in different pH values at the equivalence point, determined by whether the acid and base are strong or weak.
Acids and Bases
Acids are substances that donate protons (H+) in aqueous solutions, increasing the concentration of H+ ions. Common acids include HCl, HNO3, and HC2H3O2. Bases are substances that accept H+ ions, producing hydroxide (OH-) ions when dissolved in water. Common bases include NaOH, KOH, and Ca(OH)2. Strong acids and bases are completely ionized in solution, while weak acids and bases are only partially ionized. The strength of an acid or base determines how reactive it is.
Chapter 8 a acid-base titration
1. The document describes acid-base titrations involving strong acids and bases as well as weak acids and bases. Titrations result in titration curves that can be divided into regions based on the relative amounts of acid and base present.
2. Key regions include before equivalence, at equivalence, and after equivalence. The pH at equivalence depends on whether the acid or base is strong or weak. Indicators are used to detect the equivalence point based on their color change near the desired pH.
3. Calculations are provided to determine pH in each region based on acid or base strength, volumes reacted, and equilibrium constants. Assumptions made in the calculations must be checked for validity.
P h (titration) curves
The document discusses pH titration curves for different acid-base reactions. It explains that the equivalence point occurs when reactants are mixed in exact proportions according to the balanced chemical equation. The end point is seen by a color change in the indicator. Titration curves show a steep pH change near the equivalence point. Curves are provided for strong acid-strong base, strong acid-weak base, weak acid-strong base, and weak acid-weak base reactions. More complex curves are discussed for reactions producing multiple products.
Ph and POH
This document provides an overview of acid/base chemistry and pH. It defines pH as a measure of hydrogen ion concentration, describes the pH scale from 0-14, and explains how to calculate pH, pOH, and hydrogen or hydroxide ion concentration from other values. Sample problems demonstrate how to determine pH from concentration and vice versa, as well as the relationship between pH and pOH. Key points are that pH is a log scale measurement of acidity, and that the sum of pH and pOH equals 14 for any aqueous solution.
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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 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.
prepare and standardize N10 NaOH solution.pptx
To prepare a 0.1 N NaOH solution, 4 g of NaOH is dissolved in distilled water and made up to 1 L. This solution is then standardized against a 0.1 N oxalic acid solution prepared by dissolving 6.3 g of (COOH)2·2H2O in 1 L of distilled water. The equivalence point of the titration reaction between NaOH and oxalic acid is determined using a pH indicator to find the volume of acid needed to neutralize the base solution.
Acid Base.pptx
1. The document discusses Brønsted-Lowry acids and bases, including strong and weak acids and bases. Strong acids and bases completely dissociate in water, while weak ones only partially dissociate.
2. Examples of calculations involving titration reactions are shown, including determining the pH and volume of base needed to neutralize an acid. The autoionization of water and acid/conjugate base pairs are also covered.
3. Buffers are discussed as mixtures of a weak acid and its conjugate base that resist changes to pH when acids or bases are added. The pH of a buffer solution is close to the pKa of the acid/base pair.
Introduction to Volumetric Analysis for Basic Chemistry
Introduction to Volumetric Analysis for Basic Chemistry
Acidbasetitr
This document discusses acid-base titrations and summarizes the key steps in standardizing a sodium hydroxide (NaOH) solution and using it to determine the concentration of acetic acid in vinegar. It explains that NaOH is hygroscopic and absorbs moisture, so it must be standardized against a primary standard, potassium hydrogen phthalate (KHP), through titration. The standardized NaOH solution can then be used to titrate vinegar against it and calculate the moles and molarity of acetic acid present. It also discusses concepts like the pH at the equivalence point of a titration involving a weak acid and strong base.
Acidbasetitr 150602233035-lva1-app6891
This document discusses acid-base titrations and summarizes the key steps in standardizing a sodium hydroxide (NaOH) solution and using it to determine the concentration of acetic acid in vinegar. It explains that NaOH is hygroscopic and absorbs moisture, so it must be standardized against a primary standard, potassium hydrogen phthalate (KHP), through titration. The standardized NaOH solution can then be used to titrate vinegar against it and calculate the moles and molarity of acetic acid present. It also discusses concepts like the pH at the equivalence point of a titration involving a weak acid and strong base.
How do we find unknown concentrations of acids
The document discusses titration as a method to determine unknown concentrations of acids and bases. Titration involves slowly adding a base to an acidic solution until the pH reaches 7.00, indicating neutralization. The concentration of the base and volumes of the acid and base added can be used to calculate the concentration of the original acid using the titration equation. Indicator solutions that change color between acid and base are used to pinpoint the endpoint of neutralization.
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Introduction to Volumetric Analysis for Basic Chemistry
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Here are the steps to solve this problem using Hess's Law:
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This document provides information about buffers and how to prepare a buffer solution in the laboratory. It defines a buffer as a solution that can maintain a constant pH level when small amounts of acid or base are added. Buffers are composed of an acid and its conjugate base. The key properties that should match the desired pH are the acid's pKa value and that its conjugate base only differs by one hydrogen ion. The document outlines how to calculate the mass of salt needed using the Henderson-Hasselbalch equation and dilute the acid solution, and describes various methods for measuring the pH of buffer solutions. It also defines buffer capacity and how it can be experimentally determined.
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The document provides instructions for students to complete calculations to design a buffer of their assigned pH for a chemistry lab experiment. Students are told to use provided resources to determine the appropriate acid-salt pair to use based on the acid's pKa value relative to the assigned pH. They are then directed to use dilution equations to calculate amounts of acid and the Henderson-Hasselbalch equation to find the conjugate base concentration needed to achieve the target pH with their buffer.
16 solubility
The document provides instructions for an experiment to determine the solubility of calcium hydroxide. The experiment involves preparing a saturated calcium hydroxide solution, standardizing hydrochloric acid, and then titrating the calcium hydroxide solution with the standardized acid. Students are advised to review solubility and the solubility product constant (Ksp) in their textbook before setting up calculations involving Ksp. The titration will have two equivalence points but sometimes only one will be visible, and students are told to continue titrating until their data plateaus.
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15 buffers
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14 standardization&titration
Standardization and titration is a quantitative technique that can produce accurate and precise results when done carefully. It involves standardizing a sodium hydroxide (NaOH) solution by performing multiple titrations of a primary standard potassium hydrogen phthalate (KHP) of known molarity. This allows the determination of an unknown concentration of acetic acid by titrating samples with the standardized NaOH solution. Proper experimental procedure and tips are provided to complete the standardization and acetic acid titrations.
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06 microscale titration_fa06
Editor's Notes
- Analyte: Solution with the unknown concentration Standard: Solution with the unknown concentration You must have a BALANCED chemical equation before looking at the stoichiometry
- pH is a measurement of the H+ concentration, pH = - log (H+) pOH is a measurement of the OH- concentration, pOH = - log (OH-) pH + pOH = 14 If you start with a solution of HCl, is the pH going to be high or low? What happens to that pH as base is added? Before doing any calculations you must start with a balanced chemical equation. To predict the products of an acid-base reaction, ions change partners. Result is salt and water. Balance equation before doing calculations. Just about every exam for CHE 115 has had at least one titration calculation without a 1:1 ratio.
- The equivalence point is when the moles of H+ and OH- are equal, NOT when the moles of acid and base are equal. For an acid such as HCl and a base like NaOH, when the moles of H+ and OH- are equal, the moles of HCl and NaOH are equal. That is not the case for polyprotic acids such as H2SO4 or H3PO4 or bases such as Ba(OH)2 Endpoint is where the indicator changes color. This is *usually*, but not always after the equivalence point. It is always best to go past the end point to make sure you have completed the titration. If the indicator was chosen well, the endpoint and equivalence point will be very close to the same volume.
- The equivalence point can be determined in several ways. Measuring the pH is the most accurate method, but indicators can also be used. Indicators weak acids or bases that change color as the amount of H+ and OH- change. Typically change over 1-2 pH units and different indicators are used for to determine different pH changes. See Appendix in lab manual. The important thing is to understand why we need the equivalence point – to calculate the concentration of the unknown solution.
- Explain set-up.
- They can view the graph on their calculator or on the computer, but they should have three good titration curves before they leave the lab. What makes a good titration curve? A steep portion near the equivalence point with several data points in the region. If they have multiple outliers (i.e. pH goes up and then back down on the next drop), they should do another trial.
- Before leaving the lab, all pH sensors should be stored upright so that the tip is in the storage bottle, which is on the bottom, and the tip is immersed in the solution.
- - Students who have had calculus may be familiar with the term second derivative, others may not. It is okay if they don’t understand what it means, just follow the instructions.
- Make a big point that acid-base titration problems are just like other stoichiometry problems. Start with a balanced chemical equation, convert to moles, use mole rations, and convert to the desired units. The main difference is that instead of talking about a product and a reactant, you are looking at two reactants. If you start with a balanced chemical equation, it doesn’t matter what of the equation the substance is on.