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titration experiment of hcl and naoh

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  • Basic terms
  • Titration end point
  • End point indicators
  • End point detection
  • Equivalence point calculation
  • Titration curve calculation
  • Titration calculation
  • Back titration
  • Sample & titrant volume
  • Volumetric glassware
  • Volumetric glass cleaning
  • Glassware calibration
  • Standard substances
  • Sources of errors
  • End point (pH) indicators
  • Indicators preparation
  • Polyprotics and mixtures
  • Solutions used
  • Solutions standardization
  • HCl determination
  • H 2 SO 4 determination
  • Acetic acid in vinegar
  • H 3 PO 4 determination
  • NaOH determination
  • NaOH & Na 2 CO 3 assay
  • Potentiometric titrations
  • Permanganate titration
  • Precipitation
  • Argentometry
  • Solution standardization
  • Chlorides - Mohr method
  • Chlorides - Volhard method
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  • EDTA standardization
  • magnesium titration
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  • nickel titration
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  • Further reading
  • » Acid-base titration
  • » Hydrochloric acid titration

Titration of hydrochloric acid with sodium hydroxide

General remarks.

Determination of hydrochloric acid concentration is probably the most often discussed example of acid-base titration. Both acid and base are strong, which not only makes determination of end point easy (steep part of the curve is long), but also means that calculation of titration curve and equivalence point are pretty straightforward.

This is a simple neutralization reaction:

HCl + NaOH → NaCl + H 2 O

It is worth of noting, that - as we can assume both acid and base to be completely dissociated - net ionic reaction is just

H + + OH - → H 2 O

which is the simplest form of neutralization reaction possible.

In the reality every acid and every base - no matter how strong - have some dissociation equilibria described by dissociation constant. In this particular case K a for HCl is listed as 10 4 (which means it can be safely neglected) and dissociation constant K b for NaOH is listed as 0.6 - which means sometimes it has to be taken into account.

sample size

Depending on the titrant concentration (0.2 M or 0.1 M), and assuming 50 mL burette, aliquot taken for titration should contain about 0.26-0.33 g (0.13-0.16 g) of hydrochloric acid (7-9 or 3.5-4.5 millimoles).

end point detection

Equivalence point of strong acid titration is usually listed as exactly 7.00. That's not necesarilly the case, as it depends on the solution temperature and ionic strength of the solution, besides, slight hydrolysis of NaOH shifts pH down by about 0.02 unit. Not that it changes much - we are still very close to 7. Thus the best indicator of those listed on pH indicators preparation page is bromothymol blue . However, as we have discussed on the acid-base titration end point detection page, unless we are dealing with a diluted solution (in the range of 0.001 M) we can use almost any indicator that gives observable color change in the pH 4-10 range. Thus we can safely use the most popular phenolphthalein and titrate to the first visible color change.

Color change of phenolphthalein during titration - on the left, colorless solution before end point, on the right - pink solution after end point. Note we have to end titration at first sight of color change, before color gets saturated.

solutions used

To perform titration we will need titrant - 0.2 M or 0.1 M sodium hydroxide solution , indicator - phenolphthalein solution and some amount of distilled water to dilute hydrochloric acid sample.

  • Pipette aliquot of hydrochloric acid solution into 250mL Erlenmeyer flask.
  • Dilute with distilled water to about 100 mL.
  • Add 2-3 drops of phenolphthalein solution.
  • Titrate with NaOH solution till the first color change.

result calculation

According to the reaction equation

Hydrochloric acid reacts with sodium hydroxide on the 1:1 basis. That makes calculation especially easy - when we calculate number of moles of NaOH used it will be already number of moles of HCl titrated.

To calculate hydrochloric acid solution concentration use EBAS - stoichiometry calculator . Download determination of hydrochloric acid concentration reaction file, open it with the free trial version of the stoichiometry calculator .

Click n=CV button above NaOH in the input frame, enter volume and concentration of the titrant used. Click Use button. Read number of moles and mass of hydrochloric acid in the titrated sample in the output frame. Click n=CV button in the output frame below hydrochloric acid, enter volume of the pipetted sample, read hydrochloric acid concentration.

sources of errors

Apart from general sources of titration errors , when titrating hydrochloric acid we should pay special attention to titrant. Sodium hydroxide solutions are not stable as they tend to absorb atmospheric carbon dioxide. Hydrochloric acid is much stronger than carbonic acid, so it will slowly expel carbon dioxide from the solution, but initially presence of carbonates will mean that to reach end point we need to add axcess of titrant.

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Technical info: $section - alkalimetry, $subsection - , $page - HCl, $url_call - acid-base-titration-hydrochloric-acid

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3 Determination of the Concentration of a Solution of Sodium Hydroxide

To determine the concentration of an unknown hydrochloric acid solution using acid-base titration methods.

Expected Learning Outcomes

  • Describe and perform an acid-base titration experiment using indicators to find the end point.

Solve stoichiometry problems in the solution phase.

Textbook Reference

Tro,  Chemistry: Structures and Properties , 2nd Ed, Ch. 8.7.

Introduction

Acid-base chemistry, definitions.

There are a number of different definitions of acids and bases present. However, for most purposes, the Brønsted-Lowry definition of acids and bases are the most useful:

  • An acid is a proton (H + ) donor. Common examples include hydrochloric (muriatic) acid, citric acid, and ascorbic acid (Vitamin C).
  • A base is a proton acceptor. Among the best known of these substances are hydroxides, such as sodium hydroxide; however, other bases exist. Examples of these include sodium bicarbonate (baking soda) and sodium hypochlorite (bleach).

\textrm{CH}_3\textrm{COOH}

Acid-Base Neutralization Reactions

One of the most commonly observed and studied types of acid-base reaction is the  acid-base neutralization reaction , where an acid is reacted with a base:

\begin{equation*} \textrm{acid}+\textrm{base} \to \textrm{water} + \textrm{salt} \end{equation*}

where the identity of the salt can be determined by taking the cation from the base and taking the anion from the acid.

If I react sodium hydroxide with hydrochloric acid, the cation is Na + from the sodium hydroxide and the anion is Cl – from the hydrochloric acid:

\begin{equation*} \ce{HCl}(aq) + \ce{NaOH} (aq) \to \ce{H2O}(l) + \ce{NaCl}(aq) \end{equation*}

The Titration Setup

A titration experiment is one where one attempts to determine the concentration of a sample solution by reacting two samples in such a way that the exact volume of the two reactants are known. If the concentration of one of the two reactant solutions is known exactly, then one can determine the concentration of the second solution accurately.

A buret (a long tube with a tap at the bottom) with an Erlenmeyer flask underneath.

A typical titration setup is illustrated in the figure below. A measured volume of the analyte – typically the acid – is placed in the Erlenmeyer flask at the bottom along with an indicator.  The indicator is a compound that will change color when the reaction has completed to the end point. A volumetric pipet is typically used to measure out the exact amount of analyte needed.

A buret is used to deliver the titrant (typically the base in acid-base titrations). The titrant is added in dropwise such that the volume of titrant required to change the color of the indicator can be recorded to the nearest 0.01 mL. Therefore, we can determine the exact (or as close to exact as we can) volume of titrant required to react with the analyte.

\textrm{HC}_8\textrm{H}_4\textrm{KO}_4

From the information obtained the concentration of the provided sodium hydroxide solution can be determined. Sodium hydroxide cannot be used as a primary standard readily, as it tends to be extremely hydroscopic and therefore the mass of sodium hydroxide used is not an accurate starting point for determining the concentration of base used. For this reason, we use KHP as a primary standard, with which the concentration of the sodium hydroxide solution can be determined accurately (and therefore act as a secondary standard).

The end-point – which is when the analyte is completely reacted with the titrant – is determined using phenolphthalein as an indicator. When the end-point of the titration is reached, the color of the solution changes from colorless to pink.

You will use the Determine the concentration of an unknown HCl solution and Standardization of NaOH experiment (OLI Unit 3 – Module 11)

  • You must prepare your notebook just like a regular face-to-face experiment and record all your data there in handwriting .  At the end of the experiment, you are required to scan and submit your lab notebook entries.
  • While there are directions on the OLI webpage, we have written new ones that are designed to allow you to repeat the experiment.
  • Volumes from volumetric flasks and pipettes should have four significant figures .
  • For each titration, you may want to try a rough titration first where you go through it quickly so you have a maximum volume it might be at, and then try it more carefully by starting off with adding NaOH to within a few mL before doing the titration in full.  The rough attempt should be recorded in the notebook but can be ignored for the data analysis.

Preparation of a Solution of Potassium Hydrogen Phthalate

  • Remove the KHP, a 250 mL beaker, and scale from the stockroom.
  • Place a weigh boat on the scale and tare the balance. Be sure to record the entire reading on the scale.
  • Transfer 7-9 g of KHP to the weigh boat. Note that the transfer mode is “Realistic”. Clicking once, quickly on the “Hold to Pour” button will transfer this amount. Record the mass of the KHP.  Afterwards, transfer all of this into a 250 mL volumetric flask.  (Note: the volume of a volumetric flask has four significant figures).
  • Make up the solution to 250 mL in the volumetric flask with distilled water from the stockroom.

Standardization of the Sodium Hydroxide Solution

In this part, you will determine the exact concentration of the ~0.1 M sodium hydroxide solution using the potassium hydrogen phthalate solution you prepared earlier as a standard.

  • Fill the buret approximately to the 5.00 mL mark with the NaOH solution.
  • Using a volumetric pipet, measure out 25.00 mL of the potassium hydrogen phthlate solution and put this in a 250 mL Erlenmeyer flask.
  • Put a few drops of phenolphthalein from the stockroom into the Erlenmeyer flask.
  • Overlap the buret on top of the Erlenmeyer flask so it is placed directly above the Erlenmeyer flask and is poised to deliver the solution.
  • Add portions of the sodium hydroxide from the buret into the Erlenmeyer flask.  As you near the desired concentration, be sure to add in the sodium hydroxide dropwise until the phenolphthalein indicator turns pink.
  • Record the final volume in the buret.
  • Empty the contents of the Erlenmeyer flask.

Repeat this part at least two more times.  Do a fourth trial if the volume delivered from the first three trials are more than 0.2 mL apart.

Determination of the Concentration of the Unknown Hydrochloric Acid Solution

Repeat how you did the titration of KHP against sodium hydroxide, except using 25 mL of the unknown hydrochloric acid solution instead of 25 mL of the KHP solution.

Data Analysis

This laboratory experiment is rather complex to calculate and takes some time to do. Be sure to complete all calculations carefully. If you do not understand how to do these calculations, please be sure to consult your instructor or a mentor in the Math/Science Resource Center.

Determining the Concentration of Potassium Hydrogen Phthalate (KHP) Solution

Potassium hydrogen phthalate has a molar mass of 204.2 g/mol. Given the mass of the potassium hydrogen phthalate present and the volume of the volumetric flask (which you will need to convert into liters), you should be able to determine the molarity of the solution using the definition

\begin{equation*} \textrm{molarity (M)} = \frac{\textrm{moles of solute}}{\textrm{liters of solution}} \end{equation*}

If we dissolve 3.224 g potassium hydrogen phthalate (KHP) to form a solution of 100.0 mL, the molarity of KHP in this solution is

\begin{align*} ? \mbox{ mol KHP} &= 3.224\mbox{ g KHP} \times \frac{\mbox{mol KHP}}{204.2\mbox{ g KHP}} \\ &= 0.01579\mbox{ mol} \\ ?\mbox{ L soln} &= 100.0 \textrm{ mL} \times \frac{\mbox{L}}{1000 \textrm{ mL}} \\ &= 0.1000\textrm{ L} \\ ?\frac{\mbox{mol}}{\mbox{L}} &= \frac{0.01579\textrm{ mol}}{0.1000\textrm{ L}} \\ &= 0.1579\textrm{ M} \end{align*}

Determining the concentration of the 0.1 M sodium hydroxide solution

This can be solved as a solution stoichiometry problem. To do this, the first step is to determine the number of moles of sodium hydroxide present, and then divide by the volume of NaOH used.

10.00 mL of a 0.1523 M KHP solution requires 14.22 mL of sodium hydroxide to neutralize.  Determine the concentration of the NaOH solution.

\begin{align*} ?\mbox{mol NaOH} &= 10.00 \mbox{ mL KHP} \times \frac{\mbox{L KHP}}{1000\mbox{ mL KHP}} \times \frac{0.1523 \mbox{ mol KHP}}{\mbox{L KHP}} \times \frac{1\mbox{ mol NaOH}}{1\mbox{mol KHP}} \\ &= 0.001523 \mbox{ mol NaOH} \\ ?\mbox{ L NaOH} &= 14.22\mbox{ mL} \times \frac{\mbox{L}}{1000\textrm{ mL}} \\ &=0.01422\textrm{ L}\\ \unit[?]{M} &= \frac{0.001523\textrm{ mol}}{0.01422\textrm{ L}} \\ &= 0.1071\textrm{ M} \end{align*}

Similar approaches can then be used to find the concentration of hydrochloric acid.

Determination of the Concentration of a Solution of Sodium Hydroxide Copyright © by Yu Kay Law is licensed under a Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License , except where otherwise noted.

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Shop Experiment Acid-Base Titration Experiments​

Acid-base titration.

Experiment #24 from Chemistry with Vernier

titration experiment of hcl and naoh

Introduction

A titration is a process used to determine the volume of a solution needed to react with a given amount of another substance. In this experiment, you will titrate hydrochloric acid solution, HCl, with a basic sodium hydroxide solution, NaOH. The concentration of the NaOH solution is given and you will determine the unknown concentration of the HCl. Hydrogen ions from the HCl react with hydroxide ions from the NaOH in a one-to-one ratio to produce water in the overall reaction:

{{\text{H}}^{\text{ + }}}{\text{(aq) + C}}{{\text{l}}^{\text{ - }}}{\text{(aq) + N}}{{\text{a}}^{\text{ + }}}{\text{(aq) + O}}{{\text{H}}^{\text{ - }}}{\text{(aq)}} \to {{\text{H}}_{\text{2}}}{\text{O(1) + N}}{{\text{a}}^{\text{ + }}}{\text{(aq) + C}}{{\text{l}}^{\text{ - }}}{\text{(aq)}}

When an HCl solution is titrated with an NaOH solution, the pH of the acidic solution is initially low. As base is added, the change in pH is quite gradual until close to the equivalence point, when equimolar amounts of acid and base have been mixed. Near the equivalence point, the pH increases very rapidly. The change in pH then becomes more gradual again, before leveling off with the addition of excess base.

In this experiment, you will use a computer to monitor pH as you titrate. The region of most rapid pH change will then be used to determine the equivalence point. The volume of NaOH titrant used at the equivalence point will be used to determine the molarity of the HCl.

In this experiment, you will

  • Use a pH Sensor to monitor changes in pH as sodium hydroxide solution is added to a hydrochloric acid solution.
  • Plot a graph of pH vs . volume of sodium hydroxide solution added.
  • Use the graph to determine the equivalence point of the titration.
  • Use the results to calculate the concentration of the hydrochloric acid solution.

Sensors and Equipment

This experiment features the following sensors and equipment. Additional equipment may be required.

titration experiment of hcl and naoh

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This experiment is #24 of Chemistry with Vernier . The experiment in the book includes student instructions as well as instructor information for set up, helpful hints, and sample graphs and data.

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A thermometric titration

In association with Nuffield Foundation

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Use this class practical to practise locating the end-point of a titration by measuring the temperature change

In this experiment, students titrate sodium hydroxide solution with hydrochloric acid. By measuring the temperature change each time a portion of acid is added, students can determine the end-point of the titration, indicated by the highest temperature. They then use this information to calculate the concentration of the hydrochloric acid.

The practical takes about one hour, and is best carried out individually or in pairs.

  • Eye protection (goggles)
  • Thermometer, 0–100 °C (see note 5 below)
  • Two insulated (polystyrene) cups
  • Beaker, 250 cm 3
  • Burette, 50 cm 3
  • Burette stand
  • Clamp and stand (optional)
  • Cork, one-holed, to fit thermometer (optional)
  • Pipette, 20 cm 3 or 25 cm 3
  • Pipette safety filler
  • Hydrochloric acid, 2.00 M (IRRITANT), about 75 cm 3
  • Sodium hydroxide solution, 1.50 M (CORROSIVE), about 30 cm 3

Health, safety and technical notes

  • Read our standard health and safety guidance.
  • Wear eye protection throughout.
  • Hydrochloric acid, HCl(aq), (IRRITANT at concentration used) – see CLEAPSS Hazcard  HC047a  and CLEAPSS Recipe Book RB043. This concentration is necessary to achieve a reasonable change in temperature. The concentration of the hydrochloric acid should not be indicated on bottle available to the students.
  • Sodium hydroxide solution, NaOH(aq), (CORROSIVE at concentration used) – see CLEAPSS Hazcard  HC091a  and CLEAPSS Recipe Book RB085. This concentration is necessary to achieve a reasonable change in temperature. The concentration of the sodium hydroxide should be indicated on bottle available to the students.
  • Instead of using the thermometer to stir the titration mixture, it could be clamped in position in a cork, as shown in the diagram, and the mixture swirled after each addition of acid. Alternatively, a temperature sensor attached to a computer can be used in place of a thermometer. Data logging software could then be used to provide a detailed plot of the readings.
  • Stand an insulated cup in a beaker for support.

A diagram showing the equipment required for a thermometric titration using hydrochloric acid

Source: Royal Society of Chemistry

In this thermometric titration, students can determine the end-point of the titration using the highest temperature recorded during the experiment

  • Using a pipette and safety filler, transfer 20 cm 3 (or 25 cm 3 ) of the sodium hydroxide solution into the cup, and measure the steady temperature.
  • Using the burette, add a small portion (3–5 cm 3 ) of dilute hydrochloric acid to the solution in the cup, noting down the actual volume reading. Stir by swirling the cup and measure the highest temperature reached.
  • Immediately add a second small portion of the dilute hydrochloric acid, stir, and again measure the highest temperature and note down the volume reading.
  • Continue in this way until there are enough readings to decide the maximum temperature reached during this experiment. You will need to add at least 30 cm 3 of the acid.
  • Plot a graph of temperature against the volume of acid added, and use extrapolation of the two sections of the graph to deduce the maximum temperature reached without heat loss.
  • Use your results to calculate the concentration of the hydrochloric acid.

Teaching notes

The main concern in this experiment is the heat loss. If possible, a lid should be used. More reliable results can be achieved using two polystyrene cups (one inside the other).

With abler or older students, it is possible to discuss the extrapolation of the cooling curve to estimate the maximum temperature reached without heat loss. Creative Chemistry provide a resource on thermometric titration which includes an example of a typical plot of temperature vs volume of acid for this experiment, as well as the use of extrapolation to determine the maximum temperature change.

To reinforce the theory involved here, an indicator could also be used to show that the end-point really did occur at the highest temperature.

Additional information

This is a resource from the  Practical Chemistry project , developed by the Nuffield Foundation and the Royal Society of Chemistry. This collection of over 200 practical activities demonstrates a wide range of chemical concepts and processes. Each activity contains comprehensive information for teachers and technicians, including full technical notes and step-by-step procedures. Practical Chemistry activities accompany  Practical Physics  and  Practical Biology .

The experiment is also part of the Royal Society of Chemistry’s Continuing Professional Development course:  Chemistry for non-specialists .

© Nuffield Foundation and the Royal Society of Chemistry

  • 14-16 years
  • 16-18 years
  • Practical experiments
  • Acids and bases
  • Quantitative chemistry and stoichiometry
  • Reactions and synthesis

Specification

  • (e) simple procedures to determine enthalpy changes
  • determine the enthalpy changes for combustion and neutralisation using simple apparatus; and
  • 2.8.6 recall experimental methods to determine enthalpy changes;
  • 1.8.11 investigate the temperature change during neutralisation and demonstrate understanding that neutralisation reactions are exothermic (heat is given out);
  • 1.8.10 investigate the temperature change during neutralisation and demonstrate understanding that neutralisation reactions are exothermic (heat is given out);

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Acid-Base Titrations: Standardization of NaOH and Antacid Analysis

Acid- Base Titrations By Julia Trimble

Introduction:

A titration is a procedure in which two solutions are introduced to form a reaction that once completed, reaches an identifiable endpoint (Murphy, 2012, p.305).  During a titration, the volume of one reagent, the analyte, is predetermined while the other reagent, the titrant, is prepared in a buret and slowly introduced to the analyte solution.  The completion of this reaction can be determined via observation through the use of an indicator.  The amount of titrant released from the buret at the endpoint signifies the volume of solution necessary to complete the reaction.  From this given volume, the concentration of either titrant or analyte can be determined when equilibrium is reached between reactant and product (Murphy, 2012, p.305).

In this experiment, the reagents combined are an acid, HCl (aq) and a base, NaOH (aq) where the acid is the analyte and the base is the titrant.  The reaction between the two is as follows:

HCl (aq) + NaOH (aq) → H2O (l) + Cl – (aq) + Na + (aq)

In this case, Sodium and Chloride act as spectator ions and form into salts in a neutralization reaction.  The titration of this reaction that occurs allows one to “standardize” the concentration or value of either reagent used.  In part A of this experiment, the concentration of NaOH (aq) was determined.  The calculated value of the base can then be used in the “back-titration” of another reagent as is done in part B where an antacid, composed mainly of CaCO 3(s) , is used to neutralize the acidic HCl (aq) solution.  The reaction is as follows:

CaCO 3(s) + 2H 3 O + (aq) → Ca +2 (aq) + 2H 2 O (l) + H 2 CO 3(aq) → H 2 O (l) + CO 2(g)

For this particular experiment, back- titration occurs when an antacid is dissolved in excess HCl (aq) and the remaining acid is then neutralized by the standardized NaOH (aq) (Swartz, 2013).  From this, the change in molarity between the initial concentration of HCl (aq) and the moles neutralized will determine the mass of the active ingredient in the antacid that can later prove the effectiveness of antacid brands as buffers.

Materials / Methods:

Begin by preparing a 0.10 M of 250 mL solution of NaOH (aq) using the M I V I =M 2 V 2 formula and fill a 50.0 mL buret to the top.  In 3 separate 250 mL Erlenmeyer flasks, pipet 25 mL of 0.10 M HCl (aq) and dilute with water until the total volume reaches 75 mL.  An indicator, phenolthalein, must be added to the analyte solution to make the endpoint visible.  Release the titrant into the analyte solution until the reaction has occurred completely.  This will become apparent when the solution remains a constant pink color.  Use the calculated molarity from the volume of titrant to standardize the solution produced and prepare second half of experiment.

Prepare the standardized NaOH (aq) solution in a 50.0 mL buret just as in Part A and fill 3 separate Erlenmeyer flasks with a new solution of 25 mL 0.50 M HCl (aq) , DI water, and a crushed antacid tablet with a total volume less than 75 mL.  Heat the analyte solution for 5 minutes to finalize the HCl (aq) / CaCO 3(s) reaction.  Once cooled, add the indicator Methyl Purple.  Slowly release the titrant and observe the solution transition from purple, the initial state, to grey, the midrange, and green, the endpoint.  Once completed, use the volume of titrant released to calculate the moles of excess HCl (aq) and determine the mass of CaCO 3(s) present in the antacid tablet.

Data Summary

Part A: Standardization of NaOH
Trial 1 Trial 2 Trial 3 units
Volume HCl pipetted into flask: 25 25 25 mL
Molarity HCl pipetted: 0.11 0.11 0.11 M
Volume reading, initial (NaOH): 0 0 0 mL
Volume reading, final (NaOH): 25.78 26.19 26.18 mL
Volume of NaOH let out of buret: 25.78 26.19 26.18 mL
Molarity NaOH: 0.107 0.105 0.105 M
Average M of NaOH: 0.106 M
Average Dev. Of M NaOH: 0.001 M
Part B: Antacid Tablet Analysis
Trial 1 Trial 2 Trial 3 units
Vol. HCl added to flask total: 25 25 25 mL
M HCl: 0.524 0.524 0.524 M
Volume reading, initial (NaOH): 0 0 0 mL
Volume reading, final (NaOH): 17.08 15.8 16.38 mL
Volume of NaOH let out of buret: 17.08 15.8 16.38 mL
M NaOH: 0.106 0.106 0.106 M
Total moles HCl added to tablet: 0.0131 0.0131 0.0131 mol
Moles NaOH added from buret: 0.00181 0.00167 0.00174 mol
Moles HCl neutralized by tablet: 0.0113 0.0114 0.0114 mol
Moles of active ingredient: 0.0226 0.0229 0.0227 mol
Mass of active ingredient: 565 572 569 mg
Avg mass of active ingredient: 568 mg
Average deviation in mass: 2 mg
Percent Error 14%

From Part A of this experiment, it was found that the average volume of titrant used to complete the reaction was approximately 26.05 mL meaning that it took about 26 mL of NaOH (aq) for the moles of each reagent to equal each other.  This makes sense considering that both are stoichiometrically equivalent given their common molar coefficient.  Therefore they both reagents should require similar volumes to reach equilibrium.  Given this volume, the molarity of NaOH (aq) was calculated to be an average of 0.106 M ± 0.001.  When this standardized titrant was used in Part B of the experiment, its average volume of 16.42 mL determined the amount of HCl (aq) left unreacted from the buffer reaction with CaCO 3(aq) to be an average of 0.00174 moles.  The amount of HCl (aq) neutralized by the reaction was then used to calculate the moles of CaCO 3(aq) in the tablet.  Then it was converted to a theoretical mass of 568mg in comparison to the actual mass of 500mg.  This value is within reason due to the antacid brand’s unusually high range of uncertainty of ± 50 mg making the error a mere 14%.  A few observable results of this reaction would include the significant color changes indicating the midrange and endpoint of the reaction.  The solution would produce a vibrant purple when the indicator was added and it would abruptly shift from purple to grey and soon after, from grey to green.  Another observation would be the chemical reaction between the antacid and the HCl (aq) .  The solution produced heat, gas, and condensation along the sides of the Erlenmeyer flask as a result of the CO 2(g) and H 2 O (g) escaping the system.

Discussion:

Throughout this experiment, several environmental influences played a big role on the measured data as well as precision.  A major source or error occurred in the measurement of CaCO 3(s) in the antacid tablet.  The active ingredient of a tums tablet was measured in 1 significant figure and measured under an increment of hundreds.  Therefore, the uncertainty of the company measurement was ± half of each hundred, which is far too great of a range to truly have an accurate read on how effective the experiment was.  Another issue arose with the crushed antacid tablet.  After crushing the Tums tablet thoroughly with the mortar and pestle, when the powder was transferred to the solution, it was difficult to completely dissociate in the solution.  Even after constant stirring and swirling, a few small regions of powder settled to the bottom and formed small clumps.  While heating, this could have prevented a 100% complete reaction between HCl (aq) and the antacid despite the excessive heat.  It may have left more excess HCl than would normally have been.  The final source of error could have been due to the composition of the antacid tablet itself.  In this experiment, only the Calcium Carbonate was considered in the chemical formula however there may have been other ingredients within the tablet that could influence the concentration of substances of the reaction.  Within an uncertain concentration of CaCO 3(s) , there is an amount of uncertainty in the measured data.

Formula Copy

Formula-Copy

Works Cited

Murphy, K. (2012). Acid-Base Titrations. A Labratory Manual for General Chemistry (5th ed., pp. 305-311).

Amherst, New York: Daemen College

Swartz, D. (Director) (2013, April 3). Experiment 19A: Standardize NaOH. Chemistry II

Lab. Lecture conducted from Daemen College, Amherst.

Swartz, D. (Director) (2013, April 10). Experiment 19B: Back Titration. Chemistry II

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a student performed a titration using 12.5. ml of naoh solution with 0.5 m hcl solution. if 50.0 ml of hcl was consumed during the titration, calculate the concentration of naoh.

if 50.0 ml of hcl was consumed during the titration concentration of naoh is [tex]$0.25 \mathrm{M}$[/tex]

Explanation:

[tex]0.500 \frac{\mathrm{mol}}{\mathrm{L}} \mathrm{H}_2 \mathrm{SO}_4 \cdot 0.0125 \mathrm{~L}=0.00625 \mathrm{molH}_2 \mathrm{SO}_4[/tex]

Write a balanced equation for the reaction:

[tex]$$\begin{aligned}& \mathrm{H}_2 \mathrm{SO}_4+2 \mathrm{NaOH} \rightarrow 2 \mathrm{H}_2 \mathrm{O}+\mathrm{Na}_2 \mathrm{SO}_4 \\& 1 \text { mol } \mathrm{H}_2 \mathrm{SO} 4 \text { neutralizes } 2 \mathrm{~mol} \mathrm{NaOH} \text {, so: } \\& 0.00625 \cdot 2=0.0125 \mathrm{~mol} \mathrm{NaOH} \\& 0.0125 \mathrm{~mol} \mathrm{NaO} \frac{\mathrm{H}}{0.05} L=0.25 \mathrm{MNaOH}\end{aligned}$$[/tex]

concentration of naoh is 0.25 M

In chemistry, concentration is the abundance of a constituent divided by the total volume of a mixture. Several types of mathematical description can be distinguished: mass concentration, molar concentration, number concentration, and volume concentration.

You can determine how much solute has been dissolved in the solvent by looking at the concentration of the solution. For instance, the concentration could be stated as 1 teaspoon of salt per 2 cups of water if you add 1 teaspoon to 2 cups of water.

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Enhanced adsorption of aqueous pb(ii) by acidic group-modified biochar derived from peanut shells.

titration experiment of hcl and naoh

1. Introduction

2. materials and methods, 2.1. materials, 2.2. ambc preparation, 2.3. acid–base titration, 2.4. characterization, 2.5. adsorption test, 3. results and discussion, 3.1. characterization of adsorbents, 3.2. pb(ii) adsorption by ambc, 3.2.1. effects of contact time, 3.2.2. effects of ph value, 3.2.3. effects of the dosage of ambc, 3.2.4. effects of initial pb(ii) concentration, 3.3. comparison of pre-ambc and ambc with other adsorbents, 3.4. pb(ii) adsorption isotherms, 3.5. thermodynamic study of the adsorption, 3.6. kinetics study of the adsorption, 3.7. possible adsorption mechanisms, 4. conclusions, author contributions, data availability statement, acknowledgments, conflicts of interest.

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Click here to enlarge figure

AdsorbentS
(m /g)		V
(cm /g)		D
(nm)		Acid Density (mmol/g)
TotalPh-OH-SO H-COOH
Pre-AMBC832.63.99 × 10 2.21.260.05-1.21
AMBC329.32.67 × 10 2.12.280.330.521.43
AdsorbentsQ
(mg/g)		R
(%)
Pre-AMBC83.139.7
AMBC148.676.0
AdsorbentsQ
(mg/g)		ActivatorExperiment Conditions
pHT
(°C)		Dosage
(g/L)		Equilibration Time
(min)		Initial
Concentration
(mg/g)		Ref.
AMBC148.6H SO 5250.5120100This study
C-KOH57.5KOH525235100[ ]
SBC191.1H SO 4.5-25200[ ]
HP-BC60.9H O 525-500100[ ]
M-RH-AC134.9HNO 5.5-190180[ ]
PLAC98.4H PO 5.0250.63080[ ]
Temperature
(°C)		Langmuir ModelFreundlich Model
K
(L/mg)		q
(mg/g)		R R K
(mg/g)		1/nR
250.0716210.10.046~0.1950.9894768.1170.2110.97800
350.0786220.30.048~0.2030.9860669.6110.2150.98978
450.0826318.50.053~0.2180.9670675.7080.2800.97686
Temperature
(°C)		K ΔG
(kJ/mol)		ΔH
(kJ/mol)		ΔS
(kJ/mol.K)
255.304−4.143.2480.025
355.334−4.29
454.764−4.13
Temperature
°C		q
mg/g		k
g/(min.mg)		R
25147.05.53 × 10 0.99989
35155.36.92 × 10 0.99991
55157.58.71 × 10 0.99992
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Share and Cite

Wu, Y.; Li, C.; Wang, Z.; Li, F.; Li, J.; Xue, W.; Zhao, X. Enhanced Adsorption of Aqueous Pb(II) by Acidic Group-Modified Biochar Derived from Peanut Shells. Water 2024 , 16 , 1871. https://doi.org/10.3390/w16131871

Wu Y, Li C, Wang Z, Li F, Li J, Xue W, Zhao X. Enhanced Adsorption of Aqueous Pb(II) by Acidic Group-Modified Biochar Derived from Peanut Shells. Water . 2024; 16(13):1871. https://doi.org/10.3390/w16131871

Wu, Yumeng, Ci Li, Zhimiao Wang, Fang Li, Jing Li, Wei Xue, and Xinqiang Zhao. 2024. "Enhanced Adsorption of Aqueous Pb(II) by Acidic Group-Modified Biochar Derived from Peanut Shells" Water 16, no. 13: 1871. https://doi.org/10.3390/w16131871

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COMMENTS

  1. Titrating sodium hydroxide with hydrochloric acid

    Stage 1. Using a small funnel, pour a few cubic centimetres of 0.4 M hydrochloric acid into the burette, with the tap open and a beaker under the open tap. Once the tip of the burette is full of solution, close the tap and add more solution up to the zero mark. (Do not reuse the acid in the beaker - this should be rinsed down the sink.)

  2. Titration of a Strong Acid With A Strong Base

    This leaves the final product to simply be water, this is displayed in the following example involving hydrochloric acid (HCl) and sodium hydroxide (NaOH). From Table \(\PageIndex{1}\), you can see that HCl is a strong acid and NaOH is a strong base. Therefore, the reaction between HCl and NaOH is initially written out as follows:

  3. Titration of hydrochloric acid with sodium hydroxide

    Titrate with NaOH solution till the first color change. result calculation. According to the reaction equation. HCl + NaOH → NaCl + H 2 O. Hydrochloric acid reacts with sodium hydroxide on the 1:1 basis. That makes calculation especially easy - when we calculate number of moles of NaOH used it will be already number of moles of HCl titrated.

  4. 7.18: Titration Experiment

    Titration Experiment. In the neutralization of hydrochloric acid by sodium hydroxide, the mole ratio of acid to base is 1:1. HCl(aq) + NaOH(aq) → NaCl(aq) +H2O(l) (7.18.1) (7.18.1) HCl ( a q) + NaOH ( a q) → NaCl ( a q) + H 2 O ( l) One mole of HCl HCl would be fully neutralized by one mole of NaOH NaOH. If instead the hydrochloric acid was ...

  5. PDF Chemistry 120: Experiment 1 Preparation of a Standard Sodium Hydroxide

    Add about 6 mL (use a 10-mL graduated cylinder) of the 50% sodium hydroxide solution to a 1-liter bottle. Fill the bottle to its shoulder with the distilled water. Place the stopper or the lid on the bottle, and thoroughly mix the solution by shaking the bottle. Warning: Concentrated NaOH is very corrosive!

  6. 21.17: Titration Experiment

    A titration is an experiment where a volume of a solution of known concentration is added to a volume of another solution in order to determine its concentration. Many titrations are acid-base neutralization reactions, though other types of titrations can also be performed. In order to perform an acid-base titration, the chemist must have a way ...

  7. PDF Acid/Base Chemistry: Titration Lab

    A titration is an analytical procedure used to determine the accurate concentration of a sample by reacting it with a standard solution. One type of titration uses a neutralization reaction, in which an acid and a base react to produce a salt and water: In equation 1, the acid is HCl (hydrochloric acid) and the base is NaOH (sodium hydroxide).

  8. Titration of HCl with NaOH

    About Press Copyright Contact us Creators Advertise Developers Terms Privacy Policy & Safety How YouTube works Test new features NFL Sunday Ticket Press Copyright ...

  9. PDF Lab Report #4 Titration of Hydrochloric acid with Sodium Hydroxide

    Lab Report #4. um HydroxideSCH3U. 02Thursday, December 19, 2013Introd. tion The following lab was an acid-base neutralizing titration. A titration is a technique, in which a reagent, called a titrant, of known concentration is use. to determine the concentration of an analyte or unknown solution. Using a.

  10. Determination of the Concentration of a Solution of Sodium Hydroxide

    When the end-point of the titration is reached, the color of the solution changes from colorless to pink. Procedures. You will use the Determine the concentration of an unknown HCl solution and Standardization of NaOH experiment (OLI Unit 3 - Module 11)

  11. Titration: Practical and Calculation (NaOH and HCl)

    Must see: My new website at http://ww.acechemistry.co.uk. This video, with the help of flash animations, shows and explains how you can determine the concent...

  12. Acid-Base Titration > Experiment 24 from Chemistry with Vernier

    A titration is a process used to determine the volume of a solution needed to react with a given amount of another substance. In this experiment, you will titrate hydrochloric acid solution, HCl, with a basic sodium hydroxide solution, NaOH. The concentration of the NaOH solution is given and you will determine the unknown concentration of the HCl. Hydrogen ions from the HCl react with ...

  13. A thermometric titration

    In this thermometric titration, students can determine the end-point of the titration using the highest temperature recorded during the experiment. Using a pipette and safety filler, transfer 20 cm 3 (or 25 cm 3) of the sodium hydroxide solution into the cup, and measure the steady temperature. Using the burette, add a small portion (3-5 cm 3 ...

  14. Titration Lab (HCl + NaOH)

    Learn how titrations are used to figure out the unknown molarity of solutions.Student Work- https://docs.google.com/document/d/17TFZjASP--7vJKUVoyT-1ySilzzRx...

  15. PDF Experiment 6 Titration II

    Figure 1 is a titration curve for the titration of HCl by NaOH, a strong acid and strong base, where 25.0 mL of 0.1 M HCl is titrated with 0.1 M NaOH. NaOH titration of HCl 0 2 4 6 8 10 12 14 0 102030 mL NaOH added p H 40 Phenolphthalein Equivalence Point Methyl Red Figure 1. Titration of 25.0 mL of 0.1M HCl by 0.1 M NaOH. Blocked areas on the ...

  16. PDF HCl(aq)+NaOH(aq) H O(l)+NaCl(aq) HC H O (aq)+NaOH(aq) NaC H O (aq)+H O

    Experiment 7 Titration Curves of Strong and Weak Acids and Bases Goals To calibrate a pH probe. To become familiar with acid-base titration curves. To determine the concentration of an unknown acid. Equipment and Materials 0.10 M ammonium hydroxide solution, 0.10 M sodium hydroxide solution, 0.0074 M hydrochloric acid

  17. Performing Titrations

    HCl + NaOH NaCl + H 2 O During the course of the titration, the titrant (NaOH) is added slowly to the unknown solution. As it is added, the HCl is slowly reacted away. The point at which exactly enough titrant (NaOH) has been added to react with all of the analyte (HCl) is called the equivalence point.

  18. Titration Lab Report

    About the lab tiration titration lab report aim to find the concentration of sodium hydroxide titrating hydrochloric acid with sodium hydroxide. introduction. ... HCl+NaOH → H2O+NaCl. Predictions. When this experiment is carried out the base in the conical flask should turn colorless after adding acid to the solution. The base solution should ...

  19. Acid-Base Titrations: Standardization of NaOH and Antacid

    In this experiment, the reagents combined are an acid, HCl (aq) and a base, NaOH (aq) where the acid is the analyte and the base is the titrant. The reaction between the two is as follows: HCl (aq) + NaOH (aq) → H2O (l) + Cl -(aq) + Na +(aq) In this case, Sodium and Chloride act as spectator ions and form into salts in a neutralization ...

  20. PDF Experiment 7: Titration of an Antacid

    minus the number of moles that were neutralized by the NaOH: moles acid neutralized = (moles of HCl added) - (moles of NaOH required for back-titration) = (MHCl x VHCl) - (MNaOH x VNaOH) Eq. 4 where M = molarity and V = volume in liters. In today's experiment, half of the students in the class will analyze Tums tablets

  21. NaOH vs HCl Titration using Phenolphthalein

    In this video, I have demonstrated the whole experiment related to NaOH vs HCl titration. Also included precautions, practical copy record and important poin...

  22. 7.2: Lab

    pH Titrations. In a pH titration you measure the pH as a function of the volume of titrant added and determine the equivalence point as the point in where there is an inflection in the slope of the curve. Figure 7.2.2 7.2. 2 shows the four common types of titrations. Initially the pH is that of the pure analyte.

  23. Quantitative Analysis of Acid-Base Titrations: A Virtual

    Procedure Part I: Practicing the Titration Technique - Titration of an Acid of Known Concentration 1. In this portion of the experiment, you will practice titrating 25.00 mL of 0.3000 M HCl (your analyte) with 0.3000 M NaOH (your titrant). You will use your data to calculate the experimental concentration of HCl, with the goal to get as close to 0.3000 M as possible (since we know this is ...

  24. Copy of pH and Titration Online Lab (pdf)

    Exercise 1: Titration of HCl solutions In this exercise, we will use two HCl solutions that have been premixed with phenol red, 0.02 M HCl + phenol red and 0.04 M HCl + phenol red, to practice the titration procedure. Watch this Introduction to Exercise 1 video to view the experimental set-up that includes steps 1 and 2 in the protocol below.

  25. Put the following experimental steps for the Acid-Base Equilibria (HCL

    One pair performs a careful colorimetric titration of HCL with NaOH and the other pair performs a potentiometric titration of HCL with NaOH. Hence correct sequence is e, a, b, ... To determine the concentration of a sodium thiosulfate solution as in this experiment, a student pipetted 25.0 mL of 0.0100 M potassium iodate (KIO3) solution into a ...

  26. A Student Performed A Titration Using 12.5. Ml Of Naoh Solution With 0

    7 is the pH of a solution that results from mixing equal volumes of 0.04M ammonia and a 0.02mM solution of hydrochloric acid.. The pH of a solution that results from mixing equal volumes of 0.04 M ammonia and a 0.02 M solution of hydrochloric acid can be calculated using the formula for the pH of a mixture of weak acids and bases.. The pH of a mixture of weak acids and bases can be calculated ...

  27. Colorimetric Titration: Techniques and Equivalence Points

    Titration is a technique in which an unknown concentration of a solution (the analyte) will be investigated while reacting with a known concentration of another solution (the titrant) until it completely reacts and reaches the equivalence point. In this lab, colorimetric titration is being utilized which the equivalence point is measured by the color change.

  28. Titration NaOH vs HCl

    Titration involving Sodium Hydroxide and Hydrochloric Acid

  29. Water

    Subsequently, HCl solution (50 mL, ... V is the volume of NaOH standard solution used in the titration process, mL; 25 is the volume of NaOH standard solution used, mL; c 0 is the ... concentration (50-200 mg/L) on the adsorption performance were investigated. After the adsorption experiment, the solution was filtered, and the concentration ...