# UP Chemistry Acid Base Titration Lab Report

i need a lab report about Acid-Base Titration

Laboratory Report

Reports consist of the following sections:-Report Sheets – the completed Data and Report Sheets cut out of your lab manual.Calculations – If there are calculations, you must clearly and logically show all of the calculations to getfull credit.Discussion – Explain your results, referring to the chemistry theory where appropriate. Did the dataconform to what was expected? What does it show in relation to the experiment? If the results are notwhat was expected explain why. Possible sources of error (estimate experimental error in numericalresults). Conclusion – briefly state any final result(s), or conclusions that can be drawn from the experiment.References – any sources consulted for the writing of the report. These should be referenced (cited) inthe body of the text.

The discussion should be at least one page.

Put the calculation separately.

The conclusion should NOT be short

here all information including titration (volumetric) Experiment instruction including worksheet, and calculation links. Also attached separately table 2 and table 3 data screen shot. Table 1 can be calculated based on instruction in the lab manual. In Table 1 KHP concentration 0.001572 M to standardize NaOH in table 2. on FIELS.

GENERAL CHEMISTRY LAB SERIES

UNIVERSITY OF BRIDGEPORT-CHEMISTRY DEPARTMENT

Acid-Base Titration (Volumetric Glassware uses)

OBJECTIVES

• To standardize a NaOH(aq) solution by titration

with potassium hydrogen phthalate.

• To determine the acid concentration of an

unknown solution by titration with a standard

NaOH(aq) solution.

• To master using volumetric flasks, volumetric

pipets, and burets.

• To gain understanding of titration calculations.

BACKGROUND INFORMATION

General Overview

A titration is an analysis in which a solution with

a known reagent concentration (the standard) is

used to determine the unknown concentration of a

different reagent (the analyte). The standard and

analyte chemically interact in a reaction with known

stoichiometry. In the most common type of titration,

in an acid/base neutralization reaction occurs.

Titration is a form of volumetric analysis that

requires careful volume measurements.

The

experimental setup for a typical titration is shown in

Figure 1. A buret is a long glass tube with fine

graduated markings that allow for relatively high

precision measurement of delivered volume. A

valve (stopcock) at the bottom of the buret controls

the delivery of solution from the buret to the

receiving flask.

The reagent which is being

dispensed by the buret is known as the titrant.

Depending upon the experimental circumstances,

the titrant may be either standard or analyte.

In the more common configuration, an exactly

measured volume of analyte is transferred to the

receiving flask with a volumetric pipette.

The

standard titrant is then delivered from the buret to

the receiving flask in a controlled manner where it

reacts with and depletes the analyte. The titration is

stopped at the exact point when all of the analyte is

consumed – this is the equivalence point. At the

equivalence point the buret valve is closed, and the

volume of standard titrant delivered is read from the

buret. The known concentration of the standard

titrant, the volume of standard titrant, the volume of

analyte, and the reaction stoichiometry are then

used to calculate the concentration of the analyte.

Dr. H.

Figure 1: A typical titration apparatus

Indicators and Endpoints

So how do you know when you reach the

equivalence point? For most titrations, the analyte,

titrant, and product solutions are all colorless and

thus indistinguishable. To make the equivalence

point observable, an indicator is added to the

sample solution. The indicator (In) is a substance

that changes color when it reacts with the titrant (T).

In (Color 1) + T → In (Color 2)

Prior to the equivalence point, the analyte is in

excess and quickly consumes any added titrant and

the indicator retains its original color. However, at

the equivalence point, all of the analyte has been

consumed. Thus, the next drop of titrant will react

with the indicator and convert it into its product color.

In this experiment, we will be using phenolphthalein

as our indicator. For our titrations, phenolphthalein

will initially be colorless and turn pink at the

endpoint. Another indicator, bromothymol blue is

also available in the lab (it will turn from blue to

yellow at the equivalence point).

The point at which the analyst first notices a color

change is known as the endpoint. In an ideal

titration, the endpoint occurs at the exact same

volume as the equivalence point. In practice, it is

possible for the analyst to go past the equivalence

point before a color change is observed. When this

occurs, the difference between the equivalence point

and the endpoint is known as the titration error.

PAGE 2 OF 8

Standards and Standardization

In order to accurately determine the analyte

content of your sample, the titrant concentration

must be accurately known. A primary standard is a

pure substance from which very accurate solution

concentration can be made. Thus, the primary

standard functions as the chemical measuring stick

used to quantify the analyte. However, sometimes

the analyte cannot be directly titrated with the

primary standard. For example, in this experiment,

both the primary standard and the analyte are acids

and cannot react with each other.

In these

circumstances, a secondary standard must be used

as an intermediate measuring stick between the

primary standard and the analyte.

A secondary standard is a substance capable

of reacting with both the primary standard and the

analyte. When a secondary standard is used, the

titration

analysis

occurs

in

two

stages:

standardization and sample analysis.

In the

standardization titrations, the secondary standard

is titrated with the primary standard. From these

titrations, the concentration of the secondary

standard can be accurately determined. Then, the

sample is titrated with the secondary standard to find

the analyte concentration.

In this experiment, our primary standard will be

potassium hydrogen phthalate and our secondary

standard will be sodium hydroxide. Potassium

hydrogen phthalate is an acid salt with the formula

KHC8H4O4, MM = 204.2 g/mol. For shorthand,

chemists frequently refer to potassium hydrogen

phthalate as “KHP” where P2– = C8H4O42–. Because

KHP is a weak acid, an acid-base neutralization

reaction occurs when it is added to sodium

hydroxide:

KHP(aq) + NaOH(aq) → H2O(l) + NaKP(aq)

Or in net ionic form:

–

–

2–

HP + OH → H2O(l) + P

Our experimental procedure will consist of three

parts. First, we will be preparing a primary standard

KHP solution with an accurately known

concentration. Second, we will use this standard

KHP solution to establish the conentration of a

sodium hydroxide solution. Finally, we will use this

secondary standard NaOH(aq) to determine the acid

concentration of a unkown solution.

At this point, you may be wondering why we can’t

eliminate the entire standardization step by simply

preparing an NaOH(aq) solution with an accurate

concentration. Two problems prevent us from doing

this.

First, solid NaOH is very hygroscopic—

meaning it quickly absorbs moisture from the air. If

Dr. H.

a NaOH pellet is placed on an analytical balance in

humid air, you can actually see the NaOH pick up

mass. Because NaOH(s) pellet begins absorbing

water as soon as it is exposed to the air, it is

impossible to know how much its mass is pure

NaOH and how much is absorbed water. This

uncertainty in true NaOH(s) mass will compromise

the accuracy of the NaOH(aq) concentration.

The second problem is that NaOH also reacts

with the carbon dioxide in air:

NaOH(s) + CO2(g) → NaHCO3(s)

Although this reaction is slower, it produces the

same effect as water absorption. In both cases,

impurities will make up a significant proportion of the

mass of a “NaOH(s)” pellet and the resulting

concentrations will be inaccurate. For this reason,

NaOH cannot be used as a primary standard.

Even after your NaOH(aq) solution is prepared,

the absorption of CO2 is a cause for concern. If your

NaOH(aq) solution is left exposed to the atmosphere

for long periods of time, its concentration will

change. For this reason the NaOH(aq) reagent

bottle should be kept tightly capped while not in use.

Titration Calculations

Just as at the heart of every titration is a reaction,

at the heart of every titration calculation is a reaction

equation. Consider the general reaction.

aA + sS → pP+qQ

where A is the analyte, S is the standard, and P and

Q are products. The symbols a, s, p, and q are the

stoichiometric coefficients of the equation. These

coeffiecients give the mole ratio in which reactants

combine and products are generated.

The

relationship between the number of analyte moles

(nA) and the number of titrant moles (nT) is:

n A = nS ´

a

s

(1)

For every possible type of titration calculation,

Equation (1) is always used. Similarly, because the

titrant is always in solution form, the number of

moles of titrant is always calculated as

n = M·V

(2)

where M is the concentration and V is the volume.

The part of a titration calculation which is different

from case to case is the method of calculating the

number of moles of analyte. If the analyte is a

solution then the number of analyte moles is

calculated from equation (2). If the analyte is a

solid, then the number of analyte moles is calculated

from its mass, m, and its molar mass, MM.

PAGE 3 OF 8

n=

m

MM

(3)

While there are many different varieties of

titration calculations, all of them are based upon

some combination of Equations (1), (2), and (3).

A systematic stepwise solution using Equations (1)

and (2) is shown in the sample calculation in Box 1

below.

Streamlining Calculations 1: Combined Equation

The most common types of titration calculations

are those involving two solutions. A standard

titration calculation of this type will consist of three

steps in which equation (2) is used twice and

equation (1) is used once (see sample calculation

below). The calculation can be streamlined by

combining all these steps into one grand equation.

This is achieved through substitution of nA = MA·VA

and nS = MS·VS into equation (1):

M AV A = M SVS ´

a

s

(4)

This equation is valid for any titration involving two

solutions. It is very important to notice the a/s factor

at the end of equation (4). Many students forget this

factor in their calculations and obtain erroneous

answers when a ≠ s. The use of this equation is

shown in the streamlined titration calculation in Box

1 below.

Streamlining Calculations 2: Using Millimole Units

For typical laboratory experiments, titration

volumes are in mL rather than in L. Since the units

of molarity are mol/L, both our analyte and titrant

volumes must be divided by 1000 to calculate the

number of moles. However, a close examination of

the sample calculation below reveals that the mL→L

conversions are unnecessary. Notice that we are

dividing by 1000 in step 2 and multiplying by 1000 in

step 4. Thus, in the final answer, the two mL→L

conversion steps cancel each other out.

So can we just ignore units in our calculations?

No, no, no! Don’t succumb to this temptation which

each year leads students to commit violence to their

Chem120 grades. The proper approach is to do our

multiplication and division of 1000 in our units rather

than in our calculation. Notice what happens if we

divide both the numerator and denominator of

molarity by 1000 and substitute the milli prefix:

M =

1 mol (1/1000) mol 1 mmol

=

=

1L

(1/1000) L

1 mL

Box 1: Sample Titration Calculation

The concentration of a citric acid (H3C6H5O7) solution is being determined by titration with a standard

0.1067 mol/L NaOH solution. When 25.00 mL of the citric acid are titrated, it requires 18.34 mL of the NaOH

standard solution to reach the endpoint. What is the citric acid concentration?

Stepwise Solution

1. Start by writing a balanced equation for the reaction:

H3C6H5O7 (aq) + 3 NaOH(aq) → 3 H2O(l) + Na3C6H5O7 (aq)

2. Next, calculate the number of moles of the NaOH standard using equation (2).

nS = M SVS = 18.34 mL ´

1 L

´ 0.1067 mol L NaOH = 1.9569 ´ 10-3 mol NaOH

1000 mL

3. Now, we use equation (1) to find the moles of the citric acid analyte.

n A = nS ´

a

1 mol H 3C6 H 5O7

= 1.9569 ´ 10-3 mol NaOH ´

= 6.5229 ´ 10 -4 mol H 3C6 H 5O 7

s

3 mol NaOH

4. And finally, we use a rearranged equation (2) to find the citric acid analyte concentration.

n A 6.5229 ´ 10-4 mol H 3C6 H 5O7 1000 mL

MA =

=

´

= 0.02609 mol H 3C6 H 5O7

L

VA

25.00 mL

1 L

Streamlined Solution – Using Rearrangement of Equation (4) and Millimole Units

MA =

M SVS a 0.1067 mmol mL ´ 18.34 mL 1 mol H 3C6 H 5O7

´ =

´

= 0.02609 mol L H 3C6 H 5O7

VA

s

25.00 mL

3 mmol NaOH

PAGE 4 OF 8

Thus, another valid definition of molarity is mmol/mL.

Using this definition allows us to conveniently keep

our titration volumes in milliliters and eliminate two

conversion calculations. The use of millimole units

is shown in the streamlined in Box 1.

Concentration of an Acid Sample

Once we have established the concentration of

our sodium hydroxide solution (the secondary

standard) we will use it to determine the

concentration of an acid solution with an unknown

concentration. It is often the case that the acid

concentration is significantly higher than the

concentration of our standard solution and, if we

titrated directly, it would require an inordinate

amount of titrant. This is not a significant problem as

we can accurately dilute the acid to bring its

concentration down to a more reasonable level. The

following scenario gives an example of the

manipulations and calculations involved.

Acid Sample Concentration Analysis

A chemist needs to accurately determine the

concentration of an industrial waste sample that

contains oxalic acid (a diprotic acid) with a

concentration of about 0.7 mol/L. She has available

the KOH solution standardized in the previous

example (0.1566 mol/L or 0.1566 mmol/mL). Using

a volumetric pipet she transfers a 25.00 mL portion

of the original oxalic acid solution to a 250.0 mL

volumetric flask. After diluting and mixing she uses

a volumetric pipet to transfer 20.00 mL of this

solution into an Erlenmeyer flask. She then titrates

this sample to the endpoint with the standard KOH

solution. It takes 17.52 mL of the KOH solution to

reach the endpoint. What is the concentration of the

original unknown?

Titration Calculation

We will use the streamlined solution shown at the

end of the previous example to calculate MA, the

concentration of the diluted oxalic acid solution (the

analyte). To calculate this concentration we need to

extract the values for the following variables from the

data given in the problem.

MS

VS

VA

a

s

concentration of standard titrant

volume of standard titrant required

volume of analyte titrated

stoichiometric coefficient of analyte

stoichiometric coefficient of standard titrant

We find a and s from the balanced chemical

equation for the reaction. The formula for oxalic acid

is H2C2O4 and it reacts with potassium hydroxide

solution according to the following.

Dr. H.

H2C2O4(aq) + 2KOH(aq) → K2C2O4(aq) + 2H2O(l)

The KOH solution is the titrant so s = 2 mol KOH

and the oxalic acid solution is the analyte, so a = 1

mol H2C2O4.

The concentration of the standard KOH solution

is given as 0.1566 mol/L and 17.52 mL of this

solution was needed to reach the endpoint, thus we

have MS and VS.

Determining VA can be trickier because, with a

dilution step and a titration step, we see several

volumes associated with the oxalic acid solution.

Since we are dealing with the titration step in this

calculation we must look for the volume of solution

that is titrated.

Careful reading reveals that

20.00 mL of the dilute oxalic acid is placed in the

Erlenmeyer for titration, so VA is equal to 20.00 mL.

We now have values for all the variables so it is

just a matter of running the calculation.

MA =

(0.1566 mol KOH L )(17.52 mL ) æ 1 mol H 2C 2O 4 ö

ç

÷

è 2 mol KOH ø

20.00 mL

M A = 0.06859 mol/L H 2C 2O 4

This is not the final answer though – this calculation

gives the concentration of the dilute oxalic acid

sample. But we want to know the concentration of

the original oxalic acid solution, so we have one

more calculation to perform – a dilution calculation.

Dilution Calculation

Dilution calculations are usually presented as

follows:

M iVi = M f V f

(5)

where M represents concentration and V represents

volume (just as with the titration calculation) and the

subscripts indicate initial (i) conditions and final (f)

conditions, that is before and after the dilution.

Sometimes you may see this relation written with 1’s

and 2’s in place of the I’s and f’s, but it works the

same either way.

The original oxalic acid solution concentration is

represented by Mi, thus we need to rearrange

Equation 5 to solve for Mi.

Mi =

M fVf

Vi

(6)

The final concentration, Mf, is the concentration of

the dilute oxalic acid determined by the titration and

calculated above. As part of the dilution process the

chemist withdrew a 25.00 mL sample and placed it

in a 250.0 mL volumetric flask, thus Vi = 25.00 mL

PAGE 5 OF 8

and Vf = 250.0 mL. Substituting into equation (6),

the original solution is found to have an oxalic acid

concentration of 0.6859 mol/L.

Mi =

(0.06859 mol L )(250.0 mL ) = 0.6859 mol

25.00 mL

L

REFERENCES

For additional information on the topics covered

in the Background Information, consult the following

sections from Chemistry (3rd ed.), Nivaldo J. Tro,

Pearson, 2011.

• Solution Concentrations & Dilutions, Section 4.4

• Acid-Base Titrations, Section 4.8

PRACTICAL ASPECTS

Throughout this experiment you will need to

exercise careful laboratory technique in order to

obtain successful results (and thus a good accuracy

and precision score). Review the “Lab Equipment

Instructions” document for the proper use burets and

pipets used in this experiment.

beaker and the funnel with three small portions of

deionized water into the volumetric flask. Fill the

volumetric flask half-full of deionized water, stopper

it, and swirl to dissolve the crystals. When all solid is

dissolved, fill the volumetric flask to the mark with

deionized water and mix thoroughly by inverting 20

times or more.

2. Standardization of the Sodium Hydroxide

Solution.

Buret Preparation

Review the “Lab Equipment Instructions” section

on buret usage before proceeding further. Rinse

your burets thoroughly with tap water and deionized

water. A clean buret will drain freely without forming

drops on the inside surface.

Insure that the

stopcock functions properly; it must turn freely and

must not leak when turned off.

EXPERIMENTAL PROCEDURE

Rinse the buret with three 5-mL portions of the

NaOH solution running a portion of each rinse

solution through the stopcock and tip of the buret.

Discard the rinse solution. Fill the buret with the

NaOH solution. Place the buret cap on the buret to

reduce the rate of reaction between sodium

hydroxide and atmospheric carbon dioxide.

Chemicals:

Standardization Titrations

0.1 M sodium hydroxide (NaOH)

Potassium hydrogen phthalate (KHP),

KHC8H4O4, anhydrous

Dilute acid solutions of unknown concentration

Phenolphthalein indicator solution, or

bromothymol blue solution (an alternate

indicator)

Special Equipment (To be checked out):

50-mL buret;

buret funnel

buret cap

100 mL volumetric flask

One 10 mL volumetric pipet

One pipet bulb

1. Preparation of a Standard KHP Solution:

Refer to the “Lab Equipment Instructions”

document on volumetric flask technique before

reading further. Carefully weigh a clean, dry 50 mL

beaker as accurately as possible. Place between 3

g and 3.5 g of anhydrous potassium phthalate, KHP,

in the beaker and weigh accurately. Transfer the

weighed crystals to a clean 100 mL volumetric flask,

being careful to prevent loss of the crystals. To

ensure every bit of KHP is transferred, rinse the

Review the “Lab Equipment Instructions”

document covering the use of volumetric pipets and

burets before proceeding.

Using a 10 mL volumetric pipet, transfer 10 mL of

the KHP solution into a clean, rinsed 250 mL

Erlenmeyer flask.

Touch off the pipet tip if

necessary and rinse the inside walls of the flask with

a minimum amount of deionized water. Add enough

DI water such that the solution covers the bottom of

the flask. Add 2 drops of phenolphthalein indicator.

(Alternatively, if you have difficulty seeing pink, you

may use thymol blue as the indicator. The color

change for thymol blue is from yellow to blue.)

NOTE: be consistent from one run to the next for

good precision and accuracy.

Your first titration (Trial 0) will be a rough titration

to get an approximate volume of titrant needed to

reach the endpoint. This approximate volume,

although not used to calculate the average

concentration of your unknown, will give you practice

using the buret and will allow you to more quickly

complete subsequent trials.

Record the initial buret volume to the nearest

0.01 mL. Position the Erlenmeyer flask under the

NaOH buret. Place a sheet of white paper or a

paper towel under the flask to help you see the

indicator more easily. Begin titrating the KHP with

PAGE 6 OF 8

NaOH by opening the stopcock of the NaOH buret

almost all the way. Immediately begin stirring the

solution by swirling the flask. As the end point is

approached, the faint pink color will appear in parts

of the unmixed NaOH solution will persist for several

seconds.

For Trial 0 slow the flow somewhat (to a rapid

rate of drops instead of a continuous flow) and stop

when the pink color of the indicator persists. Record

the final buret volume and calculate the volume of

titrant used (final volume − initial volume).

For Trials 1, 2, and 3 follow the same procedure,

except slow the rate of addition when you reach

approximately 90% of the volume of Trial 0. The

addition rate should be such that you are adding

single drops and thoroughly mixing the solution after

each drop is added. The end point is reached when

a single drop (or fraction of a drop) produces a faint

pink color which lasts for at least 30 seconds after

vigorous swirling. The pink color will fade after the

solution has been standing awhile because carbon

dioxide will be absorbed from the air, and will make

the solution slightly acidic again. (You can verify the

effect of carbon dioxide by blowing on your finished

titration solutions while swirling and watching the

indicator change back to its original color.) Record

the final buret volume and calculate the volume of

titrant used (final volume − initial volume).

When the titration is finished, the neutralized

mixture in the titration flask may be disposed of by

pouring down the sink with running water.

Refill your buret with the sodium hydroxide

solution as necessary.

Complete at least three standardization titrations.

Enter your titration data into spreadsheet. Run an

additional titration (Trial 4) if the volume range

between the highest and lowest value of Trials 1, 2,

and 3 is greater than 0.2 mL.

Contact your

instructor if, after your fourth trial, this difference is

still greater than 0.2 mL.

Calculations

From your titration data, calculate the

concentration of your NaOH solution. Because of

the volumetric accuracy of your measurements, you

should be able to report your answers to four

significant digits (unless your titration volumes are

less than 10 mL). Calculate the average, standard

deviation, and relative standard deviation of your

concentration.

We will need to recycle the volumetric glassware

from this part of the experiment for use in the

second part. However, we do not want you to

dispose of your KHP solution unless we know you

Dr. H.

have (reasonably) good data. Have your instructor

or TA check your data, and, once given the OK,

discard the leftover KHP solution in the sink with the

tap water running. Rinse the volumetric flask, flask

cap, and pipet three times with regular tap water and

then do a final rinse with a small amount of DI water.

If you accidentally drew KHP solution into your pipet

bulb rinse it as well.

3. Concentration of an Acid Solution

Each student will be assigned one of the

available acid solutions of unknown concentration

for analysis. Record the letter, chemical name, and

chemical formula of your acid unknown in your

report and in the spreadsheet.

As discussed above, the acid unknowns are too

concentrated to analyze directly so we must dilute

them before analysis. There are common use pipets

associated with each of the acid unknowns. Use the

appropriate common pipet to transfer a sample of

the assigned unknown to your rinsed 100 mL

volumetric flask. NOTE: do not worry if there is

residual DI water in the volumetric flask – it will not

affect your results. Fill the flask about two-thirds full

with DI water, insert the cap, and carefully mix the

solution by inverting the flask at least ten times.

Remove the cap and CAREFULLY fill to the line with

DI water. It is a good idea to use a dropper instead

of your squirt bottle as the dropper gives you more

control – if you go over the line you will need to

dump the solution and start over. Reinsert the cap

and mix again.

Your own 10 mL volumetric pipet should be

rinsed at this point, but it probably contains residual

DI water from the rinsing. This water must be

removed otherwise it will dilute your acid and cause

erroneous results. To rinse the pipet, transfer about

20 mL of the acid from your flask into a clean, dry

50 mL beaker. Draw a small portion of this acid into

your pipet and let it rinse through the pipet into the

sink. Repeat this step with a second portion of acid.

Thought Question: why do we need to remove

residual water from the pipet but not from the

volumetric flask?

Use the rinsed pipet to transfer 10 mL of your

unknown acid from the volumetric flask into a rinsed

250 mL Erlenmeyer flask. Rinse the walls with DI

water and add enough DI water so that the solution

just covers the bottom of the flask. Add your

preferred indicator and titrate the acid with the

standardized NaOH solution to the endpoint. As

with the standardization titrations in Part 2 above,

use the first titration (Trial 0) to get an approximate

volume and then carefully titrate three more samples

(and a fourth if necessary, see below).

PAGE 7 OF 8

When the titration is finished, the neutralized

mixture in the titration flask may be disposed of by

pouring down the sink with running water.

Complete at least three unknown acid titrations.

Enter your titration data into spreadsheet. Run an

additional titration (Trial 4) if the volume range

between the highest and lowest value of Trials 1, 2,

and 3 is greater than 0.2 mL.

Calculations

Using your NaOH titration volumes and the

standardized concentration determined in the first

part of this experiment, calculate the concentration

of the acid. This is the concentration of the acid in

your volumetric flask but we want you to report the

concentration of the acid as it is in the main

container. Remember that you titrated the diluted

acid solution so you will need an additional dilution

calculation to determine the concentration of the

original acid.

Waste Disposal

Pour any unused KHP, NaOH, and unknown acid

solutions together into your largest beaker and mix

by stirring. Flush this mixture down the sink with

copious amounts of running water. Allow the tap

water for run for about a minute after pouring

unused solution in the sink.

Clean your area, replace any equipment you may

have taken from the common equipment drawer,

and make sure all reagent bottles at your station are

closed.

Before leaving the lab…

• Be sure you have cleaned and returned all of the

equipment that you checked out from the

storeroom.

• Be sure you have properly entered and saved all

of your data in the Lab06.xls spreadsheet.

• Police and clean up your area – you will not be

“signed out” until your station is clean.

• Have your laboratory report “signed out” by your

instructor or teaching assistant.

NAME:

INSTRUCTOR NAME:

1. KHP SOLUTION PREPARATION

Data (MUST be completed in ink)

Standard KHP Solution

Mass of KHP + beaker

Mass of beaker (0 if tared)

Mass of KHP

Final Dilution Volume

Calculations (should be completed in pencil)

1.

Concentration of Standard KHP Concentration

Calculate the concentration of your KHP solution in mol/L. Show your work. In your calculations, be sure to

show units, and report your answer to the correct number of significant figures.

2. NAOH STANDARDIZATION

Buret Volume Data (in ink)

* Trial 0

Trial 1

Trial 2

Trial 3

†

Trial 4

Volume KHP Delivered

NaOH buret: Final reading

Initial reading

Volume Titrant Delivered

‡

Instructor Initials for Excluded Data →

* Trial 0 is a practice run to get a rough idea of the volume of titrant needed – do not transfer this value

to the Excel spreadsheet.

†

Run a fourth trial if the volume range between the highest and lowest value of Trials 1, 2, and 3 is

greater than 0.2 mL.

‡

Instructor must confirm and initial appropriate box for excluded data, otherwise five points will be

deducted from accuracy and precision score.

Use the space below to explain: (i) if a known error was made in any of your trials; and/or (ii) your reasoning if

you decide to exclude one of your trials from your results.

Dr. H.

Calculations (should be completed in pencil)

Calculate the exact molarity of your NaOH solution to the proper number of significant digits. You must show a

complete set of calculations for Trial 1. It is not necessary to show your work for successive trials, but you must

record the final result for all trials in a table of your own construction. From your trial results, calculate the

average NaOH molarity, the standard deviation, and the percent relative standard deviation – include the results

of these calculations in your table.

3. UNKNOWN ACID SOLUTION CONCENTRATION

Unknown #: __________

Acid Name: ______________________________

Formula: ______________

Write a balanced chemical equation for your titration reaction.

__________________________________________________________________________________________

Dilution Data

Volume Initial (common pipet volume)

Titration Data

Volume Final (volumetric flask volume)

* Trial 0

Trial 1

Trial 2

Trial 3

†

Trial 4

Volume Analyte, Diluted Acid

NaOH buret: Final reading

Initial reading

Volume Titrant Delivered

‡

Instructor Initials for Excluded Data →

* Trial 0 is a practice run to get a rough idea of the volume of titrant needed – do not transfer this value

to the Excel spreadsheet.

†

Run a fourth trial if the volume range of Trials 1, 2, and 3 is greater than 0.2 mL.

‡

Instructor must confirm and initial appropriate box for excluded data, otherwise five points will be

deducted from accuracy and precision score.

Use the space below to explain: (i) if a known error was made in any of your trials; and/or (ii) your reasoning if

you decide to exclude one of your trials from your results.

Titration Calculations (should be completed in pencil)

Calculate the concentration of your diluted acid sample to the proper number of significant digits. You must show

a complete set of calculations for Trial 1. It is not necessary to show your work for successive trials, but you

must record the final result for all trials in a table of your own construction.

Dilution Calculations (should be completed in pencil)

Calculate the concentration of your original acid sample to the proper number of significant digits. You must show

a complete set of calculations for Trial 1. It is not necessary to show your work for successive trials, but you

must record the final result for all trials in a table. From your trial results, calculate the average concentration of

your unknown, the standard deviation, and the percent relative standard deviation – include the results of these

calculations in your table. Record your final results in the summary section below.

Results Summary:

Unknown Number: __________Acid Name: ______________________________

Formula: ______________

Original Solution Concentration (avg.): ___________________ +/– Standard Deviation: ___________________

Relative Standard Deviation: ____________________

Dr. H.