Jeoloji M

8/9/2019 Jeoloji M..

1/25

GENERAL CHEMISTRY

MANUAL

2009


2/25

EXPERIMENT 11-PREPARATION OF SOLUTIONS

1. PURPOSE

The purpose of the experiment is to learn how to prepare a solution with known

concentration.

2. THEORY

A solution is a homogeneous mixture created by dissolving one or more solutes in a solvent.

The chemical present in smaller amount, the solute, is soluble in the solvent (the Chemical

present in larger amount). Solutions with accurately known concentrations can be referred to

as standard (stock) solutions. These solutions are bought directly from the manufacturer or

formed by dissolving the desired amount of solute into a volumetric flask of a specific

volume.

Preparing a Standard Solution from a Solid

A solution of known concentration can be prepared from solids by two similar methods.

Although inherent errors exist with each of the methods, with careful technique either will

suffice for making solutions in General Chemistry Laboratory. In the first method, the solid

solute is weighed out on weighing paper or in a small container and then transferred directly

to a volumetric flask (commonly called a "vol flask"). A funnel might be helpful when

transferring the solid into the slim neck of the vol flask. A small quantity of solvent is then

added to the vol flask and the contents are swirled gently until the substance is completely

dissolved. More solvent is added until the meniscus of the liquid reaches the calibration mark

on the neck of the vol flask (a process called diluting to volume). The vol flask is then

capped and inverted several times until the contents are mixed and completely dissolved. The

disadvantage of this method is that some of the weighed solid may adhere to the original

container, weighing paper, or funnel. Also, solid may be spilled when it is transferred into the

slim neck of the vol flask.

In the second method the solid is weighed out first in a small beaker. A small amount of

solvent is added to the beaker and the solution is stirred until the solid is dissolved. The

solution is then transferred to the vol flask. Again, a funnel may need to be inserted into the

slim neck of the vol flask. Before adding additional solvent to the flask, the beaker, stirring

rod, and funnel must be rinsed carefully and the washings added to the vol flask making sure

all remaining traces of the solution have been transferred. Finally, the vol flask is diluted to

volume (additional solvent is added to the flask until the liquid level reaches the calibration

mark).


3/25

In general chemistry molarity is the most commonly used concentration unit.

solutionsofliters

soluteofmolesmolarity =

Percent solutions

a) Mass percent means the number of grams of solute per 100 g of solution. For example, 10

g sodium chloride in 90 g water is a 10% by mass solution.

mass percentage = (mass of solute / mass of solution)*100

= 10 g/ (10 g + 90 g) x100%

= 10%

b) Volume percentage means the number of milliliters of solute per 100 mL of solution

Diluting a Solution of Known Concentration

Dilution is the addition of more solvent to produce a solution of reduced concentration. Most

often a diluted solution is created from a small volume of a more concentrated stock solution.

To make such a solution, a volumetric pipet is used to deliver an exact amount of the stock

solution into a clean vol flask, which is then diluted to volume. To prevent extra dilution or

contamination, prerinse the vol pipet with the stock solution to remove any water droplets or

impurities. (The rinsings should be placed in an appropriate collection container.)

The diluted solutions molarity is less than the stock solution it was created from. The moles

present in the volume of stock solution delivered by the volumetric pipet is equal to the moles

present in the diluted solution created

2211

21

VMVM

NN

=

=

N: Number of moles of solute in the solutions

M: Molarity of Solution

V: Volume of solutions

3. EXPERIMENTAL PROCEDURE

1. For preparing % 10 weight percent NaOH solution, weight 10 g NaOH and solved itin 90 g water.

2. For preparing 1 M NaCl solution, calculate the required amount of NaCl then put itinto small beaker and add small amount of distillated water on it. After then put this solution

into

3. Calculate the amount of HCl for preparing 0,5 M 100 ml HCl solution from theconcentrated % 37 HCl solution with a density 1.18 g/ ml. Take required volume with the


4/25

help of pipettes then put it into volumetric flask. It is important that initially small amount of

water have to be exist in the flask and fill the volumetric flask to the volume line.

4. REFERENCES

1. Temel ve Genel Kimya Laboratuar Deney Fy; Mula nivesitesi FEF KimyaBlm; 20072. Manual for Chemistry Laboratory; Department of Chemical Eng. Ege University;2001

3. http://library.thinkquest.org/3310/nographics/experiments/titrate.html4. http://en.wikipedia.org/


5/25

EXPERIMENT 12-PREPARATION AND STANDARDIZATION OF

ACID-BASE SOLUTIONS

1. PURPOSE

The aim of the experiment is to learn how to determine the concentration of a solution also

referred to as standardizing a solution and to practice the technique of titration. Finally we

will determine the molar mass and identity of an unknown acid using titration and your

standardized solution.

2. THEORY

An acid-base reaction is a chemical reaction that occurs between an acid and a base. Several

concepts exist which provide alternative definitions for the reaction mechanisms involved and

their application in solving related problems. Despite several similarities in definitions, their

importance becomes apparent as different methods of analysis when applied to acid-base

reactions for gaseous or liquid species, or when acid or base character may be somewhat less

apparent. There are some definitions for acids and bases.

The Arrhenius definition of acid-base reactions is a more simplified acid-base concept

devised by Svante Arrhenius, which was used to provide a modern definition of bases that

followed from his work with Friedrich Wilhelm Ostwald in establishing the presence of ions

in aqueous solution in 1884, and led to Arrhenius receiving the Nobel prize in chemistry in

1903 for "recognition of the extraordinary services rendered to the advancement of chemistry

by his electrolytic theory of dissociation"

As defined at the time of discovery, acid-base reactions are characterized by Arrhenius acids,

which dissociate in aqueous solution form hydrogen or the later-termed oxonium (H 3O+)

ions, and Arrhenius bases which form hydroxide (OH-) ions. More recent IUPAC

recommendations now suggest the newer term "hydronium" be used in favor of the older

accepted term "oxonium" to illustrate reaction mechanisms such as those defined in the

Brnsted-Lowry and solvent system definitions more clearly, with the Arrhenius definition

serving as a simple general outline of acid-base character. More succinctly, the Arrhenius

definition can be surmised as;

Arrhenius acids form hydrogen ions in aqueous solution with Arrhenius bases forming

hydroxide ions.

The universal aqueous acid-base definition of the Arrhenius concept is described as the

formation of water from hydrogen and hydroxide ions, or hydronium ions and hydroxide ions

produced from the dissociation of an acid and base in aqueous solution


6/25

(2 H2O OH- + H3O

+ ),

which leads to the definition that in Arrhenius acid-base reactions, a salt and water is formed

from the reaction between an acid and a base in more simple scientific definitions, this form

of reaction is called a Neutralization reaction.

acid+ + base- salt + waterThe positive ion from a base can form a salt with the negative ion from an acid. For example,

two moles of the base sodium hydroxide (NaOH) can combine with one mole of sulfuric acid

(H2SO4) to form two moles of water and one mole of sodium sulfate.

2NaOH + H2SO4 2 H2O + Na2SO4

The Brnsted-Lowry definition, formulated independently by its two proponents Johannes

Nicolaus Brnsted and Martin Lowry in 1923 is based upon the idea of protonation of bases

through the de-protonation of acids -- more commonly referred to as the ability of acids to

"donate" hydrogen ions (H+) or protons to bases, which "accept" them. In contrast to theArrhenius definition, the Brnsted-Lowry definition refers to the products of an acid-base

reaction as conjugate acids and bases to refer to the relation of one proton, and to indicate that

there has been a reaction between the two quantities, rather than a "formation" of salt and

water, as explained in the Arrhenius definition.

It defines that in reactions, there is the donation and reception of a proton, which essentially

refers to the removal of a hydrogen ion bonded within a compound and its reaction with

another compound, and not the removal of a proton from the nucleus of an atom, which would

require inordinate amounts of energy not attainable through the simple dissociation of acids.

In differentiation from the Arrhenius definition, the Brnsted-Lowry definition postulates that

for each acid, there is a conjugate acid and base or "conjugate acid-base pair" that is formed

through a complete reaction, which also includes water, which is amphoteric.

The Lewis definition of acid base reactions, devised by Gilbert N. Lewis in 1923 is an

encompassing theory to the Brnsted-Lowry and solvent-system definitions with regards to

the premise of a donation mechanism, which conversely attributes the donation of electron

pairs from bases and the acceptance by acids, rather than protons or other bonded substances

and spans both aqueous and non-aqueous reactions.

Ag+ + 2 : NH3 [H3N:Ag:NH3]+

A silver cation reacts as an acid with ammonia which acts as an electron-pair donor, forming

an ammonia-silver adduct

In titration, we carefully add a solution of one reactant (the titrant) to a measured amount of

a second reactant. One of these is a standard (known concentration or molar mass) and the


7/25

other is unknown (the analyte). A color change (or some other distinctive change) occurs

when enough titrant has been added to consume all the reactant in the analyte. In this acid-

base titration we make use of the general reaction:

OH- + HA _ H2O + A

The equivalence point for this reaction is the state when we have added equal moles of OHand HA because they react in a 1:1 mole ratio. We can't always determine the equivalence

point accurately, and we often approximate it by the endpoint. Near the equivalence point of

an acid base titration, the pH of the solution changes rapidly with changing relative amounts

of HA and OH-, and this changes the color of an indicator. As we approach the endpoint, the

color change will occur near where the drop of titrant hits the solution, but this color quickly

disappears. We try to reach the point where addition of one drop of titrant causes a permanent

change in color after swirling. (Permanent means lasting more than 10 seconds.) This is the

endpoint. Titration is a very precise technique. Replicate runs should yield values that have arelative deviation of less than 30 parts per thousand (ppt). The indicator in this titration is

phenolphthalein (pronounced fee-nol-thay-lean). Phenolphthalein is colorless in acid solution,

but turns pink in basic solutions.

Applications of acid/base titrations include: determining the amount of acid or base in a

mixture, determining the molar mass of an acid or base known to be pure, and in acid rain

studies, determining the total amount of acid or base present in natural water.

The titrant concentration is generally determined by standardization. In this experiment you

prepare an analyte solution of sodium hydroxide (NaOH) approximately 0.10 M in

concentration. You then standardize this solution by using it to titrate a weighed amount of a

primary standard acid, HCl, (abbreviated KHP).

Primary Standard A substance of known purity for which we can determine the amount in

moles we have by simply weighing the material. In this experiment, NaOH is a primary

standard; (commercial HCl) is not.

Standardization is the process of determining the concentration of a solution. In this case we

standardize a HCl solution by determining the volume of HCl that reacts with a known

amount in moles of NaOH.

NaOH is a base, HCl is an acid, and they react in a 1:1 mole ratio. The ionic reaction is:

Na+ + OH- + H+ + Cl- Na+ + Cl+ + H2O


8/25


1. Preparation of 100 ml 0,2 N HCl solution: For this purpose firstly calculate theamount of HCl must be taken from the stock solution with 1,19 g/ml density and % 36 wt

percent. After then take enough amount of HCl solution with the help of 10 ml pipette and put

it into volumetric flask then add water to fill it to the volume line.

Important:HCl is very dangerous and corrosive material. It damages the cloths also body

so you have to be careful when you work with it. Dilution of HCl with water is extremely

exothermic so add HCl very slowly onto water.

DON T ADD WATER ONTO CONCENTRATED ACIDS

IT MAY CAUSE EXPLOSION

2. Preparation of NaOH solution with known concentration: Calculate the amount ofNaOH for preparing 100 ml 0,1 N NaOH solution. Weight it and put it into a volumetric flask

carefully. Then add little amount of water on it and solve completely. Finally fill it with water

till the volume line.

3. Standardization of HCl solution with NaOH solution: Firstly take 10 ml from theHCl solution and put it into Erlenmeyer flask. Add one or two drops of phenol phtalane as

indicator. Secondly take enough amount of NaOH solution and fill the burette. And titrate the

HCl solution with the NaOH and calculate the concentration of HCl solution.

4. REFERENCES

1. Temel Ve Genel Kimya Deney Fy; Mula 20072. Bowdon College Department Of Chemistry Lab. Manual; SPRING 20023. http://en.wikipedia.org/wiki/Acid-base_reaction_theories


9/25

EXPERIMENT 13-ACID BASE TITRATION

1.THEORY

Titration is a common laboratory method of quantitative/chemical analysis that can be used to

determine the concentration of a known reactant. Because volume measurements play a key

role in titration, it is also known as volumetric analysis. A reagent, called the titrant, of known

concentration (a standard solution) and volume is used to react with a solution of the analyte,

whose concentration is not known in advance. Using a calibrated burette to add the titrant, it

is possible to determine the exact amount that has been consumed when the endpoint is

reached. The endpoint is the point at which the titration is complete, as determined by an

indicator. This is ideally the same volume as the equivalence point - the volume of added

titrant at which the number of moles of titrant is equal to the number of moles of analyte, or

some multiple thereof (as in polyprotic acids). In the classic strong acid-strong base titration,the endpoint of a titration is the point at which the pH of the reactant is just about equal to 7,

and often when the solution permanently changes color due to an indicator. There are

however many different types of titrations.

Figure 1. Titration apparatus[2]

Before starting the titration a suitable pH indicator must be chosen. The endpoint of the

reaction, when all the products have reacted, will have a pH dependent on the relative

strengths of the acids and bases. The pH of the endpoint can be roughly determined using the

following rules:

A strong acid reacts with a strong base to form a neutral (pH=7) solution. A strong acid reacts with a weak base to form an acidic (pH7) solution.


10/25

When a weak acid reacts with a weak base, the endpoint solution will be basic if the base is

stronger and acidic if the acid is stronger. If both are of equal strength, then the endpoint pH

will be neutral.

A suitable indicator should be chosen, that will experience a change in color close to the end

point of the reaction.2. EXPERIMENTAL PROCEDURE

1. Clean and dry the burettes and beaker, and clamp the two burettes to the ring stand. Fillone of the two burettes with 1M HCl solution, and the other with the NaOH solution.

2. Use the buret to measure out 20 mL of HCl into an empty beaker. Add 2-3 drops of theindicator solution.

3. Titrate slowly with the NaOH solution, with constant swirling, until one single drop ofNaOH causes a permanent pink color that does not fade on swirling. Record the volume of

NaOH used.4. Use the formula M1V1=M2V2 to determine the concentration of the NaOH solution. Thissolution may now be used to titrate the unknown acid sample.

5. Replace the burette containing the 1M HCl with the burette containing the HCl solution ofunknown concentration. Refill the NaOH burette, and wash out the beaker.

6. Repeat the titration from steps 2-4 using 20 mL of the unknown acid solution to determinethe concentration of the HCl solution.

3. REFERENCES

1. Temel ve Genel Kimya Laboratuar Deney Fy; Mula nivesitesi FEF KimyaBlm; 2007

2. Manual for Chemistry Laboratory; Department of Chemical Eng. Ege University;2001

3. http://library.thinkquest.org/3310/nographics/experiments/titrate.html4. http://en.wikipedia.org/wiki/Titration


11/25

EXPERIMENT 14-pH and INDICATOR

1. THEORY

A pH indicator is a halochromic chemical compound that is added in small amounts to a

solution so that the pH (acidity or alkalinity) of the solution can be determined easily. Hence

a pH indicator is a chemical detector for hydronium ions (H3O+) (or Hydrogen ions (H+) in

the Arrhenius model). Normally, the indicator causes the colour of the solution to change

depending on the pH. pH values above 7.0 are basic, and pH values below 7.0 are acidic.

Solutions with a pH value of 7.0 are neutral.

pHindicator themselves are frequently weak acids or bases. When introduced into a solution,

they may bind to H+ (Hydrogen ion) or OH- (hydroxide) ions. The different electron

configurations of the bound indicator causes the indicator's color to change

pH indicator are frequently employed in titrations in analytic chemistry and biologyexperiments to determine the extent of a chemical reaction. Because of the subjective

determination of color, pH indicator are susceptible to imprecise readings. For applications

requiring precise measurement ofpH, a pH meter is frequently used.

Figure1. The color characteristics of some common indicators[2]2. EXPERIMENTAL PROCEDURE

1. Add 9 ml of water to 1 ml of 0.1 M HCl solution. Measure the pH of this solution. Prepare

a solution that has a Ph OF 4.

2. Prepare 0.01 M 100 ml CH3COOH solution by using 1 M CH3COOH solution.


12/25

3. Take 4 test tubes and mark them as 1,2,3,4. Then put 2-2.5 ml 0.1 M HCl at each of them.

Repeat the steps by taking four other test tubes and by putting same amount of 0.1 M NaOH

instead of HCl

4. Add 1 drop of methyl orange to test tubes number 1 of acid and base

5. Add 1 drop of bromotymol blue to test tubes number 2 of acid and base6. Add 1 drop of phenolphthalein to test tubes number 3 of acid and base

7. Add 1 litmus paper to test tubes number 4 of acid and base

3. REFERENCES

1. Temel ve Genel Kimya Laboratuar Deney Fy; Mula nivesitesi FEF KimyaBlm; 2007


3. http://en.wikipedia.org/


13/25

EXPERIMENT 15-BUFFER SOLUTIONS

1. PURPOSE:

The purpose of this experiment is to prepare buffer solutions and to determine their buffer

capacity.

2. THEORY

Its more difficult to prevent a known pH solution than to prepare it. If the solution is open to

air then it will be more acidic because of the carbondioxide absorption. If the solution is kept

on in glass bottle, the basic impurities from the glass may change the pH of it.

A buffer solution is one that is resistant to change in pH when small amounts of strong acid or

base are added. For example, when 0.01 mole of strong acid or base are added to distilled

water, the pH drops to 2 with the acid and rises to 12 with the base. If the same amount of

acid or base is added to an acetic acid sodium acetate buffer, the pH may only change afraction of a unit.

Buffers are important in many areas of chemistry. When the pH must be controlled during the

course of a reaction, the solutions are often buffered. This is often the case in biochemistry

when enzymes or proteins are being studied. Our blood is buffered to a pH of 7.4. Variations

of a few tenths of a pH unit can cause illness or death. Acidosis is the condition when pH

drops too low. Alkalosis results when the pH is higher than normal.

Two species are required in a buffer solution. One is capable of reacting with OH- and the

other will react with H3O+

. The two species must not react with each other. Many buffersare prepared by combining a weak acid and its conjugate (acetic acid and sodium acetate) or a

weak base and its conjugate (ammonia and ammonium chloride). In general, the pH range in

which a buffer solution is effective is +/- one pH unit on either side of the pKa. The

HendersonHasselbalch provides the information needed to prepare a buffer.

The pH of the acidic or basic solutions are calculated by following formulas:

s

a

aC

CpKpH log=

s

b

bC

CpKpOH log=

There is a limit to the amount of acid or base that can be added to a buffer solution before one

of the components is used up. This limit is called the buffer capacity and is defined as the

moles of acid or base necessary to change the pH of one liter of solution by one unit.

Buffer Capacity = (number of moles of OH- or H3O+ added)

Ca = Concentration of weak acidCs = Concentration of salt

Cb = Concentration of weak base


14/25

(pH change)(volume of buffer in L)


1. Take three erlenmayer of 100 ml and pour 2 ml of 30 % acetic acid solution, 25 ml ofdistilled water, 5 drops of 0.1 % Congo red. The solutions have to be turn into dark

blue color.2. Pour 5 ml of concentrated sodium acetate (CH3COONa) in to first and second erlen.

The color of the solutions changes from dark blue to pink (pH changes from 2.7 to

5.2). So, a buffer solution is prepared.

3. Pour 5 ml of 0.1 N HCl into first, 5 ml of 0.1 N NaOH into second erlenmayers. Thecolors of both solutions stay constant.

4. When 5 ml of 0.1 N NaOH solution is poured into third erlenmayer, you will see thechange in color.

4. REFERENCES

1. Temel ve Genel Kimya Laboratuar Deney Fy; Mula nivesitesi FEF Kimya

Blm; 2007


3. http://library.thinkquest.org/3310/nographics/experiments/titrate.html4. http://en.wikipedia.org/


15/25

EXPERIMENT 16-QUALITATIVE ANALYSIS OF GROUP-I CATIONS

AND GROUP-I ANIONS

1. THEORY

Classical qualitative inorganic analysis is a method of analytical chemistry which seeks to

find elemental composition of inorganic compounds. It is mainly focused on detecting ions in

an aqueous solution, so that materials in other forms may need to be brought into this state

before using standard methods. The solution is then treated with various reagents to test for

reactions characteristic of certain ions, which may cause color change, solid forming and

other obviously visible changes.

According to their properties, cations are usually classified into five groups. Each group has a

common reagent which can be used to separate them from the solution. The separation must

be done in the sequence specified below; otherwise, for example, some ions of 1st group can

also react with 2nd group reagent, so that the solution must not have any ions left from

previous groups to obtain meaningful results. The division and precise details of separating

into groups vary slightly from one source to another; given below is one of the commonly

used schemes.

2. 1ST

ANALYTICAL GROUP OF CATIONS

1st analytical group of cations consists of ions that form insoluble chlorides. As such, the

group reagent to separate them is hydrochloric acid, usually used at concentration of 12 M.

Concentrated HCl must not be used, because it forms a soluble complex ion - [PbCl 4]2- with

Pb2+. Consequently the Pb2+ ion would go undetected

The most important cations in 1st group are Ag+, Hg22+, and Pb2+. The chlorides of these

elements cannot be distinguished from each other by their colour - they are all white solid

compounds. PbCl2 is soluble in hot water, and can therefore be differentiated easily. To

distinguish between the other two, ammonia is used as a reagent. While AgCl dissolves in

ammonia (due to the formation of the complex ion [Ag(NH3)2]+ion), Hg2Cl2 gives a black

precipitate consisting of a mixture of chloro-mercuric amide and elemental mercury.

Furthermore, AgCl is reduced to silver under light, which gives samples a violet colour.

PbCl2 is far more soluble than the chlorides of the other two ions, especially in hot water.

Therefore, HCl in concentrations which completely sedimentize Hg22+ and Ag+, may not be

sufficient to do the same to Pb2+ and higher concentrations ca not be used for the

aforementioned reasons. Thus, a filtrate obtained after first group analysis of Pb2+ contains an


16/25

appreciable concentration of this cation, enough to give the test of the second group, viz.

formation of an insoluble sulfide. For this reason, Pb2+ is usually also included in the 2nd

analytical group.

2.1. EXPERIMENTAL PROCEDURE:

Route for analysis of silver group is given below.

General procedure foranalysis of Group 1 cations


17/25

3. DETERMINATION OF SOME ANIONS

The principles that are employed in the identification of cations can also be applied to the

analysis of anions. The qualitative detection of anions in a sample depends on the distinctive

solubility properties of particular salts of the ions and specific chemical reactions that are(ideally) unique to a particular ion. In this experiment, we will explore ways to detect the

presence of , SO4-2, Cl-,and I-. You will be testing both known and unknown solutions.

3.1. EXPERIMENTAL PROCEDURE:

Test for the Sulfate Ion

To 1 mL of the test solution, add 6 M HNO3 drop by drop until the solution is acidic. Then

add 1 mL of 0.1 M BaCl2 solution in order to produce a white precipitate of BaSO4.

SO4-2(aq) + Ba+2(aq) BaSO4(s) (1)

Test for the Chloride Ion

To 1 mL of a new test solution, add a couple drops of 6 M HNO3 as needed to make the

solution slightly acidic. Add 10 drops of 0.1 M AgNO3. No precipitate proves the absence of

Cl-, Br-, or I-. Centrifuge the mixture.

Test the clear filtrate with 1 drop of 0.1 M AgNO3 for complete precipitation. If necessary,

centrifuge again. Discard the filtrate. To this precipitate, add 1 mL of D.I. water, 2 drops of 6

M NH3, and 6 drops of 0.1 M AgNO3. The proportions are important, since we want to

dissolve ONLY AgCl.

AgCl(s) + 2 NH3(aq) Ag(NH3)2 +(aq) + Cl-(aq)

Shake the mixture well and centrifuge. Transfer the clear solution to a clean test tube, and

acidify once again with 6 M HNO3. A white precipitate of AgCl confirms the presence of Cl-.

Ag(NH3)2+(aq) + Cl-(aq) + 2H+ AgCl(s) + 2NH4+(aq)

Test for the Iodide Ion

Acidify a 2 mL sample of a new test solution by adding 6 M HCl. Add 1 mL of 0,1 M FeCl3

to oxidize any I- to I2. (Br- is not oxidized by Fe+3.) Add 1 mL of CCl4 and agitate the

mixture. A purple color of I2 in the CCl4 layer indicate I-was present in the original sample.

2I-(aq) + 2Fe+3(aq) I2(aq) + 2Fe+2(aq)

Test for the NO3-Ion

Place 10 drops of the solution to be tested in a clean, well-rinsed test tube. Make the solution

acidic by adding 3 M H2SO4 as needed. Next, add 5 drops of a freshly prepared, saturated

solution of iron (II) sulfate (FeSO4) and mix gently. Incline the test tube at a 45o angle, and, as


18/25

shown in the sketch, carefully add 5 drops of concentrated H2SO4 so the drops roll down the

side of the test tube and slide gently onto the top of the solution. DO NOT MIX the solutions!

Two separate liquid layers will be observed in the test tube. If NO3 - is present, a very faint

brown ring will be observed near the bottom of the test tube, thereby confirming the presence

of nitrate ion.4. REFERENCES

1. Temel ve Genel Kimya Lab. Fy, Mula niv., 20072. A New Approach to General Chemistry Laboratory; G.A.ktem, M. Acm; Abant

zzet Baysal niv. Yaynlar No:9; 1998

3. Genereal Chemistry Laboratory Manual; Ege niv.; 20014. General Chemistry; 8th Edition; R.H.Petrucci et al; Prentice Hall; 2002


19/25

EXPERIMENT 17-CHEMICAL EQUILIBRIUM

1.THEORY

In a chemical process, chemical equilibrium is the state in which the chemical activities or

concentrations of the reactants and products have no net change over time. Usually, this

would be the state that results when the forward chemical process proceeds at the same rate as

their reverse reaction. The reaction rates of the forward and reverse reactions are generally not

zero but, being equal, there are no net changes in any of the reactant or product

concentrations. This process is called as dynamic equilibrium.

In a chemical reaction, when reactants are mixed together in a reaction vessel (and heated if

needed), the whole of reactants do not get converted into the products. After some time

(which may be shorter than millionths of a second or longer than the age of the universe),

there will come a point when a fixed amount of reactants will exist in harmony with a fixedamount of products, the amounts of neither changing anymore. This is called chemical

equilibrium.

For any reaction such as:

aA + bB cC + dD

to be at equilibrium the rates of the forward and backward (reverse) reactions have to be

equal. In this chemical equation, A and B are reactant chemical species, C and D are product

species, and a,b,c and d are the stoichiometric coefficients of the respective reactants and

products.

Forward reaction rate: baBAk+

Backward reaction rate: dcDCk

where A, B, C and D are active masses and k+ and k are rate constants. Since forward and

backward rates are equal:

dc

ba

d

dcba

DC

BA

k

kK

DCkBAk

==

=

+

+

and the ratio of the rate constants is also a constant, now known as an equilibrium constant.

b

b

a

a

d

d

c

cp

PP

PPK = P: Partial pressure

The following equilibrium can be written for Kd and Kp:

( ) ndp

RTKK

= R: Gas Constant; T: Temperature; n: (nproducts - nreactants)


20/25


Take and wash six glass tubes by distilled water carefully. Pour 5 ml of 0.002 M KSCN into

each. Add 5 ml of 0.2 M Fe(NO3)3 into first tube. This (first) tube will be your standard. Pour

10 ml of 0.2 M Fe(NO3)3 into a graduated cylinder. Add distilled water by stirring up to 25

ml. Pour 5 ml of this diluted (concentration is 0.08 Fe+3

) solution into second tube. Throw thesolution out from graduated cylinder to the lavatory until it contains 10 ml. Fill the graduated

cylinder again by distilled water up to 25 ml. Strongly stir the solution. And pour 5 ml of this

solution into the third glass tube. Repeat this dilution procedure for 4th, 5th and 6th glass tubes.

The concentration of FeSCN2+ in the tubes will be found by comparing with the first glass

tube. Compare the darkness of the color of 1st tube with other tubes. For this procedure you

can take two tubes in hand and cover them with a white paper. If you look them from top you

can easily compare the colors. If their darkness of color are same then measure the heights of

the solutions. If they are not same, take a few of the solution from 1st glass tube into a cleandry beaker (it can be used later) and compare them again. Repeat the procedures until the

darkness are same. Measure the heights of the solutions. Repeat the procedure for all glass

tubes.

Table 1. Data Table For Chemical Equilibrium Experiment

Height of the solution Height of the Standard

2nd tube

3rd tube

4th tube

5th tube

6th tube

3. REFERENCES



3. Genereal Chemistry Laboratory Manual; Ege niv.; 20014. General Chemistry; 8th Edition; R.H.Petrucci et al; Prentice Hall; 2002


21/25

EXPERIMENT 18-EFFECT OF TEMPERATURE ON REACTION RATE

1.THEORY

The rate of a chemical reaction is the time required for a given quantity of reactant(s) to be

changed to product(s). Reaction rate usually is expressed in terms of moles per unit time.This rate is affected by several factors, including the nature of the reactants, concentration of

the reactants, temperature, pressure, and the presence of catalysts. In this experiment, you

will study the effects of temperature and concentration.

A chemical reaction is the result of effective collisions between particles of reactants.

Increasing the temperature of a system raises the average kinetic energy of the particles of the

system. This results in more collisions and, of greater importance, more effective collisions

per unit time. This affects the rate of the reaction.

At constant temperature, increasing the concentration of one or more of the reactantsincreases the number of particles present and, hence, the number of collisions. This affects

the rate of the reaction.

Figure 1. Reaction progress with respect to potential energy [ref. 5]

The rate of reaction increases by increasing temperature, generally. The Arrhenius equation

is a simple, but remarkably accurate, formula for the temperature dependence of the rate

constant, and therefore rate, of a chemical reaction. In short, the Arrhenius equation gives "the

dependence of the rate constant kof chemical reactions on the temperature T(in Kelvin) and

activation energyEa", as shown below.

RT

Ea

Aek

=

whereA is the pre-exponential factor or simply theprefactorandR is the gas constant.


22/25


Pour 5 ml of 0.0005 M KMnO4 solution and 1 ml of 0.25 M H2SO4 solution into five glass

tube. Take five more glass tube and pour 9 ml of 0.0025 M oxalic acid (C 2H2O4) in to them

with the help of a pipette. Put one of the tube containing KMnO4 and one containing C2H2O4

into the water bath at 25o

C. Wait for three minutes at this temperature. Then, pour the oxalicacid solution into the permanganate solution carefully. Record the time passed when the color

of the solution changes from pink to colorless. Be sure the temperature of the final solution

kept at 25 oC.

Repeat the same procedure for different temperatures (35, 45, 55, 65 oC) and record the time

passed for the color change.

3. RESULTS

1.Draw the graph for time versus temperature.

2.Calculate the reaction rate for each temperature.

4. REFERENCES



3. Genereal Chemistry Laboratory Manual; Ege niv.; 20014. General Chemistry; 8th Edition; R.H.Petrucci et al; Prentice Hall; 20025. http://en.wikipedia.org/


23/25

EXPERIMENT 19-SAPONIFICATION AND THE MAKING OF SOAP

1. THEORY

In today's experiment, we will perform a reaction that has been used for millennia: the making

of soap. Animal fat and vegetable oils are composed principally of esters of the long chain

fatty acids and glycerol (glycerin; 1,2,3-propantriol). Hydrolysis of these triglycerides

(triacylglycerides; TAG) in base (e.g., NaOH) yields glycerol (a carbohydrate) and the sodium

salts of the fatty acids. Because the fatty acids are ions, they are soluble in low concentrations

in water (actually they are soluble because they form micelles), but in high concentration form

insoluble aggregates called soap. You will start with a vegetable oil and will use NaOH to

hydrolyze these triglycerides. Basic hydrolysis of esters is called saponification. The reaction

for this experiment is shown:

The triglycerides most commonly used to make soap commercially are from animal sources,

such as tallow, although plant fats from coconut, palm and other vegetable oils can be used.

Pure coconut oil yields soap that is very soluble in water because it contains predominately

myristic and lauric acids (14- and 12-carbon fatty acids, respectively). Soaps made from

animal and other vegetable sources contain more 16- and 18-carbon fatty acids and are

generally harder and easy to form into shapes. To soften these harder soaps, coconut oil is

often included in the saponification reaction to make the soap softer.

The function of soaps and detergents is to remove grease and dirt by emulsifying the grease

(bringing it into suspension). Dirt adheres to clothing and to skin primarily by being glued

to these surfaces with a thin film of oil or grease; the oil (lipid) on the skin is generally

secreted during perspiration. The soap or detergent removes the oil film and the dirt can be

washed away. How do soaps and detergents dissolve non-polar substances such as fats, oils,


24/25

and greases? Molecules of soaps and detergents contain a non-polar (hydrophobic)

hydrocarbon end, and a polar (hydrophilic) end that is usually ionic. The non-polar ends of the

molecule surround the tiny oil droplets and are partially dissolved in them (like dissolves

like). The polar ends of the molecules, which are extremely soluble in water, solubilize or

emulsify the entire droplet.

2.EXPERIMENTAL PROCEDURE

Prepare a mixture of 15 mL of 20% (5 M) sodium hydroxide and 10 mL of vegetable oil in a

150-mL flask. Add a stirring bar to the flask, to prevent explosive boiling of the NaOH-oil

mixture. Turn the stirring hot plate on at any setting to get it boiling, and then switch to the

lowest setting when boiling begins. Turn on the stirrer to let the stirring bar rotate. Boil the

mixture, observing the precautions listed above. Carefully control the heating, but you should

heat the mixture high enough to maintain a constantly boiling mixture. The saponification is

complete if a wax-like solid begins to form that on further cooling becomes hard and

somewhat brittle. On the other hand, if the mixture cools to a syrupy liquid, saponification is

not complete, and heating and stirring must be resumed. It might be advisable to add more (5

mL) 20% NaOH and boil the mixture until its water is expelled. Saponification should be

complete by 30-45 min (but it may take only 15-20 min).

While the mixture is heating, prepare a concentrated salt solution by dissolving 50 g of NaCl

in 150 mL of distilled water in a 400-mL beaker. (Prepare this solution immediately so that it

is ready when your soap is ready.) When the saponification reaction is complete, remove the

flask from the heat source using HotHandsTM to hold the hot flask. Pour the reaction mixture

quickly into the saturated salt solution (you may have to scrape the solid into the NaCl

solution using a scoopula). Stir the mixture thoroughly for several minutes; then, collect the


25/25

precipitated soap on a Bchner funnel. Wash the soap twice with 10 mL of ice-cold distilled

water (cool the water with ice, but don't add the ice to the distilled water). After you have

collected and washed the soap, continue to draw air through the soap for several minutes to

help dry it. Save the soap for use in the evaluation section.

Documents

Jeoloji M