chemistry - fall 2018 · qualitative chemical analysis usually refers to the systematic procedure...

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Chemistry - FALL 2018

LAB EXPERIMENT 1: REACTIONS OF GROUP I AND GROUP II CATIONS

OBJECTIVES:

To differentiate between electrolytes and nonelectrolytes.

To get the idea of cation classification.

To get introduced to the chemistry laboratory.

To become familiar with test tube reactions.

To learn how to recognize Group I and Group II cations.

INTRODUCTION

Electrolytes and Nonelectrolytes

Electrolyte is a substance capable of conducting electricity when dissolved in water, while

nonelectrolyte dissolved in water does not conduct electricity. For example, a solution of kitchen salt in

water conducts, while the solution of sugar in water does not conduct electricity. It means that salt is

electrolyte, while sugar is not.

In order to conduct electricity, a substance has to be able to separate or dissociate into positively

(cations) and negatively (anions) charged ions in water. Electrolytes can generally be divided into two

categories, namely strong and weak electrolytes. While strong electrolytes give high ion concentration

when dissolved in water, weak electrolytes give low ion concentration. Figure 4.1 is giving an overview of

substance classification into electrolytes and nonelectrolytes.

Figure 1: Electrolytes and nonelectrolytes.

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Precipitation Reactions

If a reaction in which two liquid solutions are mixed is producing a precipitate, it is called a precipitation

reaction. The reaction product is slightly soluble in water. An example is a reaction between lead(II)

nitrate and sodium iodide, in which lead(II) iodide precipitate forms:

Pb(NO3)2(aq) + 2NaI(aq) → PbI2(s) + 2NaNO3(aq)

Precipitate formation is marked by letter s next to the insoluble compound. (In a reaction, s stands for

solid, l for liquid, g for gas and aq for aqueous solutions.)

It is possible to predict the outcome of a precipitation reaction before it is experimentally performed. The

chemists are classifying substance in the following three groups: soluble, slightly soluble and

insoluble. A substance is considered to be soluble if a large amount of it can be dissolved in water. If it is

not the case, we are talking about slightly soluble or insoluble substances. There is a set of simple rules

which are used to predict precipitation reaction outcome when such a reaction is performed at 25⁰C

(Figure 4.2).

Figure 2: Solubility of the most important ions.

Molecular, Ionic and Net Ionic Equations

The following equation:

Pb(NO3)2(aq) + 2NaI(aq) → PbI2(s) + 2NaNO3(aq)

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Represents an example of a molecular equation. It is called so because formulas of compounds are

written as neutral molecules. This type of equations is useful since it shows all compounds undergoing a

chemical reaction.

However, molecular equation is not showing what is happening at the level of particles, atoms, molecules

or ions. Therefore, all compounds should be written as they really appear in aqueous solutions, that is,

dissociated into ions, if they are electrolytes and non-dissociated, if they are nonelectrolytes. Such an

equation is referred to as ionic equation:

Pb2+

(aq) + 2NO3-(aq) + 2Na

+(aq) + 2I

-(aq) → PbI2(s) + 2Na

+(aq) + 2NO3

-(aq)

If we take a look at the left-hand side and the right-hand side of this equation, we can see that Na+

and

NO3- did not change their state during reaction. Therefore, they are called spectator ions. If we wish to

write an equation without these ions, we are emphasizing the chemical change that happened during the

course of the reaction. Such an equation is called net ionic equation:

Pb2+

(aq) + 2I-(aq) → PbI2(s)

Another example is a reaction between silver nitrate and hydrochloric acid.

Molecular equation: AgNO3(aq) + HCl(aq) → AgCl(s) + HNO3(aq)

Ionic equation: Ag+

(aq) + NO3-(aq) + H

+(aq) + Cl

-(aq) → AgCl(s) + H

+(aq) + NO3

-(aq)

Spectator ions: H+ and NO3

-

Net ionic equation: Ag+

(aq) + Cl-(aq) → AgCl(s)

Group I Cations

Group I contains the following cations: Ag+, Pb

2+ and Hg2

2+. As shown in Figure 4.2, all chloride ions are

soluble, with an exception of silver, lead and monovalent mercury chlorides. This is used to distinguish

them from the other groups of cations, which means that a group reagent here is hydrochloric acid (HCl),

which is giving enough free chloride (Cl-) ions in a solution. In order to test Group I cations, we will use

AgNO3 as a source of Ag+ ions. Lead and mercury(I) ions are not used due to their toxicity.

Group II Cations

Group II contains cations whose chlorides are soluble, but whose sulfides are insoluble, namely Hg2+

,

Cu2+

, Bi3+

, Pb2+

, Cd2+

, As3+

, As5+

, Sb3+

, Sb5+

, Sn2+

and Sn4+

. Group reagent is H2S (hydrosulfuric acid).

This group is further divided in subgroups IIa and IIb on the basis of sulfide solubility in (NH4)2S. The first

five ions belong to IIa, while the remaining six belong to IIb group. We will use Cu2+

(from Cu(NO3)2) as a

subgroup IIa representative and Sn2+

(from SnCl2) and As3+

(from As2O3) as a subgroup IIb

representative.

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SAFETY MEASURES

HCl and H2S are acids and have to be handled in chemical hood. NaOH is a strong base and is a

corrosive substance. NH4OH is very hazardous in the case of skin contact (corrosive and irritant), eye

contact (irritant) or ingestion. It may cause damage to mucous membranes and respiratory tract. Handle

carefully in the chemical hood. All chemicals are to be handled with care and according to assistant’s

instructions. Wearing protective clothes (lab coat and gloves) is mandatory throughout whole exercise. In

the case of spills, inform your instructor and clean carefully. In the case of skin or eye contact,

immediately inform your instructor and wash with a plenty of water.

MATERIALS AND METHODS

Chemicals:

Silver nitrate (AgNO3)

Hydrochloric acid (HCl)

Sodium carbonate (Na2CO3)

Hydrosulfuric acid (H2S)

Potassium chromate (K2CrO4)

Sodium hydroxide (NaOH)

Copper(II) nitrate (Cu(NO3)2)

Ammonium hydroxide (NH4OH)

Tin(II) chloride (SnCl2)

Arsenic(III) trioxide (As2O3)

Potassium iodide (KI)

Lab equipment:

Test tubes

Test tube rack

Droppers

Chemical hood

Procedure:

1. Obtain 0.1 M solutions of compounds to be tested from your assistant. During this exercise, you

will be using 1 ml of reagent per reaction. Use separate dropper for every reagent. Note that HCl,

H2S and NH4OH are available in chemical hood.

2. Take AgNO3 solution and put 1 ml in each of six test tubes. Ag+ ion will be tested by the addition

of the following reagents: HCl, Na2CO3, H2S, KI, K2CrO4 and NaOH.

3. Observe reaction outcome and record observations on report sheet attached to this handout.

4. Take Cu(NO3)2 solution and put 1 ml into each of three test tubes. Cu2+

ion will be tested on the

following substances: H2S, NaOH and NH4OH. Record your observations.

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5. Take SnCl2 and place 1 ml of this reagent in two test tubes. Sn2+

ion will be tested using H2S and

NaOH. Record your observations.

6. Place 1 ml of As2O3 into a test tube. As3+

will be tested on H2S. Record your observations.

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REPORT SHEET

PRE-LAB QUESTIONS

1. What is a precipitation reaction?

2. How can the results of a precipitation reaction be predicted?

3. Which cations belong to the Group I? How can they be distinguished from other groups?

4. Which cations belong to the Group II? What is the group reagent and why?

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EXPERIMENT RESULTS

Group I Cations: Reactions of Ag+

Reaction with: Precipitating compound Precipitate appearance

HCl

Na2CO3

H2S

KI

K2CrO4

NaOH

Group IIa Cations: Reactions of Cu2+


H2S

NaOH

NH4OH

Group IIb Cations: Reactions of Sn2+


H2S

NaOH

Group IIb Cations: Reactions of As3+

Reaction with: Resulting precipitate Precipitate appearance

H2S

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POST-LAB QUESTIONS

1. Write molecular, ionic and net ionic equation for the following reactions of Ag+ ion.

Reaction with HCl

Molecular:

Ionic:

Net ionic:

Reaction with Na2CO3

Molecular:

Ionic:

Net ionic:

Reaction with K2CrO4

Molecular:

Ionic:

Net ionic:

2. Write net ionic equation for all reactions of Cu2+

.

Reaction with H2S:

Reaction with NaOH:

Reaction with NH4OH:

3. Write net ionic equation for all reactions of Sn2+

.

Reaction with H2S:

Reaction with NaOH:

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BASIC LABORATORY EQUIPMENT

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LAB EXPERIMENT 2: REACTIONS OF GROUPS III, IV AND V OF CATIONS

OBJECTIVES:

To further understand cation classification.

To get introduced to the chemistry laboratory.

To learn how to recognize cations of Groups III, IV and V.

INTRODUCTION

Qualitative and Quantitative Chemical Analysis

The fact that certain ions will form hardly soluble precipitates of specific color when reacting under

controlled conditions is used to identify those ions. For example, Ag+ cation will form white AgCl

precipitate if HCl, NaCl or any other Cl--containing electrolyte is added to the solution. In other words, we

have used qualitative chemical analyses to identify certain substance (qualitative relates to the property

and type of substance). Qualitative chemical analysis usually refers to the systematic procedure of

proving the presence (or absence) of a substance, usually an ion.

Another important thing in chemistry is to determine the amount of substance present in a sample. In

order to do this, quantitative chemical analysis is used (quantity is the amount). It is important to note

that qualitative analysis has to precede quantitative, as is it first necessary to determine the type of

substance and then its amount.

Flame Tests

Flame test is another type of qualitative analysis. It was noted that the salts of some elements belonging

to the Group I and Group II of the periodic table (not Group I and Group II cations) color the flame in a

specific way. If an inert platinum wire is immersed in a concentrated solution of these salts and then

burned in flame, that flame will get characteristic color. For example, Ba gives yellowish-green color, Ca

gives orange-red and Sr carmine red color. This is another type of test which will be used in today’s lab.

This phenomenon happens as a consequence of the fact that when a particle absorbs light or any other

form of energy, it gets excited. In other words, if a substance is heated, enlightened or gets exposed to

mechanical energy, its own energy level gets higher. An example is (the only) hydrogen electron. In its

unexcited state, it is described by four quantum numbers: n=1, l=0, m=0 and s=1/2. If an H atom absorbs

enough energy to enable passage of its electron to a higher energy level, it will take up one of four

possible orbitals available at the higher level (2s, 2px, 2py or 2pz). If energy input was even higher, it

would pass to energy level 3 (n=3) and take up one of nine possible orbitals. If energy input was high

enough, electron will get free from nuclear attraction force and will leave an atom ionizing it (turning it into

cation).

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If energy absorption does not cause ionization, excitation state is unstable and will not last for long time

period, since an atom always tends to get stabilized by releasing excess energy. This energy is released

through light emission and released light wavelength corresponds to atomic energy levels. If the

wavelength belongs to the visible spectrum, it will be visible to the naked eye. Since excitation energy of

Group I and some Group II elements from the periodic table is relatively low, as these atoms contain one

or two electrons, light emission of these elements belongs to the visible spectrum. Different flame

coloration primarily depends on two factors; the number of valence electrons and total number of

electrons.

Group III Cations

The members of Group III are the following cations: Fe2+

, Fe3+

, Al3+

and Cr3+

. These cations can be

precipitated as hydroxides in the presence of ammonia. Therefore, their group reagent is NH4OH. This

group will be tested in today’s lab using FeCl3 (iron(III) chloride) as a source of Fe3+

cations.

Group IV Cations

The following cations belong to this group: Mn2+

, Ni2+

, Co2+

and Zn2+

. They are usually precipitated as

sulfides using (NH4)2S as a group reagent. ZnCl2 (zinc chloride) will be used as a source of Zn2+

ions to

test Group IV cations reactions.

Group V Cations

Group V cations are Ba2+

, Sr2+

and Ca2+

. These ions form soluble sulfides and insoluble carbonates in

neutral and alkaline solutions. Group reagent is (NH4)2CO3. We will use BaCl2 (barium chloride) as a

source of Ba2+

ions to test the reactions of Group V cations. Additionally, SrCO3 (strontium carbonate)

and CaCO3 (calcium carbonate) will be used for flame test.

Group VI Cations

Group VI cations are: Mg2+

, K+, Na

+ and NH4

+. These cations cannot be precipitated with any group

reagent mentioned for the previous five groups.

SAFETY MEASURES

H2SO4 is a strong acid and has to be handled in chemical hood. NaOH is a strong base and is a corrosive

substance. NH4OH and (NH4)2S are very hazardous in the case of skin contact (corrosive and irritant),

eye contact (irritant) or ingestion. They may cause damage to mucous membranes and respiratory tract.

Handle carefully in the chemical hood. All chemicals are to be handled with care and according to

assistant’s instructions. Wearing protective clothes (lab coat and gloves) is mandatory throughout whole

exercise. In the case of spills, inform your instructor and clean carefully. In the case of skin or eye contact,

immediately inform your instructor and wash with a plenty of water.

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Chemicals:

Iron(III) chloride – FeCl3

Ammonium hydroxide – HN4OH

Potassium ferrocyanide – K4[Fe(CN)6]

Sodium hydroxide - NaOH

Ammonium sulfide - (NH4)2S

Sodium hydrogen phosphate – Na2HPO4

Zinc chloride – ZnCl2

Barium chloride – BaCl2

Ammonium carbonate – (NH4)2CO3

Sulfuric acid – H2SO4

Potassium chromate – K2CrO4

Strontium carbonate – SrCO3

Calcium carbonate – CaCO3

Lab equipment:

Test tubes

Test tube rack

Droppers

Chemical hood

Watch glass

Inoculating loop

Lab burner

Procedure:

Obtain 0.1 M solutions of compounds to be tested from your assistant. During this exercise, you

will be using 1 ml of reagent per reaction. Use separate dropper for every reagent. Note that

NH4OH, (NH4)2S and H2SO4 are available in chemical hood.

Take FeCl3 solution and put 1 ml in each of five test tubes. Fe3+

ion will be tested by the addition

of the following reagents: NH4OH, K4[Fe(CN)6], NaOH, (NH4)2S and Na2HPO4.

Observe the reaction outcome and record observations on the report sheet attached to this

handout.

Take ZnCl2 solution and put 1 ml into each of three test tubes. Zn2+

ion will be tested on the

following substances: NaOH, (NH4)2S and K4[Fe(CN)6]. Record your observations.

Take BaCl2 and place 1 ml of this reagent in four test tubes. Ba2+

ion will be tested using

(NH4)2CO3, H2SO4, Na2HPO4 and K2CrO4. Record your observations.

Perform flame tests for Ba2+

, Sr2+

and Ca2+

cations according to assistant’s instructions. Note the

color of each flame on the report sheet.

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REPORT SHEET

PRE-LAB QUESTIONS

1. Differentiate between qualitative and quantitative chemical analysis.

2. What is the goal of flame test?

3. What is the common property of all Group V cations?

4. What is the only non-metal cation mentioned in labs so far?

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EXPERIMENT RESULTS

Group III Cations: Reactions of Fe3+


NH4OH

K4[Fe(CN)6]

NaOH

(NH4)2S

Na2HPO4

Group IV Cations: Reactions of Zn2+


NaOH

(NH4)2S

K4[Fe(CN)6]

Group V Cations: Reactions of Ba2+


(NH4)2CO3

H2SO4

Na2HPO4

K2CrO4

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Group V cations: Flame tests

Cation Flame color

Ba2+

Sr2+

Ca2+

POST-LAB QUESTIONS

1. Write molecular, ionic and net ionic equation for the following reactions of Fe3+

ion.

Reaction with NH4OH

Molecular:

Ionic:

Net ionic:

Reaction with K4[Fe(CN)6]

Molecular:

Ionic:

Net ionic:

Reaction with Na2HPO4

Molecular:

Ionic:

Net ionic:

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LAB EXPERIMENT 3: pH VALUE (CALCULATIONS AND EXPERIMENTAL DETERMINATION)

OBJECTIVES:

To get familiar with the concept of pH in chemistry.

To learn how to calculate pH of a solution.

To determine pH experimentally using indicator strips and pH meter.

INTRODUCTION

Acids and Bases

When talking about the concept of pH value, we usually think of strong acids and bases with toxic and

corrosive nature. However, the majority of acids and bases are weak. For example, Vitamin B and aspirin

are weak acids, while caffeine, nicotine and indigo are weak bases.

To decide if an acid or a base is strong or weak, one should consider how well it dissociates (ionizes) in

water. For example, any acid HA will react with water in the following way:

HA(aq) + H2O(l) → H3O+

(aq) + A-(aq)

If an acid is a stronger protein donor than hydronium ion (H3O+), it will give its proton to water and

dissociate completely. Such as acid is considered to be strong. On the other hand, weak acids are

weaker proton donors than hydronium ion, which means that a reaction between a weak acid and a water

molecule is reversible; acid HA will donate its proton to water to form H3O+, but H3O

+ will donate its proton

to A- to form HA at the same time.

If a substance HA is a weaker proton donor than H2O, it will receive a proton from water, as in the

following reaction:

HA(aq) + H2O(l) → H2A+

(aq) + OH-(aq)

Such compounds are termed strong bases. These bases are removing a proton from water and forming

OH- (hydroxide) ion, therefore completely ionizing in water.

The following are the general rules for classification of acids and bases:

1. Any protein donor stronger than H3O+ is a strong acid in water.

2. If a substance is a protein donor of strength between H3O+ and H2O, it is a weak acid.

3. Any protein acceptor stronger than H2O (weaker protein donor) is a weak base.

4. Any protein acceptor stronger than OH- is a strong base.

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Therefore, the concentration of H3O+ and OH

- ions is a measure of how strong acids and bases are.

pH Value

Water is an amphoteric substance, that is, it can act both as an acid and a base and is capable of self-

ionization, which is shown in the following reaction:

H2O(l) + H2O(l) → H3O+

(aq) + OH-(aq)

At 25⁰C, only two out of 109 water molecules self-ionize. The product of H3O

+ and OH

- is the ionic

product of water, Kw, which is 1.0 x 10-14

at 25⁰C:

[H3O+][OH

-] = 1.0 x 10

-14

[H3O+] = [OH

-] = 1.0 x 10

-7

If an acid is added to this solution, hydronium ion concentration goes above 1.0 x 10-7

, while hydroxide

ion concentration decreases, but never reaches zero. If base is added, OH- concentration goes above 1.0

x 10-7

, while H3O+ concentration decreases, but never reaches zero.

A good way of presenting such small concentrations of these two ions is pH value scale. pH is a

negative logarithm of hydronium ion concentration and is a dimensionless unit:

pH = -log[H3O+]

On the other hand, H3O+ concentration can be calculated if pH of a solution is known:

[H3O+] = 10

-pH mol/l

pOH is a negative logarithm of hydroxide ion concentration:

pOH = -log[OH-]

[OH-] = 10

-pOH mol/l

pKw is a negative logarithm of ionic product of water:

pKw = -logKw = -log(1.0 x 10-14

) = 14

pH + pOH = pKw

pH + pOH = 14

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Figure 3: pH value chart.

pH Calculations (Examples)

1. pH of a solution is 6.6. Calculate the concentration of H+ and OH

- ions. What is pOH of that

solution?

[H+] = 10

-pH

[H+] = 10

-6.6 = 2.5 x 10

-7 M

[OH-] = Kw / [H

+]

[OH-] = 1.0 x 10

-14 / 2.5 x 10

-7 = 0.4 x 10

-7 = 4 x 10

-8 M

pOH = 14 – pH = 14 – 6.6 = 7.4

2. What is the pH and OH- concentration in a solution of pOH = 1?

pH = 14 – pOH = 14 – 1 = 13

[OH-] = 10

-pOH = 10

-1 = 0.1 M

Measuring pH Value

pH value can be measured in different ways. Today, we will use two methods:

1. Universal pH indicator strips are using different color combinations to show pH of a solution.

2. pH meters are giving the most accurate measurements. They are using electric potential in a

glass electrode whose potential depends on hydronium ion concentration in a solution. The result

is shown on a display. When used, pH meter should first be calibrated with one or two buffers of

known pH value.

SAFETY MEASURES

All acids have characteristic unpleasant odors which is why they should be prepared in chemical hood.

Strong acids and bases are corrosive substances. Although they are available as diluted solutions, it is

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obligatory to wear lab coat and protective gloves. Any spills should be immediately reported and carefully

cleaned. In case of skin or eye contact, inform you assistant immediately.


Chemicals:

100 mM HCl

1 mM HCl

100 mM CH3COOH

1 mM CH3COOH

100 mM NaOH

1 mM NaOH

100 mM Na2HPO4

1 mM Na2HPO4

100 mM KCl

1 mM KCl

Lab equipment:

Erlenmeyer flasks

pH indicator strips

pH meter.

Procedure:

Necessary solutions of HCl (hydrochloric acid), CH3COOH (ethanoic or acetic acid), NaOH

(sodium hydroxide), Na2HPO4 (sodium hydrogen phosphate) and KCl (potassium chloride) are

available in the lab.

Measure pH of every solution using pH indicator strips. Record values on the report sheet.

Measure pH of every solution using pH meter. Record values on the report sheet.

Obtain household products available in the laboratory. Measure their pH using indicator strips.

Record pH values and acid/base nature of those products on the report sheet.

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REPORT SHEET

PRE-LAB QUESTIONS

1. Differentiate between an acid and a base.

2. Calculate the pH and pOH of 1.2 x 10-3

M HCl solution.

3. Calculate pH, pOH and [OH-] of 0.1 M HNO3 solution.

4. If a solution X has pH = 5, which of the following is true:

a. Solution X is neutral.

b. H3O+ ion concentration is higher than OH

- concentration.

c. OH- ion concentration is higher than H3O

+ concentration.

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EXPERIMENT RESULTS

Table 6.1: pH values of acid, base and salt solutions.

Solution pH by indicator strips pH by pH meter Type of solution

100 mM HCl

1 mM HCl

100 mM CH3COOH

1 mM CH3COOH

100 mM NaOH

1 mM NaOH

100 mM Na2HPO4

1 mM Na2HPO4

100 mM KCl

1 mM KCl

Table 6.2: pH values of common household products.

Product pH by indicator strips Acid/base/neutral

Apple juice

Liquid soap

Clothes detergent

Table salt

Sugar

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LAB EXPERIMENT 4: DETERMINING HCl CONCENTRATION (ACID-BASE TITRATION)

OBJECTIVES:

To get introduced to acid and base definitions by Arrhenius, Bronsted-Lowry and Lewis.

To understand what is acid-base titration.

To understand the role of indicators.

To grasp the concept of equivalence point and its position on the graph depending on the type of

acid and base used in titration.

To learn how to calculate the unknown concentration.

To perform titration.

INTRODUCTION

Definitions of Acids and Bases: Neutralization Reaction

Acids and bases are compounds which are usually studied and defined together. The first definition of

acids and bases was given by Swedish scientist Arrhenius, who explained that acids are compounds

that have extra H+ ions when dissolved in water, while bases donate extra OH

- ions in water, for

example:

HCl(aq) → H+

(aq) + Cl-(aq)

NaOH(aq) →Na+

(aq) + OH-(aq)

However, this definition gave an explanation of aqueous solutions only and did not explain the fact that H+

ion never exists alone, but is always paired with water molecule to give H3O+. Therefore, a new definition

of acids and bases was necessary.

Bronsted-Lowry definition states that both acids and bases are capable of reaction with water, whereby

acids donate a proton, while bases accept a proton:

HCl(aq) + H2O(l) → H3O+

(aq) – Cl-(aq)

NH3(aq) + H2O(l) → NH4+

(aq) + OH-(aq)

Due to a completely new definition of a base, many species that were not defined as acids/bases by

Arrhenius, were now included in these two classes. Acids and bases are connected to each other and

always act as a pair:

HCl(aq) H+

(aq) + Cl-(aq)

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acid proton + base

In this example, HCl/Cl- is an acid/base pair which is exchanging a proton and we call them conjugated

acid-base pair, where conjugated means common.

NaOH(aq) + H+

(aq) Na(H2O)+

(aq)

base + proton acid

Conjugated acid-base pair in this reaction is Na(H2O)+/NaOH, or simply Na

+/NaOH.

Neutralization is a reaction between two conjugated acid-base pairs. In a reaction between HCl and

NaOH, conjugated pairs are HCl/Cl- as a pair 1 and Na

+/NaOH as a pair 2. By consensus, acid is always

written first in an acid-base conjugated pair.

HCl(aq) + NaOH(aq) → [Na(H2O)]+ + Cl

-(aq)

acid 1 + base 1 → acid 2 + base 2

Acids and bases that can exchange more than one proton are said to be polyprotic, like H3PO4 and

H2CO3 (acids), or CO32-

and SO42-

(bases). Each ionization degree is making a new acid-base pair. For

example, carbonic acid is making two pairs: H2CO3/HCO3- and HCO3

-/CO3

2-, while phosphoric acid is

making three: H3PO4/H2PO4-, H2PO4

-/HPO4

2-, HPO4

2-/PO4

3-. Substances or ions that can act as acids and

bases are called amphoteric. From the previous examples, amphoteric ions are HCO3-, H2PO4

- and

HPO42-

.

Lewis gave an even newer definition of acids and bases in 1923 by expanding it to non-aqueous

solutions. He stated that a substance which is donating an electron pair is a base, while the

substance that is accepting an electron pair is an acid. Therefore, all compounds which have a free

electron pair and can give to an empty orbital of H+ are said to be bases (:N, :S, :Cl). In this way, many

metal ions and some oxides were recognized either as acids or bases (SOCl2, AlCl3, BF3).

:NH3 + H+

NH4+

Volumetric Analysis: Vocabulary

Volumetric analysis is a quantitative chemical method which is measuring volume in order to determine

the amount of substance. Volumetry is an analysis of solutions, mainly aqueous, in which the volume of a

solution of a known concentration necessary for a complete reaction of an analyzed substance is

measured. The basic vocabulary of volumetric analysis includes the following:

1. Standard solution – a solution of known concentration used for titration

2. Titration – careful addition of a standard solution from a burette to the analyzed substance until,

according to analyst’s opinion, the reaction is complete

3. Titrant – solution used to perform titration

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4. Equivalence point – a point in titration in which the amount of added titrant is chemically

equivalent to the amount of analyzed substance in a solution. If NaCl is reacting with AgNO3,

then:

NaCl(aq) + AgNO3(aq) → AgCl(s) + NaNO3(aq)

Equivalence point is reached when 1 mol of NaCl reacts with 1 mol of AgNO3 (that is, 1 mol of Cl- reacts

with 1 mol of Ag+). In a reaction of H2SO4 and NaOH, equivalence point is reached when 2 mol of a base

react with 1 mol of an acid:

2NaOH(aq) + H2SO4(aq) → Na2SO4 + 2H2O

5. Titration endpoint – equivalence point is a theoretical point that cannot be experimentally

determined, but can be detected by a physical change, e.g. change of color or electric signal.

When detected, this point is called titration endpoint. The difference between equivalence point

and endpoint depends on analyst’s capability to detect the change and it should be as small as

possible to decrease the possibility of titration error.

6. Indicator – a substance which is added to the analyzed solution to detect endpoint; chosen

indicator should be able to make the difference between equivalence point and endpoint as small

as possible.

7. Primary standard – a compound or a solution of high purity and stability which is used as a

reference material in all volumetric analyses. The accuracy of the method depends on these

solutions (compounds) and their properties.

Volumetric Calculations

Titration can be used to determine concentration of unknown solution due to the basic premise that the

number of moles of a base added equals the number of moles of an acid at the equivalence point (where

titration ends). In other words, the volume of a base used to neutralize an acid can be used to precisely

determine concentration of an acid and vice versa. This is valid only if an acid and a base taking part in

the reaction have the ability to exchange the same number of protons.

n1 = n2; c1V1 = c2V2

1. How many milliliters of 0.1 M NaOH solution is necessary to neutralize 300 ml of 0.1 M HCl

solution?

Firstly, write the reaction: HCl + NaOH → NaCl + H2O

c1V1 = c2V2

0.1 M * V2 = 0.1 M * 300 ml

V2 = (0.1 M * 300 ml) / 0.1 M

V2 = 300 ml

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2. 50 ml of HCl has been neutralized with 29.71 ml of 0.01963 M Ba(OH)2 solution. What is HCl

concentration?

Reaction: 2HCl + Ba(OH)2 → BaCl2 + 2H2O

From the equation, it is obvious that 2 mol of HCl are reacting with 1 mol of Ba(OH)2. Therefore:

=

; 2Ba(OH)2 = 1HCl; 2nBa(OH)2 = 1nHCl

2 * cBa(OH)2 * VBa(OH)2 = cHCl * VHCl

cHCl = (2 * 0.01963 M * 29.71 ml) / 50 ml

cHCl = 0.02333 M

Volumetric Analysis: Titration Curves

When performing an acid-base titration, there are four possible combinations: strong acid-strong base,

strong acid-weak base, weak acid-strong base and weak acid-weak base (Figure 8.1). Each of these

combinations will give a specific titration curve which will differ in the position of an equivalence point

(equivalence point is found at the halfway of a vertical line which shows a dramatic increase in pH value).

From the shape of the titration curve, the type of acid and base used can be easily determined. Note that

each of four curves below has only one equivalence point.

Figure 8.1. Four possible titration curves depending on the nature of acid and base reacting.

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However, polyprotic acids are characterized by titration curves with more than one equivalence point. The

number of equivalence points in a polyprotic acid is equal to the number of protons it can release.

Therefore, H2SO4 with have two, while H3PO4 will have three equivalence points (Figure 8.2).

Figure 8.2. Top: Titration curve of a diprotic acid. Note that 200% (2 mol) of a base are necessary to neutralize 1 mol

of an acid. Bottom: Titration curve of a triprotic acid. 3 mol of a base are necessary to neutralize 1 mol of an acid.

Experimental Determination of HCl Concentration

HCl concentration is determined by titrating a known volume of an acid with a standard NaOH solution,

which is serving as a primary standard. Given below are chemical reactions that occur during the titration.

NaOH + HCl → NaCl + H2O

Na+

(aq) + OH-(aq) + H

+(aq) + Cl

-(aq) → Na

+(aq) + Cl

-(aq) + H2O(l)

OH-(aq) + H

+(aq) → H2O(l)

In other words, strong acid-strong base titration is a reaction between protons and hydroxide ions. In the

reaction, 1 mol of HCl reacts with 1 mol of NaOH, which is an important information to calculate HCl

concentration.

Since both solutions are colorless, it is necessary to add an indicator which will show equivalence point,

that is, the end of neutralization reaction. The indicators of choice in our lab are phenolphthalein and

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methyl orange. Phenolphthalein is colorless in acidic and pink in alkaline solution. When the

concentration of H+ and OH

- ions is equal, the change in color will happen. One more drop of NaOH will

cause indicator color to turn into light purple. If this happens, it means that titration is already behind its

equivalence point and that the result is not accurate. Another indicator, methyl orange, works on the

exactly same principle, but has different color; it is red in acidic, orange in neutral and yellow in basic

solutions (Figure 8.3).

Figure 8.3. Left: Change of color for litmus, methyl orange and phenolphthalein. Right: Both indicators are suitable

for strong acid-strong base titration.

SAFETY MEASURES

HCl and NaOH are present as diluted solutions. However, you are obliged to wear a lab coat and

protective gloves. Report any spills to your assistant immediately.


Chemicals:

Standard NaOH solution of known concentration

HCl solutions 1 and 2 of unknown concentration

Phenolphthalein

Methyl orange

Distilled water.

Lab equipment:

Erlenmeyer flasks

20 ml glass pipette

Burette

Dropper

Wash bottle

Ring stand.

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Procedure:

1. Pipette exactly 20 ml of HCl solution 1 of unknown concentration. Transfer it to 250 ml

Erlenmeyer flask (wide neck).

2. Carefully pour NaOH solution in the burette. Use funnel if necessary. Put a beaker under the

burette. Slowly open the pipe and release NaOH to bring its level in burette to 0 ml. Meniscus has

to be exactly on 0 mark.

3. In HCl-containing Erlenmeyer flask, add distilled water to make liquid level 1 to 1.5 cm high.

Added water will not affect titration results, since the amount (number of moles) of acid remains

the same.

4. Add 3 drops of phenolphthalein and swirl the flask to mix.

5. Place the flask under the burette and start titration. Put a piece of white paper under the flask to

notice the change in color easier.

6. Perform titration by continuous flask mixing. When you get close to titration endpoint, add titrant

drop by drop. After each addition, observe the color of the solution.

7. Titration endpoint is reached when the solution gets stable pink color that lasts for longer than 20

seconds. Record titrant volume used on the report sheet.

8. Repeat the same titration two more times.

9. Volume of titrant is the average value of three titrations. Calculate the concentration of solution 1.

10. Obtain HCl solution 2 from your assistant. Perform three titrations using methyl orange indicator.

Calculate the concentration of solution 2.

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REPORT SHEET

PRE-LAB QUESTIONS

1. Can you determine titration endpoint using pH meter instead of color indicator? How?

2. Why is equivalence point in strong acid-weak base titration below pH 7?

3. How many equivalence points do you expect H2CO3 titration curve to have? Why?

4. Show all conjugated acid-base pairs of H2SO4. Which ion(s) is/are amphoteric?

5. How many milliliters of 0.116 M H2SO4 will be needed to titrate 25.0 ml of 0.0084 M Ba(OH)2 to

the equivalence point? Write chemical equation and perform calculations.

6. 27.0 ml of 0.310 M NaOH is titrated with 0.740 M H2SO4. How many ml of H2SO4 are needed to

reach the end point? Show the chemical reaction.

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EXPERIMENT RESULTS

Table 8.1: Titration of HCl solution 1 with NaOH.

Titration Volume of NaOH used (ml)

1

2

3

Average value

VNaOH: _____________

cNaOH: _____________

VHCl: _______________

cHCl: _______________

Show how you calculated HCl concentration below.

Table 8.2: Titration of HCl solution 2 with NaOH.

Titration Volume of NaOH used (ml)

1

2

3

Average value

VNaOH: _____________

cNaOH: _____________

VHCl: _______________

cHCl: _______________

Show how you calculated HCl concentration below.

chemistry - fall 2018 · qualitative chemical analysis usually refers to the systematic procedure...

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