University of Bahrain
Department of Chemistry
Edited by: Dr. Ali Hussain
CONTENTS
Experiment 1 Enthalpy Change of Reaction 3
Experiment 2 Molecular Geometries of Covalent
Molecules: Lewis Structure and
Vsepr Theory 7
Experiment 3 Determination of the Dissociation
Constant and Concentration of
a Weak acid 15
Experiment 4 Acid-Base Titration Curves 19
Experiment 5 Hydrolysis of Salts and the Action of
a Buffer Solution 24
Experiment 6 Determination of the Solubility-
Product Constant for a Sparingly
Soluble Salt 34
Experiment 7 Rate of Chemical Reaction:
the Iodination of Acetone 38
Experiment 8 Kinetics of the Hydrolysis of Ethyl Acetate by Sodium Hydroxide by a Conductivity Method 46
Experiment 9 Oxidation-Reduction Titration: Determination of Oxalate 48
Experiment 1:
Enthalpy Change of Reaction
Objective
To measure, using a calorimeter, the enthalpy change accompanying the following displacement reaction:
Zn (s) + Cu2+ (aq) → Cu (s) + Zn2+ (aq)
Discussion
Every chemical change is accompanied by a change in energy, usually in the form of heat. The energy change of a reaction that occurs at constant pressure is termed the heat of reaction or the enthalpy change. The symbol ΔH is used to denote the enthalpy change. If heat is evolved, the reaction is exothermic (ΔH 0); and if heat is absorbed, the reaction is endothermic (ΔH 0). In this experiment, you will calculate the enthalpy change of the above displacement reaction by adding an excess of zinc powder to a measured amount of CuSO4 (aq) and measuring the temperature change over a period of time.
This quantity of heat is measured experimentally by allowing the reaction to take place in a thermally insulated vessel called calorimeter. The heat liberated in the reaction will cause an increase in the temperature of the solution and of the calorimeter. If the calorimeter were perfect, no heat would be radiated to the laboratory. The calorimeter you will use in this experiment is shown in Figure 3.1.
Because we are concerned with the heat of the reaction and because some heat is absorbed by the calorimeter itself, we must know the amount of heat absorbed by the calorimeter. In this experiment, the temperature of the calorimeter and its contents is measured before and after the reaction. The change in enthalpy, ΔH, is equal to the negative product of the heat capacity of the calorimeter and its content times the temperature change, ΔT,:
ΔH = – (heat capacity of the calorimeter + heat capacity of content) ´ ΔT [1]
The heat capacity of the system represents the amount of heat required to raise the temperature of the system 1oC, and ΔT is the difference between the final and initial temperatures: ΔT = Tf – Ti.
Figure 1.1
A simple calorimeter
However, we will assume here that no heat is gained by the calorimeter and no heat is lost to the laboratory. Therefore,
ΔH ≈ – (heat capacity of content) ´ ΔT [2]
For a pure substance of mass m, the expression for ΔH can be written as:
ΔH = – m ´ c ´ ΔT [3]
The quantity c is known as the specific heat; it is defined as the amount of heat required to raise the temperature of one gram of a substance 1oC.
For dilute aqueous solutions, the expression for ΔH can be approximated as:
ΔH = – mwater ´ cwater ´ ΔT [4]
Where mwater is the mass of water and cwater is the specific heat of water which equals to 4.18 J/g×oC. The mass of water can be calculated by considering the density of water equals to 1.00 g/mL.
Procedure
1. Pipette 25.0 mL of cupper (II) sulfate solution into the Styrofoam cup.
2. Weigh about 3 grams of zinc powder in the weighing bottle. Since this is an excess, there is no need to be accurate.
3. Put the thermometer through the hole in the lid, stir and record the temperature to the nearest 0.1oC every half minute for 2.5 minutes.
4. At precisely 3 minutes, add the zinc powder to the cup.
5. Continue stirring and record the temperature till you get to close to room temperature. Tabulate your results on the report sheet.
Plot the temperature (y-axis, oC) against time (x-axis, min). Extrapolate the curve to 3.0 minutes to establish the maximum temperature rise as shown in Figure 2.3.
Calculate the enthalpy change for the quantity of CuSO4 solution used and for one mole of Zn and CuSO4, and write the thermochemical equation for the reaction in the report sheet.
Figure 1.2 Temperature as a function of time for the reaction Zn-CuSO4
Name …………………………………………… ID ………… Sec …..
Experiment 1: Enthalpy Change of Reaction
Report Sheet
Time/min / Temp./
oC / Time/
min / Temp./
oC / Time/
min / Temp./
oC / Time/
min / Temp./
oC / Time/
min / Temp./
oC
0.0 / 6.0
0.5 / 6.5
1.0 / 7.0
1.5 / 7.5
2.0
2.5
3.0
3.5
4.0
4.5
5.0
5.5
1. ΔT determined from your curve ______oC
2. Heat gained by solution (25.0 g ´ 4.18 J/g×oC ´ ΔT oC)
______J
3. The number of moles of CuSO4 in 25.0 mL
of 1.00 M CuSO4. ______mol
4. Heat released per mole of CuSO4 (and Zn) [(2)/(3)] ______kJ/mol
5. Complete the thermochemical equation:
Zn (s) + Cu2+ (aq) → Cu (s) + Zn2+ (aq) ΔH = ______
6. Compare your result with the accepted value of –217 kJ/mol by calculating the
percentage error in your answer:
% Error = │[(experimental value – accepted value)/accepted value]│ ´ 100%.
7. What are the main sources of error in this experiment?
17
Experiment 2:
Molecular Geometries of Covalent Molecules: Lewis Structures and Vsepr Theory
Objective
To become familiar with Lewis structures, the principles of VSEPR theory, and the three-dimensional structures of covalent molecules.
Discussion
Types of Bonding Interactions
Whenever atoms or ions are strongly attached to one another, we say that there is a chemical bond between them. There are three types of chemical bonds: ionic, covalent, and metallic. The term ionic bond refers to electrostatic forces that exist between ions of opposite charge. A covalent bond results from the sharing of electrons between two atoms. The more familiar examples of covalent bonding are found among nonmetallic elements interacting with one another. This experiment illustrates the geometric (three-dimensional) shapes of molecules and ions resulting from covalent bonding among various numbers of elements, and one of the consequences of geometric structures, polarity. Metallic bonds are found in metals such as gold, iron and magnesium. In the metals, each atom is bonded to several neighboring atoms and the bonding electrons are free to move throughout the metal.
Lewis Symbol
The electrons involved in chemical bonding are the valence electrons, those residing in the incomplete outer shell of an atom. The Lewis symbol for an element consists of the chemical abbreviation for the element plus a dot for each valence electron. The dots are placed on the four sides of the atomic abbreviation. The number of valence electrons of any representative element is the same as the group number of the element in the periodic table. The Lewis structure of oxygen, for example, is shown below
The Octet Rule
Atoms often gain, lose, or share electrons so as to achieve the same number of electrons as the noble gas closest to them in the periodic table. According to the octet rule then, atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons. There are many exceptions to the octet rule, but it provides a useful framework for many important concepts of bonding.
Covalent Bonding
Lewis reasoned that atoms might acquire a noble-gas electron configuration by sharing electrons with other atoms to form covalent bonds. The hydrogen molecule, H2, provides the simplest possible example of a covalent bond. The attraction between the nuclei and the electrons cause electron density to concentrate between the nuclei.
Lewis Structures
The formation of a covalent bond can be represented using Lewis symbols as shown below for H2:
+ →
The formation of a bond between two F atoms to give a F2 molecule can be represented in a similar way:
+ →
By sharing the bonding electron pair, each fluorine atom requires eight electrons (an octet) in its valence shell. It thus achieves the noble-gas electron configuration of neon. The structures shown here for H2 and F2 are called Lewis structures (or Lewis electron-dot structures). In writing Lewis structures, we usually show each electron pair shared between atoms as line, to emphasize it is a bond, and the unshared electron pairs as dots. Writing them this way, the Lewis structures for H2
and F2, are shown as follows:
Multiple Bonds
The sharing of a pair of electrons constitutes a single covalent bond, generally referred to simply as a single bond. In many molecules, atoms attain complete octets by sharing more than one pair of electrons between them. When two electron pairs are shred, two lines (representing a double bond) are drawn. A triple bond corresponds to the sharing of three pairs of electrons. Such multiple bonding is found in CO2 and N2.
Drawing Lewis Structures
Lewis structures are useful in understanding the bonding in many compounds and are frequently used when discussing the properties of molecules. To draw Lewis structures, we follow a regular procedure:
1- Sum the valence electrons from all atoms. Use the periodic table to help determine the number of valence electrons on each atom. For an anion, add an electron to the total for each negative charge. For a cation, subtract an electron for each positive charge.
2- Write the symbols for the atoms to show which atoms are attached to which, and connect them with a single bond (a dash, representing two electrons). Chemical formulas are often written in the order in which the atoms are connected to the molecule or ion, as in HCN. When a central atom has a group of other atoms bonded to it, the central atom is usually written first, as in or BF3. In other cases you may need more information before you can draw the Lewis structure.
3- Complete the octets of the atoms bonded to the central atom. (Remember, however, that hydrogen can have only two electrons).
4- Place any leftover electrons on the central atom, even if doing so results in more than an octet.
5- If there are not enough electrons to give the central atom an octet, try multiple bonds. Use one or more of the unshared pairs of electrons on the atoms bonded to the central atom to form double or triple bonds.
Bond Polarity and Dipole Moments
A covalent bond between two different kinds of atoms is usually a polar bond. This is because the two different atoms have different electronegativities and the electrons in the bond are then not shared equally by the two bound atoms. This creates a bond dipole moment, μ, which is a charge separation, Q, over a distance, r:
μ = Qr
The bonding electrons have an increased attraction to the more electronegative atom, thus creating an excess of electron density (δ‾) near it and a deficiency of electron density (δ+) near the less electronegative atom. We symbolize the bond dipole (a vector) with an arrow and a cross, ®, with the point of the arrow representing the negative and the cross the positive end of the dipole. Thus, in the polar covalent molecule H—Cl, because the chlorine is more electronegative than hydrogen, the bond dipole is as illustrated below:
δ+ δ‾
H ® Cl
The bond dipole is a vector (it has a magnitude and direction), and the dipole moments of polyatomic molecules are the vector sums of the individual bond dipoles. Therefore, H2O has a dipole moment (a polar molecule) and CCl4 does not (a nonpolar molecule).
VSEPR Theory
In covalent molecules, atoms are bonded together by sharing pairs of valence-shell electrons. Electron pairs repel one another and try to stay out of each other’s way. The best arrangement of a given number of electron pairs is the one that minimizes the repulsions among them. This simple idea is the basis of valence-shell electron pair repulsion theory, or the VSEPR model. Thus, as illustrated in Table 2.1, two electron pairs are arranged linearly, three pairs are arranged in a trigonal planar fashion, four are arranged tetrahedrally, five are arranged in a trigonal bipyramidal geometry, and six are arranged octahedrally. The shape of a molecule or ion can be related to these five basic arrangements of electron pairs.
Predicting Molecular Geometries
When we draw Lewis structures, we encounter two types of valence-shell electron pairs: bonding pairs, which are shared by atoms in bonds, and nonbonding pairs (or lone pairs) such as in the Lewis structure for NH3. Because there are four electron pairs around the N atom, the electron-pair repulsions will be minimized when the electron pairs point toward the vertices of a tetrahedron (see Table 2.1). The arrangement of electron pairs about the central atom of an ABn molecule is
Table 2.1 Electron-Pair Geometries as a Function of the Number of
Electron Pairs.