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UW Department of Chemistry
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Chem 162
Lab 5: Periodic Trends
Part I: (Prelab) A Computer Study and Introduction to ChemDraw
Part II: Acid-Base Properties of Period 3 and Group 5A Elemental Oxides
Part III: Oxidizing Ability of the Elemental Halides
Introduction
The focus of this lab is learning about trends among the elements in the periods and groups of
the Periodic Table. Some properties can easily be investigated in the general chemistry
laboratory, as you will see in Parts II and III of this experiment. Others are not as easily
measured within the scope of this course, but scientists have performed many of these
measurements and the results have been compiled in various databases, which can then be
used to search for the properties of elements and groups of elements.
In Part I of this lab you will use database software, called KC Discoverer, to look up and
compare several properties among elements on the Periodic Table. By using the program to
plot the data, you will be able to identify the trends within the group, period, or property you
select. The database contains information for many properties beyond those discussed in the
Zumdahl text. You will be asked to comment on the trends (e.g., “increases down a group,”
“decreases across a period,” or “increases then decreases across a period with a maximum
near X,” where X is the symbol of an element or a group or period number. You will also
sketch the graph that you see on the computer screen so that the data will be included in your
report. While this should be a quick sketch, be sure to include the axis labels so it is clear your
TA what data you have plotted. Finally, you will answer some questions and/or explain the
reasons for the trends you have identified.
In Parts II and III of this lab, you will perform experiments that will demonstrate the trends
among several elements on the Periodic Table. In Part II, you will determine whether the
oxides of the period 3 and group 5A elements are acidic or basic. The acid-base properties of
the elemental oxides will be tested using the indicators phenolphthalein and bromothymol blue.
Phenolphthalein is pink in basic solutions and colorless in acidic or neutral solutions.
Bromothymol blue is yellow in acidic solutions and blue in basic or neutral solutions. By using
a combination of both indicators, you will determine whether an oxide is acidic, basic, or
neutral in an aqueous solution.
In Part III, you will determine the relative oxidizing abilities of the elemental halides. Consider
the following reaction:
2KBr (aq) + Cl2 (aq) Æ Br2 (aq) + KCl (aq)
In this reaction Cl2 oxidizes the bromide ions into bromine, Br2. The conclusion must be that
chlorine Cl2 is a stronger oxidizing agent than bromine Br2. In this case we have both
reactions in aqueous solution. If we provide a non-aqueous phase (for example toluene,
C6H5CH3) at the time the reagents are mixed, after the reaction the non-polar products will
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preferentially distribute into the non-aqueous (non-polar) phase (like dissolves like). The above
reaction can therefore be written:
2KBr (aq) + Cl2 (aq) + (toluene) Æ Br2 (toluene) + KCl (aq)
where we indicate the non-aqueous phase by the use of brackets around the toluene. This
separation of Br2 into the toluene phase will result in an observed color change in the toluene
layer - Br2 is brown in color. It is important to clarify that the toluene is not actually changing
color, but rather that the colored Br2 product is moving from the aqueous phase into the nonpolar phase. If there is no “change in color” of the toluene layer, then you will know that the
reaction did not proceed and you still have the same species in solution that you added to the
test tube.
In order to understand the trends you will be investigating in all three parts of this experiment,
let’s review several pieces of information from the Zumdahl text covered in Chem 152 and 162.
The behavior of elements is dependent on the properties of the atoms. Remember that
reactions and bonds involve the valence electrons of the individual atoms or ions, so the
number of electrons, their proximity to the nucleus, and the resulting molecular structure all
play roles in the physical and chemical properties of the elements. Electron configuration,
atomic size, ionization energy, and electronegativity are among the commonly discussed
atomic properties and periodic trends. For this experiment, we will also consider density,
melting point, heat of fusion, polarizability, conductivity (both thermal and electrical), acid-base
properties of elemental oxides, oxidation states, and oxidizing abilities.
Before we review these various properties, recall that a major classification of elements in the
Periodic Table is into metals and nonmetals, with the metalloids (or semimetals) located along
the “staircase” that runs from the top of group 3A to the bottom of group 6A. The fundamental
chemical property of metals is their tendency to give up electrons to form cations; the
fundamental chemical property of nonmetals is their ability to accept electrons and form
anions. The metallic character of elements increases going down a group and decreases
moving from left to right across the Periodic Table. Keep this information in mind as we
discuss the following properties.
Electron configuration. This is the distribution of electrons among the energy levels and
sublevels of an atom. The valence electrons are those in the outermost energy level of an
atom and are the electrons that participate in bonding. The number and type of valence
electrons determine an atom’s chemistry, so atoms with similar electron configurations will
exhibit similar chemical behavior. The differences between one element in a column and the
one below it are systematic – outer electron configurations are similar within a group, but
different within a period. The number of valence electrons for the main group (representative)
elements can easily be determined from the group numbers. For example, carbon, with an
electron configuration of 1s22s22p2, is in group 4A and has 4 valence electrons (2s22p2).
Atomic size. This is often discussed in terms of the atomic radius, which is calculated from
the distance between the nuclei of atoms in a chemical compound. The number of electrons
per orbital, the distance of the electrons from the nucleus, and the shielding of outer electrons
by other electrons in the atom all contribute to the atomic size. In general, the atomic radius
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decreases as you move from left to right across the Periodic Table (Figure 1.). Because of
increasing effective nuclear charge (due to decreased shielding), the valence electrons are
pulled closer to the nucleus, decreasing the size of the atom. In general, the atomic radius
increases as you move from the top of a group to the bottom because of the increasing orbital
sizes in successive principle quantum levels – outer electrons in higher periods are further
from the nucleus.
Figure 1. Atomic size data for the
representative elements.
This is Figure 12.38 in the 6th Edition of the
Zumdahl text. Copyright © Houghton Mifflin
Company – all rights reserved.
The difference in atomic size as you move down group 7A from fluorine to chlorine helps
explain the unique properties of fluorine compared to the other halogens. Fluorine’s small
atomic size draws the valence electrons closer together around a relatively small nucleus. The
large electron-electron repulsions that result contribute to the weakness of the F-F bond. This
is just one example of how atomic size can affect chemical behavior.
Ionization energy. The ionization energy (IE) is the amount of energy needed to remove an
electron from an atom or ion. Since the IE is related to the removal of electrons, electrons that
are tightly bound to the nucleus will require more energy in order to be removed. The first IE is
the energy needed to remove the highest energy electron, which is in the highest energy
orbital. This IE typically increases as you move from left to right across the Periodic Table
(Figure 2.). As you move to the right, the nuclear charge increases with each additional
proton, resulting in less shielding and the valence electrons becoming more tightly bound to
the nucleus. Electrons in lower energy orbitals will partially shield the outer electrons from the
positive charge of the nucleus, but since electrons in the same energy level do not shield each
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other well, the shielding becomes less effective as the nuclear charge steadily increases
across the row. As you move from the top of a group to the bottom, each new principle
quantum number means that the valence electrons will be farther away from the nucleus and,
therefore, less tightly bound, so the first IE decreases as you move down a group.
Figure 2. Ionization energies
for Period 3 elements.
This is Table 12.6 in the 6th Edition
of the Zumdahl text. Copyright ©
Houghton Mifflin Company – all
rights reserved.
The relative magnitude of ionization energies influences the types of bonds atoms will form.
Atoms with low IE easily give up electrons, while those with high IE prefer to share or gain
electrons. This means that, based on the fundamental chemical properties of metals and
nonmetals discussed earlier, metals will have smaller IE values than nonmetals and the
smallest IE values will be found in the lower left corner of the Periodic Table where the most
metallic elements are located. Also, since it is easier to remove an electron that is farther from
the nucleus, IE and atomic size will have opposite trends – smaller atoms in the top right
corner of the Periodic Table will have higher IE values than the larger atoms in the bottom left
corner. Second and third ionization energies have been determined for many atoms. The
energy required for these successive ionizations is affected by the charge of the ion resulting
from the previous ionization as well as the relative orbital energy of the electron being
removed.
Electronegativity. This is the tendency of an atom in a molecule to attract shared electrons to
itself. Electronegativity (EN) values range from 0.7 for francium in the bottom left corner to 4.0
for fluorine in the top right corner (Figure 3.). In general, nonmetals are more EN than metals,
so the trends are that EN increases from left to right across the Periodic Table and decreases
from top to bottom within a group. The increase in EN within a period is due to the higher
effective nuclear charge resulting in shorter distances between the electrons and the nucleus.
The increase in atomic size as you move down a group results in a greater distance between
the electrons and nucleus, so the electrons experience a weaker attraction to the nucleus. In
general, EN has a trend that is the opposite of atomic size.
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Figure 3. Trends in
electronegativity values.
This is Table 13.3 in the 6th
Edition of the Zumdahl text.
Copyright © Houghton Mifflin
Company
–
all
rights
reserved.
Differences in the electronegativities of the atoms in a bond are related to the type of bond that
occurs. For atoms that are identical, as in N2 or Cl2, there is no difference in electronegativity
so the electrons are shared equally and there is no polarity in the bond – this is a covalent
bond. The greater the difference in the electronegativities of atoms bonded together, the
greater is the ionic component of the bond (e.g. the atoms in NaCl have very different EN
values: Na = 0.9 and Cl = 3.0). Polar covalent bonds exist when there is unequal electron
sharing but the differences in EN are not as large as in ionic bonds.
Density. Since density is the ratio of mass to volume, the density of an element can be
related to the ratio of atomic mass to atomic size. Atomic mass increases with the addition of
each proton and electron as the atomic number increases. Many groups show an increase in
density as you move from the top of a group to the bottom, often because the increase in
atomic mass out-paces the increase in atomic size (volume). However, you will find that the
trends in density vary quite a bit for the different groups and periods. Some trends are smooth,
or regular, while others are not.
Heat of fusion. As you learned earlier in this course, the melting point of a substance is
related to the intermolecular forces present in the solid. Metals are solids with strong metallic
bonding in crystal structures. Metalloids, or semimetals, and carbon are solids that experience
covalent bonding in extensive networks. The lighter nonmetals are gases; they are smaller
atoms with less polarizability and experience weak dispersion forces. The heavier nonmetals
are liquids or soft solids; they are larger atoms, more polarizable, and experience stronger
dispersion forces.
Molecular compounds, experiencing non-polar intermolecular interactions, have low melting
and boiling points and require less energy for phase transitions (solid to liquid and liquid to
gas). These compounds are gases, liquids, and low-melting solids at room temperature. Ionic
compounds have very strong intermolecular interactions and will typically have very high
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melting points and boiling points, as well as have large enthalpies of fusion and vaporization.
Network covalent compounds have extremely strong forces, have extremely high melting and
boiling points, and require larger amounts of energy at phase transitions than ionic compounds
do.
Polarizability. Large atoms with more electrons are more polarizable than smaller atoms with
fewer electrons. The greater the polarizability of an atom, the greater the dispersion forces the
atom experiences. If polarizability is related to atomic size and the number of electrons, then
we would expect polarizability to increase down a group as both the size and number of
electrons increase.
Conductivity. Metals are characterized by their abilities to conduct heat and electricity. As
such, we can expect the trends in conductivity to mirror those of metallic character as well as
ionization energies (see earlier discussion of IE).
Acid-Base Behavior. The metallic behavior and electronegativity of an element determine the
type of bonding within its elemental oxide, which is the basis for the acid-base behavior of the
oxide. Metals, elements with low EN values, form basic oxides that react with water to release
OH- ions and acids to form a salt and water. Nonmetals, elements with high EN values, form
acidic oxides that react with water to release H+ ions and bases to form a salt and water.
Amphoteric oxides, which react with both acids and bases, are oxides containing elements
with intermediate EN values (metalloids/semimetals). In general, the acid-base behavior of an
elemental oxide is a good indicator of metallic/non-metallic character.
Oxidizing ability. The oxidation-reduction behavior of an element is its relative ability to gain
or lose electrons when reacting with other elements. Remember from Chem 142 that oxidation
is the loss of electrons and reduction is the gain of electrons. Also, an oxidizing agent is a
reagent that accepts electrons from another reagent, while the reducing agent is the electron
donor. With this in mind, as well as the descriptions of ionization energy (the energy required
to remove an electron from an atom or ion) and electronegativity (the ability of an atom to
attract shared electrons in a bond), we can say that elements with low IE and EN values are
strong reducing agents and those with high IE and EN values are strong oxidizing agents. In
chemical compounds, the oxidation states for the individual atoms indicate the number of
electrons that have shifted away from an atom (positive oxidation number) or towards it
(negative oxidation number). The oxidizing ability of an element describes its tendency to
facilitate the oxidation of another element and is related to the relative attraction of electrons to
its atoms – the lower the EN of an element, the weaker its pull on electrons, so it has a lower
oxidizing ability.
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Helpful information
•
The information summarized here and necessary for understanding the periodic trends
is located in several sections throughout the Zumdahl text (chapters refer to 6th edition):
o Chapter 4: basics of oxidation-reduction reactions
o Chapter 12: atomic theory, electron configurations, orbitals, and periodic trends
o Chapter 13: electronegativity
o Chapter 16: intermolecular forces, types of solids, bonding, structures, and
changes of state
o Chapter 18: properties of main-group/representative elements
•
In addition to learning about several periodic trends by using the KC Discoverer
database, the prelab assignment also includes an introduction to ChemDraw, a software
program useful for creating illustrations of chemical compounds.
•
In Part II, when dissolving the bismuth oxide, Bi2O3, know that it is only slightly soluble
in water. Your entire sample will most likely not dissolve completely, but as long as
some of it dissolves you will be able determine the acidity/basicity of the solution.
Safety Considerations
•
Some reagents used in Part II of this experiment should only be handled by the TAs
because of their high reactivity: sodium oxide (Na2O), tetraphosphorus decoxide
(P4O10), and concentrated nitric acid (HNO3).
•
In Part II, when you are burning the strip of Mg to create MgO, do not look directly at the
bright flame – the ultra-violet light can damage your eyes
•
In Part II, keep the beakers containing the reactions generating SO2 and NO2 in the
hood at all times.
•
The 6M HCl used in the generation of SO2 in Part II is corrosive. If you get any on you,
immediately rinse your skin and any contaminated clothes with plenty of water.
•
In Part III, the liquid bromine (Br2) that is used is corrosive and causes severe burns.
Only perform the reaction in the hood and if you do get any on you, immediately rinse
your skin and any contaminated clothes with plenty of water.
•
All of the waste from this experiment is considered hazardous and should be collected
in the labeled waste bottles in the hood.
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