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Find a simulation like the one provided about periodic chemistry, and make a description of what it shows.

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Listen to the listed song about the Periodic Table. Write your own song about the various chemical elements.

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Develop a different way in which one could represent the Periodic Table using a variety of physical materials. You may use the following video as inspiration.

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Irish Elements - Periodic Table of Videos

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Think about a particular element, and what makes it interesting. Research that element, and prepare a short description of the chemical properties, including bonding types, that this element undergoes.

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Hydrogen bonding differs from other uses of the word " bond" since it is a force of attraction between a hydrogen atom in one molecule and a small atom of high electronegativity in another molecule. That is, it is an intermolecular force, not an intramolecular force as in the common use of the word bond.

When hydrogen atoms are joined in a polar covalent bond with a small atom of high electronegativity such as O, F or N, the partial positive charge on the hydrogen is highly concentrated because of its small size. If the hydrogen is close to another oxygen, fluorine or nitrogen in another molecule, then there is a force of attraction termed a dipole-dipole interaction. This attraction or "hydrogen bond" can have about 5% to 10% of the strength of a covalent bond.

Hydrogen bonding has a very important effect on the properties of water and ice. Hydrogen bonding is also very important in proteins and nucleic acids and therefore in life processes. The "unzipping" of DNA is a breaking of hydrogen bonds which help hold the two strands of the double helix together.

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Hydrogen bonding differs from other uses of the word " bond" since it is a force of attraction between a hydrogen atom in one molecule and a small atom of high electronegativity in another molecule. That is, it is an intermolecular force, not an intramolecular force as in the common use of the word bond.

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Develop a skit or short scene to act out with at least one other person a notable feature of the Periodic Table. This could include properties of particular elements, features of the table as a whole, the history of the table, etc.

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Find an artistic way to depict the chemical elements. Be as creative as you choose!

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Examine the repeating properties of elements exhibited in the Periodic Table, which can include atomic radii, electron configurations, and/or the charges adopted for main group elements. Provide a description of the aspect which you select.

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The elements in the periodic table are arranged in order of increasing atomic number. All of these elements display several other trends and we can use the periodic law and table formation to predict their chemical, physical, and atomic properties. Understanding these trends is done by analyzing the elements electron configuration; all elements prefer an octet formation and will gain or lose electrons to form that stable configuration.

Atomic Radius

We can never determine the atomic radius of an atom because there is never a zero probability of finding an electron, and thus never a distinct boundary to the atom. All that we can measure is the distance between two nuclei (internuclear distance). A covalent radius is one-half the distance between the nuclei of two identical atoms. An ionic radius is one-half the distance between the nuclei of two ions in an ionic bond. The distance must be apportioned for the smaller cation and larger anion. A metallic radius is one-half the distance between the nuclei of two adjacent atoms in a crystalline structure. The noble gases are left out of the trends in atomic radii because there is great debate over the experimental values of their atomic radii. The SI units for measuring atomic radii are the nanometer (nm) and the picometer (pm). 1 nm = 1 X 10⁻⁹ m; 1 pm = 1 X 10⁻¹² m.
Figure 1: (left) Covalent Radii, (middle) Ionic Radii, (right) Metallic Radii. Provided courtesy of Jessica Thornton (UCD)
To explain this trend, the concept of screening and penetration must be understood. Penetration is commonly known as the distance that an electron is from the nucleus. Screening is defined as the concept of the inner electrons blocking the outer electrons from the nuclear charge. Within this concept we assume that there is no screening between the outer electrons and that the inner electrons shield the outer electrons from the total positive charge of the nucleus. In order to comprehend the extent of screening and penetration within an atom, scientists came up with the effective nuclear charge, \(Z_{eff}\). The equation for calculating the effective nuclear charge is shown below.
Figure 3: Courtesy of Jessica Thornton (UCD)
\[Z_{eff}= Z - S\]
Figure 4: Courtesy of Jessica Thornton (UCD)
In the equation S represents the number of inner electrons that screen the outer electrons. Students can easily find S by using the atomic number of the noble gas that is one period above the element. For example, the S we would use for Chlorine would be 10 (the atomic number of Neon). Z is the total number of electrons in the atom. Since we know that a neutral atom has an identical number of protons and electrons, we can use the atomic number to define Z. For example, Chlorine would have a Z value of 17 (the atomic number of Chlorine). Continuing to use Chlorine as an example, the 10 inner electrons (S) would screen out the positive charge of ten protons. Therefore there would be and effective nuclear charge of 17-10 or +7. The effective nuclear charge shows that the nucleus is pulling the outer electrons with a +7 charge and therefore the outer electrons are pulled closer to the nucleus and the atomic radii is smaller. In summary, the greater the nuclear charge, the greater pull the nucleus has on the outer electrons and the smaller the atomic radii. In contrast, the smaller nuclear charge, the lesser pull the nucleus has on the outer electrons, and the larger atomic radii. Additionally, as the atomic number increases, the effective nuclear charge also increases. Figure 3 depicts the effect that the effective nuclear charge has on atomic radii.
Figure 5: Courtesy of Jessica Thornton (UCD)
Now we are ready to describe the atomic radius trend in the periodic table. The atomic number increases moving left to right across a period and subsequently so does the effective nuclear charge. Therefore, moving left to right across a period the nucleus has a greater pull on the outer electrons and the atomic radii decreases. Moving down a group in the periodic table, the number of filled electron shells increases. In a group, the valence electrons keep the same effective nuclear charge, but now the orbitals are farther from the nucleus. Therefore, the nucleus has less of a pull on the outer electrons and the atomic radii are larger.

We can now use these concept to explain the atomic radius differences of cations and anions. A cation is an atom that has lost one of its outer electrons. Cations have a smaller radius than the atom that they were formed from. With the loss of an electron, the positive nuclear charge out powers the negative charge that the electrons exert. Therefore, the positive nucleus pulls the electrons tighter and the radius is smaller. An anion is an atom that has gained an outer electron. Anions have a greater radius than the atom that they were formed from. The gain of an electron does not alter the nuclear charge, but the addition of an electron causes a decrease in the effective nuclear charge. Therefore, the electrons are held more loosely and the atomic radius is increased.

Ionization Energy (ionization potential)

Expelling an electron from an atom requires enough energy to overcome the magnetic pull of the positive charge of the nucleus. Therefore, ionization energy (I.E. or I) is the energy required to completely remove an electron from a gaseous atom or ion. The Ionization Energy is always positive.
Figure 6: Courtesy of Jessica Thornton (UCD)
The energy required to remove one valence electron is the first ionization energy, the second ionization energy is the energy required to remove a second valence electron, and so on.

\[Na_{(g)} \rightarrow Na^+_{(g)}+ e^-_{(g)}\]

\[Na^+_{(g)} \rightarrow Na^{2+}_{(g)} + e^-\]

Ionization energies increase relative to high effective charge. The highest ionization energies are the noble gases because they all have high effective charge due to their octet formation and require a high amount of energy to destroy that stable configuration. The highest amount of energy required occurs with the elements in the upper right hand corner. Additionally, elements in the left corner have a low ionization energy because losing an electron allows them to have the noble gas configuration. Therefore, it requires less energy to remove one of their valence electrons

Element

1st

2nd

3rd

4th

5th

6th

7th

These are the ionization energies for the period three elements. Notice how Na after in the second I.E, Mg in the third I.E., Al in the fourth I.E., and so on, all have a huge increase in energy compared to the proceeding one. This occurs because the proceeding configuration was in a stable octet formation; therefore it requires a much larger amount of energy to ionize.

Ionization Energies increase going left to right across a period and increase going up a group. As you go up a group, the ionization energy increases, because there are less electron shielding the outer electrons from the pull of the nucleus. Therefore, it requires more energy to out power the nucleus and remove an electron. As we move across the periodic table from left to right, the ionization energy increases , due to the effective nuclear charge increasing. This is because the larger the effective nuclear charge, the stronger the nucleus is holding onto the electron and the more energy it takes to release an electron.
Figure 7: Courtesy of Jessica Thornton (UCD)
The ionization energy is only a general rule. There are some instances when this trend does not prove to be correct. These can typically be explained by their electron configuration. For example, Magnesium has a higher ionization energy than Aluminum. Magnesium has an electron configuration of [Ne]3s2. Magnesium has a high ionization energy because it has a filled 3s orbital and it requires a higher amount of energy to take an electron from the filled orbital.

Electron Affinity

Electron affinity (E.A.) is the energy change that occurs when an electron is added to a gaseous atom. Electron affinity can further be defined as the enthalpy change that results from the addition of an electron to a gaseous atom. It can be either positive or negative value. The greater the negative value, the more stable the anion is.

(Exothermic) The electron affinity is positive

\[X_{(g)} + e^- \rightarrow X^- + \text{Energy}\]

(Endothermic) The electron affinity is negative

\[X_{(g)} + e^- + \text{Energy} \rightarrow X^- \]

It is more difficult to come up with trends that describe the electron affinity. Generally, the elements on the right side of the periodic table will have large negative electron affinity. The electron affinities will become less negative as you go from the top to the bottom of the periodic table. However, Nitrogen, Oxygen, and Fluorine do not follow this trend. The noble gas electron configuration will be close to zero because they will not easily gain electrons.

Figure 8: Courtesy of Jessica Thornton (UCD)

Electronegativity

Electronegativity is the measurement of an atom to compete for electrons in a bond. The higher the electronegativity, the greater its ability to gain electrons in a bond. Electronegativity will be important when we later determine polar and nonpolar molecules. Electronegativity is related with ionization energy and electron affinity. Electrons with low ionization energies have low electronegativities because their nuclei do not exert a strong attractive force on electrons. Elements with high ionization energies have high electronegativities due to the strong pull exerted by the positive nucleus on the negative electrons. Therefore the electronegativity increases from bottom to top and from left to right.
Figure 9: Courtesy of Jesscia Thornton (UCD)

Metallic Character

The metallic character is used to define the chemical properties that metallic elements present. Generally, metals tend to lose electrons to form cations. Nonmetals tend to gain electrons to form anions. They also have a high oxidation potential therefore they are easily oxidized and are strong reducing agents. Metals also form basic oxides; the more basic the oxide, the higher the metallic character.
Figure 10: Courtesy of Jessica Thornton (UCD)
As you move across the table from left to right, the metallic character decreases, because the elements easily accept electrons to fill their valance shells. Therefore, these elements take on the nonmetallic character of forming anions. As you move up the table, the metallic character decreases, due to the greater pull that the nucleus has on the outer electrons. This greater pull makes it harder for the atoms to lose electrons and form cations.

Other Trends

Melting Points: Trends in melting points and molecular mass of binary carbon-halogen compounds and hydrogen halides are due to intermolecular forces. Melting destroys the arrangement of atoms in a solid, therefore the amount of heat necessary for melting to occur depends on the strength of attraction between the atoms. This strength of attraction increases as the number of electrons increase. Increase in electrons increases bonding.

Example: Melting point of HF should be approximately -145 °C based off melting points of HCl, HBr, and HI, but the observed value is -83.6°C.

Heat and electricity conductibility vary regularly across a period. Melting points may increase gradually or reach a peak within a group then reverse direction.

Example: Third period elements Na, Mg, and Al are good conductors of heat and electricity while Si is only a fair conductor and the nonmetals P, S, Cl and Ar are poor conductors.

Redox Potentials

\[K_{(s)} \rightarrow K^+ + e^-\]

\[F_{(s)} + e^- \rightarrow F^-\]

Uses in knowing the Periodic Properties of Elements

1. Predicting greater or smaller atomic size and radial distribution in neutral atoms and ions

2. Measuring and comparing ionization energies

3. Comparing electron affinities and electronegativities

Predicting redox potential

Comparing metallic character with other elements; its ability to form cations

Predicting what reaction may or may not occur due to the trends

Determining greater cell potential (sum of oxidation and reduction potential) between reactions

Completing chemical reactions according to trends

Summary of Periodic Trends

The Periodic Table of Elements categorizes like elements together. Dmitri Mendeleev, a Russian scientist, was the first to create a widely accepted arrangement of the elements in 1869. Mendeleev believed that when the elements are arranged in order of increasing atomic mass, certain sets of properties recur periodically. Although most modern periodic tables are arranged in eighteen groups (columns) of elements, Mendeleev's original periodic table had the elements organized into eight groups and twelve periods (rows).
Figure 11: Courtesy of wikipedia for releasing this image into the public domain
On the periodic table, elements that have similar properties are in the same groups (vertical). From left to right, the atomic number (z) of the elements increases from one period to the next (horizontal). The groups are numbered at the top of each column and the periods on the left next to each row. The main group elements are groups 1,2 and 13 through 18. These groups contain the most naturally abundant elements, and are the most important for life. The elements shaded in light pink in the table above are known as transition metals. The two rows of elements starting at z=58, are sometimes called inner transition metals and have that have been extracted and placed at the bottom of the table, because they would make the table too wide if kept continuous. The 14 elements following lanthanum (z=57) are called lanthanides, and the 14 following actinium (z=89) are called actinides.

Elements in the periodic table can be placed into two broad categories, metals and nonmetals. Most metals are good conductors of heat and electricity, are malleable and ductile, and are moderate to high melting points. In general, nonmetals are nonconductors of heat and electricity, are nonmalleable solids, and many are gases at room temperature. Just as shown in the table above, metals and nonmetals on the periodic table are often separated by a stairstep diagonal line, and several elements near this line are often called metalloids (Si, Ge, As, Sb, Te, and At). Metalloids are elements that look like metals and in some ways behave like metals but also have some nonmetallic properties. The group to the farthest right of the table, shaded orange, is known as the noble gases. Noble gases are treated as a special group of nonmetals.

Useful Relationships from the Periodic Table

The periodic table of elements is useful in determining the charges on simple monoatomic ions. For main-group elements, those categorized in groups 1, 2, and 13-18, form ions they lose the same number of electrons as the corresponding group number to which they fall under. For example, K atoms (group 1) lose one electron to become K+ and Mg atoms (group 2) lose two electrons to form Mg ²⁺. The other main-group elements found in group 13 and higher form more than one possible ion.

The elements in groups 3-12 are called transition elements, or transition metals. Similar to the main-group elements described above, the transition metals form positive ions but due to their capability of forming more than two or more ions of differing charge, a relation between the group number and the charge is non-existent.

Problems

1. Arrange these elements according to decreasing atomic size: Na, C, Sr, Cu, Fr

2. Arrange these elements according to increasing negative E. A.: Ba, F, Si, Ca, O

3. Arrange these elements according to increasing metallic character: Li, S, Ag, Cs, Ge

4. Which reaction do you expect to have the greater cell potential?

A) 2Na(s) + Cl 2 (g) → 2NaCl(s) or 2Cs(s) + Cl 2 (g) → 2RbCl(s)

B) 2Na(s) + Cl 2 (g) → 2NaCl(s) or Be(s) + Cl 2 (g) → BeCl 2 (s)

5. Which equation do you expect to occur?

A) I 2 (s) + 2Br (aq) → Br 2 (l) + 2I(aq)

B) Cl2 (g) + 2I (aq) → I 2 (s) + 2Cl(aq)

Answers

1. Fr, Sr, Cu, Na, C

2. Ba, Ca, Si, O, F

3. Li, S, Ge, Ag, Cs

4. A) Second equation

B) First equation

5. A) Yes

B) No

Problems

*Highlight Answer:_____ to view answers.

1. An element that is an example of a metalloid is (a) S; (b) Zn; (c) Ge; (d) Re; (e) none of these

Answer: (c) Ge

2. In the periodic table, the vertical (up and down) columns are called (a) periods; (b) transitions; (c) families/groups; (d) metalloids; (e) none of these.

Answer: (c) families/groups

3. Why are noble gases inert (nonreactive)?

Answer: Noble gases are inert because they already have a full valence electron shell and have little tendency to gain or lose electrons.

4. What are compounds that contain a halogen called?

Answer: Salts

5. Lanthanides and Actinides are: (a) alkali earth metals; (b) transition metals; (c) metalloids; (d) alkali metals; (e) none of these

Answer: (b) transition metals

References

Cotton, F.A.; Wilkinson, G. (1988). Advanced Inorganic Chemistry (5th Edn). New York: Wiley. ISBN 0-471-84997p. 1385.

Hutchinson, John. "Journal of Chemical Education." Concept Development Studies in Chemistry (2007). Print. Outside Links

Jolly, William L. (1991). Modern Inorganic Chemistry (2nd Edn.). New York: McGraw-Hill. ISBN 0-07-112651-1.

Petrucci, Ralph H. General Chemistry. 9th ed. New Jersey: Pearson Prentice Hall, 2005.

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The elements in the periodic table are arranged in order of increasing atomic number. All of these elements display several other trends and we can use the periodic law and table formation to predict their chemical, physical, and atomic properties.

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Read the following articles on the history of the development of the Periodic Table of Elements. Write a short synopsis of the key features of this table.

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Figure \(\PageIndex{1}\). Dimitri Mendeleev

The Periodic Table is for many the symbol of Chemistry. It is a single image that contains all of the known elements in the universe combined into an easily readable table. There are many patterns present in the table as well. All of the elements seem to fit together and connect to form a readable table and in turn the image of chemistry. The idea of elements first came about in 3000 B.C. The great Greek philosopher Aristotle conceived an idea that everything on earth was made up of these elements. In ancient times, elements like gold and silver were readily accessible, however, the elements that Aristotle chose were Earth, Water, Fire, and Air.

In 1649 the idea of elements took a huge step when Hennig Brand was the first to discover a new element: Phosphorous. Brand was an alchemist in search of the Philosopher's Stone, or an object that would turn any ordinary metal into gold. In his search he tried everything, including distilling human urine. When that experiment was carried out Brand found a glowing white rock. This was the new element he would call Phosphorous. The alchemists and scientists of the enlightenment period added incredible amounts of knowledge to the ideas about elements. In 1869 there were already 63 elements that had been discovered. With each new element that was found, scientists began to realize that there were patterns developing and some started to put the elements into a table.

Scientists like John Newlands and Alexandre-Emile Béguyer de Chancourtois formed their own versions of periodic tables. These versions were relatively simple though and were also somewhat obscure and hard to read. The scientist who brought it all together was Dmitri Mendeleev (1834 to 1907). Mendeleev was a Russian born chemist and the first to publish a modern version of the periodic table. His table ordered the elements by atomic weights (molar masses). When the elements were ordered by their atomic weights, they exhibited similar chemical properties. The table that Mendeleev compiled was so good that he was able to predict elements that were not even known to him at the time. These elements included germanium, gallium, and scandium. There were some pitfalls to the table though. Since not all of the elements had been discovered at the time of Mendeleev's publishing, he left out important elements like the noble gases. After Mendeleev's publishing future scientists contributed to adding in the elements in their proper places. Sir William Ramsay added in the noble gases, and Henry Mosley discovered a way to quantitatively find the atomic number of an element and changed the order around of Mendeleev's table to be organized by atomic number. Finally, in 1945 the Manhattan Project yielded the discovery of many new radioactive elements. Glenn T. Seaborg suggested a change to the table in the form of an addition of the actinide and lanthanide series at the bottom of the table. This idea came with the discovery of Americium and Curium and their unique properties. The change was not accepted at first, but is now included in all periodic tables.

Figure \(\PageIndex{2}\). Periodic Table showing when each element was discovered

Before 1800 (36 elements): dicoveries during and before the Age of Enlightenment.

1800-1849 (+22 elements): impulse from Scientific Revolution and Atomic theory and Industrial Revolution.

1850-1899 (+23 elements): the age of Classifying Elements received an impulse from the Spectrum analysis.

1900-1949 (+13 elements): impulse from the old quantum theory, the Refinements to the periodic table, and quantum mechanics.

1950-1999 (+15 elements): Manhattan_Project and Particle physics issues, for atomic numbers 97 and above

More on the Mendeleev's Periodic Table

The similarities among macroscopic properties within each of the chemical families lead one to expect microscopic similarities as well. Atoms of sodium ought to be similar in some way to atoms of lithium, potassium, and the other alkali metals. This could account for the related chemical reactivities and analogous compounds of these elements.

According to Dalton's atomic theory, different kinds of atoms may be distinguished by their relative masses (atomic weights). Therefore it seems reasonable to expect some correlation between this microscopic property and macroscopic chemical behavior. You can see that such a relationship exist by listing symbols for the first dozen elements in order of increasing relative mass. Obtaining atomic weights, we have

Elements which belong to families we have already discussed are indicated by shading around their symbols. The second, third, and forth elements on the list (He, Li, and Be) are a noble gas, an alkali metal, and an alkaline-earth metal, respectively. Exactly the same sequence is repeated eight elements later (Ne, Na, and Mg), but this time a halogen (F) precedes the noble gas. If a list were made of all elements, we would find the sequence halogen, noble gas, alkali metal, and alkaline-earth metal several more times.

Dmitri Ivanovich Mendeleev proposed the periodic law behind his periodic table compiling. This law states that when the elements are listed in order of increasing atomic weights, their properties vary periodically. That is, similar elements do not have similar atomic weights. Rather, as we go down a list of elements in order of atomic weights, corresponding properties are observed at regular intervals. To emphasize this periodic repetition of similar properties, Mendeleev arranged the symbols and atomic weights of the elements in the table shown below. Each vertical column of this periodic table contains a group or family of related elements. The alkali metals are in group I ( Gruppe I), alkaline earths in group II, chalcogens in group VI, and halogens in group VII. Mendeleev was not quite sure where to put the coinage metals, and so they appear twice. Each time, however, copper, silver, and gold are arranged in a vertical column. The noble gases were discovered nearly a quarter century after Mendeleev's first periodic table was published, but they, too, fit the periodic arrangement. In constructing his table, Mendeleev found that sometimes there were not enough elements to fill all the available spaces in each horizontal row or period. When this was true, he assumed that eventually someone would discover the element or elements needed to complete a period. Mendeleev therefore left blank spaces for undiscovered elements and predicted their properties by averaging the characteristics of other elements in the same group.

As an example of this predictive process, look at the fourth numbered row ( Reihen). Scandium (Sc) was unknown in 1872; so titanium (Ti) followed calcium (Ca) in order of atomic weights. This would have placed titanium below boron (B) in group III, but Mendeleev knew that the most common oxide of titanium, TiO ₂, had a formula similar to an oxide of carbon CO ₂, rather than of boron, B ₂O ₃. Therefore he placed titanium below carbon in group IV. He proposed that an undiscovered element, ekaboron, would eventually be found to fit below boron. (The prefix eka means "below.") Properties predicted for ekaboron are shown in the following table. They agreed remarkably with those measured experimentally for scandium when it was discovered 7 years later. This agreement was convincing evidence that a periodic table is a good way to summarize a great many macroscopic, experimental facts.

Table \(\PageIndex{1}\). Comparison of Mendeleev's Predictions with the Observed Properties of the Element Scandium.

Formula of oxide

Eb ₂O ₃

Sc ₂O ₃

Formula of chloride

EbCl ₃

ScCl ₃

* Mendeleev used the name ''eka''boron because the blank space into which the element should fit was ''below'' boron in his periodic table.

† The modern value of the atomic weight of scandium is 44.96.

The modern periodic table differs in some ways from Mendeleev's original version. It contains more than 40 additional elements, and its rows are longer instead of being squeezed under one another in staggered columns. For example, Mendeleev's fourth and fifth rows are both contained in the fourth period of the modern table. This ends up placing gallium, not scandium underneath boron in the periodic table. This rearrangement is due to theory on the electronic structure of atoms, in particular ideas about orbitals and the relation of electronic configuration to the periodic table. The extremely important idea of vertical groups of related elements is still retained, as are Mendeleev's group numbers. The latter appear as roman numerals at the top of each column in the modern table.

Mendeleev was an extraordinary chemist that was able to compile the greatest chemical instrument of all time. He was not alone in compiling the elements, and many other great chemists contributed too. The idea of elements began over 5,000 years ago and started to finally take shape a mere 200 years ago with Mendeleev's periodic table. Yet, it was not the end of the formation of the periodic table. It has changed over time, and with continue to transform as more and more elements are discovered.

From ChemPRIME: 4.2: The Periodic Table

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Figure \(\PageIndex{1}\). Dimitri Mendeleev The Periodic Table is for many the symbol of Chemistry. It is a single image that contains all of the known elements in the universe combined into an easily readable table. There are many patterns present in the table as well.