chapter8

=Chapter 8 Periodic Relationships Among the Elements= The periodic table serves as a great reference, and its use can give both specific information, but also predict general trends.



Topics:
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 * Development of the Periodic Table
 * Periodic Classification of the elements
 * Periodic Variation in physical properties
 * Ionization Energy
 * Electron Affinity
 * Variation in Chemical Properties of the representative elements

Development of the Periodic Table
Since the 1800s, chemists have been organizing the elements based on their properties. Although this system was developed long before the discovery of subatomic particles and nuclear structure, many salient ideas were put forth. These systems of organization continued to evolve until they reached the modern state of the periodic table.

In 1864, John Newlands noticed the first major trend of the elements. When ordered my atomic mass, the elements repeated their characteristics every 8 elements. This became known as the octaves law. However, when people realized that the law failed to work past Calcium (look at a periodic table and guess why), he lost standing in science and people genreally ignored him. In 1869, Dmitri Mendeleev and Lothar Meyer (but history has mostly forgotten him) independently created a new way to organize the elements. This method organized the elements in order of their atomic mass, and then grouped them vertically by chemical properties. Below is one of his first renditions of the table. Take a moment to see how similar it is to the [|modern table].



Fig 8.1.1 Credit:[]

Despite its obvious success, the periodic table had a few major problems. One of them was that sometimes the atomic masses were not consistent with the properties of the grouping. Later, Mosely discovered the concept of atomic number (the number of protons) to be a more accurate way of organizing the table. Thus, the modern periodic table was born, and throughout history more and more elements have been added as they have been discovered.

Today, the periodic table serves not only as a excellent reference tool for almost all of the important information about an element, it can also predict information about undiscovered elements with great accuracy (don't worry, you don't need to do this).

Viola, the modern table...

Fig 8.1.2 Credit:[|http://zebu.uoregon.edu/~imamura/122/images/periodic_table.gif]

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8.2 Periodic Classification of the Elements
The periodic table can help to classify and group the elements. As such, several different classifications have been created. Broadly, there are representative elements, transition metals, noble gases, lanthanides, and actinides.

The elements are organized according to their outer shell electron configuration. These electrons are called valence electrons, and determine the chemical properties of the element. In the table, elements are arranged vertically by these configurations. As such, in each column (group), **the elements all have similar properties**. This is important. For example, this allows you to generalize your knowledge that Na is reactive in water to include all of group 1 (K, Rb, Cs, Fr). The rows are called periods and show which energy level is being filled.

Quick Note: The groups can be labeled 1-18, or 1A-8A and 1B-8B. These systems can be used almost interchangeably, and you should become familiar with both. Just check the top of the chart to figure out how they correspond. (as a weird note, Europe reverses the A's and B's in their designation, but this isn't important for you.)

Representative elements (aka main group elements) are those in groups 1A through 7A, whose outer s or p shell is incompletely filled. This excludes group 3 through 12, as their s shells are completely filled, and they have no p shell. It also excludes group 18 as all s and p levels are filled. A few representative elements include H, C, O, Li, Pb, Na, and Ba.

The transition metals are in groups 3 through 12 (1B-8B). They all have incomplete d sublevels, and therefore readily form cations by giving up these electrons. Ti should be noted that group 12 (1B, Zn, Cd, Hg) are all exceptions to this rule. Some transition metals include Sc, Ti, Pt, and W.

The noble gases are in group 18 (8A) and have completely filled orbitals. As such they are very unreactive. Some examples include He, Ne, and Xe.

The Lanthanides and Actinides are sometimes called f-block transition metals, as they have incomplete f orbitals, as opposed to d orbitals. They are often removed from the main body of the chart and placed below, but can be inserted in-between the alkali earth metals and the transition metals. Some examples include U, Sm, Dy, and Md. They are mostly rare and radioactive, so are not very common in this course.

In equations you will oftentimes have to represent the elements in their base, uncombined form. You may often use the empirical formula, the chemical symbol with a fancy name. For example, the empirical formula for Carbon is C. ( notably, carbon exists in a complex, nearly infinite network, but this makes it easier, for everyone). There are a few exceptions (as always) to this rule. The exception comes in the form of the diatomic molecules. They are (almost) always in the form of X2 (that should be a subscript). The diatomics include H, O, Br, F, I, N, and Cl. You can remember them with the device "HOBrFINCl". Know this, or you will be screwed.
 * Free elements in Equations**

Representative elements tend to form themselves into noble gas electron configurations when they form cations or anions. Thus, they become stable, in the form ns²np
 * Electron Configurations of Cations and Anions**

Meanwhile, transition metals are weird. Rather than first losing their d level electrons, they lose their s level electrons, as the d level tends to be more stable. However, when forming cations from the transition metals it is good practice to check a table, as they can form multiple different cations and be generally weird.

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8.3 Periodic Variation in Physical Properties
If position in the periodic table is related to electron configuration, it makes some sense that it is also related to physical and chemical properties. Physical properties are a little bit easier, so we will start there.

The charge on the protons in the nucleus holds the electrons in orbit around it. Therefore, the more protons, the more attractive force is present. (of course, it takes more force to hold in electrons that are further away from the nucleus, think back to physics). However, the inner electrons help to shield the outer electrons from the attractive forces of the nucleus. This is known as shielding, and the inner electrons that help to offset the charge of the nucleus are known as shielding electrons. Furthermore, shielding electrons produce repulsive forces on the other electrons that also helps to further push them from the nucleus. All in all, shielding creates the idea of effective nuclear charge. Effective nuclear charge is the charge that the nucleus seems to exert on the surrounding electrons once shielding has been taken into account. It can be put mathematically in the following form. (Effective Nuclear Charge)=(Actual Nuclear Charge)-(Shielding Constant). The shielding constant is less than the actual nuclear charge, but greater than 0.
 * Effective Nuclear Charge**

The concept of effective charge changes the amount of energy required to remove successive electron (more on this soon). For example, in He it takes only 2373 kJ of energy to remove the first electron from a mole of atoms. But, as there is no longer a shielding effect, the next mole takes a whopping 5251 kJ.

This effect can also move between shells. For example, the electrons in the 1s shell and the electrons in the 2s shell of Be shield the outermost electron. However, only the electrons in the 1s shell shield the 1 s electrons. Moreover, electrons in inner shells are more effective at shielding outer shells than they are at shielding their own shells. (and outer electrons do almost nothing to shield inner electrons). This explains why it is so much easier to remove outer electrons than inner electron.



Many physical properties such as density and melting point are determined by atomic radius. Ergo, it is important to understand atomic radius. Because measuring the size of an undefined electron cloud is difficult, scientists have determined that the atomic radius is 1/2 of the distance between two nuclei.
 * Atomic Radius**

There are two main factors in determining atomic radius. The first is the number of shells. As each successive shell is further away from the nucleus, the further down the table you go, the larger the atom gets. Furthermore, the more protons in the nucleus, the greater the attractive forces are. As such, the further to the right you go, the smaller it gets. See the chart below for the illustration.





Ionic radius is almost exactly the same as atomic radius, the only change is that rather than measuring the radius of a neutral atom, it measures the radius of an ionized atom. There are two basic rules for determining ionic radius. Anions (-) have a larger ionic radius than their corresponding atomic radii. This is because although there is the same nuclear charge, the greater number of electrons creates more repulsion and shielding. Thus, The effetive nuclear charge drops, and the radius increases. The opposite is true for cations (+). As they have fewer electrons, there is less repulsion/shielding. As such, the ion becomes smaller.
 * Ionic Radius**

An interesting comparison exists between isolectronic (same electron configuration) ions. A cation will be smaller than the anion, and the greater the charge, the greater the effect. For example, if both Na and Mg became charged into their standard ions, Mg would be smaller than Na.



Periods first. From left to right, the elements go from being very metallic to being very nonmetallic. Additionally, those elements closer to the edges tend to be more reactive.
 * Variations of Physical Properties Across a Period and Within a Group**

In groups, elements tend to be very similar. Properties that do change, such as melting point, tend to vary systematically, and you can find one value by knowing the values for the nearby elements.

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8.4 Ionization Energy
As you have hopefully realized, valence electrons determine most of an element's chemical properties. The stability of these electrons can be measured with Ionization energy, so it is critical that we study it. //Ionization energy// is the amount of energy required to remove an electron from a gaseous atom in its ground state. This is usually measured in kJ/mol, as the amount of stripping an individual electron from a single atom is miniscule. The important thing to note is that the higher the ionization energy, the harder it is to remove electrons and visa versa.

The energy to remove the first electron is known as the first ionization energy. The energy to remove the second electron is known as the second ionization energy. Go ahead and guess for the third electron. Importantly is that ionization energy continues to increase with each electron removed. This is because, without the first electron, there is less repulsion on the second, so it requires more energy to escape.

On the periodic table, ionization energy increases to the right and towards the top. In general, metals lose electrons easily, nonmetals do not. See the chart below for a visualization.

[] has a nice chart of some ionization energies.



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8.5 Electron Affinity
There is also electron affinity, which is almost the opposite of ionization energy. //Electron Affinity// is the negative of the change in energy when an atom accepts an electron. Electron affinity is also the energy required to remove an electron from an anion. As such, elements with a high electron affinity tend to gain and hold onto electrons. On the periodic table, Electron affinity is greatest towards the top, and also towards the right (towards the right is more important). This reflects these element's tendency to form anions easily. Just remember, metals have low affinity, nonmetals have high affinity.



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8.6 Variations in Chemical properties of the Representative Elements
Here it is. What this chapter has been coming to. Everything has been in preparation of this moment. Now that we know some of the forces behind the elements, we can draw conclusions and make generalizations about various elements based on their location in the periodic table.

Our first general rule is that metallic-ness increases to the left, and non-metallic-ness increases to the right.

Elements of a group share many properties. For example, Na and Mg have many similar properties (reactivity with water etc.) or Ne and Ar (nonreactivity). Of course, there is a small exception in that the first element in a group is oftentimes missing certain properties. For example, Li is not as similar to Na as Na is to Mg.
 * General Trends**

Furthermore, the strength of a property is generally magnified down a group. For example, Na and Mg are both reactive with water. However, Mg is much more violently reactive than Na is. Note that it is also lower on the periodic table.

Beyond relationships within a group or period, there are also //Diagonal relationships//. Diagonal relationships is the similarity between elements in different periods and groups (specifically adding 1 to both group and period number). For example, Li and Mg are similar in many ways, as are Be and Al, and B and Si. The reason that diagonal relations exist is because the charge density of these elements are very similar. As such, they behave very similarly in reactions.

Also note that the intra-group properties are most valid if taken with respect to metallic character. The 1A and 7A groups are very similar within their group, while group 7 is not so similar.





Hydrogen is the most abundant element in the universe, and also denies many forms of classification. As such, it is hard to properly place it on the periodic table. We place it in group 1 out of tradition, usefulness (, and a form of artistic beauty). H is unique in that it forms both a cation and anion readily. As such, it is both metallic and nonmetallic. With H serving as an anion in a compound, it will form Hmedia type="custom" key="3859129" width="17" height="15" when mixed with water, as the cation will combine with the Oxygen to form a metal oxide.
 * Hydrogen**

2NaH(s)+Hmedia type="custom" key="3859129" width="17" height="15"O(l) > 2NaOH(aq)+Hmedia type="custom" key="3859129" width="17" height="15"(g)

Moreover, when burned in Oxygen, Hydrogen will form water.

2Hmedia type="custom" key="3859129" width="17" height="15"(g)+Omedia type="custom" key="3859129" width="17" height="15"(g) > 2Hmedia type="custom" key="3859129" width="17" height="15"O

Hydrogen has the electron configuration 1s¹. Group 1A elements are known as the alkali metals. Group 1A elements have the electron configuration ns¹ (but not for n=1 aka H). Alkali metals love to lose their outermost electron, and as such form cations easily, are very reactive and metallic. Because of their great reactivity, they are almost never found in their pure forms. They react easily with water to form Hydrogen, and react with oxygen in the air to form oxides. Although all of the alkali metals form oxides, some form other compounds with oxygen such as super oxides (2 O for every metal) or peroxides (2 metal and 2 oxygen in every molecule).
 * Group 1A**

Group 2A elements are called alkali earth metals. They are not as reactive as the alkali metals. They also tend to form Mmedia type="custom" key="3859197" width="17" height="19"ions. One interesting reaction that occurs with alkali earth metals is that with acids in aqueous solutions. The charged hydrogen in the acid transfers its positive charge (technically the negative electrons from the metal move) to the metal forming a cation. The hydrogen is then free to bond with itself into a gaseous form. This releases the hydrogen gas.
 * Group 2A**

Group 3A has the electron configuration ns² np¹. They are all metals, except for B which is a metalloid. B does not react readily with oxygen or other elements ionically, so it is a bit of an outlier. Al forms a +3 ion, and readily combines with oxygen and similar elements. The rest of group 3 forms either +1 or +3 ions. In general, the further down the period that you go, the +1 ion becomes more stable than the +3 ion.
 * Group 3A (13)**

This is also a good time to notice how the metallic properties of elements varies across the periodic table. Note that Al forms AlHmedia type="custom" key="3881719" width="9" height="13". Thus, as the metal loses some of its metallic characteristics, it can be a bit easier to form with H.

Group 4A has the electron configuration ns² np². Group 4A elements are a mixed bunch. C is a nonmetal, Si and Ge are metalloids, and Sn and Pb are metals. The metals in 4A do not react with water as many metals do, but, like group 2A talked about, they react with the H in acids to liberate the H.
 * Group 4A (14)**

In group 4A the elements form either +2 or +4 ions. Near the top of the chart, the +4 oxidation state is much more stable than the +2 state. Yet, by the time you reach Pb, the +2 is much more stable than +4 oxidation state. Note that this is because it is either shedding all of the p electrons, or all electrons in the valence shell.

Group 5A has the electron configuration ns² np³. Again, we are moving towards nonmetallic-ness. N and P are nonmetals, As and Sb are metalloids, and Bi is a metal. Several of the elements have interesting properties. N is diatomic (it forms Nmedia type="custom" key="3881785" width="21" height="21"). N forms a ton of different compounds with Oxygen, you do not need to actually memorize them. N will form the Nmedia type="custom" key="3881795" width="18" height="17" ion. P forms into Pmedia type="custom" key="3881805" width="6" height="12". The metalloids form interesting 3-D structures. Bi is surprisingly unreactive.
 * Group 5A (15)**



Group 6A has the electron configuration ns² npmedia type="custom" key="3881835" width="9" height="9". O, S, and Se are nonmetals. The rest ( Te, Po) are metalloids. O is diatomic. S and Se form into groups of 8. Te and Po have really complex 3-D structures. They all form the Xmedia type="custom" key="3881851" width="13" height="14" ion.
 * Group 6A (16)**

Group 6A has the electron configuration ns² npmedia type="custom" key="3881857" width="15" height="15". They are known as the halogens. They are all diatomic and incredibly reactive. They are so reactive that they are never found on their own in nature. They attack substances such as water readily in a desperate attempt to gain that last electron. They all have large ionization energies, and large electron affinities. The halogens form 1- ions known as halides. The halides are very stable as they share the same electron configuration as the noble gas to its right. Halogens easily form with alkali metals and alkali earth metals. The halogens also form compounds among themselves. Finally, just as reactivity increased down the chart in group 1A, reactivity also increases down the chart in 7A.
 * Group 7A (17)**

Group 6A has the electron configuration ns² npmedia type="custom" key="3881911" width="11" height="15". They are known as the noble gases (or inert gases). They have completely filled electron shells and massive ionization energies. As such, they have no tendency to react with other elements. Under extreme conditions, chemists have been able to prepare a few compounds with Xe and even fewer with Kr. These are interesting as they show that the noble gases can react, but the compounds form have no practical purpose at all. In their base form the noble gases are colorless and odorless.
 * Group 8A (18)**

Both 1A and 1B (Cu, Ag, and Au) have only one electron in their s shell (one is used to complete the d shell). Thus, we would think that they have similar properties. They don't. One reason is that the ionization energies of 1B are higher than those of 1A. The ionization energies are so much lower because the d level does not shielded as well as the inner s and p levels do. The leck of reactivity makes these metals useful for making coins. Thus, giving rise to the term coinage metals. The 1A and 1B groups are somewhat similar.
 * Comparison of 1A and 1B**

Oxygen combines very readily with other elements. It also combines with almost all elements. As such, oxides are everywhere. So, we will take a moment and compare some of them.
 * Properties of Oxides Across a Period**

Na and Mg oxide are considered to be basic (not acidic) as they form hydroxide based compounds when combined with water. MgO does not dissolve very readily, but it does react well with acids much like a base would.

Al does not dissolve very well in water, but it does react with both acids and bases. This makes it amphoteric.

Silicon Dioxide also does not dissolve very well in water. However, it does react with concentrated bases. As such it can be considered something of an acid.

The other 3rd period elements from very acidic oxides.

In general, the further right on the period, the more acidic the oxide will be. This could also be said "metallic oxides are basic." In addition, to a lesser extent the further down the more basic.

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Review Questions
Note: some of these may require the use of the text or prior knowledge of chemistry. It is recommended that you have a [|periodic table] handy when you do these.


 * Highlight the area below a question to see the answer.**




 * Highlight the area below a question to see the answer.**

1. Which of the following are metals? As, Xe, Fe, Li, B, Cl, P, I, Si. Fe, Li, Ba, ., 2. Group the following species into isoelectric groups. Be2+, F-, Fe2+, N3-, He, S2-, Co3+, Ar. N-3, F-...Be2+,He...S2-,Ar...Fe2+,Co3+ 3.How does atomic change across a period? How does it change down a group? Smaller to the right, and biger towards the bottom. 4. Which has a larger atomic radius, Na or K? K 5. Which has a larger atomic radius, N or O? N 6.What is the ground-state electron configuration of Li+? 1s^2 7.Which is larger, Ca+ or Ca2+? Ca+ is larger 8. What is ionization energy? Ionization energy is the amount of energy required to remove 1 mol of electrons from 1 mol of atoms. 9. Which has a greater ionization energy, Be or O? O has a greater ionization energy. 10. Which has a greater electron affinity, Br or F? F 11. Name two elements that form a diagonal relationship from the following. Li, Mg, Al, C. Li and Mg 12. What causes diagonal relationships? Elements with diagonal relationships have similar charge densities. 13. Give an example of how chemical properties are similar within a group. Answers may vary. Both Na and K react violently with water. 14. Why are the noble gasses so stable? The noble gasses are stable because all of their electron shells are completely filled. As such they have massive ionization energies. This makes it very difficult to get them to react at all. 15. Which elements are more likely to form acidic oxides? Which are more likely to form basic oxides? Elements on the right are more acidic when made into oxides, and those on the left will form more basic oxides (so the alkali metals for alkali oxides. This might have given rise to the term alkali metals.)(note that alakali is another term for basic.) 16. Which of the following have distinct trends on the periodic table? If so, what is the trend? Molar mass, Atomic radius, First Ionization energy, number of isotopes. Molar mass shows the trend of increasing left to right and down, in a fashion similar to reading text. Atomic radius shows the trend of increasing down a group, and decreasing from left to right in a period. First Ionization energy inreases to the right and towards the top of the periodic table. The number of isotopes shows no distinct pattern, although there is a rough correlation between atomic mass and number of isotopes.

Congratulations, if you can answer these questions you have a decent grasp on the chapter. Of course, do check the book for more problems if you wish to review further.

Overview
Both elemental Na and K will be dropped into water. The reaction of one can then be used to predict the outcome of the other reaction. Now that you have read about the periodic table, it is time to see some of those relationships in action.

Materials

 * 3x 500 mL beaker
 * Distilled water
 * Li
 * Na
 * K
 * Universal Indicator
 * Forceps
 * Scupulla

Procedure
<span style="font-family: Times New Roman,Times New Roman;">
 * 1) PUT ON GOGGLES- These must be worn at ALL TIMES during the lab. The substances in this lab are very dangerous, and likely to explode.
 * 2) Put on other safety gear, including lab aprons. As stated, these chemicals are very dangerous. If you do get some Na or K on you, flush the area with plenty of water.
 * 3) Fill all beakers half-way with water. Add some universal indicator to each.
 * 4) Cut off a small piece of Li from the block. Replace the block in its container. Carefully place the Li in the water. Observe the reaction.
 * 5) Based on this predict what will happen for Na and K.
 * 6) Cut off a small piece of Na from the block. Quickly replace the source into its solution. Careful place the Na into the water. Observe the reaction.
 * 7) Based on this, guess what will happen when K is placed in water.
 * 8) Again, carefully remove a small piece of K. Be careful when handling K, it is violently dangerous and expensive. Very carefully drop it into the other beaker. Observe the reaction.
 * 9) Clean up your lab station.
 * 10) Write up the lab including detailed observations of each reaction . Answer all questions.

Questions

 * 1) Write the equations for both reactions.
 * 2) Which reaction is more violent?
 * 3) Why must Na and K be stored in oil?
 * 4) Are the products acidic or basic?
 * 5) Explain the periodic relationship between the elements used in this lab.
 * 6) Consider using chapter 18 to determine the heat of this reaction.

Video
media type="file" key="ms1.wmv"

Information credit: __Chang, Chemistry, 8th ed.__ <span style="color: rgb(0, 128, 0);">[|http://www.chalkbored.com/lessons/chemistry-11/alkali-metals-**lab**.pdf]

Image credits: [] [|http://www.doccasagrande.net/Images/Periodic_Table.jpg] [] [] [|http://www.chem.umass.edu/~botch/Chem111F04/Chapters/Ch8/IonicRadii.jpg] [] [] Zumdahl, 4th Edition

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