Periodic Table Study Guide
Modern Periodic Table
The Main Group Elements are located in Groups 1,2 and 13-18. Groups 1 and 2 are filling out the s orbital in the outermost energy level. Groups 13-18 are filling out the p orbital in the outermost energy level.Since s orbitals can hold only 2 electrons and p orbitals can hold only 6 electrons, there are only 8 Groups included in the Main Group Elements.
Ionization Energy is the energy required to separate a valence electron from an atom. All electrons are attracted to the protons in the nucleus, and energy is needed to overcome that attraction. Ionization Energy decreases as you go down a Group. Watch this Video as it examines most all the elements in Group 1, from top to bottom. The reactions occur faster and release more energy because it gets easier and easier to remove a Valence Electron.
The simple fact that electrons are easier to remove as you move down a Group totally explains the video. But why are they easier to remove? The answer is simple, as long as you consider two facts: 1. Electrons have the same (negative) charge and therefore repulse from each other. 2. As you move each step down a group, you're adding another Energy Level (or layer) that's full of Electrons between the Nucleus and the Valence Electron that's taking part in the reaction. The repulsion that occurs between a Valence Electron and the Interior Electrons is called "Electron Shielding." The repulsion from the interior Electrons help to push away the Valence Electron. A quick look at the Periodic Table shows that Cesium, the final metal tested in the Video, is in Period 6. It therefore has 6 occupied Energy Levels. But, more importantly, it has 5 interior Energy Levels with more than 50 other electrons between the Valence Electron and the Nucleus. That Valence Electron can't wait to jump off the Cesium Atom. It therefore required very little Ionization Energy. Expect to answer trends questions about Electron Shielding and Ionization Energy on the Unit Test.
Dmitri Mendeleev was credited for devising the first Periodic Table in 1869 that included all 63 elements known at that time . He arranged the Table horizontally by increasing Mass Numbers but vertically by Chemical Properties. He and other chemists had observed only eight general chemical properties, so he used eight vertical arrangements which he called Groups. In doing so, he realized that gaps existed in his table. He correctly predicted that other elements must exist that would fill those gaps. None of them were yet discovered when he published his Table. But he was proved absolutely correct about the missing elements' existence and properties.
Transition, or d-block, elements are found in Groups 3-12. They are explained by the manner in which Electron Configurations are filled: Recall from Page 97 that d orbitals don't fill until the s orbital from the next energy level fills. So starting in Period 4, the 3-d oritals do not fill until two electrons fill the 4s orbital. The process repeats through Period 7. Since a d orbital requires 10 electrons to reach capacity, there are 10 Transition Groups. It's always simple to identify the Group Number for a Transition Element if you know its Electron Configuration. Merely add the electrons from its outermost s and d orbitals. Their total matches the Group Number.
On the other hand, Ionization Energy tends to increase as you go left to right across a Period. Periods, by definition, communicate how many Energy Levels are occupied. If you stay in the same Period, there is no difference in the Electron Shielding effect. However, the number of protons is increasing by one every element along the way. So the positive nuclear charge is also increasing. All electrons, including the Valence Electrons, love positive protons. It therefore requires more energy to remove that Valence Electron as you travel across a Period.
The Period (Horizontal) Number for any Element is obvious with one quick look at its Electron Configuration. Merely take note of the last Energy Level (that's the large Coefficient Number used in the Configuration). The Element will always be found in that period. The Periodic Table above links to an interactive one. Click on it, and then click on any element or any number of elements. A lot of information will appear, including the Electron Configuration. Anyone that proves me wrong about the Period Number gets exempted from the next test with an A-Plus.
An atom's size, normally referred to as its Atomic Radius, predictably increases as you move down a Group, because you're adding layers of Energy Levels each step of the way. But what surprises many students is that Atoms tend to become much smaller as you move left to right across a Period. How can that be, since there are so many more particles like protons, electrons and neutrons? The answer once again isn't that complicated. As you move left to right across a Period, the number of Energy Levels remains unchanged. But the extra Electrons in the orbitals absolutely love the extra protons in the nucleus. The increasing positive charge in the nucleus pulls the cloud of Electrons in tighter. Therefore, Bromine is a much smaller atom than Potassium, even though Bromine has about twice the mass of Potassium. You can check these trends on the interactive Periodic Table link above. With few exceptions, you'll always see the Atomic Radius values (measured in picometers, or pm) get smaller while you move left to right across a Period, but get much larger as you move down any Group.
For a comprehensive look at all Periodic Trends, click here. Recognizing them will suffice for a passing grade on the next test. Explaining them will earn you an A.
Notice that the Lathanide and Actinide series are in Periods 6 and 7. Also note that they each contain 14 Elements. That's because they are filling the f orbitals, which have a capacity of 14 Electrons.
Electro-negativity is merely a reference to how attracted an electron is to an atom that is ALREADY BONDED TO ANOTHER ATOM. There are no units to an Electro-negativity value. Looking ahead to Page 194, you'll see values ranging as low as 0.7 for Francium (lower left corner of the Table) to as high as 4.0 for Flourine (upper right corner of the table). The raw trend is therefore this: as you go down a Group, an Atom's attraction for an Electron decreases but going left to right across a Period, the attraction increases. Why would this be? It once more goes back to the competing forces of attraction (electrons love positive protons found in an atom's nucleus) and repulsion (electrons hate other electrons found in an atom's energy levels). As you go down a Group, you are adding layers of Energy Levels full of Electrons every step of the way. The Repulsion Force magnifies from the Electron Shielding Effect. As you go across a Period, there is no change in the number of Energy Levels and therefore little change in the Repulsion Force every step of the way. However, there is a constant increase in the number of Protons, which always attract a negatively-charged Electron.
Electron Affinity is the Energy Change that occurs when a Neutral Atom gains an Electron. It therefore follows the same trends at Electro-negativity, which measures the attraction of an Electron to a bonded Atom. Electron Affinity is measured with Energy Units, but all that's important for us is to recognize and explain the Trend: Electron Affinity reduces as you go down a Group but increases as you go left to right across a Period.
Within the Main Group Elements, you should be able to identify four specific Groups:
1. Alkali Metals are in Group 1. Their Electron Configurations all end with a single s orbital electron in the outermost Energy Level. Check this out by going to the Periodic Table.
2. Alkaline Earth Metals are in Group 2. They fill out the outermost s orbital with two electrons.
3. Halogens are in Group 17. They all end with five electrons in the outermost p orbital.
4. The Noble Gases are in Group 18. They all have the outermost p orbitals completely filled with six electrons, with the exception of Helium in the 1st Period. No p orbitals exist in that first Energy Level, so Helium's Configuration shows its only s orbital filled with two Electrons.
Hydrogen's Electron Configuration ends with a single electron in the s orbiital, so it seems like it should be part of Group 1. It's not. Hydrogen is considered its own Special Group , and not part of the Alkali Metals for various reasons. Two prominent ones are that it's not a metal, and it's the most common element in the Universe, and is likewise present in most all of the common compounds we encounter, including water, food, most fuels and all lifeforms.
Hydrogen is also considered the main building block of all Elements in our Universe. Scientists have discovered evidence that Hydrogen Atoms fuse together in the core of stars, which merges their nuclei to form larger elements. This process has been performed in a limited fashion in laboratories in recent decades. Fusion is a Nuclear Reaction, and is often referred to as a Transmutation. The word mutation has a similar meaning here as it does in Biology. A new identity is formed because the proton number inside the merged nuclei increases. This process also releases an enormous amount of Energy (like we get from the Sun). Scientists hope that harnessing industrial scale Fusion Reactions will someday provide us with inexhaustible amounts of non-polluting Energy.
MOCK TEST WITH ANSWER KEY CLICK HERE.
Reflection: Refer back to Mendeleev's original observation that Chemical Reactions for all known elements seemed to fit within eight specific categories. About 150 years later, we still only know of eight specific Reaction Types. Mendeleev could only observe it, but not explain it. But since Niehls Bohr polished off the Modern Atomic Theory with his Electron Configurations, the explanation for eight types of reactions is obvious. Recall that only Valence Electrons take part in Chemical Reactions. Then go back and count the number of Valence Electrons that exist in any Atoms you choose to check on the Periodic Table. Merely click on an Atom, and its Electron Configuration will appear, along with other information. You'll find that no Atom contains more than eight Valence Electrons. Specifically, in the last Energy Level, the Configurations are lastly filling either an s orbital (which holds two electrons) or the following p orbital, which holds six more.
On a side note, US Government officials considered using Sodium Metal as a weapon in WWII. They stockpiled, but didn't use, very large amounts of it and had to get rid of it after the War. Just think how powerful the Alkali Metal bombs would have been if they could have stockpiled large amounts of Cesium instead.
Review of key Periodic Trends. As we examine elements in Vertical Groups or Horizontal Periods on the Modern Periodic Table, you'll notice tendencies exist for certain chemical and physical properties. You'll want to recognize and explain these tendencies when it comes Unit Test time.