Back To CourseMCAT Test: Practice and Study Guide
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Nicky has taught a variety of chemistry courses at college level. Nicky has a PhD in Physical Chemistry.
Today your favorite street entertainer is performing in the local town square. You rush down to watch him, but when you arrive, quite a few people are already there. You make your way to where the crowd is gathered, but you can only get in the third row. Darn, your view is not good! The people in front of you are getting in your way. From where you are sitting, you can only see glimpses of the act. This happens to me all the time: I only get a good view, or feel the effect of the performance, if I am close to the performer.
The same effect is also true of electrons and atoms. You can imagine the electrons are the audience, and the nuclear charge is the performer. In this lesson, we learn about the effective nuclear charge, and its effect on the properties of atoms. We can think of effective nuclear charge as the positive charge felt by the outermost electrons in an atom.
Let's first remind ourselves about the atom. A simple model is shown here. An atom consists of a positively charged nucleus, containing protons and neutrons, surrounded by negatively charged electrons. Electrons exist in discrete energy levels around the nucleus called orbitals. The further an electron is away from the nucleus, the higher the energy level is. Electrons always go into the lowest energy level first, until that energy level is filled.
The periodic table is the arrangement of elements in rows and columns according to atomic numbers. Each element has its own unique atomic number. Here you can see gold, and there are two numbers shown. The atomic mass is shown below the symbol; the number 79 is the atomic number. The atomic number (Z) is the number of protons in the nucleus of an atom. It is this number that gives the atom its unique identity. The atomic number also tells us the number of electrons in an uncharged atom. An uncharged, or neutral, gold atom, has 79 protons and 79 electrons. The positive charge of the protons and the negative charge of the electrons balance to give a neutral atom.
Let's do a very quick review on electrons. We have already said that there are different energy levels in an atom, and electrons always go into the lowest energy level first until the energy level is filled. One is the first energy level, then two, and so on. You may also recall these levels are further divided into sub-levels, called s, p, d, and f. These sub-levels can take 2, 6, 10, and 14 electrons, respectively. If this is new to you, you may want to review electron configurations before coming back to this lesson.
An expression that gives us the number of electrons in each energy level is the electron configuration. For example, the electron configuration for the lithium atom is 1s^2 2s^1and the electron configuration for the Fluorine atom is 1s^2 2s^2 2p^5. It is important you can write electron configurations because they tell us where the electrons are. Lithium has an atomic number of 3. It has 3 electrons. Two electrons are in the 1s energy level, and 1 electron is in the 2s energy level. The electron in the 2s energy level is further away from the nucleus and has highest energy. Fluorine has the atomic number of 9. It has 9 electrons -- two electrons in the 1s energy level, 2 electrons in the 2s energy level, and 5 electrons in the 2p energy level. Fluorine has 7 electrons at higher energy. The outer electrons at highest energy are also called valence electrons.
Ok, so now that we have reviewed the atom and electron configurations, let's think more about the interaction between the positively charged nucleus and the negatively charged electrons. Opposite charges attract, and the positive nucleus attracts the negative electrons. The closest electrons will feel the attractive effect much more than the electrons further away. It isn't just about distance: there is also something else going on. To help us understand this, we return to our street entertainer from the beginning of the lesson. Now recall you are back in the third row and your view is blocked by the people in the two rows in front of you. You are not feeling the full effect of the performance because of this; they are shielding you from full view. The two rows in front of you are completely full, and you have to stay in the third row; that is you seat. What do you think might happen if you move closer to the second row to try and see better? Well, that didn't work! They physically pushed you away.
Exactly the same situation happens in the atom. Instead of rows of people, there are rows of energy levels of electrons. The first energy level is closest to the nucleus, and just like the first row in our street performance, feels the greatest effect. Electrons in the closest, or inner, filled energy levels are the core electrons. And just like our people, these core electrons shield the outer electrons from the full nuclear charge. This is known as the shielding effect. If the outer electrons try to get closer to the nucleus, they are forced back because like charges repel.
Now we have learned that core electrons shield outer electrons from the nuclear charge, let's now take this knowledge to predict periodic trends. The two trends we will look at are atomic radius and ionization energy. There are trends going down a group and going across a period in the periodic table. Let's first look at what happens as we go down a group. Atomic radius generally increases down a group. The outer electron configuration in a group is the same, but as you go down the group there are more and more core electrons. These core electrons are held close to the nucleus and the shielding effect increases. The size of the atom increases because the outer electrons are not held so tightly. The energy needed to remove an electron is the ionization energy, and it decreases down a group. Simply put, because the outer electron is not held so tightly, it is easier to remove.
But what is happening as you go across a period? The trend is reversed. The atomic radius is decreasing. This time the atomic number and number of electrons is increasing across the period. But unlike core electrons, shielding between electrons in the same energy level is poor. Therefore, the effective nuclear charge increases with atomic number across a period. The size of the atom decreases because the outer electrons are held more tightly. Because the outer electron is held more tightly, it is more difficult to remove, and ionization energy increases across a period.
We can also calculate the approximate effective nuclear charge by Z Effective = Z - S, where Z is the atomic number and S is the number of shielding electrons. Shielding electrons are all non-valence electrons. They are in the inner shells. An easy way to remember this formula is by using the following rhyme: Zebras Ecstatically Enjoy Zesty Moving Spacemen.
We will finish by doing a quick example. What is the effective nuclear charge on the Cl atom? Chlorine has the atomic number of 17. It has the electron configuration of neon (Ne): 3s^2 3p^5. There are 7 outer electrons and 10 shielding, inner electrons. So ZE = 17 - 10 = 7 +. The higher the effective nuclear charge, the more effect the positive nuclear charge is having on the outer electrons.
So let's review. We learned that effective nuclear charge is the positive charge felt by the outermost electrons in an atom. This helps us predict periodic trends. We learned that electrons exist in discrete energy levels, and that electrons fill energy levels closest to the nucleus first. The further away the energy level is from the nucleus, the higher the energy. We learned that core electrons shield outer electrons from the nuclear charge. This has the effect of reducing the amount of attraction the outer electrons feel. However, we also learned that electrons in the same energy level do not shield effectively. This understanding of the shielding effect can be used to predict periodic trends in both ionization energy and atomic radius. We can calculate an effective nuclear charge by using Z Effective = Z - S, where Z is the atomic number and S is the number of shielding electrons.
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Back To CourseMCAT Test: Practice and Study Guide
88 chapters | 863 lessons