Effective Nuclear Charge & Periodic Trends

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  • 0:02 Introduction to…
  • 1:03 Atoms and the Periodic Table
  • 2:23 Electrons and…
  • 5:55 Effective Nuclear…
  • 8:53 Lesson Summary
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Lesson Transcript
Instructor: Nicola McDougal

Nicky has taught a variety of chemistry courses at college level. Nicky has a PhD in Physical Chemistry.

This video lesson will describe effective nuclear charge and its role in explaining periodic trends. In particular, we will learn how to predict the trends in atomic radius and ionization energy using our knowledge of shielding and effective nuclear charge.

Introduction to Effective Nuclear Charge

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.

Atoms and the Periodic Table

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.

Electrons and Shielding Effects

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.

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