Most students who first learn electron configurations often have trouble with configurations that must pass through the f-block because they often overlook this break in the table and skip that energy level. Its important to remember that when passing the 5d and 6d energy levels that one must pass through the f-block lanthanoid and actinoid series. Keeping this in mind, this "complex" problem is greatly simplified. Another method but less commonly used of writing the spdf notation is the expanded notation format.
This is the same concept as before, except that each individual orbital is represented with a subscript. The p, d, and f orbitals have different sublevels. The p orbitals are px, py , and pz, and if represented on the 2p energy with full orbitals would look like: 2p x 2 2p y 2 2p z 2.
The individual orbitals are represented, but the spins on the electrons are not; opposite spins are assumed. When representing the configuration of an atom with half filled orbitals, indicate the two half filled orbitals. The expanded notation for carbon is written as follows:. Because this form of the spdf notation is not typically used, it is not as important to dwell on this detail as it is to understand how to use the general spdf notation.
This brings up an interesting point about elements and electron configurations. As the p subshell is filled in the above example about the Aufbau principle the trend from boron to neon , it reaches the group commonly known as the noble gases. The noble gases have the most stable electron configurations, and are known for being relatively inert.
All noble gases have their subshells filled and can be used them as a shorthand way of writing electron configurations for subsequent atoms. This method of writing configurations is called the noble gas notation, in which the noble gas in the period above the element that is being analyzed is used to denote the subshells that element has filled and after which the valence electrons electrons filling orbitals in the outer most shells are written.
This looks slightly different from spdf notation, as the reference noble gas must be indicated. Vanadium is the transition metal in the fourth period and the fifth group. The noble gas in the configuration is denoted E, in brackets: [E]. Instead of 23 electrons to distribute in orbitals, there are 5.
Now there is enough information to write the electron configuration:. This method streamlines the process of distributing electrons by showing the valence electrons, which determine the chemical properties of atoms. In addition, when determining the number of unpaired electrons in an atom, this method allows quick visualization of the configurations of the valance electrons. In the example above, there are a full s orbital and three half filled d orbitals.
Unless specified, use any method to solve the following problems. A nswers are given in noble gas notation. Find the electron configurations of the following:. Scenario: You are currently studying the element iodine and wish to use its electron distributions to aid you in your work. Thought Questions:. Identify the following elements:. Without using a periodic table or any other references, fill in the correct box in the periodic table with the letter of each question.
Find the electron configuration of the following:. To find the answer we refer to part a and look at the valence electrons. We see that iodine has 5 electrons in the p orbitals.
We know that the full p orbitals will add up to 6. Using the Hund's rule and Pauli exclusion principals we can make a diagram like the following:. The first part of this question is straightforward. The second part is slightly more complicated. Because each individual's knowledge of chemistry differs, there are many answers to this question. The important aspect is that we realize that knowing electron configurations helps us determine the valence electrons on an atom.
This is important because valence electrons contribute to the unique chemistry of each atom. This should also be a straightforward question, and if it seems a little difficult refer to the body of this text about these rules and how they relate to creating an electron configuration. Remember to make logical connections! We know that the main "tools" we have in writing electron configurations are orbital occupation, the Pauli exclusion principle, Hund's rule, and the Aufbau process.
Orbitals are occupied in a specific order, thus we have to follow this order when assigning electrons. The fourth quantum number, which refers to spin, denotes one of two spin directions.
This means that in one orbital there can only be two electrons and they mus have opposite spins. This is important when describing an electron configuration in terms of the orbital diagrams. Hund's rule states that electrons first occupy the similar energy orbitals that are empty before occupying those that are half full. This is especially helpful when determining unpaired electrons. The Aufbau process denotes the method of "building up" each subshell before moving on to the next; we first fill the 2s orbitals before moving to the 2p orbitals.
We know that the noble gas has all of its orbitals filled; thus it can be used as a "shorthand" or abbreviated method for writing all of the electron configurations after 1s. Introduction Before assigning the electrons of an atom into orbitals, one must become familiar with the basic concepts of electron configurations. Rules for Assigning Electron Orbitals Occupation of Orbitals Electrons fill orbitals in a way to minimize the energy of the atom.
The order of levels filled looks like this: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, and 7p One way to remember this pattern, probably the easiest, is to refer to the periodic table and remember where each orbital block falls to logically deduce this pattern.
Pauli Exclusion Principle The Pauli exclusion principle states that no two electrons can have the same four quantum numbers. Visually, this is be represented as: As shown, the 1s subshell can hold only two electrons and, when filled, the electrons have opposite spins.
Hund's Rule When assigning electrons in orbitals, each electron will first fill all the orbitals with similar energy also referred to as degenerate before pairing with another electron in a half-filled orbital.
The Aufbau Process Aufbau comes from the German word "aufbauen" meaning "to build. Exceptions Although the Aufbau rule accurately predicts the electron configuration of most elements, there are notable exceptions among the transition metals and heavier elements. Writing Electron Configurations When writing an electron configuration, first write the energy level the period , then the subshell to be filled and the superscript , which is the number of electrons in that subshell.
Three methods are used to write electron configurations: orbital diagrams spdf notation noble gas notation Each method has its own purpose and each has its own drawbacks. Orbital Diagrams An orbital diagram, like those shown above, is a visual way to reconstruct the electron configuration by showing each of the separate orbitals and the spins on the electrons.
Example 4: Aluminum and Iridium Write the electron configuration for aluminum and iridium. This gives the following: Note that in the orbital diagram, the two opposing spins of the electron can be visualized. Another example is the electron configuration of iridium: The electron configuration of iridium is much longer than aluminum.
Example 5: Yttrium Write the electronic configuration of Yttrium. Its electron configuration is as follows: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 1 This is a much simpler and more efficient way to portray electron configuration of an atom. Noble Gas Notation This brings up an interesting point about elements and electron configurations. Now there is enough information to write the electron configuration: Vanadium, V: [Ar] 4s 2 3d 3 This method streamlines the process of distributing electrons by showing the valence electrons, which determine the chemical properties of atoms.
References Petrucci, Ralph H et al. Print Sherman, Alan, Sharon J. Sherman, and Leonard Russikoff. Basic Concepts of Chemistry Fifth Edition. Compendium of Chemical Terminology, 2nd ed.
Compiled by A. McNaught and A. Blackwell Scientific Publications, Oxford Nic, J. Jirat, B. Kosata; updates compiled by A. ISBN Scerri, Eric R. Ostrovsky, V.
On recent discussion concerning quantum justification of the periodic table of the elements. Foundations of Chemistry , 75 , Meek, T. Configuration irregularities: deviations from the madelung rule and inversion of orbital energy levels.
ScienceDirect , 5 , Problems Unless specified, use any method to solve the following problems. Find the electron configurations of the following: silicon tin lead 2. Find the electron configuration of iodine How many unpaired electrons does iodine have? Thought Questions: In your own words describe how to write an electron configuration and why it is an important skill in the study of chemistry.
Describe the major concepts Hunds, Pauli Answers 1. Find the electron configuration of the following: a silicon: [Ne] 3s 2 3p 2 b tin: [Kr] 5s 2 4d 10 5p 2 c lead: [Xe] 6s 2 4f 14 5d 10 6p 2 2. Using the Hund's rule and Pauli exclusion principals we can make a diagram like the following: The answer is one. Thought Questions: a In your own words describe how to write an electron configuration and why it is an important skill in the study of chemistry.
Period In these cases, you can use the previous noble gas to abbreviate the configuration as shown below. You just have to finish the configuration from where the noble gas leaves it:. As with every other topic we have covered to date there are exceptions to the order of fill as well.
But based on the electron configurations that are generated, these exceptions are easy to understand. In the d block, specifically the groups containing Chromium and Copper, there is an exception in how they are filled. There are lots of quizzes on electron configurations you can practice with located here. Another way to represent the order of fill for an atom is by using an orbital diagram often referred to as "the little boxes":. The boxes are used to represent the orbitals and to show the electrons placed in them.
The order of fill is the same but as you can see from above the electrons are placed singly into the boxes before filling them with both electrons. This is called Hund's Rule: "Half fill before you Full fill" and again this rule was established based on energy calculations that indicated that this was the way atoms actually distributed their electrons into the orbitals. One of the really cool things about electron configurations is their relationship to the periodic table. Basically the periodic table was constructed so that elements with similar electron configurations would be aligned into the same groups columns.
The periodic table shown above demonstrates how the configuration of each element was aligned so that the last orbital filled is the same except for the shell. The reason this was done is that the configuration of an element gives the element its properties and similar configurations yield similar properties. Let's go through some of the Periodic Properties that are influenced directly by the electron configuration:.
The size of atoms increases going down in the periodic table. This should be intuitive since with each row of the table you are adding a shell n. What is not as intuitive is why the size decreases from left to right. But again the construction of the electron configuration gives us the answer. What are you doing as you go across the periodic table?
Answer, adding protons to the nucleus and adding electrons to the valence shell of the element. What is not changing as you cross a period? Answer, the inner shell electrons. So think of it this way, the inner shell electrons are a shield against the pull of the nucleus.
As you cross a period and increase the number of protons in the nucleus you increase its pull but since you are only adding electrons to the new shell the shield is not increasing but remains the same all the way across. This means the pull on the electrons being added to the valence shell is increasing steadily all the way across.
What happens if you pull harder on the electrons? Well, they come closer to the nucleus and the size of the atom decreases. Electronegativity may be the most important of the periodic properties you can learn and understand since so many other properties are depend on its value. Electronegativity is an atoms ability to pull electrons towards itself.
Electronegativity is generally expressed by the Pauling Scale and the values were determined experimentally. The table below shows the scale values for the elements. The electronegativity values increase from left to right and bottom to top in the periodic table excluding the Noble gases.
The most electronegative element is Fluorine. From these electronegativity values we can derive the patterns of two other periodic properties: Ionization Energy and Electron Affinity.
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