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Electrochemical Cells - Calculating Voltage
Electrochemical Cells - How to build a parallel battery.
Electrochemical Cells - How to build a series battery.
Electron Dot Notation
Freezing Point Depression
Gram - Mole Conversions
Metals and the Periodic Chart
pH of Acids
Significant Digits in Chemistry
Stoichiometry Practice Problems
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by:Sarah B. and Casey K.
Orbital notation is the process in which an unoccupied orbital is represented by a line. Arrows are placed above the line. These arrows signify a magnetic field where an arrow is spinning in one direction, and the adjacent arrow is spinning in the opposite direction. Before you can fully understand orbital notation you must learn about the diagonal chart.
To understand the orbitals, arrows must be placed under each level. The S orbital, for example, holds a maximum of 2 electrons, leaving 1 arrow in place of a single electron. What does that mean? Each two electrons (arrows) are given their own orbital so that they may appear as their own electron instead of conjoined with other electron pairs. Usually, the correct amount of arrows is equal to the amount of electrons of the element. Therefore, the S orbital may only hold two arrows, an arrow facing up and down because it only has a single orbital in which to contain the electrons. Arrows are faced in opposite ways to represent the opposite attraction amongst the pair.
With orbital notation dealing with electrons, the direction of a electron orbit deals with how many electrons are within a level. We know that there are two electrons for each orbital level thus the two arrows. The arrows represent the direction an electron spins, clockwise or counter clock wise. If an atom only has one electron in its level, ex. hydrogen, then there is only one electron orbiting that atom. If you come across a atom with more than one electron, then you have multi directional spin from the electrons, both of them going opposite directions from one another.
Looking at P orbital, it holds no more than 6 electrons because it contains three orbitals, with a maximum of 2 electrons in each orbital. Look at it from a multiplication standpoint. Each orbital holds a required "full" amount of electrons, with 2 electrons for orbital, if all orbitals are full, then the entire orbital set is full. D orbital contains a maximum of 5 orbitals with no more than 2 electrons each for a maximum of 10 electrons orbital set, and F orbital set is completed by 7 orbitals with 2 electrons in each giving a full electron level of no more than 14 electrons.
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Understand the direction of orbitals yet? Heres the breakdown; an orbital set is like your Periodic Table of elements in Blocks. You have S, P, D, and F.
The amount of arrows, or electrons, placed on each line depends upon the number of electrons in the element. Even though an orbital set has a maximum outcome, the total of electrons may not always fill the oribitals. An example could be Hydrogen; an element with only one electron will have only one arrow in the first orbital. At times the lines, or orbitals, contain only one electron or are left empty. When there are not enough electrons to fill all orbitals to their fullest, there is a special rule you must follow. You must place an arrow on each line before filling the orbital with two arrows. Meaning, every orbital must have an electron placed before another electron may be added to complete the orbital.
The only difference when writing out a notation is the final product. The final product may have all levels of orbitals full or some orbitals left with one or no electrons. Remembering that all orbitals should (if possible) be filled with electrons before going to the next orbital level following the increasing energy levels. If you are still having trouble comprehending orbital notation you can view our video below.
After you have watched our video, you can try some of the practice problems below.
1.Na- 11 electrons
2.S- 16 electrons
3.Br- 35 electrons
The answers(sorry the arrows wouldn't upload so we used the numbers to tell you how many arrows are needed):
1. 1s(2) 2s(2) 2p(6) 3s(1)
2. 1s(2) 2s(2) 2p(6) 3s(2) 3p(4)
3. 1s(2) 2s(2) 2p(6) 3s(2) 3p(6) 3d(10) 4s(2) 4p(5)
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