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Electron configuration worksheet name:_____ vandenbout/labrake. Elements in orbital notation. (this is one of many possible answers) 5. Orbital notation refers to the electron configuration of an element. The notation tells you how many electrons are in the element, what orbitals of the atom they are. Chapter 5 homework answers 1. Which statement regarding the gold foil experiment is. Draw the ground-state electron configuration using the orbital notation. Home / study / science / chemistry / questions and answers / how do you find the orbital notation. Find the orbital notation of. Answer with chegg study view. To show the electron configuration for an atom, what is the advantage of using an orbital notation compared to a dot structure? - 1398667

Copyright ©Orbital notation refers to the electron configuration of an element. The notation tells you how many electrons are in the element, what orbitals of the atom they are.

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You can see that it is going to get progressively tedious to write the full electronic structures of atoms as the number of electrons increases. There are two ways around this, and you must be familiar with both.

*Shortcut 1:* All the various p electrons can be lumped together. For example, fluorine could be written as 1s 2 2s 2 2p 5. and neon as 1s 2 2s 2 2p 6 .

This is what is normally done if the electrons are in an inner layer. If the electrons are in the bonding level (those on the outside of the atom), they are sometimes written in shorthand, sometimes in full. Don't worry about this. Be prepared to meet either version, but if you are asked for the electronic structure of something in an exam, write it out in full showing all the p_{x}. p_{y} and p_{z} orbitals in the outer level separately.

For example, although we haven't yet met the electronic structure of chlorine, you could write it as 1s 2 2s 2 2p 6 3s 2 3p_{x} 2 3p_{y} 2 3p_{z} 1 .

Notice that the 2p electrons are all lumped together whereas the 3p ones are shown in full. The logic is that the 3p electrons will be involved in bonding because they are on the outside of the atom, whereas the 2p electrons are buried deep in the atom and aren't really of any interest.

*Shortcut 2:* You can lump *all* the inner electrons together using, for example, the symbol [Ne]. In this context, [Ne] means *the electronic structure of neon* - in other words: 1s 2 2s 2 2p_{x} 2 2p_{y} 2 2p_{z} 2 You wouldn't do this with helium because it takes longer to write [He] than it does 1s 2 .

On this basis the structure of chlorine would be written [Ne]3s 2 3p_{x} 2 3p_{y} 2 3p_{z} 1 .

*The third period*

At neon, all the second level orbitals are full, and so after this we have to start the third period with sodium. The pattern of filling is now exactly the same as in the previous period, except that everything is now happening at the 3-level.

1s 2 2s 2 2p 6 3s 2 3p 6 *4s 2*

There is strong evidence for this in the similarities in the chemistry of elements like sodium (1s 2 2s 2 2p 6 *3s 1* ) and potassium (1s 2 2s 2 2p 6 3s 2 3p 6 *4s 1* )

The outer electron governs their properties and that electron is in the same sort of orbital in both of the elements. That wouldn't be true if the outer electron in potassium was 3d 1 .

*s- and p-block elements*

The elements in Group 1 of the Periodic Table all have an outer electronic structure of ns 1 (where n is a number between 2 and 7). All Group 2 elements have an outer electronic structure of ns 2. Elements in Groups 1 and 2 are described as s-block elements.

Elements from Group 3 (the boron group) across to the noble gases all have their outer electrons in p orbitals. These are then described as p-block elements.

*Note:* If you use the current IUPAC (international Union of Pure and Applied Chemistry) system for group numbering, you will probably know what I call Group 3 as Group 13. My reasons for not using the IUPAC system are discussed on this page in the Questions and Comments section.

We are working out the electronic structures of the atoms using the Aufbau ("building up") Principle. So far we have got to calcium with a structure of 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 .

The 4s level is now full, and the structures of the next atoms show electrons gradually filling up the 3d level. These are known as d-block elements.

Once the 3d orbitals have filled up, the next electrons go into the 4p orbitals as you would expect.

d-block elements are elements in which the last electron to be added to the atom using the Aufbau Principle is in a d orbital.

The first series of these contains the elements from scandium to zinc, which at GCSE you probably called transition elements or transition metals. The terms "transition element" and "d-block element" don't quite have the same meaning, but it doesn't matter in the present context.

*If you are interested:* A transition element is defined as one which has *partially filled* d orbitals either in the element or any of its compounds. Zinc (at the right-hand end of the d-block) always has a completely full 3d level (3d 10 ) and so doesn't count as a transition element.

Some UK syllabuses use a more restrictive definition which defines a transition metal as one which has one or more stable ions with partly filled d orbitals. You don't need to worry about this until you do some transition metal chemistry.

d electrons are almost always described as, for example, d 5 or d 8 - and not written as separate orbitals. Remember that there are five d orbitals, and that the electrons will inhabit them singly as far as possible. Up to 5 electrons will occupy orbitals on their own. After that they will have to pair up.

Notice in what follows that all the 3-level orbitals are written together - with the 4s electrons written at the end of the electronic structure.