Just to show you what it looks like, here is the electron configuration for silicon:

1s

^{2}2s^{2}2p^{6}3s^{2}3p^{2 }^{ }

You see, electrons only exist in certain energy states. When the electrons absorb or emit a specific amount of energy, it will instantly jump to another orbital. The energy level is basically that amount that an electron of an atom can have. This is called "

*n".*

The energy difference between 2 energy levels is called the

__quantum__of energy.Atoms can exist in 2 states:

**all electrons of the atom are at their lowest possible energy level**

__Ground State:__

__Excited State:__one or more electron is in an energy level other than the lowest possible one

**Here are some other things that you should know:**

__Orbital:__the actual amount of space that is being occupied by an electron

__Shell__: the set of all orbitals that have the name n-value

__Subshell__: a set of orbitals of the same type (like 2s and 2p)

s, p, d, and f refer to four types of orbitals.

An atom's electron shells are filled according to the following theoretical constraints:

- Each s subshell holds at most 2 electrons
- Each p subshell holds at most 6 electrons
- Each d subshell holds at most 10 electrons
- Each f subshell holds at most 14 electrons

There can only be a max of 2 electrons per orbit, so the p subshell would need 3 orbitals, each with 2 electrons. d subshell would need 5 orbitals, each with 2 electrons. f subshell would need 7 orbitals, each with 2 electrons. Sometimes an orbital will only have one electron is there is an odd number of electrons.

1s

^{2}2s^{2}2p^{6}3s^{2 }3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}4f^{14}5d^{10}6p^{6}7s^{2}5f^{14}6d^{10}7p^{6}Above is the order in which electrons are filled. For example, a neutral oxygen atom has 8 electrons, so the electron configuration would be:

**1s**

^{2}

**2s**

^{2}

**2p**

^{4 }

^{ }

^{Notice that the p only has a 4, and not a 6. This is alright. The superscripts should add up to the number of electron in the atom. In this case, it adds up to 8, which is correct. }

When writing electronic configurations for negative ions, simply add the appropriate number of electrons from where the neutral atom left off.

When writing electronic configurations for positive ions, take away electrons from the outermost shell first. (aka. the higher energy shell/larger n-value)

Another method of writiting configurations is

__Core Notation__.**The core is the part of the configuration of the nearest noble gas that comes before the given element. (Ex/ Ar is the noble gas that comes before Br)**

The outer part consists of all the electrons outside of the core.

The core notation for Rubidium would be... [Kr]5s^1

Do you see how all the numbers and letters that should've been before 5s^1 was replaced by Kr?

Another example... the core notation for Chlorine would be... [Ne]3s^2 3p^5

In electron configuration, valence electrons are those that aren't in the core, or in d- and f- subshells.

Ex/ Determine number of valence electrons by looking at core notation of Selenium.

[Ar] 4s^2 3d^10 4p^4

There are 6 valence electrons because of the 4s^2 and the 4p^4 part of the configuration. 3d^10 doesn't count because its a d- subshell.

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