Electronic+Configuration

=Topic #3 - Electron Configurations= The Bohr model of the atom was developed in 1913 by the Danish physicist Niels Bohr. It tried by explain why, when hydrogen atoms were excited (given extra energy), they emitted light of only certain wavelengths. According to classical physics, they should emit light of all wavelengths.
 * == Bohr model ==

The central proposition of the Bohr model was that the electron in a hydrogen atom moves around the nucleus only in certain allowed circular orbits.

Bohr developed an equation for showing the energy levels available to the hydrogen atom:

E = -2.178 x 10^18 J (Z^2/n^2) 

Many scientists had high hopes for Bohr’s theory, but unfortunately it could not be made to work for any atoms other than hydrogen, and was eventually discarded in favor of the quantum mechanical model of the atom.

** When a sample of hydrogen gas receives a spark (which contains a lot of energy), the H2 molecules absorb energy, and some of the H-H bonds are broken. This produces hydrogen atoms that are excited, meaning that they contain excess energy. The atoms release this energy by emitting light of various wavelengths, to produce something called an emission spectrum** Now, when white light is passed through a diffraction grating such as a prism, it is split up and you end up with a continuous spectrum containing all wavelengths of visible light. However, when light-emitting hydrogen gas (hot gas) is placed in front of a prism, you end with only a few lines of light of different colors, each corresponding to a discrete wavelength. This is called a **line spectrum or emission spectrum.** You can also end up with something called an absorption spectrum.
 * == Emission spectra ==



The attempt to explain this phenomenon to led to many advances in physics, such as the Bohr model of the atom.

In the quantum mechanical model of the atom, each orbital has a number of properties, which are described by the four quantum numbers.
 * == Quantum numbers ==

The __principle quantum number__ is n. N is an integer with the values 1, 2, 3….. It is related to the size and energy of an orbital. A high value for n means a larger, higher-energy orbital.

The __angular momentum quantum number__ is l. l has integer values from 0 to n-1 for each value of n. It is related to the shape of the orbitals.

The __magnetic quantum number__ is m. It is an integer with a value between l and -l, including 0. It is related to the orientation of the orbital in space related to the other orbitals in the atom.

The __electron spin number__ is ms. It can be either ½ or -1/2.

Value of l (__angular momentum quantum number__) and Letter Used 0-->s 1-->p 2-->d 3-->f 4-->g
 * == S,p,d,f orbitals ==







**#1 Pauli's exclusion principle** – In a given atom, __no two electrons can have the same set of four quantum numbers__. This means that since electrons in the same orbital have the same first three quantum numbers, they must have different electron spin values. Since there are only two electron spin values, an orbital can only have two electrons.
 * == Three rules to fill orbitals ==

**#2 Aufbau principle** – Aufbau is German for “building up.” As protons are added one by one to the nucleus to build up the elements, electrons are simply added to orbitals as in hydrogen,

**#3 Hund's rule** – The lowest energy configuration is the one having the maximum number of unpaired electrons allowed by the Pauli exclusion principle in a particular set of degenerate orbitals. (“Degenerate” means “having the same energy”)

=﻿Review Questions=
 * 1) 1 Give the electron configurations for sulfur (S), cadmium (Cd), hafnium (Hf), and radium (Ra) using the periodic table. (p. 308)


 * 1) 2 What are the possible values for the quantum numbers //n//, //l//, and //ml//? (pg. 322)


 * 1) 3 Which of the following orbital designations are incorrect: 1//s,// 1//p//, 7//d, 9s, 3f, 4f, 2d. (//pg. 323)

//#//4 Which of the following sets of quantum numbers are not allowed in the hydrogen atom? For the sets of quantum numbers that are incorrect, state what is wrong in each set. a. //n = 3, l = 2, ml = 2// b. //n = 4, l = 3, ml = 4// c. //n = 0, l = 0, ml = 0// d. //n = 2, l = -1, ml = 1 (//pg. 323)

a. //n = 3, l = 3, ml = 0, ms = -1/2// //b. n = 4, l = 3, ml = 2, ms = -1/2// //c. n = 4, l = 1, ml = 1, ms = +1/2// //d. n = 2, l = 1, ml = -1, ms = -1// //e. n = 5, l = -4, ml = 2, ms - +1/2// //f. n = 3, l = 1, ml = 2, ms = -1/2// (pg. 323)
 * 1) 5 Which of the following sets of quantum numbers are not allowed? For each incorrect set, state why it is incorrect.


 * 1) 6 How many orbitals in an atom can have the designation 5p, 3d(z^2), 4d, n=5, n=4? (pg. 323)


 * 1) 7 How many electrons in an atom can have the designation 1p, 6d(x^2-y^2)