Atomic+Structure

Part I. Electromagnetic Radiation
Electromagnetic radiation is energy radiated as a wave as a result of the motion of electric charges.

Electromagnetic waves are created when an electric field couples with a magnetic field. There are different types of electromagnetic waves in the electromagnetic spectrum. In order from longest wavelength to shortest wavelength, the different types of electromagnetic waves include: radio waves, microwaves, infrared, visible light, ultra violet, x-rays, and gamma rays. Wavelength is the distance between one peak (or crest) of a wave to the next of that same wave, and is generally denoted by the Greek letter lambda, λ.

Retrieved from http://science.hq.nasa.gov/kids/imagers/ems/waves3.html

In the visible light portion of the electromagnetic spectrum, red has the longest wavelength and violet has the shortest wavelength. Waves with longer wavelengths have shorter frequencies, and vice versa. Frequency is the number of peaks that would pass through a given stationary point in a certain amount of time, and is denoted by the letter ‘v’. Typically, waves with shorter wavelengths and higher frequencies contain more energy (which would make Gamma rays the most powerful type of wave in the diagram above). Wavelength and frequency can be connected through the equation c =λv, with ‘c’ being the speed of light (2.998 x 108 m/s) with ‘λ’ being wavelength, in meters (m) with ‘v’ being frequency, in l/s or Hz

Another useful equation in relation to wavelengths is Plank’s equation:

E=hv with ‘E’ being the energy of a photon (a photon is defined as a ‘dicrete bundle of electromagnetic (light) energy’ and does not have mass or an electric charge but does have energy and momentum and is in constant motion) with ‘h’ being Plank’s constant, 6.26 x 10-34 Joules/second with ‘v’ being frequency

If an ‘n’ is added into the equation, making it E=nhv, you can then calculate the amount of energy in any number of photons of a certain type of light, rather than in just one photon like in the original equation.

The two equations mentioned above, c =λv and E=nhv can be combined into one equation:

E= (hc) / λ, with all of the variables defined above. This combinatory equation streamlines the process of solving for any of these variables. It can also be noted that this equation can show that, as mentioned above, the longer the wavelength the lower the amount of energy.

BOHR ATOM
Niels Bohr (7 October 1885 – 18 November 1962) was a Danish physicist. He proposed a model of the atom that consisted of seven different energy levels. His model looked something like this:



Retrieved from http://www.faculty.virginia.edu/consciousness/images/bohr%20model.gif

It is similar to the planetary orbits around the sun. The model consists of a small, positively charged atom surrounded by electrons orbiting around the nucleus on 7 different energy levels. Bohr also proposed that when an atom is heated or excited in any other manner, it will ‘jump’ from one energy level to a higher one, then fall back down to its original energy level and in this process emits wavelengths of light. Through use of this model gases and vapors can be identified because there is only a specific amount of energy each can emit and this specific amount gives off a particular ray of light. Here are the specific emissions of a few atoms:

Retrieved from http://astronomy.nmsu.edu/jojohnso/images/model-bohr-3.jpg

Part II. The rydberg equation

 * E=-Rh(1/Nf^2-1/Ni^2)**

Part III. **Orbitals** (pictures from [|http://www.d.umn.edu/~pkiprof/ChemWebV2/AOs/ao4.html]) an orbital is an area around the nucleus that represents the probable location of an electron

There are 4 types of orbitals, each of which can be arranged in a number of different ways the first is the s orbital, which holds 2 electrons, and can only be arranged in one way the s orbital looks like a sphere around the nucleus of an atom the second is the p orbital, which has two different regions, and 3 different arrangements, holding 6 electrons the p orbitals 3 parts consist of pairs of regions that exist on opposite sides of eachother the d orbitals has 5 different orbital types, and holds 10 electrons 4 of the d orbitals are made up of 4 seperate regions which make an X, while the 5th is two polar opposite regions with a ring around them the f orbital has 7 different orbitals, holding 14 electrons ..f orbitals are really weird the more electrons an atom has, the more orbitals it fills, but the electrons will fall first into orbitals which have the lowest levels of energy

the first s orbital has the least energy of all, but the energy levels do not fall in place strictly by letter from least energy to most the order is as follows : 1s,2s,2p,3s,3p,4s,3d,4p,5s,4d,5p,6s,4f,5d,6p,7s,5f,6d

the way electrons fill the orbitals can be though of as people filling a hotel electrons, like people, will always prefer their own "bed", meaning they will take a single place of an orbital instead of doubling up unless all spots are taken. the most convenient spot would also be on the lowest level for a person at a hotel, or an orbital with the lowest level for an electron

for john leelike - V. Quantum numbers a. definitions of n, l, ml, ms b. quantum numbers and orbital energy

for james kelly- you smell like doodoo

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Part IV. Electron Configuration
Through the application of the **//Aufbau Principle//**, the electron configuration of an atom can be determined. The atom continues to "build up" (translation of german word aufbau is construction) as more and more electrons are added.

The Principle states that electrons will fill into the lowest energy orbital possible, before filling in higher energy subshells. For example, an electron will fill the 2s subshell before the 2p and then join the 3s after that.

Using the information of the **//Aufbau Principle//**, the //Aufbau Diagram// is created:

The //**Orbital Energy Diagram**// shows the way in which the electrons will begin to fill up the orbitals. Here is the diagram: Different elements will fill in the diagram differently. For instance, hydrogen, with one electron, will have one half arrow in the 1s orbital, written as 1s^1. Helium, on the other hand, with two electrons, will have two half arrows in the 1s orbital, written as 1s^2. The same trend continues as bigger elements have more electrons.


 * //Hund's Rule//**: every orbital in a subshell is singly occupied with one electron before any one orbital is occupied twice, and all electrons in singly occupied orbitals have the same spin.

Developed by a German scientist, Frederick Hund, //Hund's Rule// can predict the order in which electrons occupy suborbitals. Electrons space themselves as far apart as possible, occupying all available vacant suborbitals before pairing up with another electron. The unpaired  electrons all have the same spin quantum number. You can think of Hund's rule as if the electrons are people and the suborbitals are seats on a bus. Everyone will occupy an empty seat before they decide to sit with someone else.

The **//Pauli Exclusion Principle//**, formulated by Wolfgang Pauli in 1925, states that 2 electrons in the same region of space must have different spins. This causes the successive buildup of orbitals around the nucleus, preventing matter from collapsing into incredibly dense matter.

The principle applies to any two fermions, such as protons, electrons, or neutrons, stating that they cannot occupy the same quantum state. The **//Modern Periodic Table//** is arranged by electron configuration. As the electrons increase across a period, the electron configuration of the elements increase. For instance, Hydrogen is 1s^1, Helium is 1s^2, Lithium is 1s^2 2s^1, etc. (shown below) Below are the elements with their electron configuration in the periodic table: 1 ||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||  || [|2He] 2 || + 2s: 2p: || [|3Li] 1 - || [|4Be] 2 - ||||||||||||||||||||||||||||||||||||||||||||||||  || [|5B] 2 1 || [|6C] 2 2 || [|7N] 2 3 || [|8O] 2 4 || [|9F] 2 5 || [|10Ne] 2 6 || + 3s: 3p: || [|11Na] 1 - || [|12Mg] 2 - ||||||||||||||||||||||||||||||||||||||||||||||||  || [|13Al] 2 1 || [|14Si] 2 2 || [|15P] 2 3 || [|16S] 2 4 || [|17Cl] 2 5 || [|18Ar] 2 6 || + 4s: 3d: 4p: || [|19K] 1 - - || [|20Ca] 2 - - ||||||||||||||||||||||||||||  || [|21Sc] 2 1 - || [|22Ti] 2 2 - || [|23V] 2 3 - || [|24Cr] 1 5 - || [|25Mn] 2 5 - || [|26Fe] 2 6 - || [|27Co] 2 7 - || [|28Ni] 2 8 - || [|29Cu] 1 10 - || [|30Zn] 2 10 - || [|31Ga] 2 10 1 || [|32Ge] 2 10 2 || [|33As] 2 10 3 || [|34Se] 2 10 4 || [|35Br] 2 10 5 || [|36Kr] 2 10 6 || + 5s: 4d: 5p: || [|37Rb] 1 - - || [|38Sr] 2 - - ||||||||||||||||||||||||||||  || [|39Y] 2 1 - || [|40Zr] 2 2 - || [|41Nb] 1 4 - || [|42Mo] 1 5 - || [|43Tc] 2 5 - || [|44Ru] 1 7 - || [|45Rh] 1 8 - || [|46Pd] - 10 - || [|47Ag] 1 10 - || [|48Cd] 2 10 - || [|49In] 2 10 1 || [|50Sn] 2 10 2 || [|51Sb] 2 10 3 || [|52Te] 2 10 4 || [|53I] 2 10 5 || [|54Xe] 2 10 6 || + 6s: 4f: 5d: 6p: || [|55Cs] 1 - - - || [|56Ba] 2 - - - || [|57La] 2 - 1 - || [|58Ce] 2 1 1 - || [|59Pr] 2 3 - - || [|60Nd] 2 4 - - || [|61Pm] 2 5 - - || [|62Sm] 2 6 - - || [|63Eu] 2 7 - - || [|64Gd] 2 7 1 - || [|65Tb] 2 9 - - || [|66Dy] 2 10 - - || [|67Ho] 2 11 - - || [|68Er] 2 12 - - || [|69Tm] 2 13 - - || [|70Yb] 2 14 - - || [|71Lu] 2 14 1 - || [|72Hf] 2 14 2 - || [|73Ta] 2 14 3 - || [|74W] 2 14 4 - || [|75Re] 2 14 5 - || [|76Os] 2 14 6 - || [|77Ir] 2 14 7 - || [|78Pt] 1 14 9 - || [|79Au] 1 14 10 - || [|80Hg] 2 14 10 - || [|81Tl] 2 14 10 1 || [|82Pb] 2 14 10 2 || [|83Bi] 2 14 10 3 || [|84Po] 2 14 10 4 || [|85At] 2 14 10 5 || [|86Rn] 2 14 10 6 || + 7s: 5f: 6d: 7p: || [|87Fr] 1 - - - || [|88Ra] 2 - - - || [|89Ac] 2 - 1 - || [|90Th] 2 - 2 - || [|91Pa] 2 2 1 - || [|92U] 2 3 1 - || [|93Np] 2 4 1 - || [|94Pu] 2 6 - - || [|95Am] 2 7 - - || [|96Cm] 2 7 1 - || [|97Bk] 2 9 - - || [|98Cf] 2 10 - - || [|99Es] 2 11 - - || [|100Fm] 2 12 - - || [|101Md] 2 13 - - || [|102No] 2 14 - - || [|103Lr] 2 14 - 1 || [|104Rf] 2 14 2 - || [|105Db] 2 14 3 - || [|106Sg] 2 14 4 - || [|107Bh] 2 14 5 - || [|108Hs] 2 14 6 - || [|109Mt] 2 14 7 - || [|110Ds] 1 14 9 - || [|111Rg] 2 14 9 - || [|112Uub] 2 14 10 - || [|113Uut] 2 14 10 1 || [|114Uuq] 2 14 10 2 || [|115Uup] 2 14 10 3 || [|116Uuh] 2 14 10 4 || [|117Uus] 2 14 10 5 || [|118Uuo] 2 14 10 6 ||
 * ~ 1 H [|hydrogen] : 1s1 ||
 * 1s1 ||  ||   ||   ||   ||   ||   ||   ||   ||   ||   ||   ||   ||   ||   ||   ||   ||   ||   ||
 * 1 ||||  ||||||   ||||||||   ||||||||   ||||||   ||||   ||
 * ~ 2 He [|helium] : 1s2 ||
 * 1s2 ||  ||   ||   ||   ||   ||   ||   ||   ||   ||   ||   ||   ||   ||   ||   ||   ||   ||   ||
 * 2 ||||  ||||||   ||||||||   ||||||||   ||||||   ||||   ||
 * ~ 3 Li [|lithium] : 1s2 2s1 ||
 * 1s2 || 2s1 ||  ||   ||   ||   ||   ||   ||   ||   ||   ||   ||   ||   ||   ||   ||   ||   ||   ||
 * 2 |||| 1 ||||||  ||||||||   ||||||||   ||||||   ||||   ||
 * 1s: || [|1H]
 * [He]
 * [Ne]
 * [Ar]
 * [Kr]
 * [Xe]
 * [Rn]

Part V. Quantum numbers
//**m**//**l** is the "magnetic quantum number"
 * //n//** is the "principle quantum number"
 * possible values = 1, 2, 3, 4, ...
 * larger values of //n// are higher in energy
 * larger values represent a larger (further from the nucleus) probability density
 * //l//** is the "angular momentum quantum number"
 * possible values = 0, 1, 2, 3, ... //n// – 1
 * The values have alternate names you also need to know:
 * //l// = 0 is //s//
 * //l// = 1 is //p//
 * //l// = 2 is //d//
 * //l// = 3 is //f//
 * larger values of //l// represent slightly larger values of energy
 * //l// corresponds to the shape of the probability density, the orbital shapes are shown above
 * possible values = –//l// ..., –2, –1, 0, +1, +2, ... +//l//
 * all //m//l values have the same energy
 * the values correspond to the orientation of probability density
 * The important part of this is knowing how many orientations are possible:
 * For //l// = 0, //m//l = 0—there is only one way the sphere is oriented or one //s// orbital
 * For //l// = 1, //m//l = –1, 0, or +1—there are three //p// orbitals
 * For //l// = 2, //m//l = –2, –1, 0, +1, +2—there are five //d// orbitals
 * For //l// = 3, //m//l = –3, –2, –1, 0, +1, +2, +3—there are seven //f// orbitals
 * //m//****s** is the "spin quantum number"
 * possible values = +1/2 or –1/2
 * both values are the same energy; however, it is lower if electrons are not paired and spinning in the same direction
 * the value shows the spin

You should now be able to use the quantum numbers to describe and compare electrons.

Here's an example:

Rank the following electrons with quantum numbers (//n//, //l//, //m//l, //m//s) from lowest energy to highest energy.


 * (2, 1, 1, +1/2)
 * (1, 0, 0, –1/2)
 * (4, 1, –1, +1/2)
 * (4, 2, –1, +1/2)
 * (3, 2, –1, +1/2)
 * (4, 0, 0, +1/2)
 * (2, 1, –1, +1/2)
 * (3, 1, 0, +1/2)

Lowest energy has the lowest //n// + //l// value. Electrons 3 and 5 have the same value (5), so the one with the lowest //n// (5) is lower in energy. Electrons 8 and 6 also have the same //n// + //l// value, and 8 is lower in energy. Electrons 1 and 7 have the same values of //n// and //l// so each electron has the same energy.
 * Solution:**


 * low energy** 2 < 1 = 7 < 8 < 6 < 5 < 3 < 4 **high energy**