Atomic+Theory+and+the+Periodic+Table

=__Atomic Theory and Periodic Trends__=
 * 1) Dalton's Theory
 * 2) Periodic Table and Periodicity
 * 3) Electron Theories and Electron Configuration
 * 4) Intermolecular Forces
 * 5) Electromagnetic Spectrum
 * 6) Flame Test Colors

= = =Dalton's Theory = In 1808, John Dalton developed the atomic model of matter that underlies modern chemistry. Three of the main postulates of modern atomic theory, all of which Dalton suggested in a somewhat different form, are stated below.
 * 1) //An element is composed of tiny particles called atoms//. All atoms of a given element show the same chemical properties. Atoms of different elements show different properties.
 * 2) //In an ordinary chemical reaction, atoms move from one substance to another, but no atom of any element disappears or is changed into an atom of another element.//
 * 3) //Compounds are formed when atoms of two or more elements combine.// In a given compound, the relative number of atoms of each kind are definite and constant. In general, these relative numbers can be expressed as integers or simple fractions.

On the basis of Dalton's theory, the atom can be defined as the smallest particle of an element that can enter into a chemical reaction.

[|Go here] for descriptions of other atomic models and theories

media type="custom" key="3957545"Travis DePriest. Learn All About Bohr's Atomic Model...and how to back 180 a stair set!!!

=Periodic Trends = First periodic table was proposed by Dmitri Mendeleev in 1869. This primitive version was ordered by atomic mass as opposed to the modern periodic table being organized by atomic number(created by Henry Mosely). This table contains numerous "trends" or repetitions of chemical properties: >
 * Atomic Radius
 * Ionic Radius
 * Ionization Energy
 * Electronegativity

Atomic Radius
Measured as half the distance from one nucleus to the next(the size of the atom). The value of the atomic radius becomes smaller as you move from left to right across a period. This occurs because as the nucleus gains protons, its attractive forces become stronger, and it pulls the outer electrons closer to the center of the atom. Atomic Radius also becomes larger the farther you move down a group because the more electrons and energy levels an atom has, the greater its size. For these reasons, Francium, or Fr, has the largest atomic radius, and Helium, He, has the smallest atomic radius.

Ionic Radius
This is defined as the size of an elements ions(charged atoms resulting from a change in electrons). Ionic Radius shares the same trends as atomic radius. Positive ions are smaller than their parent atoms because an unequal number of protons and electrons favors pull of nucleus. This often results in the loss of an entire energy level. On the other hand, negative ions are larger than their parent atoms and because there are more electrons present, more can pull away and repel each other, causing the size of the ion to increase.

Ionization energy
Defined as the amount of energy that is required to remove the outermost electron from an atom. Since energy is absorbed when an electron is removed, the value is always positive and the larger the value, the harder it is to remove an electron. Ionization energy increases as you more farther right on a period, because the closer you get to the noble gasses, the closer elements become to having a full 8 valence electrons. In addition, it decreases as you move down a group. This happens because shielding effects inner electrons and shields outer electrons from the pull of the nucleus. Therefore, Helium, He, has the highest ionization energy.

Electronegativity
The tendency of an atom to attract electrons to itself while bonding. Electronegativity shares the same trends as Ionization energy with one exception. When you get to the group of noble gases, the electronegativity sharply decreases. This occurs because once an element has a stable octet, it no longer has an attraction to other electrons. The electronegativity decreases as you move down a group because shielding increases. Therefore, Fluorine, F, is the most electronegative element.

=Electron Theories and Electron Configuration =
 * __Quantum Numbers;__** refer to the exact address of electrons. For example, four quantum numbers are used to specify the quantum state of an electron orbiting the nucleus of an atom: one characterizes its basic orbital energy level (principal or first quantum number), one the shape of its orbit (its azimuthal orbital or second quantum number), one the orientation of its orbit relative to other orbits (magnetic quantum number), and one its spin (spin magnetic number).

Go to [|this Site]

orbital- the area in which an electron is allowed to move.

n= energy level- distance from the nucleus. There are 7 energy levels. An electron can have 1-7 energy levels. l= Angular quantum number (sublevel)- determines the shape of the orbital. The general formula is 0...n-1. If n=1 then l=0 - the shape is a sphere If n=2 then l=0,1 - is like the shape of a peanut If n=3 then l=0,1,2 - is like a shape of a daisy If n=4 then l=0,1,2,3 - is like the shape of a firework

ml= magnetic quantum number. This refers to which way the orbital is facing, or the orientation of the orbital in space. If l=0 then ml=0 (1 orientation) If l=1 then ml= -1, 0, 1 (3 orientations) If l=2 then ml= -2, -1, 0, 1, 2 (5 orientations) If l=3 then ml= -3, -2, -1, 0, 1, 2, 3 (7 orientations)

ms= spin quantum number. This indicates the direction of the spin of the electrons within the orbital. The direction of the spin can either be + 1/2 or - 1/2. + 1/2 indicates a spin up - 1/2 indicates a spin down

How to read an electron configuration for Helium: 1s² 1. The large number "1" refers to the principle quantum number "n" which stands for the energy level. 2. The letter "s" stands for the angular momentum quantum number "l". Therefore there are two electrons of the helium electron occupy an "s" or spherical orbital. 3. The exponent "2" refers to the total number of electrons in that orbital or sub-shell. This shows that there are two electrons in the first energy level.

Orbitals and sublevels: An orbital is a space that can have up to two electrons. Each type of sublevel can have different numbers or orbitals, and therefore, a different number of electrons. The s sublevels have one orbital, which can hold up to two electrons. The p sublevels have three orbitals, each of which can hold 2 electrons, for a total of 6. The d sublevels have 5 orbitals, for a possible total of 10 electrons. The f sublevels, with 7 orbitals, can hold up to 14 electrons. When you write an electron orbital for an element, it is in its ground state. This means that the electrons are in the lowest energy state. Electrons can also be excited, and this is when electrons absorb enough energy to move to a higher energy level.

[|Go to this site] for an explanation on how to write electron configurations 1st level: 2 electrons 2nd level: 8 electrons 3rd level: 18 electrons
 * Any orbital can hold no more than 2 electrons* THe maximum number of electrons in each energy level 2n².

**Pauli Exclusion Principle-** No two electrons can have the same 4 quantum numbers; however, the first 3 numbers can be the same (n, l, ml). 2 electrons in the same orbital can not spin in the same direction. The last quantum number (ms- spin) can not be identical if the first 3 numbers are the same for 2 elements.

There are exceptions: chromium and copper and their electron configurations are: Copper: 1s22s22p63s23p64s13d10 Chromium: 1s22s22p63s23p64s13d5
 * Aufbau Principle-** Electrons will full up the orbitals with the lowest amount of energy before they fill in the orbitals that have more energy.
 * Hunds’s Rule-** Electrons will not share an orbital of the same energy if there is an empty orbital with that energy is available. Electrons will split up before they pair up, electrons will not share an orbital unless they have to.

Sample electron diagrams: []

Example drawing contributed by Zack Kattwinkel.

= Electromagnetic Spectrum = Definition: The Electromagnetic Spectrum is the distribution of electromagnetic radiation according to energy. The visible part of the spectrum can be subdivided by color with red at the long wavelength and violet at the short wavelength.

**__The Spectra__** v Absorption (black parts- element absorbs the wavelengths) v Emission (light is emitted when electrons move to an excited state to a ground state)  Light Emission and Absorption Tutorial Go to [|this site] Electromagnetic waves(radiation) are produced by the movement of electrically charged particles. They travel through empty space through air and other substances. Besides acting as waves, electromagnetic radiation can act like a path of particles called photons which have no mass. The higher the energy of a photon the shorter the wavelength.
 * __Electromagnetic Radiation__ **


 * = ||  ||= ** SYMBOL ** ||   ||= ** DEFINITION ** ||

||  ||= The distance between two consecutive crests or troughs (measured in meters or nanometers) ||
 * WAVELENGTH ** ||  ||= l
 * WAVELENGTH ** ||  ||= l

** FREQUENCY ** ||  ||= n  ||   ||= The number of wave cycles that pass a given point in unit time (measured in hertz- Hz, 1/sec, sec-1) ||

** AMPLITUDE ** ||  ||=  y  or A    ||   ||= The height of the waves or the difference between the center and the crest. ||

(measured in distance/time)
 * SPEED OF LIGHT ** ||  ||= C ||   ||= Frequency × wavelength
 * SPEED OF LIGHT ** ||  ||= C ||   ||= Frequency × wavelength


 * C= 2.998×108 m/s ** ||

__Wavelengths vs. Frequency__ v INVERSELY related- as wavelength increases, the frequency decreases v Shorter wavelength = Higher Frequency = More Energy (E) v You MUST CONVERT when wavelength is in nm ( 109 nm = 1 meter)

Relationship between wavelength, frequency, and speed of light (formula) ln= C   ** Planck’s Equation- relates energy with wavelength and frequency** ** E =h **** n **** or Ehc / **** l ** =Planck's constant= 6.626 x 10^-34 Joules x seconds  Bohr’s Equation - is used to find the energy of an electron in a specific energy level: **En =** = -RH/ n2 =Energy level **RH**= quantity of Rydberg constant (2.180×10-18 J)
 * E= Energy**
 * H**
 * En=** Energy of an electron **n**

=Flame Test Colors =

Go here for colors of elements in aqueous solutions, click on descriptive chemistry

What are they used for?
Flame Tests are used to identify the presence of a relatively small number of metal ions in a compound.

What happens when heat is added? The flame colors are produced from the movement of electrons in the metal ions of a compound. The heat cause electrons to gain energy and "jump" to higher energy levels. These "jumps" each have a specific amount of energy emitted(released) as light energy.

[|Go here] to watch a flame test demonstration

Gas flame || Gas flame seen through cobalt glass || Flame test on copper sulfate || Flame test on copper sulfate seen through cobalt glass || Flame test on sodium carbonate || Flame test on sodium carbonate seen through cobalt glass || <span style="border-color: rgb(0, 0, 255); border-width: 0px; display: inline-block; font-size: 0px; background-image: none; vertical-align: middle; width: 1px; height: 1px;"> Flame test on a lithium salt || <span style="border-color: rgb(0, 0, 255); border-width: 0px; display: inline-block; font-size: 0px; background-image: none; vertical-align: middle; width: 1px; height: 1px;"> Flame test on a potassium salt || = HOW DO YOU DO A FLAME TEST?? = =<span style="color: rgb(0, 0, 255);"> <span style="font-size: 80%; color: rgb(0, 0, 255);">//You can try a classic wire loop method// = =<span style="font-size: 80%; color: rgb(0, 0, 255);">First you need a clean wire loop of platinum or nickel-chromium. The loop must be cleaned between tests. Dip the clean loop in either a powder or solution of an ionic metal salt. The loop with the sample is then placed in the clear or blue part of a flame and a color can be observed. Identify the element with its color using the charts above. The test will only work, however, if you dip the loop in cleansing solution and use distilled water. //Wooden Splint or Cotton Swab Method// = =<span style="font-size: 80%; color: rgb(0, 0, 255);">Wooden splints or cotton swabs can be used instead of the wire loop and are much more inexpensive. Soak the splint overnight in distilled water. Pour out the water and rinse the splints with clean water, being sure you do not contaiminate the splint in any way. Take a damp splint or cotton swab that has been moistened in distilled water, dip it in the sample to be tested, and wave the splint or swab through the flame. Do not hold the sample in the flame or you are risking igniting the splint on fire. Use a new splint or swab for each test. = =<span style="font-size: 80%; color: rgb(0, 0, 255);">contribution: Jennifer Samuels = =<span style="font-size: 80%; color: rgb(0, 0, 255);"> __Citations__= __http://en.wikipedia.org/wiki/Flame_test__ Chemistry: Principles and Reactions by Masterton & Hurley http://www.clickandlearn.org/Gr9_Sci/atoms/modelsoftheatom.html http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch6/quantum.html#aufbau http://www.mpcfaculty.net/mark_bishop/complete_electron_configuration_help.htm http://www.fordhamprep.org/gcurran/sho/sho/lessons/lesson36.htm
 * ~ Symbol ||~ Name ||~ Color ||
 * As || [|Arsenic] || [|Blue] ||
 * B || [|Boron] || [|Bright green] ||
 * Ba || [|Barium] || [|Pale/Apple green] ||
 * Ca || [|Calcium] || Brick red ||
 * Cs || <span style="color: rgb(0, 0, 238);">__Cesium__ || [|Blue] ||
 * Cu(I) || [|Copper](I) || [|Blue] ||
 * Cu(II) || [|Copper](II) (non-[|halide]) || [|Green] ||
 * Cu(II) || [|Copper](II) ([|halide]) || [|Blue-green] ||
 * Fe || [|Iron] || [|Gold] ||
 * In || [|Indium] || [|Blue] ||
 * K || [|Potassium] || [|Lilac] ||
 * Li || [|Lithium] || [|Red] ||
 * Mn(II) || [|Manganese](II) || [|Yellowish green] ||
 * Mo || [|Molybdenum] || [|Yellowish green] ||
 * Na || [|Sodium] || [|Intense yellow] ||
 * P || [|Phosphorus] || [|Pale bluish green] ||
 * Pb || [|Lead] || [|Blue] ||
 * Rb || [|Rubidium] || [|Red-violet] ||
 * Sb || [|Antimony] || [|Pale green] ||
 * Se || [|Selenium] || [|Azure blue] ||
 * Sr || [|Strontium] || [|Crimson] ||
 * Te || [|Tellurium] || [|Pale green] ||
 * Tl || [|Thallium] || [|Pure green] ||
 * Zn || [|Zinc] || [|Bluish green] ||
 * [[image:http://upload.wikimedia.org/wikipedia/commons/thumb/4/47/Flametest--.swn.jpg/72px-Flametest--.swn.jpg width="72" height="120" caption="external image 72px-Flametest--.swn.jpg" link="http://en.wikipedia.org/wiki/File:Flametest--.swn.jpg"]]
 * [[image:http://upload.wikimedia.org/wikipedia/commons/thumb/e/e5/Flametest--Na.swn.jpg/72px-Flametest--Na.swn.jpg width="72" height="120" caption="external image 72px-Flametest--Na.swn.jpg" link="http://en.wikipedia.org/wiki/File:Flametest--Na.swn.jpg"]]