Chemistry Helpers
Atomic Theory Definitions Atomic Theory History
Atomic Structure


Subatomic Particles: An atom is the smallest unit of an element. It consists of three major particles:

1 amu
1 amu
No charge
1/1836 amu

Note:        a.m.u.= atomic mass unit
Atomic Models:
1)     Dalton’s Theory
a) All elements are composed of indivisible atoms
b) All atoms of a given element are identical
c) Atoms of different elements are different; that is, they have different masses
d) Compounds are formed by the combination of atoms of different elements    
2) JJ Thomsonused a cathode ray tube to show smaller units that make up an atom. The ray was deflected a certain way by a magnetic field, so he concluded that the ray was formed by particles and that the particles were negatively charged. The only source available for the particles was the atoms present. Therefore, Thomson theorized that an atom contains small, negatively charged particles called electrons.   This theory is referred to as the Plum Pudding Model. In this model, the mass of the rest of the atom was evenly distributed and positively charged, taking up all of the space not occupied by the electrons.
3) Rutherford’s Gold Foil Experiment: If electrons are present in atoms, what makes up the rest of the atom? Ernest Rutherford (with other scientists) directed positively charged alpha particles at a piece of gold foil. If the plum pudding theory were correct, all of the alpha particles would pass through the foil with just a few being slightly deflected. What they had expected did in fact occur, but some of the alpha particles were greatly deflected and some even bounced back. Rutherford concluded that:
a)        Majority of the volume of an atom is empty space
b)        Atoms have a dense positively charged central core
4) The Bohr Model             1 or K-shell = max 2 e-
                                                 2 or L-shell = max 8 e-
                                                 3 or M-shell = max 18 e-
                                                 4 or N-shell = max 32 e- 
5) The Orbital Model:
            *Important Definitions:
a)     Principle Energy Level: Region around the nucleus in which electrons can be found. (The closer to the nucleus, the lower the energy)
b)     Quanta: Small amount of energy that a(n) electron can absorb or release as it moves through principle energy levels
c)      Ground Stateall electrons fill lowest energy levels before higher energy levels begin to fill
d)     Excited Stateone or more electrons fills a higher energy level before the lower ones are filled
e)     Spectral Lines: As electrons at higher energy levels (excited electrons) fall back to their normal energy levels (ground state) they release energy in the form of the spectrum. ROYGBIV
Electron Configurations
Looking at the periodic table of elements, you will notice numbers at the bottom of each element. These numbers represent the electron configuration of the element (the address of the electrons). The _Period___ represents the number of principle energy levels (orbitals) present. The _Group_ represents the sublevel for each principle energy level.
Period 1 = 1 shell                       Group 1 & Group 2 = ‘s’ sublevel (max – 2 electrons)
Period 2 = 2 shells                     Group 13 – Group 18 = ‘p’ sublevel (max – 6 electrons)
Period 3 = 3 shells                     Group 3 – Group 12 = ‘d’ sublevel (max – 10 electrons)
Period 4 = 4 shells                     Lanthanum & Actinum Series = ‘f’ sublevel (max – 14 e)
Period 5 = 5 shells   
Period 6 = 6 shells   
Example: Write the electron configuration for the following:
Na       2-8-1
S         2-8-6
Kr        2-8-18-8
1)     Valence Electrons and Electron Dot Diagrams
a) Valence Electrons are electrons that fill the outermost principle energy level of an atom
Example: Mg 2-8-2 has 2 valence electrons
                  Ne 2-8 has 8 valence electrons
Valence electrons are largely responsible for an element’s chemical and physical properties.
b) Electron Dot Diagrams: the term Kernel refers to all of the non-valence electrons and the nucleus of an atom. The Kernel is represented by the element’s symbol, valence electrons are represented by dots.
Differences between atoms
1)     Atomic Number: the number of Protons in the nucleus of an atom. (Atom by definition is an electrically neutral particle, so this must also be equal to the number of Electrons).
2)     Mass Number: the number of protons plus neutrons
Question: Why are there fractional mass numbers on the periodic table? (ex: Na, O2, …) Answer: Due to the existence of isotopes.
Note: Atomic Symbols: One or two letters, 1st is always capital, the 2nd is always lower case.
3) Isotope: Atoms if the same element having the same number of protons, but different number of neutrons 
            Example: Isotopes of Hydrogen

Mass Number

Calculating Isotopes (weighted atomic mass)
·        1)   Take the percent of each isotope and convert it back to a decimal (¸ 100)
·        2)   Multiply the decimal by the mass number
·        3)  Add the numbers together to get the Weighted Atomic Mass
            a.         C-12   99%                C-14     1%
                        (.99 x 12) + (.01 x 14) = 12.02                     or         .99 x 12 =       11.88
                                                                                                            .01 x 14 =        0.14
            b.         Mg-26   1.75 %          Mg-24 98.25%
                        (0.0175 x 26) + (.9825 x 24) = 24.035
            c.         Cl-35    75%               Cl-37   25%
                        (.75 x 35) + (.25 x 37) = 35.5
            **         This is why we can’t round Chlorine to 35
4) Ion: Atoms of the same elements having the same number of protons, but different number of electrons
            Example: Na ->  Na+1                                                F ->  F-1
a)     Cations: positive charge formed by losing electrons
Example: sodium loses its outermost valence electron to become positively charged
b)     Anions: negative charge formed by gaining electrons
Example: fluorine gains an electron to fill its outermost shell becoming negatively charged

Write Your Questions Below

what is quantum theory of radiation?????
July 14, 2012

Max Planck defined Quantum Theory of Radiation.  Radiant energy is absorbed or released in the form of quanta whose energy is dependent on the frequency. The amount of energy can be calculated using the frequency and Planck's constant in the formula       Energy = (Planck's constant) (frequency)

July 16, 2012 -  Replied By Expert

the element with highest ionisation enthalpy is---------?????
July 13, 2012

Ionization enthalpy decreases as you go down and increases as you go across the periodic table.

July 14, 2012 -  Replied By Expert

what is ionisation enthalpy??????????
July 13, 2012

The energy required to remove the last electron from an atom.

July 14, 2012 -  Replied By Expert

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