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Atomic Theory
http://www.skanschools.org/webpages/rallen/
The Evolution of the Atomic Model
Atom Basic building block of matter
Cannot be broken down chemically
A single unit of an element
Dalton (1803)Known as the founder of the atomic theory
Dalton’s Postulates:1. All matter is composed of indivisible particles called atoms2. all atoms of a given element are identical in mass and properties. Atoms of different elements have different masses and different properties3. Compounds are formed by a combination of 2 or more atoms4. Atoms cannot be created, destroyed, or converted into other kinds of atoms during chemical reactions
1. Spherical
2. Uniform Density
Cannonball Theory/Model
J.J. Thomson (1897)
Used a cathode ray tube with charged particle field (+/-)
Cathode ray deflected by negative electrode toward positive electrode
Discovered subatomic particle called the ELECTRON
SmallNegatively charged
J.J Thomson
Plum Pudding Theory/Model
Positive “pudding”
Negative electrons embedded (just like raisin bread)
Rutherford (1909)
Conducted the GOLD FOIL EXPERIMENT where he bombarded a thin piece of gold foil with a positive stream of alpha particles
Expected virtually all alpha particles to pass straight through foil
Most passed through but some were severely deflected
Conclusion led to…
Rutherford’s Experiment
RutherfordNuclear Theory/Model
1. The atom is mostly empty space
2. At the center of the atom there is a DENSE, POSITIVE CORE called the NUCLEUS
**Provided no information about electrons other than the fact that they were located outside the nucleus**
Neils Bohr (1913)
Bohr Model or Planetary Model
1. Electrons travel AROUND the nucleus in well defined paths calls ORBITS (like planets in the solar system)
2. Electrons in different orbits possess different amounts of energy
Neils Bohr 3. Absorbing/Gaining a certain amount of energy causes electrons to jump to a higher energy level or an excited state
When excited electrons emit/lose a certain amount of energy which causes electrons to fall back to a lower energy level or the Ground State
Wave-Mechanical/Cloud Model
Modern, present day model
Electrons have distinct amounts of energy and move in areas called orbitals
Orbital= an area of high probability for finding an electron (not necessarily a circular path)
Developed after the famous discovery that energy can behave as both waves & particles
Many scientists have contribute to this theory using X-Ray diffraction
Vocabulary(of the Periodic
Table)Atomic#= the number of protons in every atom of the element (NEVER CHANGES!)
Atomic Mass= average mass of all the isotopes of an element
Element Symbol= the letter(s) used to identify an element
Subatomic
Particle
Charge
Relative Mass
Location
Symbol How to Calcula
te
Proton +1 1 amu Nucleus
or Look at atomic
#
Neutron 0 1 amu Nucleus
Mass # - Atomic
#
Electron -1 1/1836 amu
Outisde nucleus
P= e(in
neutral atoms)
**Nucleons= Protons and Neutrons (any subatomic particle found within the nucleus)
Determining Subatomic
Particles (p, n, e)# of Protons = sum of the protons and neutrons in an atom of an element (Atomic Number)
# of Neutrons = Mass # - Atomic # or Mass # - # Protons
# of Electrons = sum of the protons and neutrons in an atom of an element ( equal to # protons in a neutral atom)
Determining Subatomic
Particles (p, n, e)Atomic Number:1. Look at the element symbol and locate on the periodic table
2. same as the # of protons or the nuclear charge
Mass Number = # Protons + # Neutrons
Nuclear Charge = # of protons or the atomic #P = nuclear charge
Atoms vs. IonsVocabulary
TermDefinition Example
Neutral Atom An atom with the same number of protons and electrons
P = e (no charge indicated)
Ion Two Types
Anion and Cation
Protons and electrons are not the same
Atoms vs. Ions
Anion Cation
aNion
An atom that has GAINED one or more electrons
e > p
NEGATIVE ION
ca+ion
An atom that has LOST one or more electrons
p > e
POSITIVE ION
Ions
Isotope
Atoms of the same element with different mass numbers; same atomic number, same number of protons, different number of neutrons
Isotopes
p = p = p = e = e = e = n = n = n =
Example 1: Carbon (12, 13, 14)
Isotopes
p = p = e = e = n = n =
Example 2: Uranium (238, 240)
U-240
Practice
1. Two different isotopes of the same element must contain the same number of
2. Two different isotopes of the same element must contain a different number of
a. protons b. neutrons c. electrons*
a. protons b. neutrons c. electrons*
Practice
3. Isotopes of a given element have
a. the same mass number and a different atomic numberb. the same atomic number and a different mass numberc. the same atomic number and the same mass number
*
Calculating Atomic Mass (for any
element)Atomic mass = the weighted average of an element’s naturally occurring isotopes
% abundance of isotope 1 x (mass of isotope 1)
% abundance of isotope 2 x (mass of isotope 2)
+ % abundance of isotope 3 x (mass of isotope 3)
Average Atomic Mass of the Element
Calculating Atomic Mass
Chlorine has two naturally occurring isotopes, Cl-35 (isotopic mass 34.9689 amu) and Cl-37 (isotopic mass 36.9659 amu). In the atmosphere, 32.51% of the chlorine is Cl-37, and 67.49% is Cl-35. What is the atomic mass of atmospheric chlorine?
Step 1: Multiply the mass of each separate isotope by its percent abundance
Example 1: The exact mass of each isotope is given
CL-35 = 34.9689 amu x = CL-37 = 36.9659 amu x =
(.6749)
(.3251) 12.0176
23.6005
Calculating Atomic Mass
Step 2: Add the products of all the calculated isotopes together from step 1
23.6005
+12.0176
35.6181
***This is your average atomic mass***
Calculating Atomic Mass
The element Carbon occurs in nature as two isotopes. Calculate the average atomic mass for Carbon based on the information below
C-12 = 98.89% C-13 = 1.11%
**Since the mass numbers were not given for either isotope, use the mass number
instead**
C-12 = 12
C-13 = 13
Calculating Atomic Mass
Step 1: Multiply the mass by the percent abundance
C-12 = 12 x (.9889) =
C-13 = 13 x (.0111) =
Step 2: Add the products together
11.8668
0.1443
*These are weighted masses*
11.8668+ 0.1443
12.0111
Calculating Atomic Mass
Practice: The element Boron occurs in nature as two isotopes. Calculate the average atomic mass for Boron, using the information below.
Isotope Mass % abundance
Boron-10 10.0130 amu 19.9 %
Boron-11 11.0093 amu 80.1 %
Average Mass of Boron = 10.8104
Practice: The element Hydrogen occurs in nature as three isotopes. Calculate the average atomic mass of Hydrogen.
Isotope % abundance
Protium 99.0%
Deuterium 0.6%
Tritium 0.4%
Average Mass= 1.014
Mass Number Atomic Mass
The MASS of ONE isotope of a given
element
The AVERAGE MASS of ALL isotopes of a
given element
Electron Configurations
The dashed chain of numbers found in the lower left hand corner of an element box
Tells the number of energy levels as well as the number of electrons in each level(how the electrons are arranged around the nucleus)
Electron Configuration
Is the representation of the arrangement of electrons distributed among the orbitals
Used to describe the orbitals of an atom in the ground state
Can be used to describe ionized atoms (cations and anions)
Many of the chemical and physical properties of elements can be correlated to their electron configuration
Electron Configurations
All electron configurations on the Periodic Table are NUETRAL ( p = e)
For IONS, add or subtract electrons from the LAST NUMBER in the electron configuration only
Electron Configuration
Orbitals
There are four types (s, p, d, and f)
They have different shapesEach orbital can hold a maximum of two electrons
P, d and f orbitals have different sublevels
Electron Configuration
Is unique to an element’s position on the periodic table
Energy level is determined by the period
Number of electrons is given by the atomic number
Electron Configuration
Orbitals on different energy levels are similar to each other but they occupy different areas in space
ex. 1s and 2s orbitals both have s orbital characteristics ( radial nodes, spherical volume, can only hold two electrons) but since they are found in different energy levels they occupy different spaces around the nucleus
Electron Configuration
Electrons fill orbitals to minimize energy
Electrons fill the principal energy levels in order of increasing energy
Electrons are getting further from the nucleus
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p
Electron Configuration
A single orbital can hold a maximum of two electrons
These electrons must have opposing spins
One electron spins up and the other spins down
Ensures they have different quantum numbersPauli Exclusion Principal
Electron Configuration
S subshell 1 orbital that can hold 2 electrons
P subshell 3 orbitals that can hold up to 6 electrons
D subshell 5 orbitals that can hold up to 10 electrons
F subshell 7 orbitals that can hold up to 14 electrons
Principal Energy Level (n)
Electron energy levels consisting of orbitals which designated s, p, d, or f.
Electron Configurations
Valence Electrons:
Electrons found in the OUTERMOST shell or orbital
Kernel Electrons:
INNER electrons (all non-valence electrons)
Orbital Notation
Orbital Notation
When filling in the orbital
Electrons fill the lowest vacancy levels first
When there’s more than one subshell at a particular energy level (ex. 3p or 4d) only one electron fills each subshell until each subshell has one electron. Then electrons start pairing in each subshell
Hund’s Rule
Every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin.
Energy Level Diagram: Oxygen.
Oxygen is atomic number 8.
8 protons How many electrons?
8 electrons
Bohr Diagrams
A method for showing electron location in an atom/ion
All electrons must be drawn
Look up the electron configuration and the period (row) the element is in on the periodic table
Draw a circle (nucleus) write in the number of protons and neutrons
Draw the shells (energy levels) around the nucleus
Bohr Models
Add the electrons
The first shell only holds 2 electrons
Now add the rest of the electronsAdd one at a time starting on the right side and going counter clockwise
Shell (energy level) Maximum number of electrons
1 22 83 184 325 506 727 98
2n2 n=Formula for maximum # of electrons per quantum number (or energy level)
Bohr ModelsExample:
Draw the Bohr Model for Carbon
Bohr ModelsExampleDraw the Bohr Model for Oxygen
Bohr ModelsExample
Draw the Bohr Model for Na+
E- configuration 2-8
Lewis Dot Diagrams
Only illustrates valence electrons
Write the element symbol
Look up the electron configuration (use the last number in the configuration- # of valence electrons)
Use either an X or dot to represent the electrons
Place that many electrons around the symbol at (12, 3, 6 and 9)
Lewis Dot Diagram
Example: Carbon e- configuration 2-4
Practice
The number of unpaired electrons is equal to the number of BONDS that an element can form with other elements
When determining the number of bonds an element can form, arrange the valence electrons so that you have the MAXIMUM number UNPAIRED
Example: Carbon
How many bonds can carbon form?4
Ground State vs. Excited State
Ground State electrons are in the lowest energy configuration possible (the configuration found on the periodic table)
Excited State electrons are found in a higher energy configuration (any configuration not listed on the periodic table)
Ground State Excited State
EXCITED
EXCITED
EXCITED
EXCITED
EXCITED
GROUNDGROUND
GROUND
The greater the distance from the nucleus, the greater the energy of the electron
When ground state electrons absorb energy they jump to a higher energy level or an excited state
This is very unstable/temporary condition
Excited electrons fall rapidly to a lower energy level
When excited electrons fall from an excited state to a lower energy level, they release energy in the form of light
Ground ExcitedEnergy is absorbedDark line spectrum is produced
Excited Ground Energy is releasedBright line spectrum is produced
Dark Lines Absorbed
Bright LinesEmitted
Balmer Series: electrons falling from an excited state down to the second (2nd) energy level give off visible light (Bright Line Spectrum or Visible Light Spectrum)
Different elements produce different colors of light or spectra
These spectra are unique for each element Spectral lines are used to identify different elements
*
Gas A and Gas D