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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 atoms 2. all atoms of a given element are identical in mass and properties. Atoms of different elements have different masses and different properties 3. Compounds are formed by a combination of 2 or more atoms 4. Atoms cannot be created, destroyed, or converted into other kinds of atoms during chemical reactions Cannonball Theory/Model 1. Spherical 2. Uniform Density 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 Small Negatively 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 Rutherford Nuclear 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 **Nucleons= Protons and Neutrons (any subatomic particle found within the nucleus) Subatomic Charge Relative Particle Mass Location Symbol How to Calculate or Look at atomic # Proton +1 1 amu Nucleus Neutron 0 1 amu Nucleus Mass # Atomic # Electron -1 1/1836 amu Outisde nucleus P= e (in neutral atoms) 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. Ions Vocabulary Term Neutral Atom Ion Definition An atom with the same number of protons and electrons P=e Two Types Anion and Cation Protons and electrons are not the same Example (no charge indicated) Atoms vs. Ions Ions Anion aNion Cation ca+ion An atom that has An atom that has GAINED one or more LOST one or more electrons electrons e>p p>e NEGATIVE ION POSITIVE ION Isotope Atoms of the same element with different mass numbers; same atomic number, same number of protons, different number of neutrons Isotopes Example 1: Carbon (12, 13, 14) p= e= n= p= e= n= p= e= n= Isotopes Example 2: Uranium (238, 240) U-240 p= e= n= p= e= n= Practice 1. Two different isotopes of the same element must contain the same number of * a. protons b. neutrons c. electrons 2. Two different isotopes of the same element must contain a different number of a. protons * b. neutrons c. electrons Practice 3. Isotopes of a given element have * a. the same mass number and a different atomic number b. the same atomic number and a different mass number c. 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 Example 1: The exact mass of each isotope is given 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 CL-35 = 34.9689 amu x (.6749) = 23.6005 CL-37 = 36.9659 amu x (.3251) = 12.0176 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) = 11.8668 C-13 = 13 x (.0111) = *These are weighted masses* Step 2: Add the products together 11.8668 + 0.1443 12.0111 0.1443 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 shapes Each 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 numbers Pauli 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 electrons Add one at a time starting on the right side and going counter clockwise Shell (energy level) 1 2 3 4 5 6 7 2 2n n= Maximum number of electrons 2 8 18 32 50 72 98 Formula for maximum # of electrons per quantum number (or energy level) Bohr Models Example: Draw the Bohr Model for Carbon Bohr Models Example Draw the Bohr Model for Oxygen Bohr Models Example 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 GROUND GROUND EXCITED GROUND EXCITED EXCITED 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 Excited Energy is absorbed Dark line spectrum is produced Excited Ground Energy is released Bright line spectrum is produced Dark Lines Absorbed Bright Lines Emitted 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