UNIT 2 - ATOMIC THEORY VOCABULARY: Allotrope Electron Configuration Nuclear Charge Anion Element Nucleons Atom Excited state Nucleus Atomic Mass Ground state Orbital Atomic Mass unit (a.m.u.) Ion Proton Atomic number Isotope Quantum Theory Bohr model Kernel electron(s) Valence electron(s) Cation Lewis Dot Diagram Wave-mechanical model Compound Mass number Electron Neutron OBJECTIVES: Upon completion of the unit you will be able to do the following: Discuss the evolution of the atomic model Relate experimental evidence to models of the atom Identify the subatomic particles of an atom (electron, proton, and neutron) Know the properties (mass, location, and charge) of subatomic particles Determine the number of protons, electrons, and neutrons in a neutral atom and an ion Calculate the mass number and average atomic weight of an atom Differentiate between an anion and a cation Distinguish between ground and excited state Identify and define isotopes Write electron configurations Generate Bohr diagrams Differentiate between kernel and valence electrons Draw Lewis Dot Diagrams for an element or an ion
THE EVOLUTION OF THE ATOMIC MODEL Atom 1.) basic building block of matter 2.) cannot be broken down chemically 3.) a single unit of an element Dalton (cannonball model) 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 MODEL: SPHERICAL UNIFORM DENSITY J.J.Thomson (plum pudding model) EXPERIMENT: Used a CATHODE RAY TUBE with charged electrical field (+/-) o Cathode ray deflected BY NEGATIVE electrode TOWARD POSITIVE electrode Discovered SUBATOMIC PARTICLE called the ELECTRON: 1. SMALL 2. NEGATIVELY CHARGED PLUM PUDDING MODEL: POSITIVE PUDDING NEGATIVE ELECTRONS EMBEDDED (just like raisin bread)
Rutherford (nuclear model) EXPERIMENT: 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 (see diagram above) NUCLEAR MODEL: The ATOM is MOSTLY EMPTY SPACE At the center of the atom 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 (BOHR MODEL) BOHR MODEL or PLANETARY MODEL: Electrons travel AROUND the nucleus in well-defined paths called ORBITS (like planets in a solar system) Electrons in DIFFERENT ORBITS possess DIFFERENT AMOUNTS OF ENERGY 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 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 o 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 contributed to this theory
VOCABULARY (of the Periodic Table) Atomic Mass = AVG. mass of all the isotopes of an element Oxidation # s = possible charges an atom of an element can have Atomic # = the number of protons in EVERY ATOM of the element Element Symbol = the letter(s) used to identify an element Electron configuration = number of electrons in each energy level; add the numbers to get the total # of electrons SUBATOMIC PARTICLES Subatomic Charge Relative Mass Location Symbol How to Calculate Particle p Proton +1 1 amu Nucleus Neutron 0 1 amu Nucleus ---------- 1 p 1 n ---------- 1 n 0 Look at the atomic # Mass # - Atomic # Electron -1 1/1836 th amu Outside e- or e P = e- in neutral atom (negligible) nucleus
- DETERMINING SUBATOMIC PARTICLES (p, n, e) DIRECTIONS: Use the information below to complete the chart that follows. Atomic number = the # of protons in an atom of an element (p) Mass # = sum of the protons and neutrons in an atom of an element (p + n) = mass # Nuclear Charge=charge w/in the nucleus; = to the # of protons or the atomic # nuclear charge=# of p Nucleons= any subatomic particles found w/in the nucleus protons and neutron Symbol # # # Atomic Mass Nuclear Nuclear Protons Neutrons Electrons Number Number Charge Symbol 35 Cl-35 17 18 17 17 35 17 Cl 17 15 16 C-14 8 16 8 O 8 Ar-40 18 22 20 40
ATOMS (neutral) VS. IONS (charged) Vocabulary Term Definition Example/Diagram Neutral Atom An atom with the same number of protons and electrons p = e 12 6 or C C (no charge indicated) Ion anion An atom that has GAINED one or more electrons e > p 19 9 F -1 the -1 indicates a negative charge with a magnitude of one Ca+ion An atom that has LOST one or more electrons p > e 3 7 +1 Li the +1 indicates a positive charge with a magnitude of one
ISOTOPE = atoms of the same element with different mass # s but the same atomic #; (same number of protons, different number of neutrons) p = p = p = n = n = n = e = e = e = Example 2: Isotopes of Uranium (U-238, U-240) p = p = n = n = e = e = Calculating Atomic Mass (for any element): Atomic Mass = the weighted average of ALL the element s naturally occurring isotopes (% abundance of isotope in decimal form) x (mass of isotope 1) +(% abundance of isotope in decimal form) x (mass of isotope 2) + (% abundance of isotope in decimal form) x (mass of isotope 3) Average Atomic Mass of the Element
Example 1: The exact mass of each isotope is GIVEN to you. Chlorine has two naturally occuring 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 in DECIMAL FORM (move the decimal 2 places to the left) Cl-35 = 34.9689 x (.6749) = 23.6005 Cl-37 = 36.9659 x (.3251) = 12.0176 ** These are the weighted masses ** Step 2: Add up the products of all the calculated isotopes from step 1. 23.6005 + 12.0176 35.6181 ** This is your average atomic mass **
Example 2: The exact mass for each isotope is NOT GIVEN to you. The element Carbon occurs in nature as two isotopes. Calculate the average atomic mass for Carbon based on the information for the isotopes below. 12 C = 98.89% 13 C = 1.11% **Since the exact masses were NOT given for either of the isotopes, just use the mass number instead. For 12 C, the mass would be 12 amu, and for 13 C, the mass would be 13. Step 1: Multiply the mass of each separate isotope by its percent abundance in DECIMAL FORM (move the decimal 2 places to the left) 12 C = 12.9889) = 11.8668 13 C = 13 x (.0111) = 0.1443 ** These are the weighted masses ** Step 2: Add up the products of all the calculated isotopes from step 1. 11.8668 + 0.1443 12.01 ** This is your average atomic mass **
Practice 1: The element Boron occurs in nature as two isotopes. In the space below, calculate the average atomic mass for Boron. Isotope mass percent abundance Boron-10 10.0130 amu 19.9% Boron-11 11.0093 amu 80.1%
Practice 2: The element Hydrogen occurs in nature as three isotopes. In the space below alculate the average atomic mass for Boron. Isotope of Hydrogen 1 H Protium 99.0% 1 2 H Deuterium 0.6% 1 3 H Tritium 0.4% 1 Percent Abundance Challenge Practice 3: Chlorine has two naturally occuring isotopes, Cl-35 (isotopic mass 34.9689 amu) and Cl-37 (isotopic mass 36.9659 amu). If chlorine has an atomic mass of 35.4527 amu, what is the percent abundance of each isotope? Mass Number The MASS of ONE ISOTOPE of a given element. Atomic Mass The AVERAGE MASS of ALL ISOTOPES of a given element
ELECTRON CONFIGURATIONS = a dashed chain of numbers found in the LOWER LEFT CORNER of an element box (see below); tells us the number of ENERGY LEVELS as well as the number of ELECTRONS in each level (tells us how the electrons are arranged around the nucleus) Electron configuration **All electron configurations on the Periodic Table are NEUTRAL (p=e) SUBSTANCE ELECTRON CONFIGURATION Magnesium 2-8-2 Mg +2 2-8 Bromine Br -1 Barium *Lead * shortcut allows you to cut out the first two orbitals to shorten the address Valence Electrons: electrons found in the OUTERMOST shell or orbital the LAST number in the electron configuration Kernel Electrons: INNER electrons (all non-valence electrons) Ex: Practice 1: Practice 2: Chlorine Nitrogen Sodium # valence e - = 7 # valence e - = # valence e - = # kernel e - = 10 # kernel e - = # kernel e - =
Principle Energy Level (n) = electron energy levels that contain a certain number of SUBLEVELS (s, p, d, f); each sublevel contains one a set number of ORBITALS Maximum # of electrons in an energy level = 2n 2 where n = quantum # (or period #) Principle Energy Level (n) Maximum number of electrons (2n 2 ) 1 2 3 4 BOHR DIAGRAMS (one method for expressing electron location in an atom or ion); ALL ELECTRONS MUST be drawn BOHRing 1. Look up electron configuration of element at hand on Periodic Table (if you are working with an ion, add/subtract the proper amount of electrons from outer shell(s) of configuration) Example: Oxygen is 2-6 2. Draw nucleus (with a square) and notate correct amount of protons and neutrons inside P = 8 N = 8 3. Using rings or shells, place the proper number of ORBITS around your nucleus.
4. Use either and x or a dot to represent your electrons, place the correct number of electrons in the area that would correspond to the number 12 on the face of a clock in the ORBIT CLOSEST TO THE NUCLEUS ONLY. Remember, you can have a maximum of 2 e - in the first orbit. x x P = 8 N = 8 5. Place one x or one dot at a time around your 2 nd or valence orbit in the areas that would correspond to the numbers 12, 3, 6, and 9 on the face of a clock. x x x x P = 8 N = 8 x 6. If there are any electrons remaining in the configuration, pair them up with electrons you have already placed. You may have no more than 2 e - in any of the 4 spots, and no more than 8 e - total in the 2 nd or valence orbit. In this case, we have 2 e - left to place. x P = 8 N = 8 x x
Carbon Fluorine Beryllium Al Li Ca 2+ Na + S 2-
LEWIS (ELECTRON) DOT DIAGRAMS (Only illustrates VALENCE ELECTRON CONFIGURATION) 1. Write the element s symbol 2. Retrieve electron configuration from Periodic Table. The last number in the configuration is the NUMBER OF VALENCE ELECTRONS 3. Arrange the valence electrons (DOTS) around the symbol using the following rules: Only two electrons maximum per side of the symbol (therefore no more than 8 total surrounding symbol 8 is great!) Always pair the first two If you have more than 2 valence electrons, deal them one at a time to the other three sides until you run out 8 1 5 2 OR 8 1 2 3 4 X 6 5 X 6 3 7 4 7 Ex: Using a dot or an x place the valence electrons around the symbols for carbon below in order according to each of the models above. C OR C 4. If you are working with an ION you must adjust the valence electrons (add or subtract electrons) in the configuration before constructing your Dot Diagram. Your final diagram must include brackets and the charge on the ion. Ex 1: S -2 ADD 2 e - to the 6 that S normally has in its valence shell. Ex 2: K +1 REMOVE 1 e - from the valence shell of K. *NEGATIVE IONS always end up with EIGHT VALENCE e - *POSITIVE IONS always end up with ZERO VALENCE e -
QUICK NOTE: The number of UNPAIRED VALENCE 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. Ex: Carbon can be arranged in either of the two ways shown on the previous page. Draw the Lewis dot structure for carbon that gives it the maximum number of UNPAIRED valence electrons. How many bonds can an atom of carbon form? The alternate Lewis structure for carbon with only 2 unpaired valence electrons represents the LOWEST ENERGY STATE.
Ground State vs. Excited State *Notice that one electron from the 2 nd orbital has moved to the 3 rd orbital Ground State = electrons in LOWEST ENERGY CONFIGURATION possible (the configuration FOUND ON PERIODIC TABLE) ground state electron configuration for Li is 2-1 Excited State = electrons are FOUND IN A HIGHER ENERGY CONFIGURATION (ANY configuration NOT FOUND ON PT) excited state electron configuration for Li could be 1-2, 1-1-1 Distinguish between ground state and excited state electron configurations below: Bohr Electron Configuration Ground (G) or Excited (E) state? 2-1 Ground 2-0-1 Excited 1-1-1 2-7-3 2-8-2 2-8-8-2 2-8-17-6 2-8-18-8 2-6-18-1 2-5-18-32
***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 a very UNSTABLE/TEMPORARY condition EXCITED ELECTRONS rapidly FALL BACK DOWN or DROP to a LOWER energy level When excited electrons fall from an excited state to 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 show the specific wavelengths of light being ABSORBED by the electrons (becoming excited) BRIGHT LINES show the specific wavelengths of light being EMITTED by the electrons (falling down) Balmer Series: electrons falling from an EXCITED STATE down to the 2 ND ENERGY LEVEL give off VISIBLE light (AKA Bright Line SPECTRUM or Visible Light SPECTRUM ) Different elements produce different colors of light or SPECTRA. These spectra are UNIQUE for each element (just like a human fingerprint is unique to each person). We use spectral lines to identify different elements.