# UNIT 2 - ATOMIC THEORY

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1 UNIT 2 - ATOMIC THEORY VOCABULARY: Allotrope Anion Atom Atomic Mass Atomic Mass unit (a.m.u.) Atomic number Bohr model Cation Compound Electron Electron Configuration Element Excited state Ground state Ion Isotope Kernel electron(s) Lewis Dot Diagram Mass number Neutron Nuclear Charge Nucleons Nucleus Orbital Proton Quantum Theory Valence electron(s) Wave-mechanical model OBJECTIVES: Upon completion of the unit you will be able to do the following: Understand that the modern model of the atom has evolved over a long period of time through the work of many scientists 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 Identify what element the amu unit is derived from Define the term orbital 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 1

2 THE EVOLUTION OF THE ATOMIC MODEL Atom = Democritus = Dalton (1803) o Known as the of the atomic theory o 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 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 ****** ****** J.J. Thomson (1897) (Click link to see animation of Cathode Ray Tube Experiment) o Used a with charged electrical field (+/-) o Cathode ray deflected electrode electrode o Discovered called the : ****** ****** (just like raisin bread) J.J. Walker (1975) (Click link to see one of my favorite TV catch phrases) 2

3 Rutherford (1909) * (Click link to see animation of Gold Foil Experiment) o Conducted the where he a thin piece of with a o Expected virtually all alpha particles to pass straight through foil o Most passed through, but some were severely deflected (see diagram above) o Conclusion led to ****** ****** 1. The is 2. At the center of the atom is a called the *Provided no information about electrons other than the fact that they were located outside the nucleus. Neils Bohr (1913) (Click link to see animation of the Bohr Model) ****** or ****** 1. Electrons travel the nucleus in well-defined paths called (like planets in a solar system) 2. Electrons in possess 3. a certain amount of causes electrons to to a or an When electrons a certain amount of causes electrons to to a or the 3

4 Wave-Mechanical/Cloud Model (Modern, present-day model) o Electrons have distinct amounts of energy and move in areas called ORBITAL = an area of for finding an (not necessarily a circular path) o Developed after the famous discovery that energy can behave as both & o have contributed to this theory PRACTICE: 1. Which of the following did Rutherford s Gold Foil experiment prove? a. That the atom was a uniformly dense sphere. b. That the atom is mostly empty space with a dense, positive core. c. That most the atom consists of a uniform positive pudding with small negative particles called electrons embedded throughout. d. That electrons travel around the nucleus in well-defined paths called orbits. 2. J.J. Thomson s Cathode Ray Tube experiment led to the discovery of a. the positively charged subatomic particle called the electron b. the positively charged subatomic particle called the proton c. the positively charged subatomic particle called the electron d. the negatively charged subatomic particle called the electron 3. According to the Bohr Model, a. electrons are found in areas of high probability called orbitals b. electrons travel around the nucleus in circular paths called orbits c. electrons are found in areas of high probability called orbits d. electrons travel around the nucleus in random paths called orbitals 4. According to the Wave-Mechanical Model, a. electrons are found in areas of high probability called orbitals b. electrons travel around the nucleus in circular paths called orbits c. electrons are found in areas of high probability called orbits d. electrons travel around the nucleus in random paths called orbitals 4

5 5. Which of the following is NOT one of Dalton s Postulates? a. All atoms of a given element are identical b. All atoms look like a simple sphere c. The mass of an atom is evenly distributed d. Atoms of different elements have different properties and masses 6. When electrons in an atom gain or absorb enough energy, they can a. jump to the ground state b. fall back to the ground state c. jump to an excited state d. fall back to an excited state 7. When excited electrons lose or emit enough energy, they can a. jump to the ground state b. fall back to the ground state c. jump to an excited state d. fall back to an excited state 5

6 VOCABULARY (of the Periodic Table) SUBATOMIC PARTICLES Subatomic Particle Charge Relative Mass Location Symbol How to Calculate Proton Neutron Electron 6

7 Mass number = Nuclear Charge = Nucleons = PRACTICE: Calculating Subatomic Particles in Different Atoms: Symbol # Protons # Neutrons # Electrons Atomic Number Mass Number Nuclear Charge Nuclear Symbol Cl Cl C O 8 7

8 ATOMS (neutral) VS. IONS (charged) Vocabulary Term Definition Example/Diagram Neutral Atom Ion anion Ca+ion 8

9 ISOTOPE = Example 1: Isotopes of Carbon (C-12, C-13, & C-14) p = p = p = n = n = n = e = e = e = Example 2: Isotopes of Uranium (U-238, U-240) p = p = n = n = e = e = Sample Questions: 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 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 9

10 Calculating Atomic Mass (for any element): Atomic Mass = the weighted average of an 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 Ex: The element Carbon occurs in nature as two isotopes. Calculate the average atomic mass for Carbon. 12 C = 98.89% 13 C = 1.11% 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 x (.9889) = C = 13 x (.0111) = ** These are the weighted masses ** Step 2: Add up the products of all the calculated isotopes from step ** 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 amu 19.9% Boron % 10

11 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 1 2 H Deuterium 1 3 H Tritium 1 Percent Abundance 99.0% 0.6% 0.4% Challenge Practice 3: Chlorine has two naturally occuring isotopes, Cl-35 (isotopic mass amu) and Cl-37 (isotopic mass amu). If chlorine has an atomic mass of amu, what is the percent abundance of each isotope? Mass Number The of of a given element. Atomic Mass The of of a given element 11

12 ELECTRON CONFIGURATIONS = a dashed chain of numbers found in the of an element box (see below); tells us the number of as well as the number of 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 Mg 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 shell or orbital the number in the electron configuration Kernel Electrons: 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 - = 12

13 Principle Energy Level (n) = electron energy levels that contain a certain number of ; each sublevel contains one a set number of Maximum # of electrons in an energy level = where n = # (or period #) Principle Energy Level (n) Maximum number of electrons (2n 2 ) Sample question: What is the maximum number of electrons that can occupy the 3 rd principal energy level in any atom? PRACTICE: 1. Give the correct electron configuration for an atom of oxygen. 2. Give the correct electron configuration for the K + ion. 3. How many valence electrons are there in atom of fluorine? 4. How many valence electrons are there in a F - ion? 5. How many kernel electrons are there in a phosphorus atom? 6. How many principal energy levels are there in a Bromine atom? 7. How many principal energy levels are there in a Bromine atom? 13

14 BOHR DIAGRAMS (one method for expressing electron location in an atom or ion); 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. Draw nucleus (with a ) and notate correct amount of protons and neutrons inside 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. 14

15 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. 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. Carbon Fluorine Beryllium Al Li Ca 2+ Na + S 2-15

16 LEWIS (ELECTRON) DOT DIAGRAMS (Only illustrates ) 1. Write the element s symbol 2. Retrieve electron configuration from Periodic Table. The last number in the configuration is the - 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 X 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. 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. C 2 6 OR OR X C 3 6 Ex 1: S -2 to the 6 that S normally has in its valence shell. Ex 2: K +1 from the valence shell of K. S -2 K +1 * always end up with * always end up with 16

17 Draw LEWIS ELECTRON-DOT DIAGRAMS for the following: Argon Phosphorus Silicon Beryllium Oxygen Aluminum Sodium N 3- Na 1+ O 2- QUICK NOTE: The number of is equal to the number of 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 number. 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 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. 17

18 PRACTICE: 1. What is the advantage to using the Bohr model as opposed to the Lewis model? a. The Bohr model provides more information than the Lewis model. b. The Bohr model is less bulky than the Lewis model. c. The Bohr model shows the valence electrons. d. The Bohr model shows how many bonds atoms of an element can form. 2. What is the advantage to using the Lewis model as opposed to the Bohr model? a. The Lewis model provides more information than the Bohr model. b. The Lewis model is less bulky than the Bohr model. c. The Lewis model shows the valence electrons. d. The Lewis model shows how many bonds atoms of an element can form. 3. What is the maximum number of electrons an atom or an ion can have in its valence shell? a. 2 c. 6 b. 4 d The number of bonds an atom of an element can form is the same as the number of a. electrons in its valence shell. b. paired electrons in its valence shell. c. unpaired electrons in its valence shell. 5. An atom of which of the following elements can form the most bonds? (Hint: Look at your Lewis diagrams) a. Phosphorus c. Sodium b. Oxygen d. Carbon 18

19 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 possible (the configuration ) ground state electron configuration for Li is 2-1 Excited State = electrons are ( configuration ) excited state electron configuration for Li could be 1-2, Distinguish between ground state and excited state electron configurations below: Bohr Electron Configuration Ground or Excited state? 2-1 Ground Excited

20 ***The greater the distance from the nucleus, the greater the energy of the electron When they jump to a energy level or an. This is a very condition rapidly or to a energy level When excited electrons fall from an excited state to lower energy level, they release energy in the form of. Energy is is produced Energy is is produced show the specific wavelengths of light being by the electrons (becoming excited) show the specific wavelengths of light being by the electrons (falling down) Balmer Series: electrons falling from an down to the give off light (AKA Bright Line or Visible Light ) Different elements produce different colors of light or. These spectra are for each element (just like a human fingerprint is unique to each person). We use spectral lines to identify different elements. 20

21 Practice: 1. Which of the following is a ground state electron configuration? a c b d Which of the following is an excited state electron configuration? a c b d Give one possible example of an excited state electron configuration for oxygen: 4. When atoms of an element in the ground state absorb enough energy, some of their electrons may a. fall back to the ground state and give off energy in the form of light. b. fall back to an excited state and give off energy in the form of light c. jump to the ground state. d. jump to an excited state. 5. When excited electrons lose enough energy, they will a. fall back to the ground state and give off energy in the form of light. b. fall back to an excited state and give off energy in the form of light c. jump to the ground state. d. jump to an excited state. 6. A dark-line spectrum is produced by a. electrons emitting energy and falling back to the ground state. b. electrons absorbing energy and falling back to the ground state. c. electrons absorbing energy and jumping to an excited state. d. electrons emitting energy and jumping to an excited state. 21

22 7. A bright-line spectrum is produced by a. electrons emitting energy and falling back to the ground state. b. electrons absorbing energy and falling back to the ground state. c. electrons absorbing energy and jumping to an excited state. d. electrons emitting energy and jumping to an excited state. 8. Based on the known bright line spectra produced by the four gases below, which gases are present in the unknown mixture? Gases present: 22

23 What to Study for the Atomic Exam (Unit 2) Structure of the atom (role/nature of the nucleus and electrons etc.) Gold Foil Experiment (know what was observed and the conclusions that were drawn from these 2 observations) Orbit vs. orbital (what is the difference?) Location, mass, charge of protons, neutrons, and electrons in an atom (see chart in notes) Isotopes (same element; different number of neutrons or masses ex: C-12 & C-14) be able to identify them Nucleons (the subatomic particles found in the nucleus) Nuclear Charge (charge inside the nucleus; always positive and directly dependent on # protons) Chemical notations: where do we find the atomic #, mass #, atomic mass etc. Determining the number of protons, neutrons, and electrons (pne) in an atom or ion (use mass # or atomic mass rounded to whole number) remember that the # protons ALWAYS tell us the symbol/element we have Atomic mass weighted average mass of an element s naturally occurring isotopes (ex: C is ) Ground state vs. excited state and the energy absorbed/released during electron movement (remember: light is emitted or seen when an electron jumps from an excited state back down to ground state or a lower energy level) Principle energy levels (orbitals) and the maximum # electrons found in each use 2n 2 formula Calculating average atomic mass given isotope abundances and masses The nucleus makes up pretty much all the mass in an atom therefore (mass #) = (#protons) + (#neutrons) Who performed the cathode ray experiment? What subatomic particle was discovered in this experiment? Atom vs. Ion (what is the difference?) Electron Configuration found below the atomic number on PTable (know how to write them for an atom or ion) Drawing Bohr models show all the electrons within an atom/ion and their location (construct nucleus & use shells or circles to place electrons) Lewis Dot Diagrams write element symbol and distribute dots around (draw only the valence electrons) 23

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