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Unit-4 Atomic Structure and Quantum Theory 1. Bohr was the first to propose that the electrons were located in energy levels. A lower case “n” is used to denote these principle energy levels (also called principle quantum numbers). The principle energy levels are numbered, so that the level closest to the nucleus is labeled n = 1. The next level is labeled n = 2 and so forth. Each principle energy level had a certain energy value associated with the level. The closer the level was to the nucleus, the lower the energy of the level. The farther away from the nucleus, the higher the energy of that level in the atom. As long as the electrons were in these levels, the electrons do not give off energy. The dark circle below represents the nucleus. The rings around the nucleus represent the principle energy levels. Number the principle energy levels starting with the one closest to the nucleus: n = 1, n = 2, n = 3 etc. 2. Electron Configuration and the Periodic Table Each principle energy level can only hold so many electrons before the level is full. A quick and easy way to determine the maximum number of electrons (max e-) that a principle energy level can hold is given by the following: max e- = 2 n2. First square the principle energy level number (n) then multiply by 2. Energy Level (n) Maximum number of electrons (max e- = 2 n2) 1 2 3 4 5 6 Electrons are arranged around the nucleus by filling up the first principle energy level (n=1), then the second energy level, etc. This is the electron configuration given on your periodic table. The number of electrons are listed for each level with a dash between levels: for oxygen (O) which has a total of 8 electrons, the configuration is 2–6 (2 electrons are located in the first principle energy level and 6 electrons are located in the second principle energy level. Look up the electron configuration on the periodic table for the element given and fill in the chart. Ca is done as an example. Element Ca Na F B Al C H n=1 2 n=2 8 n =3 8 n=4 2 3. Drawing Bohr Diagrams of Atoms: 1) A circle is used for the nucleus- the # protons (# p or +) and the # of neutrons (#n) are placed in the circle. 2) A ring is drawn around the nucleus for each energy level. 3) The electrons for each energy level are placed in pairs symmetrically around the nucleus For F: atomic # = _____________ atomic mass = _____________ electron configuration: ______________ # p = _________ # n =____________ For Al : atomic # = _____________ atomic mass = _____________ electron configuration: ______________ # p = _________ # n =____________ Going Backwards: Determining the identity of an element from the Bohr diagram: # p = _____________ # n =______________ atomic # = _____________ atomic mass = # p + # n = ________________ electron configuration: # p = _____________ # n =______________ atomic # = _____________ atomic mass = # p + # n = ________________ electron configuration: ____________________________________ Isotopic Notation: ____________________________________ Isotopic Notation: 2. Ground and Excited States The lowest possible energy state that an electron can occupy is called the __________ ____________. This is a very __________ condition. The principle energy levels, which are occupied match those predicted by the electron configuration on the periodic table. When electrons gain energy, the electrons move to higher principle energy levels then they would normally occupy. This unstable situation is called the _____________ _______________. The electrons will release the absorbed energy, often seen as the bright line spectrum of the element, and fall back to the ground state. I) How to tell when energy will be absorbed or released The Principle Energy Level (n) changes: If the number of the principle energy level (n) goes up, then energy is _____________ or ______________ n = 1 to n = 3 OR n = 3 to n = 4 If the number of the principle energy level (n) goes down, then energy is _____________ or ______________ n = 2 to n = 1 OR n = 5 to n = 3 Determine if energy is added/absorbed (+E) or released/emitted (-E) for the following transitions: 1) n = 1 to n = 2 ______________ 6) n = 1 to n = 5 ______________ 2) n = 4 to n = 3 ______________ 7) n = 4 to n = 2 ______________ 3) n = 2 to n = 1 ______________ 8) n = 2 to n = 3 ______________ 4) n = 2 to n = 4 _____________ 9) n = 3 to n = 6 ______________ 5) n = 5 to n = 3 ______________ 10) n = 6 to n = 2 ______________ II) How do you tell the excited and ground state apart from the electron configuration?? Ground State: Matched the predicted electron configuration found on the periodic table. In other words, it follows the order given Ground State for Oxygen (O) on PT= 2 – 6 (8 total electrons) Possible Excited State for Oxygen = 1 – 7 (still 8 total electrons) The first energy level is not filled before moving into the second energy level. The KEY here is that the configuration does not MATCH the one on the PT. Another possible excited state for oxygen: 2 – 5 – 1 (still 8 total electrons) For the following elements, fill in the chart and determine if the electron configuration is in the GROUND STATE (GS) or EXCITED STATE (ES). Element Symbol Electron Configuration C 2-4 F 1-8 Cl 2-8-7 B 2-1-1-1 Na 2-7-2 O 2-6 S 1-8-7 Zn 2-8-18-2 Br 1-7-17-8-1-1 Protons Neutrons Electrons Principle Energy Levels Ground or excited state? Completely Filled vs. Occupied Principle Energy Levels Occupied means that there is at least one electron in the Principle Energy Levels Li: 2 – 1 has 2 occupied Principle Energy Levels Completely Filled means that each level has its maximum number of electrons which can be determined by the 2n2 rule. Li: 2 – 1 has only 1 Completely Filled Principle Energy Level To help you review the 2n2 rule complete the following chart. Principle energy levels Max number of e- 1 2n2 = 2( )2 = 2 2n2 = 2( )2 = 3 2n2 = 2( )2 = 4 2n2 = 2( )2 = 5 2n2 = 2( )2 = 6 2n2 = 2( )2 = 7 2n2 = 2( )2 = (2n2) a) Copy the electron configuration from the Periodic Table b) Determine the number of Occupied Principle Energy Levels (PEL) c) Determine the number of Completely Filled Principle Energy Levels Element Electron Configuration # Occupied PEL # Completely Filled PEL C Na O Cl He F Ne Si Zn Au 3. Bright Line Spectra and Introduction to Light Every atom has many electron energy levels. Electrons can drop from the excited state in many different pathways. Each different drop in energy level gives off light with different amounts of energy, and therefore wavelength. The electrons will fall back down in the different ways, giving off different wavelengths of light. If you look at the element sample in the flame, it will appear to be glowing with a certain color. If the light is broken up by a prism (as raindrops can break sunlight up into a rainbow) and projected onto a white surface, the individual energies of light can be seen as bright lines of different color. These colored bright lines are different from element to element, making each spectrum unique to a particular element. This process of observing atomic and molecular spectra is called ______________ and the device used to observe the spectra is called a _____________________ . Example bright-line spectrum for hydrogen (Numbers in parenthesis indicate the drop in electron energy level that gives off the light, numbers in brackets indicate the wavelength in nanometers (10-9 m) of the photons making up that bright line): violet violet-blue alpha) (62) (52) [410.2 nm] [434.0 nm] green (42) [486.1 nm] red (hydrogen(32) [636.5 nm] In order for the electron to return to a lower and more stable energy level, the added energy must be given off. When the electrons return to the lower energy levels this decreases the PE because the added energy is given off and the colors of the bright line spectra are seen. BRIGHT LINE SPECTRUM BRIGHT LINE SPECTRA are produced when “electrons in the EXCITED STATE” fall back to lower energy levels of the GROUND STATE. Unlike the continuous spectrum of sunlight, only certain colors will be present in the BRIGHT LINE SPECTRA. The BRIGHT LINE SPECTRUM is like a “fingerprint” of the element that produced the spectrum. Like a fingerprint, the BRIGHT LINE SPECTRA can be used to identify the element. When viewed with a spectroscope, the individual bands of colors in the BRIGHT LINE SPECTRUM can be seen and the wavelength of each band determined. 1. Below are the BRIGHT LINE SPECTRA of three elements. From the position of the lines determine which element is the unknown. (HINT: Match up the lines present in the unknown with the three known elements.) Unknown element = ____________ Element X Element Y Element Z Unknown 2. Which of the two elements above are present in the BRIGHT LINE SPECTRUM given below? (HINT: Match up the lines present with the three known elements. Only two patterns should match perfectly.) ____________ and _____________ Introduction to Light Visible Light (energy we see with): part of the Electromagnetic Spectrum 1. Two theories to explain light’s behavior: Waves Particles of Packets of Energy There was evidence for both models so the two theories were put together!! Light: QUANTUM THEORY OF LIGHT a) packets or bundles of energy called _________________ or ______________ b) travel in wave-like fashion c) produced when electrons drop from ______________ energy levels to _____________ energy levels (the greater the drop, the greater the energy the light has) 2. Properties of Light Wavelength ( ) - ____________________________________________ Frequency (F) - ______________________________________________ (units: Cycles / second OR Hertz) Energy (E) - __________________________________ Speed (velocity) – same for all electromagnetic radiation _______________ Relationships: Frequency and Energy: Type _________________ F ______, E ________ or F ______, E ________ Frequency and Wavelength: Type _________________ F ______, ________ or F ______, ________ Wavelength and Energy: Type _________________ ______, E ________ or ______, E ________ 3. The Rainbow: A Continuous Spectrum R Long Low F Low E O Y G B I V LONG STEM RED ROSES: All “L’s” go together with RED Short High F High E B. Continuous Spectrum when radiation from the sunlight passes through a prism, a rainbow – a spectrum of colors – is seen the colors are not separated from one another but blend together due to the overlap of the line spectra of the 67 different elements in the sun lightbulb C. Bright Line Spectrum when radiation from an excited atom (element) passes through a prism, the radiation is separated into various wavelengths and colors Colors are not blended – spectrum is discontinuous – and you observe lines of color at different locations Flame Updating The BoHR MoDeL: QuAnTuM oR WaVe MeChAnIcaL MoDeL The Bohr Model was very good at explaining many things about atoms and electrons and how they behaved. Just like your wardrobe needs updating, so did the model of the atom. A revolution in physics occurred in the early 1900’s when experiments showed that matter, just like light energy, could have a dual nature…it can act as a particle or a wave. 1. The Wave Mechanical Model 1.__________________________________________________________________ 2.__________________________________________________________________ 3.__________________________________________________________________ Using complicated math, this math showed that electrons were not moving in definite fixed orbits like planets but had distinct amounts of energy. a. These electrons were found in “regions” or “areas” around the nucleus called orbitals. b. An orbital is defined as a region in which an electron with a particular amount of energy is most likely to be found The probability of finding an electron near the nucleus is _____________ The probability of finding an electron far away from the nucleus is __________ B. Four “Quantum Numbers” are used to describe the location of the electron. a. Principle Energy Levels (PEL) get divided into sublevels. (PEL are also called the PRINCIPLE QUANTUM NUMBER. ) b. These sublevels are designated: s, p ,d and f c. The number of sublevels is determined by the number of the princple energy level (n): n = # sublevels d. Within the same principle energy level: the sublevels have different energies: s p d f lowest energy highest energy e. The sublevels are divided into orbitals (rooms) where electrons are found. Each oribital can hold 2 electrons. f. Orbitals are “regions” or “areas” around the nucleus g. The number of orbitals depends on the sublevel type: Sublevel # orbitals x2= # of electrons in sublevel s x2= p x2= d x2= f x2= There are two ways to show the Quantum Atom Electron Configuration. One way shows the number of electrons in the Principle Quantum Level (PEL) and Sublevel type (letter). The other way, known as BOX DIAGRAMS, show all four quantum numbers by using boxes for orbitals and arrows for electrons to show the opposite spin of the electrons. Using the electron configuration from the periodic table, we will do both types. For Hydrogen: EC form PT is 1 Quantum Atom Box Diagram 1s1 PEL 1s # of electrons SUBLEVEL PEL ORBITAL SUBLEVEL ARROW SHOW SPIN RULES TO REMEMBER WHEN DRAWING BOX DIAGRAMS; * S has one orbital so draw 1 box * P has 3 orbitals so draw 3 boxes, * Each “P” box in a sublevel must have an electron before pairing up because each orbital has the same energy as the other orbitals (Boxes) * Electrons spin in opposite directions! Principle Energy Level Sublevel(s) Present Orbitals Present Total # of e per PEL 1 2 3 4 n = _____________________ & n = ___________________________ Element Elec. Config. From P T Quantum Atom E. C. Box Diagrams H He Li Be B C N O F Ne Na Mg Al Si P After argon (Ar) the energy levels begin to overlap and things get complicated. We will not be doing elements above Calcium. However…. The Kernel, Valence Electrons and Lewis (Electron) Dot Diagrams A very powerful tool for showing how different elements bond together, called Lewis (Electron) Dot Diagrams, was developed by a chemist named Lewis. Many times in chemical bonding, only the electrons on the outside or in the highest energy level, called the VALENCE SHELL, actually become involved in bonding to form compounds. We are going to look at the Lewis (Electron) Dot Diagrams and elements in this next section. Valence Shell: Outermost energy level of an element (highest number) Oxygen: O 2-6 the 2nd principle energy level is the valence shell Valence Electrons (VE): electrons located in the outermost energy level - these are the electrons which are the furthest to the right in the electron configuration found on the PT - Oxygen: O 2-6 the 6 electrons are the valence electrons - In a dot diagram, the VE are represented by dots: Kernel: This is the nucleus and all the other electrons located in the inner energy levels In a dot diagram, the chemical symbol represents the kernel: O The valence shell always contains only 4 orbitals. These orbitals are represented by the four sides around the symbol. The top side represents one orbital which has less energy than the other three orbitals. The other three orbitals have the same amount of energy. Frist 2 electrons are always placed first in this top orbital. Going around clockwise, one electron is placed on each remaining side until you need to pair the electrons up. Orbital lower energy than others O Orbitals with the same energy O Orbitals with the same energy The Periodic Table: 1) Look up and write down the electron configuration, then circle the valence electrons 2) Try your hand at drawing dot diagrams Symbol Electron Configuration Li Be B C N O F Ne Li Be B C N O F Ne Dot Diagram Guided Practice: Complete the following chart Symbol Electron Configuration (Circle Number of VE the Valence Electrons (VE) Dot Diagram S Se Na K Si Ge Cl Br Do you see any patterns in the relationship between the diagrams and the element’s position on the periodic table?