CH 115 1st Edition Lecture 6Unit 1: Atoms, Isotopes, and IonsOutline of Last Lecture I. Wave Particle DualityII. Heisenberg Uncertainty Principle III. Quantum Numbersa. Electron Shells and Subshells b. Principal, Subsidiary and Magnetic Quantum NumbersOutline of Current Lecture VI. Shapes of Quantum Numbersa) Principalb) Subsidiaryc) Magnetic V. Energy Levels of Electronsa) Rulesi. Aufbau Principleii. Pauli Exclusion Principleiii. Hund’s Ruleb) Orbital Notationi. Box vs. Energy DiagramsCurrent LectureShapes of Quantum Numbers- ReviewPrincipal Quantum Number–Represented by “n”, this number shows theenergy and size. The letter “n” shows an energy level (n=1; n=2; n=3; etc.): the larger the value for “n”, the bigger the size and higher the energy of the orbital.Subsidiary Quantum Number – Represented by “l”, this number shows the shape and some energy. The letter “l” is equal to n-1 so if n=3 , l=2.Magnetic Quantum Number – Represented by “m”, this number shows the orientation. The letter “m” shows something like a range of values for the subsidiary quantum number “l”. The letter “m” is equal to –l,…,0,…,l. Thus, if m=3 , l= -3, -2, -1, 0, 1, 2, 3. (It might help to think of this stupid acronym l0l – Little 0rigin Large – not laugh out loud). If you get stuck remembering what values for “l” are possible given for the value of “m” just think:These notes represent a detailed interpretation of the professor’s lecture. GradeBuddy is best used as a supplement to your own notes, not as a substitute.o Little – The first value for “l” is going to be the negative of the value for “m”, which will also be the smallest (hence, the word “little”) of all possible values. So if m = 2, then the first and smallest possible value for l would be -2.o 0rigin – When you’ve identified the smallest possible value for “m”, literally count down until you reach zero, or the origin if you’re ironicallylooking at a graph. From here you should have, for m = 2, l = -2, -1, 0.o Large – Now, simply count up until you get the original value for “m”, which will also be the “largest” possible value for “l”. So after this extremely bad acronym you should have l = -2, -1, 0, 1, 2 (for the value of m = 2)Even though this acronym is awful, it is simple to remember and will ALWAYS get you the correct possible values for “l” for any value of “m” if followed correctly. Spin Quantum Number– Represented by “ml”, this number shows the spin. This number can literally only be one number or the other: ½ or -½.This is also known as “spin up” or “spin down”.- Drawing Shapes The principal quantum number at the lowest energy level (which is n = 1; l = 0) is simply the shape a sphere. The n = 1 energy level is also known as the s – orbital. When drawing the s – orbital, simply draw a circle. It can beshaded or unshaded. As “n” increases, the size gets bigger and starts to have different shapes (determined by the subsidiary quantum number). When n = 2 and l = 1, the shape resembles the number “8” or a dumbbell. This is now the p – orbital.  When drawing the p – orbital, simply draw the number “8”.  Make sure one half of the 8 is shaded and the other half is not. The half you choose does not matter.  Pay attention to axis labels: there are 3 orientations for the p – orbital: px, py, and pz. If you are sketching px, then you must ensure that the shape of the figure 8 is sideways (like an infinity symbol) so that it lies along the x-axis. You must also label the axis with an x because the orbital you are sketching is px. You must also draw another axis perpendicular to the x axis but it does not matter what you label the axis. When sketching py and pz, you must put both ofthe along the y axis Examples are below: When n = 3 and l = 2, the energy level has reached the d – orbital. When lequals 2, the values for the magnetic quantum numberare -2, -1, 0, 1, 2. The shape is like a clover with 2 shaded portions. Since there are 5 values for l, there are 5 different shapes. Three of them look the same, but have different axis. The last two are different as pictured below: When labeling the axis for a d – orbital, the first letter of the orbital you’re drawing will be on the horizontal axis and the second letter will be on the vertical axis for the first three d – orbitals. Also, make sure that when drawing the first three orbitals that you have the lobes (or “leaves” of the “clover”) intersecting the horizontal and vertical axis. In other words, draw the first three orbitals like an “X”. You must also ensure that 2 lobes are shaded and the other 2 are clear. The 2 clear lobes must be across from each other and the 2shaded lobes must be across from each other.  For the 4th d – orbital (the one that looks like a lowercase “t”), follow the same rule ofputting the first letter of the orbital on the horizontal axis and the second letter on the vertical axis. Only this time, make sure the lobes are ALONG the horizontal and vertical axis. Again, 2 lobes must be shaded and 2 must be clear (doesn’t matter which 2) and they MUST be across from each other. For the 5th d – orbital (the one with a ring around it), there is only one letter (“z”) and it will always be along the vertical axis. The horizontal axis will either be x or y, as shown in the picture, however Jared (SI Leader) strongly encourages that “y” is used to label the horizontal axis. Patterson most likely will not count off for either one. Make sure to remember that the 5th d – orbital IS NOT DRAWN LIKE THE REST of the d – orbitals. It has 2 lobes with a ring around it. The shading of the lobes does not matter AS LONG AS THE RING IS OPPOSITE from the color of the lobes. Soyou could shade the lobes but the ring must be clear. Or you could shade the ring, but the lobes must be clear. Both versions will be counted as correct.Energy Levels of Electrons- RULES – THEY MUST BE FOLLOWED AND MEMORIZED!1. Aufbau Principle– Orbitals are filled to give the lowest total energy for the atom. Orbitals increase in energy with increasing “n”. 2. Pauli Exclusion Principle – no two electrons can have the same 4 quantum numbers. This rule means that you cannot have 2 electrons in an orbital that have the same spin. It takes