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Atomic Emission Spectra, Electron Configuration, Periodicity Ch.5 Coral Gables Senior High Ms. Kiely Pre-IB Chemistry I Bell-Ringer #8 Today 1. Individual teacher-student meetings with rebuild 2. Unit 2: Ch.6 Introduction Questions DUE TODAY 1. State the group number(s) for each of the following groups of elements: the alkali metals, the alkaline metals, the halogens, the noble gases, the transition metals. Take note that your book refers to these groups as 1A-8A, and 3B-7B. TRANSLATE THESE INTO GROUPS 1-18! 2. Which group numbers correspond to the representative elements, and why are they given that name? 3. Define atomic radius. What is the trend across a period for the atomic radius of an atom? 4. Define ionization energy. Does ionization energy pertain to cations or anions? 5. Define electronegativity. Explain why electronegativity increases left to right across a period. Does electronegativity pertain to cations or anions? Bell-Ringer #9 Turn in Unit 2 Introduction Questions Bell-Ringer #9 Answer: B Bohr’s Experiment Bohr excited the atoms of hydrogen gas with an electric current (energy). He then noticed a blue light emitting from the gas sample. He had this blue light pass through a prism. The prism showed that the blue light was composed of four specific wavelengths of visible light: red, blue-green, blue-violet, and violet. He was observing hydrogen’s Atomic Emission Spectra This is hydrogen’s Atomic Emission Spectra. When the electrons of an atom absorb a particular amount of energy, a quanta, they are able to “jump” to higher energy levels in the atom. We call this an electron’s excited state. The excited state is, however, unstable and the electron soon falls back to its original, lower energy position in the atom, we call this its ground state. As it travels back towards its ground state it releases the same amount of energy it originally absorb, however it releases it in the form of electromagnetic radiation, or light. Bohr’s Model of the Atom This lead Bohr to propose that electrons travel along well defined circular orbits, which he referred to as energy levels, that are located around the nucleus of an atom. He figured since hydrogen always emitted the same spectra of electromagnetic radiation, that it was because electrons were located in specific locations in the atom. Although the Bohr model of the atom was able to explain the emission spectrum of hydrogen with great success, it fails to predict the emission spectrum of atoms with more than one electron. Why was Bohr’s model wrong? 1. Wave-Particle Duality: electrons actually behave as both particles and waves. Waves are in continuous motion, and therefore are not localized to a specific position in space. 2. Heisenberg’s Uncertainty Principle: we cannot know where an electron is at any given moment in time - the best we can hope for is a probability picture of where the electron is likely to be. The possible positions of an electron are spread out in space in the same way that a wave is spread across a water surface. Recall that the mere act of locate an object in space (whether via naked eye or instrument) requires light to reflect off of that object. We cannot find an electron even if it is behaving as a particle. The subatomic particles are so small that light reflecting off of them will cause their trajectory to change. Why can’t we observe an electron directly? Why do we believe electrons have particle-wave duality? VIDEO: QUANTUM WORLD Figure: beam of electrons going through a slit produces a diffraction pattern, a wave pattern, instead of a straight line. Evidence for particle-wave duality of subatomic particles. Schrödinger and the Electron Cloud -Based on Heisenberg’s uncertainty principle, Erwin Schrödinger in 1926 developed a mathematical equation to figure out where an electron might be located within the electron cloud based on its level of energy. -He knew that an electron’s location within the electron cloud could not be known for certain- but a probability of where an electron might be could be calculated! -We call areas of high probability of where an electron might be an atomic orbital. Atomic orbitals describe the three-dimensional areas where there is a high probability that the electron will be located. They are regions around an atom’s nucleus in which there is a 90% probability of finding the electron. Shapes of orbitals will depend on the energy of the electron and its motion in that area. When an electron is in an orbital of higher energy, it means it is most likely far from the nucleus. s orbitals are spherical p orbitals are dumbbell-shaped. There are three of them, located on an x, y, and z axis. Despite not being able to directly detect the location of an electron directly, scientists Electron Configuration Diagram like Schrödinger have provided mathematical formulas and theoretical models for us that will help us determine where an electron MOST LIKELY is, as well as where an electron may be traveling to when it absorbs energy and travels away from the nucleus. The formulas can be summarized in something called electron configuration. 1s_ We will fill in diagonal arrows to help us follow the Aufbau Principle. 2s_ 2p_ _ _ 3s_ 3p_ _ _ 3d_ _ _ _ _ 4s_ 4p_ _ _ 4d_ _ _ _ _ 4f_ _ _ _ _ _ _ 5s_ 5p_ _ _ 5d_ _ _ _ _ 5f_ _ _ _ _ _ _ Electron Configuration Rules 1. Aufbau Principle: lowest energy orbitals are filled first 2. Pauli Exclusion Principle: up to two electrons can occupy an orbital, however they must have opposite spins so as to reduce the repulsion of like charges. (Up arrow means clockwise, down arrow means counter-clockwise.) 3. Hund’s Rule: electrons are placed into orbitals of a sublevel one electron at a time, so as to minimize the amount of repulsion in one given orbital. Once all orbitals in a sublevel are filled with one electron, then if more electrons remain a second electron with opposite spin can be added to completely fill the orbital. *Transition metals (metals located in groups 3-12) are exceptions to these rules. To be discussed later. 1. Using the diagram, determine the electron configuration of the following elements: Li, Mg, & Ar 2. State the full electron configuration of vanadium (V) and deduce the number of unpaired electrons. Answers: 1. Li: 1s²2s Mg: 1s²2s²2p⁶3s² Ar: 1s²2s²2p⁶3s²3p⁶ 2. V: 1s²2s²2p⁶3s²3p⁶3d³4s² There are three unpaired electrons. Argon’s and Vanadium’s electron configurations give us a great example for a different way of writing out an atom’s electron configuration, we call it condensed electron configuration. Ar: 1s²2s²2p⁶3s²3p⁶ V: 1s²2s²2p⁶3s²3p⁶3d³4s² Notice that Vanadium’s electron configuration is similar to Argon’s since Argon only has 18 electrons and Vanadium has 23. We can therefore condense or abbreviate the electron configuration of Vanadium like this: [Ar]3d³4s² To use this form, you must abbreviate according to the noble gas that corresponds to the specific element. The electron configurations of the first 30 elements: