Slide 1 / 184 Slide 2 / 184 New Jersey Center for Teaching and Learning Progressive Science Initiative This material is made freely available at www.njctl.org Electron Configurations and the and is intended for the non-commercial use of students and teachers. These materials may not be Periodic Table used for any commercial purpose without the written permission of the owners. NJCTL maintains its website for the convenience of teachers who wish to make their work available to other teachers, participate in a virtual professional learning community, and/or provide access to course materials to parents, students and others. Click to go to website: www.njctl.org www.njctl.org Slide 3 / 184 Slide 4 / 184 The Problem with the Nuclear Atom The Problem with the Nuclear Model The nucleus of an atom is small, 1/10,000 the size of the atom. If the Rutherford model of the atom were correct, the atom should The electrons are outside the nucleus, moving freely within the vast empty atom. The nucleus is positive; the electron is emit energy as the orbit of the electron decays. negative Since the electron would speed up as it decays, the amount of energy released should be of an increasingly higher frequency. This There is an electric force, F E = kq 1 q 2 /r 2 , pulling would create what is called a continuous spectrum representing all the electrons towards the nucleus. There is frequencies of light. no other force acting on the e- electrons;they feel a net force towards the nucleus emits energy continuous spectrum Why don't the electrons fall in... why doesn't the atom collapse into its nucleus? Slide 5 / 184 Slide 6 / 184 The Problem with the Nuclear Model The Problem with the Nuclear Model When energy is added to atoms, atoms do release energy in the form Our observations tell us the nuclear model is insufficient of light. 1. Most atoms are stable and do not release energy at all Electrons in atoms can absorb If electrons were continuously orbiting the nucleus in uniform circular energy from collisions with photons motion, they would be accelerating, and accelerating charges release or other particles and become energy. This is not observed. "excited." The excited electrons move from their initial state farther from the nucleus. Electron Absorbing Energy Then they emit energy in the form of light as they return to their original state. Electron Releasing Energy http://imagine.gsfc.nasa.gov/docs/teachers/lessons/xray_spectra/background-atoms.html
Slide 7 / 184 Slide 8 / 184 The Problem with the Nuclear Model Emission Spectra and the Bohr Model A scientist named Niels Bohr 2. When energized atoms do emit energy, a continuous spectrum interpreted these observations and created a new model of the is not produced; instead, an emission spectrum is produced atom that explained the existence displaying emitted light at specific wavelengths and frequencies. of emission spectra and provided a framework for where the electrons can exist around the e- nucleus. e- External energy added light energy (electricity, light, etc.) emitted Emission Spectrum nucleus Slide 9 / 184 Slide 10 / 184 Emission Spectra and the Bohr Model Emission Spectra and the Bohr Model No one knew what "n" was. Bohr proposed that "n" referred to a Bohr knew that the wavelengths seen in the emission spectra of particular orbit around the nucleus where an electron could be. hydrogen had a regular pattern. Each series was named after the scientist who observed these particular spectral lines. Bohr proposed that electrons could orbit the Lyman Series (spectral lines in the UV range) nucleus, like planets orbit the sun...but only in certain specific orbits. He then said that in these orbits, they wouldn't radiate Balmer Series (spectral lines in the visible and UV range) energy, as would be expected normally of an accelerating charge. These stable orbits would somehow violate that rule. Paschen Series (spectral lines in the infrared range) Each orbit would correspond to a different energy level for the electron. Slide 11 / 184 Slide 12 / 184 The Bohr Atom 1 An accelerating charge emits light energy. True The lowest energy level is called the ground state; the others are False excited states. Answer n 5 4 3 2 1
Slide 13 / 184 Slide 14 / 184 3 Why was the Nuclear Model insufficient? 2 When hydrogen atoms are energized by electricity, what is observed? A It could not explain the existence of emission spectra B It could not account for the stability of the atom Answer Answer C It required the electrons to be in the nucleus and the protons in orbit around the nucleus D A and B A A continuous spectrum of light E A, B, and C B An emission spectrum of specific colors only. C Neither a nor b Slide 15 / 184 Slide 16 / 184 Emission Spectra and the Bohr Model 4 In the Bohr model of the atom an electron in its lowest energy state Since atoms do not normally emit radiation, Bohr believed that the electrons existed in discrete stable orbits (n) around the nucleus A is in the ground state which varied in energy relative to their distance from the nucleus. B is farthest from the nucleus Answer n = 3 Bohr was able to calculate the C is in an excited state n Increasing energy = 2 energy of each of these orbits. n = 1 D emits energy n Energy (J) + 1 -2.178 x 10 -18 E both a and b 2 -5.445 x 10 -19 3 -2.417 x 10 -19 Slide 17 / 184 Slide 18 / 184 Emission Spectra and the Bohr Model Emission Spectra and the Bohr Model Bohr reasoned that each spectral line was being produced Interestingly, the energy differences between the Bohr orbits were by an electron "decaying" from a high energy Bohr orbit to found to correlate exactly with the energy of a particular spectral a lower energy Bohr orbit. lines in the emission spectra of Hydrogen! n = 3 Hydrogen atom n = 2 Hydrogen atom n = 4 n = 3 n = 1 n = 2 + n = 1 + Energy of n = 3 = -2.417 x 10 -19 J Hydrogen emission spectrum Energy of n = 2 = -5.445 x 10 -19 J Red line wavelength ( )= 656.3 nm Since only certain frequencies of light were produced, only certain E = (-2.417 x 10 -19 J) - (-5.445 x 10 -19 J) E = h/ orbits must be possible. E = 3.03 x 10 -19 J E = 3.033 x 10 -19 J EQUAL!!!
Slide 19 / 184 Slide 20 / 184 Emission Spectra and the Bohr Model Emission Spectra and the Bohr Model According to Bohr's model, first an electron is excited These possible energy states for atomic electrons were quantized – from its ground state by absorbing energy. only certain values were possible. The spectrum could be explained as transitions from one level to another. n = 4 n = 3 Electrons would only radiate when they moved between orbits, n = 2 not when they stayed in one orbit. photon n = 1 e- upper + upper e- lower lower Slide 21 / 184 Slide 22 / 184 Review: Emission Spectrum of Hydrogen Emission Spectra and the Bohr Model Hydrogen atoms have one proton and one electron. The emission spectrum of hydrogen shows all of the different possible Once an electron is excited, it can take any number of routes wavelengths of visible light emitted when an excited electron back to its ground state, so long as it is releasing energy in returns to a lower energy state. discrete quantitized packets. Transition light emitted 6 2 410 nm Here we see 2 separate emissions coming from the same electron. The electron can either go from n=3 right to n=1 or it 4 2 486 nm can go from n=3 to n=2 to n=1. + n = 4 n = 4 n = 3 n = 3 n = 2 n = 2 3 2 656 nm n = 1 n = 1 + + Click here for Bohr model animation Both are acceptable and both will occur. Slide 23 / 184 Slide 24 / 184 Emission Spectra and the Bohr Model Emission Spectra and the Bohr Model The difference in energy between the orbits decreases as Due to the differing numbers of protons in the nucleus and one moves farther from the nucleus. number of electrons around them, each atom produces a unique emission spectrum after being energized. n = 3 n = 2 wavelength of Transition spectral line Energy (J) n = 1 produced (nm) + Since the emission spectrum 3 --> 2 656 3.03 x 10 -19 of each element is unique, it 2 --> 1 122 1.63 x 10 -18 can be used to identify the presence of a particular element.
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