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The concept of ionization energy, its trends in the periodic table, and its significance in determining element reactivity. It also includes several learning objectives and questions for further study.
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Learning Objectives
Ionization Energy Ionization is the process by which atoms become ions, specifically positive ions (cations, pronounced cat-ions, like a charged feline). If there is enough energy available, an electron will be removed from an atom. The energy required to remove an electron is called ionization energy , and it varies from one element to another, but is very predictable with the trends in the periodic table. The energy required to remove an electron from a neutral atom is known as the first ionization energy, and is given by the general formula model
The energy required to remove two electrons follows (a second electron is removed),
…and so on… This energy is usually expressed in kJ/mol (kilojoules per mole), which means the amount of energy it takes for all the atoms in a mole of the substance to lose one electron each. Trends in ionization energy Ionization energy is closely related to Coulomb’s Law. If a negatively-charged electron is closer to the positively-charged nucleus, it will have a greater attractive force, and therefore be held more tightly by the nucleus. As a result, elements with smaller atoms tend to have higher ionization energies. Contributing to this feature, atoms in a certain period (horizontal row on the periodic table) with more protons tend to have a greater force holding on to electrons in the same energy level. Therefore, Ionization energy increases from left to right across the periodic table. If an electron is further away from the nucleus, the attractive force decreases, as described and predicted by Coulomb’s Law. The outer electrons (known as the valence electrons) are the first
noble Simplified charts show just the noble gases:
Similar trends appear for Second Ionization Energies (IE2), third ionization energies (IE3), and so on: Ionization Energies (kJ/mol) IE 1 IE 2 IE 3 IE 4 IE 5 IE 6 IE 7 IE 8 H 1312 He 2372 5250 Li 520 7297 11810 Be 899 1757 14845 21000 B 800 2426 3659 25020 32820 C 1086 2352 4619 6221 37820 47260 N 1402 2855 4576 7473 9442 53250 64340 O 1314 3388 5296 7467 10987 13320 71320 84070 F 1680 3375 6045 8408 11020 15160 17860 92010 Ne 2080 3963 6130 9361 12180 15240 20000 23100 Na 496 4563 6913 9541 13350 16600 20113 25666 Mg 737 1450 7731 10545 13627 17995 21700 25662 As more electrons are removed, it becomes much more difficult to remove remaining electrons because nuclear attraction increases greatly when shielding is not present. Each successive ionization energy shows a similar trend, but shifted. Second Ionization Energy is greatest in the first group (alkali metals) and lowest in the second group (alkaline earth metals), Third Ionization Energy is greatest in the second group, and lowest in group 13 (Boron family)
7 ) Which group would generally have the lowest first ionization energy? a. Transition Metals (Groups 3 - 12) b. Alkali Metals (Group 1) c. Noble Gases (Group 18) d. Alkaline Earth Metals (Group 2) e. Halogens (Group 17) Justify your response. 8 ) Sulfur has a first ionization energy of 999.6 kJ/mol. Rubidium has a first ionization energy of 403 kJ/mol. What bond do they form when chemically combined? a. Covalent b. Polar Covalent c. Ionic 9 ) Low first ionization energy is considered a property of a. Metals b. Nonmetals Justify your response The following are “higher order” questions for Honors, but CP are encouraged to try as well. 10 ) Ionization energy, when supplied to an atom, results in a(n) a. Anion and a proton b. Cation and a proton c. Cation and an Electron d. Anion and an electron 11 ) Gallium has a first ionization energy of 580 kJ/mol, and Lithium 520 kJ/mol. Justify this based on atomic structure and periodic trends.