Anomaly #1 By following the atomic radius trend on the periodic table, one would assume that hydrogen would have a greater atomic radius than helium. This notion would coincide with the ionization energy trend as well, since helium should have a higher attraction to its own electrons, according to ionization energy principles. Helium’s higher ionization energy would suggest a smaller size than hydrogen’s due to hydrogen having a lower ionization energy level. But, as shown on the “Atomic Radius and Ionization Energy vs. Atomic Number” graph, the atomic radius of helium is instead greater than the atomic radius of hydrogen. Although not conclusive, this anomaly may be explained through valence shell electron pair repulsion otherwise known as, “VSEPR”. After …show more content…
Instead, the ionization energy level of nitrogen is actually greater than oxygen. The anomaly can be examined more thoroughly through orbital diagrams of both nitrogen and oxygen. Since oxygen is the eighth element, it has 8 electrons, which would make its electron configuration 1s22s22p4 (University of Maryland, 2016), As for nitrogen, it is the seventh element with 7 electrons. Nitrogen’s electron configuration is 1s22s22p3(University of Maryland, 2016). The orbital diagrams for both of these elements are very similar with the difference of 1 electron. In nitrogen’s case, the 2p orbitals are all half-filled, whereas oxygen carries one extra electron in one of its 2p orbitals. Due to the extra electron in oxygen’s 2p orbital, oxygen may want to dispose of the extra electron in order to fulfill a more stable configuration- one that resembles nitrogen’s. Thus, nitrogen may possess a higher ionization energy level due to its more stable form- where it would rather keep all of its half-filled
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Atoms are electrically neutral; the electrons that bear the negative charge are equal in number to the protons in the nucleus
In 1907, Einstein used Planck’s hypothesis of quantization to explain why the temperature of a solid changed by different amounts if you put the same amount of heat into the material. Since the early 1800’s, the science of spectroscopy had shown that different elements emit and absorb specific colors of light called “spectral lines.” In 1888, Johannes Rydberg derived an equation that described the spectral lines emitted by hydrogen, though nobody could explain why the equation worked. This changed in 1913 when Danish physicist Niel Bohr applied Planck’s hypothesis of quantization to Ernest Rutherford’s 1911 “planetary” model of the atom, which affirmed that electrons orbited the nucleus the same way that planets orbit the sun. Bohr offered an explanation for why electrical attraction does not make the electrons spiral into the nucleus. He said that electrons in atoms can change their energy only by absorbing or emitting quanta. When an electron absorbs a quantum it moves quickly to orbit farther from nucleus. When an electron emits a quantum the electron jumps to a closer
An atom, by definition, is the smallest part of any substance. The atom has three main components that make it up: protons, neutrons, and electrons. The protons and neutrons are within the nucleus in the center of the atom. The electrons revolve around the nucleus in many orbitals. These orbitals consist of many different shapes, including circular, spiral, and many others. Protons are positively charged and electrons are negatively charged. Protons and electrons both have charge of equal magnitude (i.e. 1.602x10-19 coulombs). Neutrons have a neutral charge, and they, along with protons, are the majority of mass in an atom. Electron mass, though, is negligible. When an atom has a neutral charge, it is stable.
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“The half-life of a radioisotope is the time required for half the atoms in a given sample to undergo radioactive decay; for any particular radioisotope, the half-life is independent of the initial amount of...
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Ionic compounds, when in the solid state, can be described as ionic lattices whose shapes are dictated by the need to place oppositely charged ions close to each other and similarly charged ions as far apart as possible. Though there is some structural diversity in ionic compounds, covalent compounds present us with a world of structural possibilities. From simple linear molecules like H2 to complex chains of atoms like butane (CH3CH2CH2CH3), covalent molecules can take on many shapes. To help decide which shape a polyatomic molecule might prefer we will use Valence Shell Electron Pair Repulsion theory (VSEPR). VSEPR states that electrons like to stay as far away from one another as possible to provide the lowest energy (i.e. most stable) structure for any bonding arrangement. In this way, VSEPR is a powerful tool for predicting the geometries of covalent molecules.