Day 14: Covalent Bonding (Sharing electrons, Naming, Identifying); Day 15: Polar or Non-Polar (Solutions, solubility, anesthetics, ethics and application); Day 14 & 15 - Reactions - Slides 106 - 129 The term "electronegativity" was introduced by Jöns Jacob Berzelius in 1811, though the concept was known even before that and was studied by many chemists including Avogadro. For Berzelius, all chemical compounds contained two electrically opposing constituents: the acidic, or electronegative, and the basic, or electropositive. In spite of its long history, an accurate scale of electronegativity was not developed until 1932, when Linus Pauling proposed an electronegativity scale, which depends on bond energies, as a development of valence bond theory. The opposite of electronegativity is electropositivity: a measure of an element's ability to donate electrons. As it is usually calculated, electronegativity is not a property of an atom alone, but rather a property of an atom in a molecule. Properties of a free atom include ionization energy and electron affinity. It is to be expected that the electronegativity of an element will vary with its chemical environment, but it is usually considered to be a transferable property, that is to say that similar values will be valid in a variety of situations. The first paper from Pauling on electronegativity "The Nature of the Chemical Bond. IV. The Energy of Single Bonds and The Relative Electronegativity Of Atoms" published in 1932 gives a different scale with H = 0.0 and F = 2.0. Later, he changed the values to 2.1 for H and 4.0 for F [http://scarc.library.oregonstate.edu/coll/pauling/bond/notes/sci5.001.14notes-01.html]. As only differences in electronegativity are defined, it is necessary to choose an arbitrary reference point in order to construct a scale. Hydrogen was chosen as the reference, as it forms covalent bonds with a large variety of elements: its electronegativity was fixed first at 2.1, later revised to 2.20. But people still find the scale Pauli used from lithium having an electronegativity around 1.0 to fluorine having an electronegativity of 4.0 the easiest to remember. Values in the change in electronegativity between 0.4 and zero are considered to have non-polar covalent bond characteristics while electronegativity differences of 0 indicate a completely non-polar covalent bond Polar and nonpolar Bond polarity is important when considering bond strength. If there is a single bond sigma bond, the smaller atoms in molecules such as H2 form stronger bonds than the larger atoms in molecules such as the carbon-carbon bonds in many organic compounds. However, the nonpolar HF bond is much stronger than the single bond in hydrogen gas. The carbon-carbon double bond is stronger than the carbon-carbon single bond in most organic compounds, but not by two folds. However, the polar carbon-oxygen double bond is more than two folds stronger than that of the carbon-oxygen bond in most organic compounds. We will find out that this is because of the freedom in space pi electrons have versus the directional sigma electrons. Valence electrons and the periodic table Valance electrons are the electrons which are not the core electrons. For main group elements, the trends are usually seen as group trends. A) Group 1: the alkali metals (ashes), ns1 silvery, luster, high thermal and electrical conductivity, colorless except with colored anions, melt point decrease, density increase, radius increases, first ionization energy decrease (up to down the periodic table), reacts with oxygen and water, forms plus one cations. B) Group 2: the alkaline earth metals, ns2 harder and less reactive than group 1 elements, colorless or white except with colored anions, magnesium needs steam to react while calcium will react with liquid water, melt point decrease, density increase, radius increases, first ionization energy decrease (up to down the periodic table). C) Hydrogen can act as a metal or nonmetal, 1s1 1s2, hydrogen cation; 1s0, hydrogen anion; 1s1, atomic hydrogen; reacts with metals to form hydrides and is the most common atom. D) Group 13, the boron group, ns2np1 E) Group 14, the carbon group, ns2np2 F) Group 15, the pnictides (older name, now used for a sub-group within the pnictogens), ns2np3 A pnictogen is a member of the nitrogen group of elements, Group 15 of the periodic table (formerly numbered as Group V or Group VA). G) Group 16, the calcogens, ns2np4 melting point increase (except polonium), density increase, radius increases, first ionization energy decrease (up to down the periodic table). H) Group 17, the halogens, ns2np5 melting point increases, density increase, radius increases, first ionization energy decrease (up to down the periodic table). I) Group 18, the noble gas, ns2np6 mono-atomic gases, generally inert, boiling point increases, density increase, radius increases, first ionization energy decrease (up to down the periodic table), xenon was shown to react with fluorine in 1963, other reactions followed. Size of atom As we have calculated and observed, the periodic table seems to have some order to it related to quantum mechanics. This is a modern idea. John Newlands "thought up" the "Law of Octaves." This is somehow related to the idea of closed orbitals for period two. By 1869, Mendeleev in Russia and Meyers in Germany agreed with the periodic arrangement of the elements. Mendeleev offered experiment prediction of new elements or eka elements. He examined physical properties such as atomic weight, density, specific heat, melting point, color, oxides, chlorides and boiling points. By 1913, a young man who died in World War I named Moseley, developed the idea of atomic number. As we examine periodic families from up to down, the number of electrons increase and so does the atomic radii for main group elements. Examining the periodic periods for the main group elements from left to right, the size of the atomic radii decreases even though the number of electrons increase. In general, the transition elements decrease at first and then increases. It takes a detailed knowledge of quantum mechanics to examine each exception to our simple rules. Lanthanides are about the same but mostly decreases. Find a book to look at the many different charts on this subject. Nuclear Charge There are electron shells in atoms. The atomic number corresponds to the number of protons in the nucleus. Each proton has a plus one charge. Examining the periodic periods for the main group elements from left to right, the nuclear charge increases. This is why the atomic radii decease. As we examine the noble gases from up to down, the K shell is drawn closed to the nucleus. Helium has only the K shell. Neon has the K shell and the L shell. So the K shell for neon is closer to the nucleus than the K shell for Helium. However, Neon has a larger atomic radius than helium. Argon is even bigger. Ionization energy The ionization energy for gases ought to be absent of many atomic interactions. We would really like to measure the ionization energy of an isolated atom. However, we cannot do that at the present moment. The ionization energy should be endothermic because the electron must leave the atom to obtain the first ionization energy. If we include quantum effects, then some changes in simple periodic predictions occur. For example, the first ionization energies of hydrogen and helium increase while it decreases for lithium. Yet, It also decreases when we compare the atomic pairs of beryllium and boron and of carbon and nitrogen. Normal periodic trends of the first ionization energy would be an increase of energy from left to right across a period and a decrease of energy up to down a main group family. Hund’s rule clearly explains the deviation from the general rule. With beryllium and boron, the boron atom has only one p-orbital electron. This electron is easily ionized so that a closed sub-shell can form with an electron pair in the 2s-orbital. Beryllium already has a closed sub-shell, and the electron is more difficult to ionize even though the nuclear charge is less for beryllium than for boron. Nitrogen has a half filled sub-shell with an electron in each orbital while oxygen has one coupled pair. It is easier for the oxygen to loss one electron so that the spaciously stable half filled sub-shell can be obtained. Electron affinity From general periodic trends, we would expect electron affinities to be exothermic because an electron attaches to the atom in the gases phase. By the octette rule, the halogens should have the strongest electron affinity and the noble gases should have the lowest. However, due to quantum effects, the group two elements have the lowest electron affinity with beryllium having the lowest for the atoms. As excepted, nitrogen had a low electron affinity which is around zero. In general, the electron affinity increases for main group elements as one examines from left to right across a period. A new effect called charge density, charge/atomic radii, causes the electron affinity of fluorine to be lower than the electron affinity of chlorine in the gas phase. Metal vs. nonmetal Metallic elements versus nonmetallic elements: A) Distinguishing luster versus various colors B) Malleable and ductile solids versus brittle solids C) Thermal and electrical conductivity versus poor conductors D) Oxides (basic, ionic solids) versus oxides (acidic, molecular solids) E) Aqueous solution, cations versus aqueous solutions anions or oxyanions Metal Rxn A) Low oxidation energy--easy to oxidize B) Mostly solids at room temperature except for liquid mercury, cesium and gallium melt at slightly above room temperature C) Metals + oxygen ÷ metal oxide 2Ni(s) + O2(g) ÷ 2NiO(s) D) Metal oxide + water ÷ metal hydroxide CaO(s) + HOH(l) ÷ Ca(OH)2(aq) E) Metal oxide + acid ÷ salt + water MgO(s) + 2HCl(aq) ÷ MgCl2(aq) + HOH(l) Non-metal Rxn A) Nonmetal commonly forms hydrides, oxides, halides, low melting solids, gases and liquids B) Metal + nonmetal ÷ salt 2Al(s) + Br2(l) ÷ 2AlBr3(s) C) Nonmetal oxide + water ÷ acid CO2(g) + HOH(l) ÷ H2CO3(aq) D) Metal oxide + base ÷ salt + water SO3(aq) + 2KOH(aq) ÷ K2SO4(aq) + HOH(l) Metalloid Rxn The general periodic trends with metals are an increase in metallic character as one looks up to down a group or family. With nonmetals, an increase in nometallic character is seen as one looks down to up. As one looks across a period from left to right, an increase in nonmetallic character is seen, And as one looks across a period from right to left, an increase in metallic character is seen. The properties and reactions of metalloids are generally in between those of metals and nonmetals. This leads to the developing field of electronics and many other things.
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