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Fluorine has the greatest electronegativity of all the elements, and the heavier alkali metals such as potassium, rubidium and cesium have the lowest electronegativities. It should be noted that carbon is about in the middle of the electronegativity range, and is slightly more electronegative than hydrogen.
The halogens are one electron short of a valence shell octet, and are among the most reactive of the elements. In their chemical reactions halogen atoms achieve a valence shell octet by capturing or borrowing the eighth electron from another atom or molecule. The alkali metals are also exceptionally reactive, but for the opposite reason. These atoms have only one electron in the valence shell, and on losing this electron arrive at the lower shell valence octet. As a consequence of this electron loss, these elements are commonly encountered as cations (positively charged atoms).
1-3 Charge Distribution If the electron pairs in covalent bonds were shared absolutely evenly there would be no fixed local charges within a molecule. Although this is true for diatomic elements such as H2, N2 and O2, most covalent compounds show some degree of local charge separation, resulting in bond and / or molecular dipoles. 1-3-1 electronegativity Different atoms have different affinities for nearby electrons. The ability of an element to attract or hold onto electrons is called electronegativity.
1-1 Electronic Configurations Electron Configurations in the Periodic Table
The periodic table shown here is severely truncated. There are, of course, over eighty other elements.
Chapter I Structure & Bonding 1-1 Electron Configurations of Atoms 1-2 Chemical Bonding & Valence 1-3 Charge Distribution in Molecules 1-4 The Shape of Molecules 1-5 Isomers 1-6 Resonance 1-7 1-8 Atomic and Molecular Orbitals The study of organic chemistry must at some point extend to the molecular level, for the physical and chemical properties of a substance are ultimately explained in terms of the structure and bonding of molecules. This module 模块 introduces some basic facts and principles that are needed for a discussion of organic molecules.
1-2-3 Valence
The number of electrons an atom gain or lose to achieve a valence octet is called valence.
The valences here represent the most common form 普通形式 in organic compounds. Many elements, such as chlorine, bromine and iodine, are known to exist in several valence states in different inorganic compounds.
1-2-2 Covalent Bonding
Water is a liquid at room temperature; carbon dioxide and carbon tetrafluoride are gases. None of these compounds is composed of ions. A different attractive interaction between atoms, called covalent bonding, is involved here.
1-3-2 Polar Covalent Bonds
When two different atoms are bonded covalently, the shared electrons are attracted to the more electronegative atom of the bond, resulting in a shift of electron density toward the more electronegative atom. Such a covalent bond is polar, and will have a dipole. The degree of polarity and the magnitude of the bond dipole will be proportional to the difference in electronegativity of the bonded atoms. Thus a O–H bond is more polar than a C–H bond, with the hydrogen atom of the former being more positive than the hydrogen bonded to carbon.
1-2 Chemical Bonding and Valence 1-2-1 Ionic Bonding When sodium is burned in a chlorine atmosphere, it produces sodium chloride. This has a high melting point (800 º C) and dissolves in water to give a conducting solution. Sodium chloride is an ionic compound, in which an electron of sodium atom was transferred to a chlorine atom and generates a sodium cation and a chloride anion. Electrostatic attraction results in these oppositely charged ions packing together in a lattice. The attractive forces holding the ions in place can be referred to as ionic bonds.
The shift of electron density in a covalent bond toward the more electronegative atom or group can be observed in several ways. For bonds to hydrogen, acidity is one criterion. If the bonding electron pair moves away from the hydrogen nucleus the proton will be more easily transfered to a base (it will be more acidic). Methane is almost non-acidic, since the C–H bond is nearly non-polar. The O–H bond of water is polar, and it is at least 25 powers of ten more acidic than methane. H–F is over 12 powers of ten more acidic than water as a consequence of the greater electronegativity difference in its atoms. Electronegativity differences may be transmitted through connecting covalent bonds by an inductive effect. This inductive transfer of polarity tapers off as the number of transmitting bonds increases, and the presence of more than one highly electronegative atom has a cumulative effect. For example, trifluoro ethanol, CF3CH2– O–H is about ten thousand times more acidic than ethanol, CH3CH2–O–H.
Covalent bonding occurs by sharing of valence electrons, rather than an outright electron transfer. Similarities in physical properties (they are all gases) suggest that the diatomic elements H2, N2, O2, F2 & Cl2 also have covalent bonds. Carbon dioxide is notable because it is a case in which two pairs of electrons are shared by the same two atoms. This is an example of a double covalent bond.
Likewise, C–Cl and C–Li bonds are both polar, but the carbon end is positive in the former and negative in the latter. The dipolar nature of these bonds is often indicated by a partial charge notation (δ+/–) or by an arrow pointing to the negative end of the bond.
Fluorine has the greatest electronegativity of all the elements, and the heavier alkali metals such as potassium, rubidium and cesium have the lowest electronegativities. It should be noted that carbon is about in the middle of the electronegativity range, and is slightly more electronegative than hydrogen.
The halogens are one electron short of a valence shell octet, and are among the most reactive of the elements. In their chemical reactions halogen atoms achieve a valence shell octet by capturing or borrowing the eighth electron from another atom or molecule. The alkali metals are also exceptionally reactive, but for the opposite reason. These atoms have only one electron in the valence shell, and on losing this electron arrive at the lower shell valence octet. As a consequence of this electron loss, these elements are commonly encountered as cations (positively charged atoms).
1-3 Charge Distribution If the electron pairs in covalent bonds were shared absolutely evenly there would be no fixed local charges within a molecule. Although this is true for diatomic elements such as H2, N2 and O2, most covalent compounds show some degree of local charge separation, resulting in bond and / or molecular dipoles. 1-3-1 electronegativity Different atoms have different affinities for nearby electrons. The ability of an element to attract or hold onto electrons is called electronegativity.
1-1 Electronic Configurations Electron Configurations in the Periodic Table
The periodic table shown here is severely truncated. There are, of course, over eighty other elements.
Chapter I Structure & Bonding 1-1 Electron Configurations of Atoms 1-2 Chemical Bonding & Valence 1-3 Charge Distribution in Molecules 1-4 The Shape of Molecules 1-5 Isomers 1-6 Resonance 1-7 1-8 Atomic and Molecular Orbitals The study of organic chemistry must at some point extend to the molecular level, for the physical and chemical properties of a substance are ultimately explained in terms of the structure and bonding of molecules. This module 模块 introduces some basic facts and principles that are needed for a discussion of organic molecules.
1-2-3 Valence
The number of electrons an atom gain or lose to achieve a valence octet is called valence.
The valences here represent the most common form 普通形式 in organic compounds. Many elements, such as chlorine, bromine and iodine, are known to exist in several valence states in different inorganic compounds.
1-2-2 Covalent Bonding
Water is a liquid at room temperature; carbon dioxide and carbon tetrafluoride are gases. None of these compounds is composed of ions. A different attractive interaction between atoms, called covalent bonding, is involved here.
1-3-2 Polar Covalent Bonds
When two different atoms are bonded covalently, the shared electrons are attracted to the more electronegative atom of the bond, resulting in a shift of electron density toward the more electronegative atom. Such a covalent bond is polar, and will have a dipole. The degree of polarity and the magnitude of the bond dipole will be proportional to the difference in electronegativity of the bonded atoms. Thus a O–H bond is more polar than a C–H bond, with the hydrogen atom of the former being more positive than the hydrogen bonded to carbon.
1-2 Chemical Bonding and Valence 1-2-1 Ionic Bonding When sodium is burned in a chlorine atmosphere, it produces sodium chloride. This has a high melting point (800 º C) and dissolves in water to give a conducting solution. Sodium chloride is an ionic compound, in which an electron of sodium atom was transferred to a chlorine atom and generates a sodium cation and a chloride anion. Electrostatic attraction results in these oppositely charged ions packing together in a lattice. The attractive forces holding the ions in place can be referred to as ionic bonds.
The shift of electron density in a covalent bond toward the more electronegative atom or group can be observed in several ways. For bonds to hydrogen, acidity is one criterion. If the bonding electron pair moves away from the hydrogen nucleus the proton will be more easily transfered to a base (it will be more acidic). Methane is almost non-acidic, since the C–H bond is nearly non-polar. The O–H bond of water is polar, and it is at least 25 powers of ten more acidic than methane. H–F is over 12 powers of ten more acidic than water as a consequence of the greater electronegativity difference in its atoms. Electronegativity differences may be transmitted through connecting covalent bonds by an inductive effect. This inductive transfer of polarity tapers off as the number of transmitting bonds increases, and the presence of more than one highly electronegative atom has a cumulative effect. For example, trifluoro ethanol, CF3CH2– O–H is about ten thousand times more acidic than ethanol, CH3CH2–O–H.
Covalent bonding occurs by sharing of valence electrons, rather than an outright electron transfer. Similarities in physical properties (they are all gases) suggest that the diatomic elements H2, N2, O2, F2 & Cl2 also have covalent bonds. Carbon dioxide is notable because it is a case in which two pairs of electrons are shared by the same two atoms. This is an example of a double covalent bond.
Likewise, C–Cl and C–Li bonds are both polar, but the carbon end is positive in the former and negative in the latter. The dipolar nature of these bonds is often indicated by a partial charge notation (δ+/–) or by an arrow pointing to the negative end of the bond.