Covalent bond
The atomic bond ( covalent bond, or electron pair binding homopolar bond ) is a form of chemical bonding, and as such is responsible for the cohesion strength of atoms in molecular assembly chemical compounds. Atomic bonds are relevant for all atomic systems with the exception of noble gases, which have completely filled electron shells. Atomic bonds are characteristic between the atoms of non-metals, such as in molecules. You are spatially oriented and primarily localized on the atomic distance of nearest neighbors. In semiconductor systems, this is reflected in characteristic crystal structures. In ionic crystals, however, act predominantly ionic and metallic bonds in metals. Here is the difference between metallic bond quantum mechanically only a gradual nature and reflects the spatial extent of the involved electron orbitals. Metallic bonds are gradually delocalized and the spatial direction by the delocalization weak.
When atomic bonds, the interaction of the outer electrons ( valence electrons ) with the atomic nuclei of the atoms involved plays the major role. The atoms forming between them of at least a pair of electrons. This electron pair holds two (two- center bond ) or more ( multi -center bond ) atoms together, so it is binding and is therefore called bonding electron pair. In addition to a bonding electron pair ( single bond ), two ( double bond) or three ( triple bond) and very rarely even four ( quadruple bond ) or five ( quintuple bond ) can act electron pairs. An atomic bond has a certain effect direction, that is a directed binding and thus determines the geometric structure of a compound. The strength of bonding will be described by the binding energy. In the chemical reaction of corresponding substances with each other socializing or separation takes place one or more atomic bonds.
- 3.1 dipole moment
- 5.1 Spatial orientation
- 5.2 bond length
- 5.3 Geometry of multiple bonds
- 5.4 Conjugated double bonds and aromatic bonds
Basics
From experience it is known that does not form molecules in chemical reactions in any combination of types of atoms and atomic numbers. The electron shells of atoms of the same or of different elements must be suitable to form bonds with each other. However, a detailed description of the electron shells is only with more sophisticated mathematical methods (see molecular orbital theory, Valenzstrukturtheorie ). An important and powerful tool for the understanding of bonding relationships is the less complicated noble gas rule. It allows the graphical representation of many chemical compounds as Valenzstrichformeln in which bonding electron pairs are indicated by dashes between the element symbols.
Noble gas rule
According to Lewis and Kossel (1916 ) chemical compounds are particularly stable when the atoms involved in the periodic table the nearest noble gas configuration reach ( noble gas rule). The nearest noble gas hydrogen is helium with only two electrons. Is hydrogen, the control is therefore filled with two electrons and is, accordingly, present as the molecule, which is held together by atomic bonding (H -H).
In many cases, atoms in compounds achieve a valence shell with four pairs of electrons, so they have an octet of electrons and satisfy the so-called octet rule. The octet rule is true for most compounds of main group elements. This rule of thumb grips good for the second period elements such as carbon, nitrogen and oxygen, which are the important elements of numerous organic compounds.
Physics of atomic bonding
Necessary for binding of two atoms (or 2 bodies in general ) at a fixed distance is the balance between an attractive and a repulsive force. If we consider the simple case of the hydrogen molecule results in the following: The repulsive force is the electromagnetic force between the positive nuclei, which is getting stronger ever closer together the nuclei. The attractive force is caused by the exchange interaction of two electrons: the wave function of two electrons results in the symmetric case a probability of electrons between the nuclei. Thus, the electrons screen the nuclear charges from each other and are even largely confined between the cores. The result is a bonding electron pair. A detailed description can be found in Article molecular orbital theory, Section hydrogen.
Bonding electron pairs
By atomic bonds ( covalent bonds ) exist molecular weight substances, such as oxygen (O2) or carbon dioxide ( CO2), but also materials such as diamond ( CDiamant ) or silicon dioxide (SiO2), which do not form molecules but atomic lattice. Complex ions, ie molecules that carry electric charges are held together by atomic bonds. While these ions form salts ( SO42 - ) are, however, held together by covalent bonds, by ionic bonds, which atoms of the complex ions such as ammonium (NH4 ) or sulfate.
In a formula, the valence electrons of nonmetal atoms can be represented pictorially, where electrons are distributed among four possible positions around the element symbol. Points represent individual electrons, while strokes lone pairs of electrons (also: lone pair nonbonding electron pair ) symbolize. The formulas electrons of the atoms can be combined to form molecules of many known chemical compounds and predict with known atomic composition of small molecules, the molecular structure of a compound. In order to arrive at a valence-bond of a molecule, lone (points) are combined to form bonding electron pairs ( dashes) or electron pairs (dashes ) between atoms shifted so that the octet rule is satisfied. It also double bonds and triple bonds between two atoms are possible (see also: Lewis notation).
The formal assignment of bonding and nonbonding electron pairs to represent a chemical compound occasionally leads to a so-called formal charge. It is the difference between the positive nuclear charge and atomic allocated negative electrons and is often given as a superscript plus or minus sign in a circle icon. Formal charges are in carbon monoxide for example.
Is the coordination number of the atom is the number of nearest neighbor atoms, and is, for example, with respect to the carbon atom of carbon monoxide in 1, wherein carbon dioxide and methane 2 4
Polarity of atomic bonds
→ Main article: Polar Covalent Bonding
The electron-withdrawing forces ( electronegativity, En) are a measure of the ability of an atom to pull in a chemical bond, the bonding electrons to itself. The electronegativity of binding partners is exactly the same only with element molecules and only here are ideal atomic bonds before. Strictly speaking, only these bonds can be called non-polar or homöopolar.
Are there differences between the binding partners in their electronegativity, on the other hand are polar or polar -called hetero atom bonds before. The bonding electrons are more or less non-uniformly distributed between the binding partners. Their focus is shifted towards the more electronegative partner. The atom with the greater electronegativity pulls the bonding electrons closer to itself. This gives these binding partners a negative charge, which is symbolized by δ -. The electron shell of the atom at the other end of the binding according to the negative charge density is depleted and the atom is replaced by a positive charge ( δ ). They are called atomic bonds polar bonds because poles arise at different part loads.
For very polar atomic bonds bonding electrons can be largely associated with a binding partner. There is the limiting case, to ionic bonds and, in some cases it is useful to describe the compound as ionic.
Dipole moment
Polar compounds may result in the entire molecule is polar The molecule then carries a dipole moment and is present as a molecule before dipole. If a molecule has a ( measurable ) dipole moment, however, depends not only on the polarity of the compounds, but also on the molecular structure. The dipole moments of various bonds in the molecule add up depending on the direction ( vectorial) and can therefore cancel each other out. Hydrogen fluoride contributes as diatomic, heteronuclear compound is a dipole moment. Carbon dioxide has a total dipole moment of zero, since the bond dipoles are oriented in opposite directions and cancel each other out. Water has a larger overall dipole as hydrogen fluoride, although the polarity of the O -H bond is smaller than that of the H -F bond. This is due to the addition of the two O-H bond dipoles, (see below ) are in a bond angle of about 105 ° to each other.
Bonding electron pairs of lone pairs of electrons
Lone electron pairs of a connection can take over the role of nonbonding electron pairs in one reaction. This type of bonds is coordinative bond: (also called a dative covalent bond ) and occurs in compounds such as ammonium cation and in complex compounds. Coordination bonds have a similarity with the weak bonds, which occur for example in the hydrogen bonding.
Geometry
Spatial orientation
→ Main article VSEPR model
Three interconnected atoms in a lattice of atoms, molecule or complex are arranged in a particular bond angle. Knowledge of bond angle design allows the structural formula of a compound. From knowledge of bonding and nonbonding electron pairs in a compound to bond angle can be estimated by using the electron pair repulsion model. The bond angles result from an arrangement of the electron clouds of each other in a large distance. An electron cloud can be a single electron ( in radicals ), a non-bonding electron pair or single bonds. For a simple estimate of double and triple bonds can be conceptually understood as a single cloud.
Final deviations can occur between an estimate of a bond angle with the aid of the electron cloud model and real molecules. The actual bond angle in the water molecule is not 109.47 °, 104.45 ° but due to the lower repulsive effect of the lone pairs on the bonding pairs, and the smaller size of sq- bonding orbitals that contain the proton.
Bond length
The interatomic distances in molecules and complexes with covalent bonding can be determined experimentally by analyzing the rotational spectra. The bond lengths depend on the size of the atoms. The larger its radius, the greater the distance between them.
When bonds between similar atoms the distance between them is also on the number of bonding electron pairs depends: act The more bonding electron pairs, the shorter the bond length. The shape of the binding potential can be described by the Morse potential.
Geometry of multiple bonds
Although single bonds determine the bond angles between atoms, but are rotated in itself. A molecule such as n- butane can easily turn into itself and is therefore in different conformations. All conformations describe the same compound. Multiple bonds can be, however, do not turn in to. Have meaning here the double bond, especially in organic compounds. Hydrocarbons such as 2-butene exist as two different chemical compounds, namely cis -2 and trans-2- butene. The rigidity of the double bond leads generally to the so-called cis-trans isomerism. Can be described by the Carter - Goddard - Malrieu - Trinquier model further distortions of multiple bonds.
Conjugated double bonds and aromatic bonds
If, as a molecule alternately double and single bonds on, the interatomic distances of the single bond shorter ( solid ) than single bonds without double bonds in the neighborhood. On the other hand, multiple bonds will be affected lengthening. This phenomenon is called conjugation and can hardly be explained by the binding models described here, simple.
A special case exists when the aromaticity: Here are just a formal sequences of double and single bonds, but the bond lengths are all the same short. A simple aromatic compound is the circular molecule of benzene ( C6H6). Valenzstrichformeln this connection lead to two possible representations, which are referred to in the figure as resonance structures. Both Valenzstrichformeln lead to the correct assumption that benzene is a flat ( planar ) molecule, since the geometry of trigonal planar orientations calls. Any carbon-carbon bond can be depicted as a double or single bond. In reality, the double bonds are not at fixed locations before, but are distributed over the whole ring ( delocalized ). All aromatic compounds, ie compounds with delocalized double bonds must satisfy the so-called Hückel rule that is quantum mechanically justified.
Creating resonance structures with the binding models described here also allows simple estimates of rather complicated bonding situations. The figure at right shows the peptide bond in two limiting structures. From the resonance structure 1 can be a CNC bond angle of 109 ° suspect ( tetrahedral ), while resonance structure 2 to an angle of 120 ° indicates ( trigonal planar). In reality, a bond angle of 122 ° is present, as it results rather from resonance structure 2 with formal charges. The CN distance of the possible double bond is 133 pm between a C -N single bond (147 pm ) and a C = N double bond (130 pm ).
Binding energy
The binding energy (also: dissociation, bond cleavage energy of binding, or Bindungsdissoziationsenthalpie Valenzenergie ) is equal to the energy that is required to a cleavage of an atomic bond and a compound ( A-B) in two radicals converted:
This dissociation is called homolytic cleavage. The Bindungsdissoziationsenthalpie can be measured in simple molecules and assess in more complex molecules by measurements and calculations. It depends - as the bond length (see above) - the size of the bound atoms from. The larger the radius of the binding partner, the greater their distance and the smaller their binding energy. Even when bonds between identical atoms can be seen that the distance between them increases with the number of bonding electron pairs is low, however, their binding energy increases.