What is the most stable ionic compound
Chemical bond is the name for the cohesion of the smallest particles in chemical substances. The smallest particles can be atoms, anions, cations or molecules. By loosening and tying chemical bonds in a chemical reaction, substances are converted into one another. The products can have completely different properties than the starting materials.
|Binding type||Attachment partner|
|Ionic bond (synonym is electrovalent bond, heterovalent bond)||Non-metals with metals|
|Metal bond||Metals with metals|
|Atomic bond (synonyms are covalent bond, electron pair bond)||Non-metals with non-metals|
Octet rule and valence of the elements
Walther Kossel (1915) and Gilbert Newton Lewis (1916) developed the octet rule to explain the numerical proportions of the elements in chemical bonds. Accordingly, the elementary atoms strive to achieve the closest noble gas configuration in the periodic table through chemical bonding by releasing or accepting the corresponding number of electrons. The determining property is therefore the value of the elements.
- Fluorine absorbs an electron and receives it as F.− the configuration of the neon.
- Calcium gives off two electrons and receives as Ca2+ thus the configuration of the argon.
The designation Octet rule is derived from the eight valence electrons of the noble gases.
However, this rule only applies in the 1st and 2nd period of the main group elements without restriction. Other configurations can also be achieved for the main group elements of the other periods. The sulfur in sulfuric acid has 12 valence electrons (this assumption is only valid as a first approximation. According to modern calculations and in MO theory, the electron octet is not exceeded! Instead, the bond is partially ionic.). The subgroup elements sometimes achieve other, relatively stable configurations.
Metal character in the periodic table
Within a period, the metallic character of the atoms in the periodic table decreases from left to right and increases from top to bottom. Accordingly, there are smooth transitions between the three types of bond in which metals and non-metals are involved. The same principle applies in reverse for the ionic character. Here are some examples from the 3rd period:
Within the compounds of chlorine with the elements of the 3rd period of the periodic table, the ionic character of the bonds decreases more and more and the covalent character increases.
Within the compounds of sodium with the elements of the 3rd period, the metallic character of the bonds decreases more and more and the ionic character increases more and more.
Within the metal lattice or molecules of the elements of the 3rd period, the metal character decreases more and more and the covalent character increases.
According to the octet rule, a chemical bond is formally created when non-metal atoms take up electrons as binding partners and metals give up electrons. This is known as the donor-acceptor principle.
A Ionic bond is formed between the metal and non-metal atom in that the metal atom completely releases its valence electrons to the non-metal atom. This creates a cation from the metal atom and an anion from the non-metal atom. Due to the electrostatic attraction between these ions, an ion lattice is created. Lattices only form in the solid state. In the liquid state, the lattice collapses, the particles can be moved more easily against each other, but the binding character is retained. Due to the high electrical forces of attraction between their particles, ionic compounds have a salt-like character: They have high melting temperatures, conduct electrical current only in melt and solution (2nd order conductor, electrolytes) and are very brittle.
Example: Formal bond formation of potassium iodide
Since with one Metal bond all binding partners are metals, all atoms also give off valence electrons. The resulting metal cations are held together by the now freely moving electrons (the so-called electron gas), creating a metal lattice. In contrast to the ionic bond, the lattice generally does not form stoichiometrically with a metal bond between different elements. Metals are therefore all electrically conductive (first-order conductor), easily deformable (ductile), good heat conductors and have a metallic sheen (see figure copper).
Example: Formal bond formation of sodium as a metal lattice
Since with one Atomic bond (also Covalent bond, Covalent bond or Electron pair binding) all binding partners are non-metals, all atoms also accept valence electrons. This creates molecules or atomic lattices that are held together by bonding pairs of electrons. The electrons are located in so-called molecular orbitals (MO) between the two binding partners (binding electron pairs), and the atomic group formed in this way is called a molecule. That is why one speaks of a molecular or covalent bond (from Latin valens, of being of value - see valency (chemistry)), also called electron pair bond or atomic bond. This can be polar or non-polar, depending on whether the electrons are distributed symmetrically or asymmetrically in the molecule. Molecular substances usually have low boiling temperatures, are electrical non-conductors (insulators) and are either in volatile form (small molecules, example: water, oxygen, hydrogen chloride) or plastic and diamond-like (giant molecules, polymers, example: polyethylene, starch, boron nitride) ).
Example: formal bond formation with hydrogen chloride
Weak ties (Synonym: Van der Waals interaction in the broader sense) usually develop between molecules and influence the specific physical properties such as boiling point and fixed point. In macromolecules (e.g. polypeptides) they also appear as intra-molecular bonds. In the case of very weak bonds, instead of the term bond, the term interaction or Intermolecular Force used.
|Type of bond / interaction||Emergence||Explanation|
|Dipole-dipole interaction||with polar molecules||Unequal partial charges within a molecule create a dipole. The molecules align themselves in such a way that dipoles of different names can stick to one another. Example: hydrogen bond.|
|Dipole-ion interaction||when dissolving salts in water||The water dipole molecules form a hydration shell around the ions.|
|Van der Waals interaction in the narrower sense (London dispersion interaction)||with non-polar molecules||The disordered movement of the particles within the molecule leads to the random formation of dipoles that interact with each other.|
In the case of weak bonds, a complete electron transfer or the formation of bonding electron pairs cannot be formulated. Here only a shift of negative charge takes place within a molecule, creating electrical dipoles that can attract other dipoles or ions (see polar atomic bond).
In proteins, all types of weak interactions, as well as ion and atom bonds, can also occur within a single polypeptide molecule.
All chemical bonds and interactions can be traced back to electrostatic attraction between opposing charges.
In molecules and atomic lattices, the spatial alignment of the binding partners depends on the geometry of the atomic orbitals. (For more details see under atomic bond.)
In metal and ion grids, the spatial structure depends on the size of the binding partners, which are arranged on an imaginary spherical surface. (see also spherical packing, crystal structure).
Permanent or induced dipole molecules align with one another in such a way that their oppositely charged molecular parts point to one another and the parts with the same partial charge are as far apart as possible. (see VSEPR model)
This is the distance between the centers of the atoms or ions in chemical bonds.
With crystalline solids with Ion or metal grids the distances between the lattice modules can be determined by X-ray or electron diffraction. Since different distances between the lattice planes can be measured in crystal lattices, the smallest distance is usually given as the bond length in tables.
In calcium fluoride, the distance between the calcium cations is Ca2+ and the fluoride anions F− 235 pm (picometers). In metal lattices, the distance is between 200 pm and 500 pm, depending on the atomic size.
To Atomic bond look there
Hydrogen bonds have distances between 120 pm and 300 pm depending on the degree of polarization.
Bond strength and bond energy
A bond is stronger the more energy is released during its formation. The reverse is also true: the stronger a bond, the more energy has to be expended to break it and the less reactive the bond or element is.
The binding energy is at Ionic compounds the lattice enthalpy is given, that is the enthalpy that has to be expended in order to transfer a solid crystal into the gas phase, in which the ions can move freely.
The lattice enthalpy depends on the one hand on the size of the ions involved: the larger the ions, the smaller the lattice energy, since the forces of attraction decrease as the distance between the positive nuclei and the negative electron shell of the binding partner increases.
Examples: Lattice enthalpy of the fluorides of the alkali metals at 25 ° C in kJ per mol:
On the other hand, the lattice energy depends on the electrical charge of the ions involved: the larger the charges, the greater the forces of attraction and the greater the lattice energy.
Examples: Lattice enthalpy at 25 ° C in kJ / mol (in the examples the ion radius changes only slightly):
Aluminum oxide Al has the highest lattice enthalpy2O3 (Al3+ and O2−) with 15157 kJ / mol.
For the force of attraction K between two oppositely charged ions with the amount of charge e1 and e2 at a distance r, Coulomb's law applies:
As a measure of the strength of the bond in the Metal bond the melting temperature can be used: the higher the melting temperature, the stronger the binding forces. These again depend both on the distance between the metal cations and the number of electrons released: the more valence electrons are released and the smaller the lattice spacing, the greater the binding forces and thus the melting temperatures.
The enthalpy of binding of the Atomic bond is defined by the change in enthalpy when molecules dissociate into their atoms in the gas phase. Like the bond length (see above), it depends on both the size of the bonded atoms and the number of bonding electron pairs: the larger the radius of the bonding partners, the greater their distance and the smaller the bonding energy. In the case of bonds between atoms of the same type, it can be seen that their distance also depends on the number of bonding electron pairs:
|Surname||formula||binding||Bond length in pm||Binding enthalpy in kJ / mol|
|Ethene||C.2H4||C = C||134||614|
For delocalized atomic bonds it applies accordingly that they are lower in energy than a multiple bond, but more energetic than a single bond. The enthalpy of binding in benzene is 147 kJ / mol.
- The intermolecular interactions only have 10% of the bond strength of the strong bonds. Nevertheless, they have a strong influence on the fixed and boiling points of the fabrics:
In the Hydrogen bond the enthalpy of binding is at least 40 kJ / mol with strong polarization of the binding partners, and at most 20 kJ / mol with weak polarization. It is responsible for the fact that the boiling point of water is 100 ° C, while the boiling point of hydrogen sulfide is −83 ° C (see boiling point anomaly) Hydrogen bonds are usually intermolecular interactions. In ice there is a molecular lattice below 0 ° C, just as in sugar (sucrose) at room temperature (Granulated sugar).
Dipole-Dipole Interactions occur between polar molecules that do not meet the conditions for hydrogen bonding. Example: Ether: H3C-O-CH3. The strength of the dipole-dipole interaction results from the formula
- µ1,2: Dipole moments
Dipole-ion interactions occur, among other things, when salts are dissolved in water. The water dipoles surround the ions as a hydration shell and thus prevent cations and anions from reuniting to form a lattice structure. The strength of the ion-dipole interaction results from the formula
- e1: Ion charge
- µ1: Dipole moment
Van der Waals interactions arises between non-polar molecules that polarize each other when they approach, induced dipoles are created (in contrast to permanent dipole molecules such as water and hydrogen sulfide). Example: atom assemblies in liquid noble gases, molecular lattice of iodine at room temperature, π-complex in the bromination of ethene. The binding enthalpy of the Van der Waals interaction is of the order of 1 kJ / mol. Their amount depends on the dipole moment of the particles. Since we are dealing here with induced dipoles, the polarizability of the initially non-polar atoms also plays a role: large “soft” atoms are easier to polarize than small “hard” ones. (The definitions “hard” and “soft” are taken from the HSAB concept.) This can be seen from the boiling points of the noble gases, which, with increasing size, develop increasingly stronger van der Waals interactions and thus increasingly require more energy to overcome these attractive forces and to pass into the gas phase.
However, the boiling points of non-polar molecules also depend on the surface with which they can exert van der Waals interactions with neighboring molecules. The boiling point of the linear, unbranched n-pentane is 36.1 ° C, while the isomeric 2,2-dimethyl-propane with the same molar mass has a boiling point of 9.5 ° C because it is almost spherical and therefore a has a smaller “contact area” with neighboring molecules.
Chemisorption and physisorption
In the case of low binding energies, which mainly come about through electrical attractive forces, one speaks of physisorption. The physisorption includes, for. B. the Van der Waals bond or the hydrogen bond.
In the case of higher binding energies, one speaks of chemisorption, in which the participating electron orbitals overlap and thus lead to a bond. Chemisorption includes the covalent atomic bond and the complex bond.
One speaks of physisorption for binding energies in the meV range, of chemisorption in the eV range and greater. A precise boundary between the two is often not possible.
Chemical bond theory
- J. Reinhold, Quantum theory of molecules, 3rd ed., Teubner, 2006, ISBN 3835100378.
- L. Pauling, The nature of the chemical bond, 2nd post d. 3rd ed., Wiley-VCH, 1988, ISBN 9783527252176.
Category: Chemical Bond
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