How does carbon affect the bond angle
1. Structure and binding of organic molecules
1.1 atomic structure
The Atomic number (or atomic number / atomic number) is the number of protons in the atomic nucleus. The Mass number (or mass number) is the sum of protons and neutrons combined.
All atoms in an element have the same atomic number, but may have different mass numbers, e.g.
|H atomic number =||1||Isotopes =||1H||2H||3H|
|99,98 * ||0.015 ||10-16||C atomic number = ||6||Isotopes = ||12C.||13C.||14C. ||1.1 ||10-10||N atomic number = ||7||Isotopes = ||14N||15N||0.37|
* Natural frequency in atomic%
The average mass number is called Atomic weight (Atomic weight) and is usually not an integer, e.g.
|H atomic weight =||1.008|
|C atomic weight =||12.01|
|N atomic weight =||14.01|
The electrons are particularly important for chemistry. Where are the electrons located? From quantum chemistry we know that the square of the wave function (& # 9682) for each point of the space around the nucleus describes the probability of finding an electron at a certain point (The Schrödinger equation combines the function & # 968, which Wave function of the electron, with its energy and the space coordinates, which are necessary to describe the system).
So - in a somewhat casual expression - the electrons are in certain orbitals.
Atomic orbitals have a characteristic shape.
e.g. for carbon atoms are only s - and p- Orbitals important:
The energy ratios for the different s, p, and d Orbitals can be described in an energy level scheme:
|Quantum numbers characterize the main energy and distance of the electron from the atomic nucleus (n), "orbital shape" (l), orientation (m) and spin (s):|
n = principal quantum number
l = minor or orbital quantum number (l = 0.1, ... n-1)
m = mag. Quantum number (m = -l..0 .. + l)
s = spin quantum number (s = ± ½)
1.2 Electron configuration
When "filling up" the possible energy levels, the following must be observed: 1) The orbitals with the lowest energy are filled first: 2) Each AO can be filled with a maximum of two electrons, which must have opposite spin (Pauli principle); 3) If two or more orbitals with the same energy are unoccupied, all orbitals are first occupied with an electron, with all electron spins aligned in parallel until all orbitals are half occupied (Hund's rule).
The electron configurations for various elements can now be described. E.g .:
It is easy to see why electron octets and duets represent particularly stable configurations (The noble gas rule). These numbers of electrons result in closed configurations with completely occupied orbitals (cf. He, Ne, Ar, Kr, Xe, Rn).
1.3 Ionic and covalent bonds
Why do atoms form bonds with one another? Roughly speaking, because the resulting product is more stable (has less energy) than the separated atoms.
A simple one Ionic bond arises between one electronegative Atom (electronegative atom means an atom that has a tendency to attract electrons) and one electropositive Atom (i.e. an atom that has a tendency to transfer electrons).
A Ionic bond comes about through electrostatic forces of attraction between the charged ions.
Elements on the far right and on the far left side of the periodic table thus form ionic bonds through the loss or gain of electrons.
The carbon atom (1s2 2s2 2p2) but cannot 4e- win or lose to achieve a noble gas configuration. Therefore, carbon atoms form bonds with other atoms by the atoms that are bonded to each other sharing electrons with each other. Each atom thus formally receives an external electron octet.
Such bonds are known as covalent single bonds. In the Lewis drawings, the valence electrons are represented by dots. Simply put, atoms that are bound together are symbolized by straight lines between the atoms. Lone electron pairs (lone pairs of electrons) are represented as colons or simply omitted (Kekulé structures).
The number of bonds that an atom can form is determined by the number of its external electrons in conjunction with the noble gas rule
1.4 The covalent bond
How do covalent bonds arise? Bonds are created by the in-phase overlap of atomic orbitals.
The in-phase overlap of the two 1s orbitals creates a new orbital with lower energy than the two original atomic orbitals - that binding molecular orbital.
If you bring two hydrogen atoms so close together that a bond is made between them, an energy of 435 KJ / mol is released. Accordingly, energy is required to break such a bond, also 435 KJ / mol (104 Kcal / mol). This energy is known as Bond dissociation energy.
Dissociation energies relate exclusively to homolytic splits. Depending on the type of bond and the atoms connected to one another, they have a characteristic value.
What do the covalent bonds look like in more complex molecules, e.g. in methane?
1.5 Hybridization of orbitals
The bond formation in the H2 Molecules are easy to describe, the situation becomes more difficult in complex organic molecules with tetravalent carbon. This is to be demonstrated using the example of methane.
As an atom of the 4th main group of the PSE, carbon has four electrons in its valence shell (2s2, 2p2) and can thus enter into four covalent bonds, e.g. with hydrogen.
In this electronic configuration, C is able to enter into four chemical bonds. In Lewis structures:
How can one describe the C-H bonds? Since carbon uses electrons from two types of orbitals (2s and 2p) to form bonds, one might assume that methane has two types of C-H bonds. In fact, all four C-H bonds are identical and point to the corners of a regular tetrahedron. How can one explain this fact?
The answer was given by Linus Pauling in 1931 by showing that one s orbital and three p orbitals combined (hybridized) can be used to obtain four equivalent atomic orbitals with tetrahedral orientation. These tetrahedral orbitals are called sp3 Hybrid orbitals, whereby the exponent 3 indicates how many atomic orbitals of the respective type are involved in the formation of the hybrid orbital (formally s1p3 -> sp3).
The concept of hybridization now describes how Carbon forms four equivalent tetrahedral bonds, but does not explain Why. The form of the hybrid orbital provides the answer. If one s orbital hybridizes with three p orbitals, the resulting hybrid orbitals are asymmetrical with respect to the atomic nucleus. One of the two lobes becomes larger than the other and can therefore better overlap with another orbital to form a bond. This makes it clear that sp3 Hybrid orbitals can form stronger bonds than unhybridized s or p orbitals.
If the four identical orbitals of a sp3 hybridized carbon atom with the 1s orbitals of four hydrogen atoms, then four identical C-H bonds are formed and a methane molecule is present.
The C-H bond in methane has a bond dissociation energy of 105 kcal / mol, a bond length of 1.1 Å. The H-C-H bond angle, also called the tetrahedral angle, is 109.5O.
1.6 The structure of ethane
The SP3-Hybrid orbitals can also overlap with other orbitals. Overlapping with four chlorine p orbitals creates carbon tetrachloride CCl, for example4. C-C bonds are also formed by the overlap of two hybrid orbitals:
In ethane, the C-C bond length is 1.54 Å, the C-C bond dissociation energy is 88 Kcal / mol and the bond angles are almost all 109.5O.
1.7 The structures of ammonia and water
Which types of orbitals can be used to describe the bond in ammonia and in the water molecule?
From the electron configuration 1s2 2s2 2p3 of nitrogen, there is its trivalent nature. Three covalent bonds are necessary to achieve an electron octet. To agree with the observed molecular structure (an approximate tetrahedron), one takes a sp3Hybridization on nitrogen.
In the fourth sp3-Orbital is the lone pair of electrons. The N-H bond dissociation energy is 391 KJ / mol (93 Kcal / mol), the bond length 1.0 Å, and the HNH bond angle 107.3O.
Correspondingly, the bond in the water molecule can best be determined via sp3- Describe hybrid orbitals.
O-H bond dissociation energy = 463 KJ / mol (110 Kcal / mol)
O-H bond length = 0.96 Å
HOH bond angle = 104.5O.
1.8 sp2-Hybrids for the representation of trigonal structures
In ethylene (= ethene) the two carbon atoms share four Electrons. They are linked by a double bond:
A sp3-Hybridization cannot explain this structure. Instead, the 2s orbital at C is hybridized with only two 2p orbitals, by three sp2 Hybrid orbitals to build. A p orbital remains unchanged:
If two sp2-hybridized carbon atoms can approach each other by overlapping two sp2-Orbitals create a & # 963 bond. At the same time, the two non-hybridized p orbitals approach from the side and thereby form a new one pi (& # 960) binding.
The overlapping of p orbitals creates pi (& # 960) bonds. The overlap of the four remaining sp2-Hybrid orbitals with 1s orbitals of the hydrogen then result in the ethylene molecule:
Ethylene has a flat structure with HCH and HCC bond angles of 120O. The C-H bond length is 1.076 Å and the C-H bond dissociation energy is 103 Kcal / mol. The C = C bond length is 1.33 Å (see ethane 1.54 Å) and the bond dissociation energy is 152 Kcal / mol (see ethane 88 Kcal / mol). As expected, the double bond is thus shorter and stronger than a C-C single bond.
The carbonyl group in aldehydes, ketones and carboxylic acids contains a C = O double bond. The carbonyl-C is also sp2-hybridizes and forms three bonds. The fourth valence electron remains in a p orbital and is involved in a pi bond to oxygen:
Like alkenes, carbonyl groups have a flat structure and bond angles of about 120O.
1.9 sp hybrids for the representation of acetylene
Carbon can also form triple bonds. In acetylene, two carbon atoms share six electrons. Here other hybrid orbitals arise through the mixture of the 2s orbital and a 2p orbital. Two sp hydride orbitals arise and form a linear geometry. The other p orbitals are perpendicular to each other:
When two sp-hybridized carbon atoms approach, the overlap of the sp-hybrid orbitals creates a strong-bond and the overlap of the two p-orbitals creates two-bonds (i.e. a triple bond). The remaining sp hybrid orbitals can again form bonds with the 1s orbitals of hydrogen.
Due to the sp hybridization, acetylene is a linear molecule with an HCC bond angle of 180O, a C-H bond length of 1.06 Å, a C-C bond length of 1.2 Å and bond dissociation energies of 125 Kcal / mol (C-H) and 200 Kcal / mol (C-C).
1.10 polarity of covalent bonds
If two identical atoms are linked by a bond (H-H, CH3-CH3 etc.) so the pair of binding electrons is distributed evenly to both atoms. If, on the other hand, the linked atoms are different (e.g. H-Cl, CH3-Cl), an unsymmetrical electron cloud is created because one of the atoms attracts the bonding electrons more strongly than the other.
The effort of an atom to attract binding electrons is called Electronegativity. In the periodic table, electronegativity increases from "left to right" and from "bottom to top":
The Pauling electronegativities relate to the most electronegative atom of fluorine, to which the value 4 is arbitrarily assigned. E.g. the C-Cl bond in chloromethane is one polar covalent bond:
Metallic elements on the left side of the periodic table are less electronegative than C:
Polarization effects that are caused by atoms or groups of atoms that attract or repel electrons and are transmitted via über bonds are called inductive effects. Depending on whether the "key atom" i.e. the electron-attracting or -rejecting atom receives a negative or positive partial charge, one speaks of -I or + I effects. The "key atoms" are also often called & # 963 acceptors (electron attracting) or & # 963 donors (electron-repellent) denotes e.g.
With an increasing number of bonds, i.e. with increasing distance from the key atom, the effect of the inductive effect decreases very strongly.
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