Bond angle in the context of Water (molecule)

In a tetrahedral molecular geometry, a central atom is located at the center with four substituents that are located at the corners of a tetrahedron. The bond angles are arccos(−⁠1/3⁠) = 109.4712206...° ≈ 109.5° when all four substituents are the same, as in methane (CH₄) as well as its heavier analogues. Methane and other perfectly symmetrical tetrahedral molecules belong to point group T_d, but most tetrahedral molecules have lower symmetry. Tetrahedral molecules can be chiral.

In chemistry, the ball-and-stick model is a molecular model of a chemical substance which displays both the three-dimensional position of the atoms and the bonds between them. The atoms are typically represented by spheres, connected by rods which represent the bonds. Double and triple bonds are usually represented by two or three curved rods, respectively, or alternately by correctly positioned sticks for the sigma and pi bonds. In a good model, the angles between the rods should be the same as the angles between the bonds, and the distances between the centers of the spheres should be proportional to the distances between the corresponding atomic nuclei. The chemical element of each atom is often indicated by the sphere's color.

In a ball-and-stick model, the radius of the spheres is usually much smaller than the rod lengths, in order to provide a clearer view of the atoms and bonds throughout the model. As a consequence, the model does not provide a clear insight about the space occupied by the model. In this aspect, ball-and-stick models are distinct from space-filling (calotte) models, where the sphere radii are proportional to the Van der Waals atomic radii in the same scale as the atom distances, and therefore show the occupied space but not the bonds.

In chemistry, transferability is the assumption that a chemical property that is associated with an atom or a functional group in a molecule will have a similar (but not identical) value in a variety of different circumstances. Examples of transferable properties include:

• Electronegativity
• Nucleophilicity
• Chemical shifts in NMR spectroscopy
• Characteristic frequencies in Infrared spectroscopy
• Bond length and bond angle
• Bond energy

Transferable properties are distinguished from conserved properties, which are assumed to always have the same value whatever the chemical situation, e.g. standard atomic weight.

The linear molecular geometry describes the geometry around a central atom bonded to two other atoms (or ligands) placed at a bond angle of 180°. Linear organic molecules, such as acetylene (HC≡CH), are often described by invoking sp orbital hybridization for their carbon centers.

According to the VSEPR model (Valence Shell Electron Pair Repulsion model), linear geometry occurs at central atoms with two bonded atoms and zero or three lone pairs (AX₂ or AX₂E₃) in the AXE notation. Neutral AX₂ molecules with linear geometry include beryllium fluoride (F−Be−F) with two single bonds, carbon dioxide (O=C=O) with two double bonds, hydrogen cyanide (H−C≡N) with one single and one triple bond. The most important linear molecule with more than three atoms is acetylene (H−C≡C−H), in which each of its carbon atoms is considered to be a central atom with a single bond to one hydrogen and a triple bond to the other carbon atom. Linear anions include azide (N=N=N) and thiocyanate (S=C=N), and a linear cation is the nitronium ion (O=N=O).

Dibromine trioxide is the chemical compound composed of bromine and oxygen with the formula Br₂O₃. It is an orange solid that is stable below −40 °C. It has the structure Br−O−BrO₂ (bromine bromate). It was discovered in 1993. The bond angle of Br−O−Br is 111.7°, the bond angle of O−Br=O is 103.1°, and the bond angle of O=Br=O is 107.6°. The Br−OBrO₂ bond length is 1.845 Å, the O−BrO₂ bond length is 1.855 Å and the Br=O bond length is 1.612 Å.

In organic chemistry, ring strain is a type of instability that exists when bonds in a molecule form angles that deviate from their normal values as a result of being part of a ring. This type of strain is most commonly discussed for small rings such as cyclopropanes and cyclobutanes, whose internal angles are substantially smaller than the idealized value of approximately 109°. Because of their high strain, the heat of combustion for these small rings is elevated.

Ring strain results from a combination of angle strain, conformational strain or Pitzer strain (torsional eclipsing interactions), and transannular strain, also known as van der Waals strain or Prelog strain. The simplest examples of angle strain are small cycloalkanes such as cyclopropane and cyclobutane.

In organic chemistry, a ring flip (also known as a ring inversion or ring reversal) is the interconversion of cyclic conformers that have equivalent ring shapes (e.g., from a chair conformer to another chair conformer) that results in the exchange of nonequivalent substituent positions. The overall process generally takes place over several steps, involving coupled rotations about several of the molecule's single bonds, in conjunction with minor deformations of bond angles. Most commonly, the term is used to refer to the interconversion of the two chair conformers of cyclohexane derivatives, which is specifically referred to as a chair flip, although other cycloalkanes and inorganic rings undergo similar processes.

In chemical bonds, an orbital overlap is the concentration of orbitals on adjacent atoms in the same regions of space. Orbital overlap can lead to bond formation. The general principle for orbital overlap is that, the greater the overlap between orbitals, the greater the bond strength. Linus Pauling explained the importance of orbital overlap in the molecular bond angles observed through experimentation; it is the basis for orbital hybridization. As s orbitals are spherical (and have no directionality) and p orbitals are oriented 90° to each other, a theory was needed to explain why molecules such as methane (CH₄) had observed bond angles of 109.5°. Pauling proposed that s and p orbitals on the carbon atom can combine to form hybrid orbitals (sp in the case of methane) which are directed toward the hydrogen atoms. The carbon hybrid orbitals have greater overlap with the hydrogen orbitals, and can therefore form stronger C–H bonds.

A quantitative measure of the overlap of two atomic orbitals Ψ_A and Ψ_B on atoms A and B is their overlap integral, defined as