Give one example of electron deficient covalent hydride

\( \newcommand\) • • • • • • The combination of hydrogen with another element produces a hydride, E xH y. The formal charge or oxidation state of the hydrogen in these compounds is dependant on the relative electronegativity of the element in question. Ionic hydrides Hydrogen compounds with highly electropositive metals, i.e., those in which the metal has an electronegativity of less than 1.2, are ionic with the hydrogen having a s 2 configuration (H -). Typical ionic metal hydrides are those of the Group 1 (IA) metals and the heavier Group 2 (IIA) metals. The ionic radius of the hydride ion is in between that of fluoride and chloride and the same as oxide (Table \(\PageIndex\] WARNING Group 1 and 2 metal hydrides can ignite in air, especially upon contact with water to release hydrogen, which is also flammable. Hydrolysis converts the hydride into the analogous hydroxide, which are caustic bases. In practice, most ionic hydrides are dispensed as a dispersion in oil, which can be safely handled in air. Covalent Hydrides The most common binary compounds of hydrogen are those in which hydrogen bonds have covalent bond character. The E-H bond is usually polar ranging from those in which the hydrogen is polarized positively (e.g., those with non-metals such as F, O, S, and C) to where it is negative (e.g., those with metals and metalloids such as B, Al, etc). Magnesium hydride is intermediate between covalent and ionic since it has a polymeric solid similar to AlH 3, but react...

Borane and diborane Borane (BH 3) formed in the gaseous state from decomposition of other compounds, (6.4.2) but cannot be isolated except as a Lewis acid-base complex (6.4.1). As such many borane adducts are known. \[ \text_3 \] Borohydride The borohydride anion (or more properly the tetrahydridoborate anion), BH 4 -, can be considered as the Lewis acid-base complex between borane and H -. A typical synthesis involves the reaction of a borate ester with a hydride source, (6.4.9). \[\text\] Sodium borohydride is a stable white crystalline solid that is stable in dry air and is non-volatile. The boron in borohydride (BH 4 -) is tetrahedral. Although it is insoluble in Et 2O, it is soluble in water (in which it reacts slowly), THF, ethylene glycol, and pyridine. Interestingly, NaBH 4 reacts rapidly with MeOH, but dissolves in EtOH. Sodium borohydride has extensive uses in organic chemistry as a useful reducing agent in which it donates a hydride (H -). Higher Boranes Higher boron hydrides contain, in addition to the bridging B-H-B unit, one or more B-B bonds. The higher boranes are usually formed by the thermal decomposition of diborane, (6.4.10) and (6.4.11). \[ \text\). Tetraborane, or to be more precise tetraborane(10) or arachno-B 4H 10, is a foul-smelling toxic gas. Pentaborane (9) is a toxic liquid (with a distinctive garlic odor) that can detonate in air, and like decaborane(14) was at one time considered as a potential rocket fuel. Because simple boron compounds burn...

Electron deficient compounds: • Compounds wherein the central atom lacks eight electrons in its outer shell or possesses eight electrons but can widen the valency owing to the existence of unoccupied d – orbitals. • In brief, electron-deficient compounds lack enough electrons to fulfill the octet of the central atom. Examples: • Electron-deficient compounds are those with less than 8 electrons in their valence shells, such as B 2 F 6 , Al 2 Cl 6 , and others. Boron family example of Electron deficient compounds : • Boron, for example, with the electronic structure [ He ] 2 s 2 2 p 1 , creates compounds known as electron-deficient compounds. • Boron has just three valence electrons, however, four orbitals are available to house these electrons. • Other boron family members, such as Aluminum, Gallium, Indium, and Thallium, have a proclivity to lose valence electrons and generate ionic M 3 + entities in their compounds.

In Another type of exception to the Lewis approach to bonding is the existence of compounds that possess too few electrons for a Lewis structure to be written. Such compounds are called electron-deficient compounds. A prime example of an electron-deficient compound is diborane, B 2H 6.… molecular orbital theory • In …outstanding problem is that of electron-deficient compounds, as typified by B 2H 6. Such molecules are classified as electron deficient because, in Lewis terms, there are fewer than two electrons available per bond. However, a consequence of delocalization is that the bonding influence of an electron pair is distributed over all the…

\( \newcommand\) • • • • • • • • • • • • • • • • • • Introduction to Lewis structures A Lewis structure is a way to show how atoms share electrons when they form a molecule. Lewis structures show all of the valence electrons in an atom or molecule. The valence electrons are the electrons in the outermost shell.For representative elements, the number of valence electrons equals the group number on the periodic table.To draw the Lewis structure of an atom, write the symbol of the atom and draw dots around it to represent the valence electrons. Note that hydrogen is often shown in both group 1A and group 7A, but it has one valence electron – never seven. Also, helium is shown in group 8A, but it only has two valence electrons. Representing a Covalent Bond Using Lewis Structures Nonmetals can form a chemical bond by sharing two electrons. Each atom contributes one electron to the bond. For example, two hydrogen atoms can form a bond, producing a molecule of H 2. Using Lewis structures, we can represent this as follows: Two fluorine atoms can form a molecule of F 2 in the same fashion. Note that each atom must contribute one electron to the bond. Atoms can form more than one bond. In a water molecule, an oxygen atom forms two bonds, one to each hydrogen atom. Chemists normally represent a bond using a line instead of two dots. The structures of H 2, F 2, and H 2O would usually be drawn as follows: Only the bonding electrons are shown using lines.Nonbonding electrons are always ...

Covalent hydrides: Covalent hydrides are compounds containing either Hydrogen or non-metal compounds that share electron pairs with another element through equivalent electronegativity between them. Types of covalent hydrides: • Electron precise: There are exact number of electrons to form covalent bond. For example, Methane ( CH 4) • Electron deficient: There are less number of electrons to form covalent bond. For example, Diborane ( B 2 H 6) • Electron rich: There are one or more lone pair of electrons around the central atom. For example, water ( H 2 O) Note: Due to their low boiling and melting points, covalent hydrides are typically gases or volatile liquids.

\( \newcommand\) • • • • Learning Objectives • Identify which elements are unable to achieve octets whenforming covalent molecules. • Draw Lewis structures of covalent molecules that contain electron-deficient atoms. • Write the chemical formulas and chemical names of covalent molecules that contain electron-deficient atoms. As discussed previously, electrons are most stable when they exist in octet configurations. However, some elements are unable to fill their valence shellswhen bonding, due to inherent deficiencies in their atomic electron configurations. Other elements are able to achieve octets when bonding, but are also able to expand their valences to contain more than eight electrons. Elements that exhibit each of these exceptionsto the "octet rule" can still be used to formstable covalent molecules, as will bediscussed in greater detail in the current and following sections of this chapter, respectively. Elements that Cannot Achieve Octets When Bonding Recall that covalent bonds are produced when unpaired electrons found within two atoms, which must be classified as either non-metals or metalloids,interact to form a shared pair of electrons. After successfully pairing all of their unpaired electrons, most of the elements that are able to bond covalentlyachieve octet configurations through a combination of bonded, or shared,and lone, or unshared, pairs of electrons. For example, carbon and silicon, which are both found in Group 4A, each contain four valence electro...