Covalent Bond
A covalent bond is a type of chemical bond formed when two atoms share one or more pairs of electrons. This sharing of electrons allows both atoms to achieve a more stable electron configuration. Covalent bonds are typically found in molecules and are characterized by the strong attraction between the shared electrons and the positively charged nuclei of the bonded atoms.
8 Key excerpts on "Covalent Bond"
eBook - ePub- J Fisher, J.R.P. Arnold, Julie Fisher, John Arnold(Authors)
- 2020(Publication Date)
- Taylor & Francis(Publisher)
...Thus, this is more of an interaction than a bond as there is no extra build up of electron density at the mid-point between the two atom centers. Covalent Bonds A Covalent Bond is formed between elements that in general are not readily ionizable; they are neither strongly electropositive nor electronegative. Atoms involved in Covalent Bonds share the bonding electrons and thus attain an apparent complete valence shell. In Covalent Bonds electron density is distributed between the atom centers. Lewis structures The Lewis structure of a molecule depicts the electrons in the valence shell of atoms involved in bonding. Each electron is depicted as a dot and this is a straightforward way of keeping check of electron balance in a molecule or reaction. Molecular orbital theory The molecular orbital theory is an approach to describe Covalent Bonding in molecules, by considering the combination of atomic orbitals. Bonding, antibonding, and nonbonding molecular orbitals are introduced and, following on from this, the shapes of simple molecules may readily be described. Coordination compounds Many elements are known to have fixed valences; however, a number of elements can exhibit two or three stable valences. These can readily be explained by consideration of the relevant atomic and molecular orbitals. A number of metals have combining powers that cannot readily be reconciled in terms of covalent or ionic bonding. Such metal-containing compounds are now referred to as coordination complexes, as groups or ligands around the central metal donate a lone pair of electrons to unoccupied orbitals of the metal. Related topics (A1) The periodic table (G1) The early transition metals The octet rule Atoms whose outer (valence) electron shells hold a complete set of electrons are substantially more stable than those that do not. Consequently, helium and neon are extremely stable, indeed they are inert (Section A1)...
eBook - ePub- Atef Korchef(Author)
- 2022(Publication Date)
- CRC Press(Publisher)
...There are three types of chemical bonds: A Covalent Bond results from sharing electrons between the atoms. The Covalent Bond is usually found between non-metals. The Covalent Bond can be polar or non-polar. Examples of molecules showing polar Covalent Bonds: HCl, HBr and H 2 O Examples of molecules showing non-polar Covalent Bonds: Homonuclear diatomic molecules such as I 2, Br 2, Cl 2, O 2 and F 2, and hydrocarbons (C n H 2n + 2). A non-polar Covalent Bond in I 2 and a polar Covalent Bond in HCl. An ionic bond results from the transfer of electrons from a metal to a non-metal. Example: The bond between sodium (Na) and chlorine (Cl) in NaCl is an ionic bond. The NaCl compound obtained is an ionic compound. A metallic bond is formed between positively charged atoms in which the free electrons are shared among a structure of positively charged ions. Metallic bonding is the main type of chemical bond that forms between metal atoms. Lewis dot representation or electron dot diagram is a simplistic way of showing the valence electrons of an atom that uses dots around the chemical symbol of the element with no more than two dots on a side. To represent the Lewis dot diagram of an atom, carry out the following steps: Write out the electronic configuration of the atom or the ion. Find the number of valence electrons. Draw the valence electrons as dots around the chemical symbol of the atom or the ion. Example: The seven valence electrons of Cl are drawn in the following configuration: Elements in the same group of the periodic table of elements have similar Lewis electron dot diagrams because they have the same valence shell electron configuration. Example: 7 N and 15 P belong to the same group (Group 15) and have similar Lewis dot diagrams: Electron dot diagrams for ions are like those of atoms, except that some electrons have been removed from or added to a neutral atom to obtain an ion...
eBook - ePub- Jeffrey Gaffney, Nancy Marley(Authors)
- 2017(Publication Date)
- Elsevier(Publisher)
...But, elements with similar electronegativities (those closer together in the periodic table) cannot completely transfer electrons from one atom to another. They can still achieve a closed valence shell by sharing their electrons. For example, two hydrogen atoms have the same electronegativity and cannot transfer electrons from one hydrogen atom to the other to form an ionic bond. But they can both achieve the closed valence shell of a helium atom (1 s 2) if both atoms share their single valence electron to form a shared electron pair. This can be represented by an electron dot structure as; The pair of electrons that is shared between the two hydrogen atoms forms a Covalent Bond between them resulting in a diatomic molecule of hydrogen with the formula H 2. Similarly, two fluorine atoms ([He]2 s 2 2 p 5) can both achieve the closed valence shell of neon ([He]2 s 2 2 p 6) by sharing one valence electron each. The shared electron pair, which forms the Covalent Bond between the two fluorine atoms, is known as bonding pair electrons. The three pairs of valence electrons around each fluorine atom that are not shared are known as lone pair electrons. Covalent Bonds are formed when two or more valence electrons are attracted by the positively charged nuclei of both atoms and so are shared between the two atoms. This description of Covalent Bonds as involving shared pairs of valence electrons that are localized in a bond between two atoms is known as valence bond theory. Fig. 3.5 shows the potential energy between two atoms of hydrogen as they approach each other. In a similar manner as in the formation of an ionic bond (Fig. 3.4), the potential energy decreases as the hydrogen atoms get closer together and the attraction between the valence electrons and the nuclei of each atom becomes stronger. The Covalent Bond is formed when the two hydrogen atoms reach the lowest potential energy, which is the Covalent Bond length...
- Kevin Reel, Derrick C. Wood, Scott A. Best(Authors)
- 2014(Publication Date)
- Research & Education Association(Publisher)
...Chapter 6 Chemical Bonding Intramolecular Forces: Bonds Between Atoms Bonds are the forces of attraction that hold atoms together. There are many types of bonding including ionic, metallic, and Covalent Bonds. You can figure out the difference between the bonding types if you look at what role the valence electrons are playing in the chemical bond—because bonding is all about the valence electrons. Many of the electrons in an atom have no impact on bonding because they are located close to the nucleus, and thus are called core electrons. In general, the valence electrons are the outermost s-shell and p-shell electrons in an electron configuration. For transition metals, the outermost d-shell electrons will also play a role. Elements will typically form bonds in order to have eight electrons in the valence shell, which is called the octet rule. The vast majority of chemical bonds that occur obey the octet rule, although a significant number of exceptions to the octet rule exist; these exceptions will be covered later in this chapter. Ionic Bonds • Ionic bonding is a bond between a cation and an anion held together by electrostatic attractions. Coulomb’s law dictates that oppositely charged particles are attracted to one another, and this is the fundamental principle behind ionic bonding. • In an ionic bond, an electron is removed from the least electronegative atom to form a positively charged ion (cation). This electron is then transferred to a more electronegative atom to form a negatively charged ion (anion). • Ionic bonds form in order to fulfill the octet rule for the elements involved with the bond. The metal loses electrons to have a filled shell. The nonmetal gains electrons to have a filled shell. • Ionic bonds usually occur between a metal and a nonmetal that have a difference in electronegativity greater than or equal to 1.7...
eBook - ePubChemistry
Concepts and Problems, A Self-Teaching Guide
- Richard Post, Chad Snyder, Clifford C. Houk(Authors)
- 2020(Publication Date)
- Jossey-Bass(Publisher)
...A special type of Covalent Bond is known as the coordinate Covalent Bond. In a coordinate Covalent Bond, both electrons in a shared pair of electrons come from the same atom. The bonding between the nitrogen atom in ammonia (NH 3) and the boron atom in boron trifluoride (BF3) is an example of a coordinate Covalent Bond. In the above equation, which atom supplied both electrons to form the coordinate Covalent Bond, H, N, B, or F? ___________ Answer: The N atom supplied both electrons to form a coordinate Covalent Bond. (Notice that in NH 3 on the left side of the equation, the N is already shown with eight electrons.) A coordinate Covalent Bond is like any other Covalent Bond except for the origin of the shared electron pair. Does the molecule below conform to the octet rule (with the exception of the hydrogen atoms)? ___________ Answer: Yes, all the atoms in the molecule (except the hydrogen atoms) have eight outer shell electrons. Review your understanding of ionic, covalent, and coordinate Covalent Bonding. A bond in which one or more electrons are removed from one atom and taken by another is called a(n) ____________ bond. A bond in which two atoms share a pair of electrons with one electron coming from each atom is called a(n) ____________ bond. A bond in which two atoms share a pair of electrons with both electrons coming from one atom is called a(n)____________ bond. Answer: (a) ionic; (b) covalent; (c) coordinate covalent Up to this point we have assumed that the electrons in a Covalent Bond are shared equally between the bonding atoms, that is, the electrons are midway between the bonding atoms. In the frames that follow we discuss what happens if the electrons are not equally shared, that is, if the electron pair is closer to one atom than the other. POLAR BONDS When all the atoms in a covalent molecule are identical, the attraction each atom has for shared electrons is equal...
eBook - ePub- Gavin Whittaker, Andy Mount, Matthew Heal(Authors)
- 2000(Publication Date)
- Taylor & Francis(Publisher)
...It also describes the geometric arrangement of the bonds, and so the shapes of molecules. The more sophisticated valence theories yield information about the electrical, magnetic, and spectroscopic properties of molecules. Elementary valence theories invoke two principal bond types. In ionic bonding, electrostatic interactions generate bonds between ions formed by electron transfer from one element to the other. In Covalent Bonding two elements are held together by shared electrons in order that both may adopt an energetically favorable electron configuration. In reality, both are extreme forms of the same bonding phenomenon. Pure Covalent Bonds are formed by elements with identical electronegativities, with more ionic bonding character being introduced to the bond as the electronegativity difference between the elements increases (see Topic H4). Even in extreme cases of ionic bonding, the degree of covalent character may still be quite high. Two complementary theories were originally developed to explain the number and nature of Covalent Bonds (Lewis theory) and the shapes of molecules (VSEPR theory). More sophisticated theories have superseded these approaches for detailed investigations, but they remain useful in semi-empirical and non-rigorous discussions of molecular bonding. Lewis theory The Lewis theory of Covalent Bonding may be regarded as an elementary form of valence bond theory. It is nonetheless useful for describing covalent molecules with simple Covalent Bonds, and works successfully in describing the majority of, for example, organic compounds. Lewis theory recognizes both the free energy gains made in the formation of complete atomic electron shells, and the ability of atoms to achieve this state by sharing electrons...
eBook - ePub- Tony Cox(Author)
- 2004(Publication Date)
- Taylor & Francis(Publisher)
...C1 E LECTRON PAIR BONDS Key Notes Lewis and valence structures A Lewis structure shows the valence electrons in a molecule. Two shared electrons form a single bond, with correspondingly more for multiple bonds. Some atoms may also have nonbonding electrons (lone-pairs). Valence structures show the bonds simply as lines. Octets and ‘hypervalence’ In most stable molecules and ions of the elements C–F, each of these atoms has eight electrons (an octet) in its valence shell. Expansion of the octet and increased valency is possible with elements in periods 3 and below. Resonance When several alternative valence structures are possible, the bonding may be described in terms of resonance between them. Formal charges Formal charges are assigned by apportioning bonding electrons equally between the two atoms involved. They can be useful to rationalize apparent anomalies in bonding, and to assess the likely stability of a proposed valence structure. Limitations Many covalent molecules and ions cannot be understood in terms of electron pair bonds between two atoms. They include electron-deficient boron hydrides and transition metal compounds. Related topics Electronegativity and bond type (B1) Molecular shapes: VSEPR (C2) Introduction to nonmetals (F1) Lewis and valence structures A single Covalent Bond is formed when two atoms share a pair of electrons. Double and triple bonds can be formed when two or three such pairs are shared. A Lewis structure is a representation of a molecule or complex ion that shows the disposition of valence electrons (inner shells are not drawn) around each atom. 1–4 show Lewis structures of CH 4, H 2 O, O 2 and N 2, the last two molecules having a double and triple bond, respectively. These representations are entirely equivalent to the valence structures (1’–4’) in which each bonding pair of electrons is represented by a line. A molecule such as H 2 O has nonbonding or lone-pair electrons localized on one atom rather than shared...
eBook - ePub- Dov M. Gabbay, Paul Thagard, John Woods(Authors)
- 2011(Publication Date)
- North Holland(Publisher)
...In a non-molecular polar substance like potassium chloride, the bonding is electrostatic and therefore radially symmetrical. Hence an individual ion bears no special relationship to any one of its neighbours. Polar bonding is non-directional, and so cannot be represented by the lines connecting atoms in classical structural formulae. The distinction between polar and non-polar bonding is still present in his next paper [ Lewis, 1916 ], and ‘roughly’ but not exactly coextensive with that between inorganic and organic chemistry [1916, 764]. But it is now a matter of degree rather than of kind, and relational, because it depends on the environment: a non-polar substance may be polarised by a polar solvent [1916, 765]. Bonds arise from atoms with incomplete electron shells filling them either by sharing electrons (in what came to be known as covalent, or shared-electron bonds) or transferring them (the ions and electrostatic bonding of polar compounds). Because shared electrons may not be shared equally, giving rise to partial charges on the bonded atoms, Lewis saw pure covalent and ionic bonds as two ends of a continuum, and offered the pairing of electrons as a unifying explanation of bonding in both polar and non-polar compounds. In his influential textbook Valence, Lewis again presented his theory as a unification, this time of the two great theories of chemical affinity of the nineteenth century [1923/1966, 20]. The electrochemical theory saw affinity as arising from the transfer of electricity between atoms: attraction between the resulting opposite charges would explain the stability of the compound. As Lewis pointed out [1923/1966, 20], this account found support in the fact that electrolysis demonstrated an intimate link between electricity and chemical combination, but foundered on the existence of homonuclear species like H 2 and N 2...
