Chemistry
Polar and Non-Polar Covalent Bonds
Polar covalent bonds occur when atoms with different electronegativities share electrons unequally, resulting in a partial positive and partial negative charge on the atoms. Non-polar covalent bonds form when atoms with similar electronegativities share electrons equally, leading to no net charge on the atoms. These types of bonds are important in understanding the behavior of molecules and their interactions.
Written by Perlego with AI-assistance
Related key terms
1 of 5
10 Key excerpts on "Polar and Non-Polar Covalent Bonds"
- Frederick Bettelheim, William Brown, Mary Campbell, Shawn Farrell(Authors)
- 2019(Publication Date)
- Cengage Learning EMEA(Publisher)
In a nonpolar covalent bond, electrons are shared equally. In a polar covalent bond, they are shared unequally. It is important to realize that no sharp line divides these two categories, nor, for that matter, does a sharp line divide polar covalent bonds and ionic bonds. Nonetheless, the rule-of-thumb guidelines given in Table 3.6 will help you decide whether a given bond is more likely to be non-polar covalent, polar covalent, or ionic. Nonpolar covalent A covalent bond between two atoms whose difference in electronegativity is less than 0.5 Polar covalent A covalent bond between two atoms whose difference in electronegativity is between 0.5 and 1.9 CHEMICAL CONNECTIONS 3B Ionic Compounds in Medicine Many ionic compounds have medical uses, some of which are shown in the table. ■ Test your knowledge with Problems 77, 78, and 79. Drinking a “barium cocktail,” which contains barium sulfate, makes the intestinal tract visible on an X-ray. Formula Name Medical Use AgNO 3 Silver nitrate Antibiotic BaSO 4 Barium sulfate Radiopaque medium for X-ray work CaSO 4 Calcium sulfate Plaster of Paris casts FeSO 4 Iron(II) sulfate Treatment of iron deficiency KMnO 4 Potassium Anti-infective permanganate (external) KNO 3 Potassium nitrate Diuretic (saltpeter) Li 2 CO 3 Lithium carbonate Treatment of bipolar disorder MgSO 4 Magnesium sulfate Cathartic (Epsom salts) NaHCO 3 Sodium bicarbonate Antacid (baking soda) NaI Sodium iodide Iodine for thyroid hormones NH 4 Cl Ammonium chloride Acidification of the digestive system (NH 4 ) 2 CO 3 Ammonium carbonate Expectorant SnF 2 Tin(II) fluoride To strengthen teeth (external) ZnO Zinc oxide Astringent (external) Charles D. Winters © Suttha Burawonk/Shutterstock.com 3.6 A Covalent Bond | 75 Copyright 2020 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. Due to electronic rights, some third party content may be suppressed from the eBook and/or eChapter(s).
eBook - PDF- Ageetha Vanamudan(Author)
- 2023(Publication Date)
- Delve Publishing(Publisher)
A triple bond is formed as a result of the two nitrogen atoms. Bonds with polar covalent nature is formed when atoms with varying degrees of electronegativity unite and share electrons unequally, they create this type of covalent connection. Electrons will be more attracted to more electronegative atoms. Even if it’s greater than 0, the electronegative difference between the atoms is less than 2.0. Thus, the shared electron pair of that atom will be in closer proximity (Huang & Zou, 2010). Hydrogen bonds may occur as a result of an imbalance in the electrostatic potential. Fluorine, hydrogen, and oxygen all react with hydrogen atoms in this situation. Bonds that are nonpolar (Nonpolar) Covalent bonds are formed when two atoms have an equal number of electrons sharing one atom’s valence shell. The electronegativity difference between any two atoms is a 0. This is the case when two diatomic elements (atoms with the same electron affinity) come together (diatomic elements), Nonpolar covalent bonds are found in gas molecules such as hydrogen and nitrogen. The Polarization of Chemical Bonds: It is always the more electronegative atom that has the electron cloud closest to it in sigma bonds. An ever-present dipole results from the bond’s being polarized. What Is the Difference Between Covalent and Ionic Bonds? Covalent and ionic Introduction to Molecular Science 122 bonds are two types of atomic bonding to consider. The architecture and behaviors of these bonds differ greatly. If the electron pairs of two atoms are connected in the same direction, covalent bonds are created. When two ions come into contact, they establish ionic bonds.
eBook - PDFFoundations of Chemistry
An Introductory Course for Science Students
- Philippa B. Cranwell, Elizabeth M. Page(Authors)
- 2021(Publication Date)
- Wiley(Publisher)
You are unlikely to be expected to memorise these values. Source: By Anne Helmenstine. Retrieved from https://sciencenotes.org/electronegativity-definition-and-trend/ 58 Chemical bonding 2.3.2 Polar bonds Consider a covalent bond formed between two atoms A and B. The single bond is composed of a pair of electrons. If atoms A and B have the same electroneg-ativity, then on average, the pair of electrons will be located evenly between the two atoms, as shown in Figure 2.21. This type of bond is called a pure covalent bond and is formed when the elements at the end of the bond are the same: for example, Cl 2 or H 2 . However, if one of the atoms (say, atom B) is more electronegative than the other, the electrons will be pulled towards that atom, and the distribution will no longer be evenly spread. Atom B will have a greater share of electrons than atom A, as shown in Figure 2.22. The result of this is that atom B becomes slightly negatively charged compared to atom A. We represent this charge by a δ -sign next to atom B. Because atom B takes up a slight negative charge, atom A must become slightly positively charged to balance the overall charge, and this is represented by the symbol δ +. A bond where there is an uneven distribution of charge, such as that shown in Figure 2.23, is called a polar covalent bond. It occurs between atoms where there is a reasonable difference in electronegativity. Examples of polar covalent bonds are found in hydrogen chloride, HCl, where the chlorine atom is more electro-negative than hydrogen; and carbon dioxide, CO 2 , where oxygen is more elec-tronegative than carbon. If there is a large difference in electronegativity between the atoms in a bond, then the electrons are concentrated on the more electronegative element. In the extreme case, the more electronegative element pulls the electrons completely towards itself and becomes a negatively charged ion (an anion), and the bonding is ionic.
eBook - PDF- R Sanderson(Author)
- 2012(Publication Date)
- Academic Press(Publisher)
FIVE Polar Covalence I: Electronegativity Equalization, Partial Charge, and Bond Length THE PRINCIPLE OF ELECTRONEGATIVITY EQUALIZATION The usual definition of electronegativity as the power of an atom in a molecule to attract electrons to itself leaves much to be desired in both clarity and significance. Electronegativity as listed is the property of an isolated atom. It changes when the atom is placed in a variety of conditions. A very simple model of the act of formation of a heteronuclear covalent bond can be very helpful in visualizing some of the fundamental consequences of chemical combination. Let us begin with atoms A and B y Β initially more electronegative than .4. Let us assume that each atom possesses at least one half-filled outermost orbital, giving it the capacity to form a covalent bond. When the two atoms come in contact, the single electron of A finds the vacancy of Β available and the single electron of Β finds the vacancy of A available. Thus both electrons come under the considerable influence of both nuclei and are shared between the two atoms. The attraction between the two atoms that results from this mutual sharing of the two bonding electrons is called a covalent bond. Since here the atoms are not alike, the bond is called heteronuclear. In the fact of sharing the same two electrons between two nuclei, the heteronuclear bond is just like a homonuclear bond. However, a difference arises from the fact that Β is initially more electronegative than A. This implies that the bonding electrons will be more strongly attracted to the nucleus of Β than to the nucleus of A. 75 76 5. Polar Covalence I A stable system cannot result unless in effect the bonding electrons are able to adjust to a condition of essentially equal attraction to both nuclei. Electrons are free to move according to the forces acting upon them, and they are certainly not to be expected to remain evenly distributed if they are not evenly attracted.
eBook - PDFIntroductory Chemistry
An Active Learning Approach
- Mark Cracolice, Edward Peters, Mark Cracolice(Authors)
- 2020(Publication Date)
- Cengage Learning EMEA(Publisher)
The two electrons joining the atoms in the H 2 molecule are shared equally by the two nuclei. A bond in which bonding electrons are shared equally is a nonpolar covalent bond. A more formal way of saying this is that the charge density is centered in the region between the bonded atoms. A bond between identical atoms, as in H 2 or F 2 , is always nonpolar. The fluorine atom in an HF molecule has a stronger attraction for the bonding electron pair than the hydrogen atom has. The bonding electrons, therefore, spend Copyright 2021 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. Due to electronic rights, some third party content may be suppressed from the eBook and/or eChapter(s). Editorial review has deemed that any suppressed content does not materially affect the overall learning experience. Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it. 451 12.5 Polar and Nonpolar Covalent Bonds more time nearer the fluorine atom nucleus than the hydrogen atom nucleus. A bond in which bonding electrons are shared, but shared unequally, is a polar covalent bond. The charge density is shifted toward the fluorine atom and away from the hydrogen atom. Because the negative charge density is closer to the fluorine atom, that atom acts as a negative pole, and the hydrogen atom acts as a positive pole (Figure 12.14). When the charge density shift is extreme, the bonding electrons are effectively trans- ferred from one atom to another, and an ionic bond results (Figure 12.15). Bond polarity in covalent bonds may be described in terms of the electronega- tivities of the bonded atoms. American chemist Linus Pauling (Figure 12.16) origi- nally described the electronegativity of an element as the ability of an atom of that element in a molecule to attract bonding electron pairs to itself.
eBook - PDF- Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson(Authors)
- 2019(Publication Date)
- Openstax(Publisher)
The charges of anions formed by the nonmetals may also be readily determined because these ions form when nonmetal atoms gain enough electrons to fill their valence shells. 7.2 Covalent Bonding Covalent bonds form when electrons are shared between atoms and are attracted by the nuclei of both atoms. In pure covalent bonds, the electrons are shared equally. In polar covalent bonds, the electrons are shared unequally, as one atom exerts a stronger force of attraction on the electrons than the other. The ability of an atom to attract a pair of electrons in a chemical bond is called its electronegativity. The difference in electronegativity between two atoms determines how polar a bond will be. In a diatomic molecule with two identical atoms, there is no difference in electronegativity, so the bond is nonpolar or pure covalent. When the electronegativity difference is very large, as is the case between metals and nonmetals, the bonding is characterized as ionic. 7.3 Lewis Symbols and Structures Valence electronic structures can be visualized by drawing Lewis symbols (for atoms and monatomic ions) and Lewis structures (for molecules and polyatomic ions). Lone pairs, unpaired electrons, and single, double, or triple bonds are used to indicate where the valence electrons are located around each atom in a Lewis structure. Most structures—especially those containing second row elements—obey the octet rule, in which every atom (except H) is surrounded by eight electrons. Exceptions to the octet rule occur for odd-electron molecules (free radicals), electron-deficient molecules, and hypervalent molecules. 7.4 Formal Charges and Resonance In a Lewis structure, formal charges can be assigned to each atom by treating each bond as if one-half of the electrons are assigned to each atom. These hypothetical formal charges are a guide to determining the most appropriate Lewis structure. A structure in which the formal charges are as close to zero as possible is preferred.
eBook - PDFChemistry
The Molecular Nature of Matter
- Neil D. Jespersen, Alison Hyslop(Authors)
- 2014(Publication Date)
- Wiley(Publisher)
The electron density at both ends of the bond is the same, because the electrons are equally attracted to both nuclei. However, when different kinds of atoms combine, as in HCl, one nucleus usually attracts the electrons in the bond more strongly than the other. Polar and Nonpolar Bonds The result of unequal attractions for the bonding electrons is an unbalanced distribution of electron density within the bond. For example, chlorine atoms have a greater attraction for electrons in a bond than do hydrogen atoms. In the HCl molecule, therefore, the elec- tron cloud is pulled more tightly around the Cl, and that end of the molecule experiences a slight buildup of negative charge. The electron density that shifts toward the chlorine is removed from the hydrogen, which causes the hydrogen end to acquire a slight positive charge. These charges are less than full 1+ and 1- charges and are called partial charges, which are usually indicated by the lowercase Greek letter delta, d (see Figure 8.8). Partial charges can also be indicated on Lewis structures. For example, HOCl a C d + d - A bond that carries partial positive and negative charges on opposite ends is called a polar covalent bond, or often simply a polar bond (the word covalent is understood). The term polar comes from the notion of poles of equal but opposite charge at either end of the bond. Because two poles of electric charge are involved, the bond is said to be an electric dipole. The polar bond in HCl causes the molecule as a whole to have opposite charges on either end, so the HCl molecule as a whole is an electric dipole. We say that HCl is a Practice Exercise 8.9 Practice Exercise 8.10 (a) (b) H δ+ δ– H H Cl Figure 8.8 | Equal and unequal sharing of electrons in a covalent bond. Each of the diagrams illustrates the distribution of electron density of the shared electron pair in a bond. (a) In H 2 , the electron density in the bond is spread equally over both atoms.
eBook - PDFChemistry
An Atoms First Approach
- Steven Zumdahl, Susan Zumdahl, Donald J. DeCoste, , Steven Zumdahl, Steven Zumdahl, Susan Zumdahl, Donald J. DeCoste(Authors)
- 2020(Publication Date)
- Cengage Learning EMEA(Publisher)
An example of this type of bond occurs in molecules of HF (hydrogen fluoride). When a sample of hydrogen fluoride gas is placed in an electric field, the molecules tend to orient themselves as shown in Fig. 3.3, with the fluoride end closest to the positive pole and the hydrogen end clos- est to the negative pole. This result implies that the HF molecule has the following charge distribution: HOF d 1 d 2 where d (lowercase delta) is used to indicate a fractional charge. The most logical explanation for the development of the partial positive and nega- tive charges on the atoms (bond polarity) in molecules such as HF is that the electrons in the bonds are not shared equally. For example, we can account for the polarity of the HF molecule by assuming that the fluorine atom has a stronger attraction for the shared electrons than the hydrogen atom. Because bond polarity has important chemical im- plications, we find it useful to quantify the ability of an atom to attract shared elec- trons. In the next section we show how this is done. 3.2 Electronegativity The different affinities of atoms for the electrons in a bond are described by a property called electronegativity: the ability of an atom in a molecule to attract shared elec- trons to itself. The most widely accepted method for determining values of electronegativity is that of Linus Pauling (1901–1995), an American scientist who won the Nobel Prizes for both chemistry and peace. To understand Pauling’s model, consider a hypothetical Ionic and covalent bonds are the extreme bond types. - δ + δ + + H No electric field Electric field F δ - δ - a b FIGURE 3.3 The effect of an electric field on hydrogen fluoride molecules. (a) When no electric field is present, the molecules are randomly oriented. (b) When the field is turned on, the molecules tend to line up with their negative ends toward the positive pole and their positive ends toward the negative pole.
eBook - PDFChemistry
Structure and Dynamics
- James N. Spencer, George M. Bodner, Lyman H. Rickard(Authors)
- 2011(Publication Date)
- Wiley(Publisher)
4.17 The Difference between Polar Bonds and Polar Molecules The difference between the electronegativities ( EN) of chlorine (EN 2.87) and hydrogen (EN 2.30) is 0.57 and is sufficiently large that the bond in HCl is said to be polar. As we saw in Section 4.11, the partial charge on the hydrogen is positive and that on the chlorine is negative. Because it contains only this one bond, the HCl molecule can be described as polar. The polarity of a bond in a structure is represented with an arrow hav- ing a plus sign at the positive end (the and shown in the above struc- ture could be replaced by the | S symbol). The arrow points in the direction of the more electronegative atom, and the plus sign is next to the least elec- tronegative atom. The polarity of a molecule can be predicted by considering the polarities of the individual bonds, the location of nonbonding pairs, and the three-dimen- sional shape of the molecule. The magnitude of the polarity of a molecule is measured using a quantity known as the dipole moment, . The dipole moment for a molecule depends on two factors: (1) the magnitude of the charge and (2) the distance between the negative and positive poles of the molecule. Dipole moments are reported in units of debye (D). The dipole moment for HCl is small, 1.08 D (Table 4.4). This can be understood by noting that the separation of charge in the HCl bond is relatively small because the H¬Cl bond is rela- tively short. C¬Cl bonds ( EN 0.33) are not as polar as H¬Cl bonds ( EN 0.57), but they are significantly longer. As a result, the dipole moment of CH 3 Cl (1.89 D) shown in Figure 4.22 is greater than that for HCl. At first glance, we might expect a similar dipole moment for carbon tetra- chloride (CCl 4 ), which contains four polar C¬Cl bonds. The experimental value of the dipole moment for CCl 4 , however, is zero. This can be understood by con- sidering the structure of CCl 4 , shown in Figure 4.22.
eBook - PDF- R Sanderson(Author)
- 2012(Publication Date)
- Academic Press(Publisher)
W e i n s t e i n (A9). It m a y b e s t a t e d : When t w o or m o r e a t o m s u n i t e t o f o r m a c o m p o u n d , t h e i r e l e c t r o n e g a t i v i t i e s b e c o m e adjusted t o t h e s a m e i n t e r m e d i a t e v a l u e within t h e 38 POLAR COVALENCE Figure 3:1 Representation of Electronegativity Equalization Atom A has relatively few outer electrons, corresponding to: 1) a relatively small effective nuclear charge 2) a relatively large nonpolar covalent radius 3) a relatively diffuse electronic cloud and therefore 4) a low electronegativity. Atom Β has a nearly filled outer octet, corresponding to: 1) a relatively large effective nuclear charge 2) a relatively small nonpolar covalent radius 3) a relatively compact electronic cloud and therefore 4) a relatively high electronegativity. When the two atoms join by a covalent bond, A-B, the bonding electrons are attracted initially more strongly to B, giving it a surplus (or partial negative charge), and leaving A with a deficiency of electrons (or partial positive charge). With its electron population reduced, a) the effective nuclear charge of A increases b) causing the cloud to contract to smaller radius c) and A to be more electronegative. With an increase in electron population, a) the effective nuclear charge of Β decreases b) causing the cloud to expand to larger radius c) and Β to be less electronegative. These adjustments cease when the two atoms have become equal in electronegativity, by virtue of acquiring partial charge. 3 The Theory of Polar Covalence 39 compound. A c o r o l l a r y is, t h e i n t e r m e d i a t e e l e c t r o n e g a t i v i t y within t h e c o m p o u n d i s t h e g e o m e t r i c m e a n of all t h e a t o m i c e l e c t r o n e g a t i v i t i e s . N o t e t h a t t h e d i f f e r e n t k i n d s of a t o m s b e c o m e equal in e l e c t r o n e g a t i v i t y b y unequal s h a r i n g of t h e b o n d i n g e l e c -t r o n s .
Index pages curate the most relevant extracts from our library of academic textbooks. They’ve been created using an in-house natural language model (NLM), each adding context and meaning to key research topics.
