Manual Inorganic Chemistry Review: Chemical Bonding (Quick Review Notes)

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One technologically important set of isoelectronic solids are the p-block semiconductors. The group 14 element Si is the most widely used semiconductor for electronics, but, as we will discuss later, it is not a good light emitter. Light-emitting diodes LEDs , which are used in lasers, high efficiency lighting, and display technologies, are made from compounds that are isoelectronic with Si and Ge, especially GaAs, GaP, AlAs, and GaN all contain four valence electrons per atom.

Like Si and Ge, these compounds have tetrahedrally bonded structures in the solid state and absorb light across most of the solar spectrum, as we will discuss in more detail in Chapters 8 and The isoelectronic principle is also a powerful tool in materials research, because it provides guidance about where to look for new materials with similar and perhaps improved properties. For example, the discovery that 8.

To use the VSEPR model, one begins with the Lewis dot picture to determine the number of lone pairs and bonding domains around a central atom. For example, in either the hypervalent or octet structure of the I 3 - ion above, there are three lone pairs on the central I atom and two bonding domains. We then follow these steps to obtain the electronic geometry: The molecular geometry is deduced from the electronic geometry by considering the lone pairs to be present but invisible.

The most commonly used methods to determine molecular structure - X-ray diffraction , neutron diffraction , and electron diffraction - have a hard time seeing lone pairs, but they can accurately determine the lengths of bonds between atoms and the bond angles. The table below gives examples of electronic and molecular shapes for steric numbers between 2 and 9. We are most often concerned with molecules that have steric numbers between 2 and 6. From the Table, we see that some of the molecules shown as examples have bond angles that depart from the ideal electronic geometry. We can rationalize this in terms of the last rule above.

The lone pair in ammonia repels the electrons in the N-H bonds more than they repel each other. This lone pair repulsion exerts even more steric influence in the case of water, where there are two lone pairs. For some molecules in the Table, we note that there is more than one possible shape that would satisfy the VSEPR rules. For example, the XeF 2 molecule has a steric number of five and a trigonal bipyramidal geometry.

There are three possible stereoisomers: The observed geometry of XeF 2 is linear, which can be rationalized by considering the orbitals that are used to make bonds or lone pairs in the axial and equatorial positions. There are four available orbitals, s, p x , p y , and p z. If we choose the z-axis as the axial direction, we can see that the p x and p y orbitals lie in the equatorial plane.

We assume that the spherical s orbital is shared equally by the five electron domains in the molecule, the two axial bonds share the p z orbital, and the three equatorial bonds share the p x and p y orbitals. We can then calculate the bond orders to axial and equatorial F atoms as follows: Because fluorine is more electronegative than a lone pair, it prefers the axial site where it will have more negative formal charge. In general, by this reasoning, lone pairs and electropositive ligands such as CH 3 will always prefer the equatorial sites in the trigonal bipyramidal geometry.

Electronegative ligands such as F will always go to the axial sites. In the case of the BrF 4 - anion, which is isoelectronic with XeF 4 in the Table, the electronic geometry is octahedral and there are two possible isomers in which the two lone pairs are cis or trans to each other. In this case, lone pair - lone pair repulsion dominates and we obtain the trans arrangement of lone pairs, giving a square planar molecular geometry.

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Orbital hybridization The observation of molecules in the various electronic shapes shown above is, at first blush, in conflict with our picture of atomic orbitals. For an atom such as oxygen, we know that the 2s orbital is spherical, and that the 2p x , 2p y , and 2p z orbitals are dumbell-shaped and point along the Cartesian axes Fig. Given the relative orientations of the atomic orbitals, how do we arrive at angles between electron domains of However, we still impose the constraint that our hybrid orbitals must be orthogonal and normalized.

For sp hybridization , as in the BeF 2 or CO 2 molecule, we make two linear combinations of the 2s and 2p z orbitals assigning z as the axis of the Be-F bond:. Here we have simply added and subtracted the 2s and 2pz orbitals; we leave it as an exercise for the interested student to show that both orbitals are normalized i. What this means physically is explained in Fig. By combining the 2s and 2p z orbitals we have created two new orbitals with large lobes high electron probability pointing along the z-axis.

These two orbitals are degenerate and have an energy that is halfway between the energy of the 2s and 2p z orbitals. For an isolated Be atom, which has two valence electrons, the lowest energy state would have two electrons spin-paired in the 2s orbital. However, these electrons would not be available for bonding. By promoting these electrons to the degenerate 2sp z hybrid orbitals, they become unpaired and are prepared for bonding to the F atoms in BeF 2.

This will occur if the bonding energy in the promoted state exceeds the promotion energy. The overall bonding energy , i. We can similarly construct sp 2 hybrids e. The three hybrids are:. These orbitals are again degenerate and their energy is the weighted average of the energies of the 2s, 2p x , and 2p y atomic orbitals.

Finally, to make a sp 3 hybrid, as in CH 4 , H 2 O, etc. The lobes of the sp 3 hybrid orbitals point towards the vertices of a tetrahedron or alternate corners of a cube , consistent with the tetrahedral bond angle in CH 4 and the nearly tetrahedral angles in NH 3 and H 2 O. Similarly, we can show that we can construct the trigonal bipyramidal electronic shape by making sp and sp 2 hybrids, and the octahedral geometry from three sets of sp hybrids. The picture that emerges from this is that the atomic orbitals can hybridize as required by the shape that best minimizes electron pair repulsions.

This is consistent with the fact that the energy difference between s and p orbitals stays roughly constant going down the periodic table, but the bond energy decreases as the valence electrons get farther away from the nucleus. In compounds of elements in the 3rd, 4th, and 5th rows of the periodic table, there thus is a decreasing tendency to use s-p orbital hybrids in bonding.

For these heavier elements, the bonding energy is not enough to offset the energy needed to promote the s electrons to s-p hybrid orbitals. Linus Pauling introduced the concept of electronegativity in order to explain the extra stability of molecules with polar bonds. This definition, while directly relevant to the strength of chemical bonds, requires thermochemical input data from many compounds, some of which were not available at the time. Mulliken [10] [11] and later Pearson [12] developed a scale of electronegativities based on the average of the electron affinity and ionization energy of the free A and B atoms, which they correlated with thermochemical data and the Pauling scale.

Carbon and hydrogen have intermediate electronegativities 2. The general trend see table below is that electronegativities increase going up and to the right in the periodic table. The first of these can be explained by the very high metal-metal bond energy of elements such as Mo and W, which can use all six of their valence electrons in bonding, as we will discuss in Chapter 6. The second however occurs with more weakly bonded noble metals such as Pt and Au, and is responsible for their low position in the activity series , [13] as well as their extraordinary properties as catalysts.

Table of Pauling electronegativities. The polarity of bonds is determined by electronegativity differences.

As a guideline we define bonds as:. The polarity of bonds helps us understand non-covalent forces between molecules, such as hydrogen bonding and dipole-dipole interactions. It also helps us interpret the reactivity of molecules. Similarly, electrophilic substitution reactions occur more readily on Si-H and P-H compounds than they do on C-H compounds. There is also a correlation between the strength of a chemical bond and the bond length, longer bonds being weaker because of weaker orbital overlap.

Pauling introduced an empirical formula relating bond length to bond strength. For a given pair of atoms for example, two carbon atoms: D 1 in this case would be the length of a C-C single bond, which we can obtain from the average bond length in alkanes 1. In a related form the Pauling formula can be used to calculate bond lengths when the single bond length D 1 is not available: Here n and m represent two different bond orders between the same kinds of atoms.

This tells, for example, that the difference in length between a triple and double bond, D 2 -D 3 , should be - 0. Some bond lengths and bond energies are anomalous. For example, the F-F bond length in F 2 is 1. By putting the extra bond length into the Pauling formula, we calculate that the bond order in the F 2 molecule is only 0.

The physical reason for this is that the F-F bond is "stretched" by repulsion of the lone pairs on the F atoms. This crowding is caused by the fact that the [He] 1s 2 core orbital, as well as the valence orbitals of the fluorine atoms, are contracted by the high nuclear charge. The Cl 2 atom, with its larger [Ne] 1s 2 2s 2 2p 6 core, in contrast, has a "normal" single bond length 1.

The weak bond in F 2 is responsible for the high electronegativity of fluorine, as well as the legendary reactivity of elemental fluorine gas, which reacts explosively with hydrogen and powdered metals. Because of the instability of elemental fluorine and the polar nature of its bonds with more electropositive elements, fluorine compounds tend to be very stable. For example, the noble gases Xe and Kr react with fluorine to make covalent compounds, whereas other halogens do not react.

Fluorocarbon compounds contain strong C-F bonds and have high thermal and chemical stability.

8.S: Basic Concepts of Chemical Bonding (Summary) - Chemistry LibreTexts

The extraordinary hydrophobicity of perfluorocarbons arises from the fact that -CF 2 - and -CF 3 groups are "fatter" than -CH 2 - and -CH 3 groups; dissolving them in water is therefore more disruptive to the hydrogen bonding network than is dissolving a hydrocarbon. Join a group of classmates from your inorganic chemistry course. Each group should have 3 or 4 students.

Arrange a time to meet for an hour or so. Prior to the meeting, read part I pp. At the meeting, discuss the following questions: Please include the names of all students in your group on the cover page; all of you will receive the same grade on this assignment. Anonymous excerpts from a few of these essays will be shared with the class. The N-N bond distance is 1. Can you draw octet structures for these compounds? Why would these molecules be unstable? Consider the compounds NH 3 and PH 3. The H-N-H bond angle in ammonia is o close to the tetrahedral angle, Why is the angle in PH 3 closer to 90 o than it is to the tetrahedral angle?

Which one is more stable? Krypton difluoride, KrF 2 , decomposes at dry ice temperature to Kr and F 2. Consider the molecule ClF 3 O 2 with Cl the central atom.

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How many isomers are possible? Which is the most stable? The Br-F bond distance in the interhalogen compound BrF is 1. Use this information to estimate the average bond lengths in BrF 3 and BrF 5. Explain why the N-N bond in hydrazine is so weak, and why this effect is not seen in N 2. Show that a set of three sp 2 orbitals satisfies the following criteria: The atomic connectivity in the ring is: From Wikibooks, open books for an open world.

Lewis, "The atom and the molecule," J. Silvi, "The octet rule and hypervalence: Kanatzidis, All-solid-state dye-sensitized solar cells with high efficiency, Nature , Electroaffinity, Molecular Orbitals and Dipole Moments". Intermolecular hydrogen bonding increases boiling point of the compound and also its water solubility. Intramolecular hydrogen bonding decreases the boiling point of the compound and also its water solubility.

A covalent bond is formed by overlapping of valence shell atomic orbital of the two atoms having unpaired electron.

Introduction to Inorganic Chemistry/Review of Chemical Bonding

When covalent bond is formed by overlapping of atomic orbitals along the same axis it is called s - bond. Such type of bond is symmetrical about the line joining the two nuclei e. This type of bond is formed by the sidewise or lateral overlapping of two half filled atomic orbitals. The strength of a bond depends upon the extent of overlapping of half-filled atomic orbitals. The extent of overlapping is between two atoms is always greater when there is end to end overlapping of orbitals than, when there is sidewise overlapping of oritals. Hence s-bond is always stronger than p-bond.

The same amount of energy is released in formation of one mol of particular bond. T he mixi ng of dissimilar orbital of similar energies to form new set of hybrid orbital.

In this case, one s and one p orbital mix together to form two sp hybrid orbitals and are oriented in a linear shape. The remaining two unhybridised p orbitals overlap with another unhybridised p orbital leading to the formation of triple bond as in HC CH. The distance between the nuclei of two atoms bonded together is called bond length.

Bond length decreases with increase in s-character since s-orbital is smaller than a p — orbital. Bond angel is the angle between two adjacent bonds at an atom in a molecule made up of three or more atoms. Bond angles mainly depend on the following three factors: Bond angle depends on the state of hybridization of the central atom. Generally s- character increase in the hybrid bond, the bond angle increases.

Bond angle is affected by the presence of lone pair of electrons at the central atom. A lone pair of electrons at the central atom always tries to repel the shared pair bonded pair of electrons. Due to this, the bonds are displaced slightly inside resulting in a decrease of bond angle. If the electronegativity of the central atom decreases, bond angle decreases. In simple homonuclear diatomic molecules the order of MO's based on increasing energy is. Dear , Preparing for entrance exams? Register yourself for the free demo class from askiitians. Studying in Grade 6th to 12th?