Chapter 8 Covalent Bonding 8.3 Bonding Theories

Chapter 8 Covalent Bonding 8.3 Bonding Theories
8.1 Molecular Compounds 8.2 The Nature of Covalent Bonding 8.3 Bonding Theories 8.4 Polar Bonds and Molecules Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

CHEMISTRY & YOU What is the problem with a flat 2-dimensional map? It does not show an accurate picture of what is really going on or what something truly looks like. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

What do scientists use the VSEPR theory for?
Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

VSEPR Theory To explain the three-dimensional shape of molecules, scientists use Valence-Shell Electron-Pair Repulsion theory (VSEPR theory). VSEPR theory states that the repulsion between electron pairs causes molecular shapes to adjust so that the valence electron-pairs stay as far apart as possible. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

(electron dot structure)
VSEPR Theory The structural formula and electron dot structure of methane (CH4) show the molecule as if it were flat and two-dimensional. They fail to reflect the true three-dimensional shape of the molecule. Methane (electron dot structure) (structural formula) Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

All of the H–C–H angles are 109.5°, the tetrahedral angle.
VSEPR Theory In methane, the hydrogen atoms are at the four corners of a geometric shape called a tetrahedron. All of the H–C–H angles are 109.5°, the tetrahedral angle. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

VSEPR Theory Unshared (lone) pairs of electrons are also important in predicting the shapes of molecules. In ammonia (NH3) one of the pairs of electrons are unshared. Nonbinding pairs repel electrons more than shared pairs. Ammonia (NH3) Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

VSEPR Theory Because no atom needs these electrons; they are held closer than bonding pairs. The lone pair pushes the bonding pairs closer together changing the angle and shape. The H—N—H bond angle is only 107°, rather than the tetrahedral angle of 109.5°. Unshared electron pair 107° Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Thus, the water molecule is bent but flat
VSEPR Theory Water’s two lone pairs of electrons push the bonding pairs father around the central oxygen. Thus, the water molecule is bent but flat With two unshared pairs, the H—O—H bond angle is compressed to about 105°. Unshared pairs 105° Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Thus, CO2 is a linear molecule.
VSEPR Theory The carbon in a carbon dioxide molecule has no unshared electron pairs. The double bonds joining the oxygen atoms to the carbon are farthest apart and the O=C=O bond angle is 180°. Thus, CO2 is a linear molecule. Carbon dioxide (CO2) No unshared electron pairs on carbon 180° Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Here are some common molecular shapes.
VSEPR Theory Here are some common molecular shapes. Linear Trigonal planar Bent Pyramidal Tetrahedral Trigonal bipyramidal Octahedral Square planar Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

What causes valence electron-pairs to stay as far apart as possible?

What causes valence electron-pairs to stay as far apart as possible?
The repulsion between electron pairs due to their negative charges causes valence electron-pairs to stay as far apart as possible. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Molecular Orbitals How are atomic and molecular orbitals related?

Molecular Orbitals Just as an atomic orbital belongs to a particular atom, a molecular orbital belongs to a molecule as a whole. A molecular orbital that can be occupied by two electrons of a covalent bond is called a bonding orbital. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Molecular Orbitals When two atoms combine, their atomic orbitals overlap to produce molecular orbitals, or orbitals that apply to the entire molecule. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Molecular Orbitals Covalent bonding results from an imbalance between the attractions and repulsions of the nuclei and electrons involved. The nuclei and electrons attract each other. Nuclei repel other nuclei. Electrons repel other electrons. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

The result is a stable diatomic molecule.
Molecular Orbitals In a bonding molecular orbital, the attractions between the nuclei and the electrons are stronger than the repulsions. The balance of all the interactions between the atoms is thus tipped in favor of holding the atoms together. The result is a stable diatomic molecule. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Its symbol is the Greek letter sigma (σ).
Molecular Orbitals Sigma Bonds When a molecular orbital forms that is symmetrical around the axis connecting two atomic nuclei, a sigma bond molecular orbital is formed. Its symbol is the Greek letter sigma (σ). s atomic orbital Bond axis Sigma-bonding molecular orbital  represents the nucleus Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Molecular Orbitals Sigma Bonds Just like s orbitals, p orbitals can also overlap to form molecular orbitals. When p orbitals overlap end to end the result is a sigma bond. + + Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Molecular Orbitals Sigma Bonds There is a high probability of finding a pair of electrons between the positively charged nuclei of a sigma bond. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Molecular Orbitals Pi Bonds The side-by-side overlap of p orbitals produces what is called a pi bond molecular orbital. Its symbol is the Greek letter pi (). p atomic orbital Pi-bonding molecular orbital  represents the nucleus Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Molecular Orbitals Pi Bonds In a pi bond, the bonding electrons are most likely to be found in sausage-shaped regions above and below the bond axis of the bonded atoms. Because orbitals in pi bonding overlap less than in sigma bonding, pi bonds tend to be weaker than sigma bonds. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

In these drawings, show where an electron is most likely to be found?
Sigma Pi Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Hybrid Orbitals What is orbital hybridization? Hybrid Orbitals

Hybrid Orbitals Orbital hybridization is the combining of two different types of atomic orbitals. In hybridization, several atomic orbitals mix to form the same total number of equivalent hybrid orbitals. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

This creates four sp3 hybrid orbitals.
Hybridization of Orbitals Recall that the carbon atom promotes one of the 2s electrons to a 2p orbital to give it four bonding sites. This creates four sp3 hybrid orbitals. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Hybrid Orbitals Hybridization of Orbitals The energy of a Hybrid orbital is the average of all orbitals involved.  an “s” orbital has less energy than a “p”  when 1 “s” combines with 3 ‘p” orbitals, the energy level of the hybrids will be closer to the energy of a “p” orbital  this allows them to fill equally Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Hybrid Orbitals Hybridization of Orbitals Sulfur has 6 valence electrons leaving room for 2 flourine atoms to combine creating a SF2 molecule. So why do we get SF6? S F F Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Hybrid Orbitals Hybridization of Orbitals Hybridization of orbitals helps explain why some atoms can violate the octet rule. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

How many hybridized orbitals form when one 2s orbital is hybridized with two 2p orbitals?

How many hybridized orbitals form when one 2s orbital is hybridized with two 2p orbitals?
Three sp2 orbitals form. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Bond Polarity Covalent bonds differ in terms of how the bonded atoms share the electrons. The characteristics of the molecule depends on the kind and number of atoms joined together. These features, in turn, determine the molecular properties. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Bond Polarity Bond Polarity Charge distribution is based on the difference in electronegativity values found in a covalent bond. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Bond Polarity The bonding pairs of electrons are pulled between the nuclei of the atoms sharing the electrons. The nuclei of atoms pull on the shared electrons, much as the knot in the rope is pulled toward opposing sides in a tug-of-war. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

* polar means “having opposite ends”
Bond Polarity A polar covalent bond, (AKA) polar bond, is a bond between atoms in which the electrons are shared unequally. * polar means “having opposite ends” (one side of the molecule has a slight negative charge and the other side has a slight positive charge) Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Bond Polarity Describing Polar Covalent Bonds Hydrogen’s electronegativity is 2.1, and bromine’s electronegativity is 2.8. These values are significantly different, so the covalent bond in hydrogen bromide is polar. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Bond Polarity Describing Polar Covalent Bonds Hydrogen’s electronegativity is 2.1, and bromine’s electronegativity is 2.8. These values are significantly different, so the covalent bond in hydrogen bromide is polar. The bromine atom, with its higher electronegativity, acquires a slightly negative charge. The hydrogen atom acquires a slightly positive charge. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Bond Polarity Describing Polar Covalent Bonds The lowercase Greek letter delta (δ) denotes that atoms in the covalent bond acquire only partial charges, less than 1+ or 1–. δ+ δ– H—Br The minus sign shows that bromine has a slightly negative charge. The plus sign shows that hydrogen has acquired a slightly positive charge. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Bond Polarity Describing Polar Covalent Bonds A polar bond may also be represented by an arrow pointing to the more electronegative atom. H—Br Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

The H—O bonds in a water molecule are also polar.
Bond Polarity Describing Polar Covalent Bonds The H—O bonds in a water molecule are also polar. The highly electronegative oxygen partially pulls the bonding electrons away from hydrogen and becomes slightly negative leaving hydrogen slightly positive. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Bond Polarity When the atoms in a bond pull equally (occurs when identical atoms are bonded), the bonding electrons are shared equally, and each bond formed is a nonpolar covalent bond. remember the DIATOMIC 7 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Bond Polarity Describing Polar Covalent Bonds The electronegativity difference tells you what kind of bond is likely to form. Electronegativity Differences and Bond Types Electronegativity difference range Most probable type of bond Example 0.0–0.4 Nonpolar covalent Se—I (0.1) 0.4–1.0 Moderately polar covalent δ+ δ– H—Cℓ (0.9) 1.0–2.0 Very polar covalent δ+ δ– H—F (1.9) >2.0 Ionic Na+Cℓ– (2.1) Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

There is no sharp boundary between ionic and covalent bonds.
Bond Polarity Describing Polar Covalent Bonds There is no sharp boundary between ionic and covalent bonds. As electronegativity differences increases, the polarity of the bond increases. If the difference is more than 2.0, the electrons will likely be pulled away completely by one of the atoms. In that case, an ionic bond will form. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Sample Problem Identifying Bond Type Which type of bond (nonpolar covalent, moderately polar covalent, very polar covalent, or ionic) will form between the following pairs of atoms? a. Cℓ and P b. O and O c. Si and F d. Be and N Use Table 6.2 on page 181 for electronegativities. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Sample Problem Analyze Identify the relevant concepts. 1 In each case, the pairs of atoms involved in the bonding pair are given. The types of bonds depend on the electronegativity differences between the bonding elements. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Electronegativities based on Table 6.2 on page 181.
Sample Problem Solve Apply concepts to this problem. 2 Electronegativities based on Table 6.2 on page 181. a. Cℓ(3.0), P(2.1) b. O(3.5), O(3.5) c. Si(1.8), F(4.0) d. Be(1.5), N(3.0) Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Calculate the electronegativity difference between the two atoms.
Sample Problem Solve Apply concepts to this problem. 2 Calculate the electronegativity difference between the two atoms. The electronegativity difference between two atoms is expressed as the absolute value. So, you will never express the difference as a negative number. a. Cℓ(3.0), P(2.1); 0.9 b. O(3.5), O(3.5); 0.0 c. Si(1.8), F(4.0); 2.2 d. Be(1.5), N(3.0); 1.5 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

moderately polar covalent nonpolar covalent very polar covalent ionic
Sample Problem Solve Apply concepts to this problem. 2 Based on the electronegativity difference, determine the bond type using Table 8.4. a. Cℓ(3.0), P(2.1); 0.9; b. O(3.5), O(3.5); 0.0; c. Si(1.8), F(4.0); 2.2; d. Be(1.5), N(3.0); 1.5; moderately polar covalent nonpolar covalent very polar covalent ionic Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Bond Polarity Describing Polar Covalent Molecules The presence of a polar bond in a molecule often makes the entire molecule polar. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Bond Polarity Describing Polar Covalent Molecules In the hydrogen chloride molecule, the hydrogen and chlorine atoms become electrically charged regions, or poles. A molecule that has two poles is called a dipolar molecule, or dipole. The hydrogen chloride molecule is a dipole. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Describing Polar Covalent Molecules
Bond Polarity Describing Polar Covalent Molecules When polar molecules are placed between oppositely charged plates, they tend to become oriented with respect to the positive and negative plates. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Bond Polarity Describing Polar Covalent Molecules The effect of polar bonds on the polarity of an entire molecule depends on the shape of the molecule and the orientation of the polar bonds. A carbon dioxide molecule has two polar bonds and is linear and therefore nonpolar. O C O Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

The water molecule also has two polar bonds.
Bond Polarity Describing Polar Covalent Molecules The water molecule also has two polar bonds. However, the water molecule is bent rather than linear. Therefore, the bond polarities do not cancel and a water molecule is polar. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

What is the difference between an ionic bond and a very polar covalent bond?

What is the difference between an ionic bond and a very polar covalent bond?
Two atoms will form an ionic bond if their electronegativity differences are great enough to cause a transfer of electrons—a difference of more than 2.0. Or if there is a metal and a non-metal involved. Remember, there is no sharp boundary between ionic and covalent bonds. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Intermolecular attractions are weaker than either ionic or covalent bonds. * These attractions are responsible for determining if a molecular compound is a gas, a liquid, or a solid at a room temperature. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Van der Waals Forces The two weakest attractions between molecules are collectively called Van der Waals’ forces, named after the Dutch chemist Johannes van der Waals. Van der Waals forces consist of dipole interactions and dispersion forces. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Van der Waals Forces Dispersion forces, the weakest of all molecular interactions, are caused by the motion of electrons. They occur in polar and nonpolar molecules. Aka. = London forces Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Van der Waals Forces Dispersion forces, the weakest of all molecular interactions, are caused by the motion of electrons. When electrons are momentarily on one side of a molecule, they force the neighboring molecule’s electrons to be momentarily on the opposite side. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Van der Waals Forces Dispersion forces, the weakest of all molecular interactions, are caused by the motion of electrons. The strength of dispersion forces generally increases as the number of electrons in a molecule increases. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Van der Waals Forces Fluorine and chlorine have few electrons and are gases at room temperature because of their weak dispersion forces. Bromine is a liquid at room temperature because it has more electrons and the dispersion forces are stronger. Iodine, with a still larger number of electrons, is a solid at ordinary room temperature. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Van der Waals Forces Dipole interactions occur when polar molecules are attracted to one another. The attraction occurs between the oppositely charged regions of molecules. Dipole interactions are similar to, but much weaker than, ionic bonds. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Van der Waals Forces The slightly negative region of a polar molecule is weakly attracted to the slightly positive region of another polar molecule. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

The strong dipole interactions in water produce an attraction between water molecules. Each O—H bond in the water molecule is highly polar, and the oxygen acquires a slightly negative charge because of its greater electronegativity. The hydrogens in water molecules acquire a slightly positive charge. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Hydrogen Bonds Hydrogen bonds are attractive forces in which a hydrogen covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another atom. The other atom may be in the same molecule or in a nearby molecule. Hydrogen bonding always involves hydrogen. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Hydrogen Bonds A hydrogen bond has about 5 percent the strength of the average covalent bond. Hydrogen bonds are the strongest of the intermolecular forces. They are extremely important in determining the properties of water and biological molecules. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Why are hydrogen bonds important?

Why are hydrogen bonds important?
Hydrogen bonds are the strongest of the intermolecular forces and are extremely important in determining the properties of water and biological molecules such as proteins. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Intermolecular Attractions and Molecular Properties
At room temperature, some compounds are gases, some are liquids, and some are solids. The physical properties of a compound depend on the type of bonding it displays, in particular, on whether it is ionic or covalent. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

The diversity of physical properties among covalent compounds is mainly because of widely varying intermolecular attractions. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

The melting and boiling points of most molecular compounds are low compared with those of ionic compounds. In most molecular solids, only the weak attractions between molecules {intermolecular forces} need to be broken. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

A few solids that consist of molecules do not melt until the temperature reaches 1000°C or higher. Most of these very stable substances are network solids (or network crystals) in which all of the atoms are covalently bonded to each other. Melting a network solid would require breaking covalent bonds throughout the solid. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Diamond is an example of a network solid. Each carbon atom in a diamond is covalently bonded to four other carbons, interconnecting carbon atoms throughout the diamond. Diamond does not melt; rather, it vaporizes to a gas at 3500°C and above. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Characteristics of Ionic and Molecular Compounds
Intermolecular Attractions and Molecular Properties Differences between ionic and covalent (molecular) substances. Characteristics of Ionic and Molecular Compounds Characteristic Ionic Compound Molecular Compound Representative unit Formula unit Molecule Bond formation Transfer of one or more electrons between atoms Sharing of electron parts between atoms Type of elements Metallic and nonmetallic Nonmetallic Physical state Solid Solid, liquid, or gas Melting point High (usually above 300°C) Low (usually below 300°C) Solubility in water Usually high High to low Electrical conductivity of aqueous solution Good conductor Poor to non-conducting Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Why do network solids take so much more heat to melt than most covalent compounds?

Why do network solids take so much more heat to melt than most covalent compounds?
Melting a network solid requires breaking covalent bonds throughout the solid. Melting most covalent compounds only requires breaking the weak attractions between molecules. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Chapter 8 Covalent Bonding 8.3 Bonding Theories

Similar presentations

Presentation on theme: "Chapter 8 Covalent Bonding 8.3 Bonding Theories"— Presentation transcript:

Similar presentations

About project

Feedback

Log in

Auth with social network:

Chapter 8 Covalent Bonding 8.3 Bonding Theories

Similar presentations

Presentation on theme: "Chapter 8 Covalent Bonding 8.3 Bonding Theories"— Presentation transcript:

Similar presentations

About project

Feedback