Why shape of molecule is important

It is important to note that electron-pair geometry around a central atom is not the same thing as its molecular structure. The electron-pair geometries shown in Figure 3 describe all regions where electrons are located, bonds as well as lone pairs. Molecular structure describes the location of the atoms , not the electrons. Figure 4. The molecular structure of the methane molecule, CH 4 , is shown with a tetrahedral arrangement of the hydrogen atoms.

We differentiate between these two situations by naming the geometry that includes all electron pairs the electron-pair geometry. The structure that includes only the placement of the atoms in the molecule is called the molecular structure.

The electron-pair geometries will be the same as the molecular structures when there are no lone electron pairs around the central atom, but they will be different when there are lone pairs present on the central atom. VSEPR structures like the one in Figure 4 are often drawn using the wedge and dash notation, in which solid lines represent bonds in the plane of the page, solid wedges represent bonds coming up out of the plane, and dashed lines represent bonds going down into the plane.

For example, the methane molecule, CH 4 , which is the major component of natural gas, has four bonding pairs of electrons around the central carbon atom; the electron-pair geometry is tetrahedral, as is the molecular structure Figure 4. On the other hand, the ammonia molecule, NH 3 , also has four electron pairs associated with the nitrogen atom, and thus has a tetrahedral electron-pair geometry. One of these regions, however, is a lone pair, which is not included in the molecular structure, and this lone pair influences the shape of the molecule Figure 5.

Figure 5. As seen in Figure 5, small distortions from the ideal angles in Figure 6 can result from differences in repulsion between various regions of electron density. VSEPR theory predicts these distortions by establishing an order of repulsions and an order of the amount of space occupied by different kinds of electron pairs.

The order of electron-pair repulsions from greatest to least repulsion is:. This order of repulsions determines the amount of space occupied by different regions of electrons. A lone pair of electrons occupies a larger region of space than the electrons in a triple bond; in turn, electrons in a triple bond occupy more space than those in a double bond, and so on.

The order of sizes from largest to smallest is:. Consider formaldehyde, H 2 CO, which is used as a preservative for biological and anatomical specimens Figure 1. This molecule has regions of high electron density that consist of two single bonds and one double bond.

Because electrostatic interactions fall off rapidly with increasing distance between molecules, intermolecular interactions are most important for solids and liquids, where the molecules are close together. These interactions become important for gases only at very high pressures, where they are responsible for the observed deviations from the ideal gas law at high pressures.

Substances with the highest melting and boiling points have covalent network bonding. This type of interaction is actually a covalent bond. In these substances, all the atoms in a sample are covalently bonded to the other atoms; in effect, the entire sample is essentially one large molecule. Many of these substances are solid over a large temperature range because it takes a lot of energy to disrupt all the covalent bonds at once.

One example of a substance that shows covalent network bonding is diamond Figure 4. Diamond, a form of pure carbon, has covalent network bonding. For interactions between different molecules, the strongest force between any two particles is the ionic bond , in which two ions of opposing charge are attracted to each other. Thus, ionic interactions between particles are an intermolecular interaction. Substances that contain ionic interactions are strongly held together, so these substances typically have high melting and boiling points.

Sodium chloride Figure 4. Solid NaCl is held together by ionic intermolecular forces. Many substances that experience covalent bonding exist as discrete molecules and do not engage in covalent network bonding. Thus, most covalently bonded molecules will also experience intermolecular forces. These intermolecular forces are weaker than those found in ionic interactions and depend on the polarity of the covalent bond. Recall that in polar covalent bonds, the electrons that are shared in a covalent bond are not shared equally between the two atoms in the bond.

Typically, the atom displaying higher electronegativity attracts the electrons more strongly than the other, leading to the unequal sharing of electrons in the bond.

This idea is illustrated in Figure 4. The electrons in the HF molecule are not equally shared by the two atoms in the bond. Because the fluorine atom has nine protons in its nucleus, it attracts the negatively charged electrons in the bond more than the hydrogen atom does with its one proton in its nucleus.

Thus, electrons are more strongly attracted to the fluorine atom, leading to an imbalance in the electron distribution between the atoms. Such a bond is called a polar covalent bond. The fluorine atom attracts the electrons in the bond more than the hydrogen atom does.

The result is an unequal distribution of electrons in the bond, favoring the fluorine side of the covalent bond. A covalent bond that has an unequal sharing of electrons is called a polar covalent bond. A covalent bond that has an equal sharing of electrons, as in a covalent bond with the same atom on each side, is called a nonpolar covalent bond.

A molecule with a net unequal distribution of electrons in its covalent bonds is a polar molecule. HF is an example of a polar molecule. The charge separation in a polar covalent bond is not as extreme as is found in ionic compounds, but there is a related result: oppositely charged ends of different molecules will attract each other.

This type of intermolecular interaction is called a dipole-dipole interaction. If the structure of a molecule is polar, then the molecule has a net dipole moment. Molecules with net dipole moments tend to align themselves so that the positive end of one dipole is near the negative end of another and vice versa, as shown in part a in Figure 4. These arrangements are more stable than arrangements in which two positive or two negative ends are adjacent Figure 4. Hence dipole—dipole interactions, such as those in part b in Figure 4.

Because molecules in a liquid move freely and continuously, molecules always experience both attractive and repulsive dipole—dipole interactions simultaneously, as shown in Figure 4. On average, however, the attractive interactions dominate. Because of this strong interaction, hydrogen bonding is used to describe this dipole-dipole interaction.

The physical properties of water, which has two O—H bonds, are strongly affected by the presence of hydrogen bonding between water molecules. The presence of hydrogen bonding in molecules like water can have a large impact on the physical properties of a substance. These are very weak intermolecular interactions and are called London dispersion forces. Since electrons naturally orbit the nucleus of the atom, there are momentary dipoles that are present in the atom as the electrons are shifting from one side to the other.

If other atoms are in close proximity, the electrons of the other atoms will orbit in concert with the neighboring atom, i. The actual molecule is an average of structures 2 and 3 , which are called resonance structures. Structure 1 is also a resonance structure of 2 and 3 , but since it has more formal charges, and does not satisfy the octet rule, it is a higher-energy resonance structure, and does not contribute as much to our overall picture of the molecule.

The real molecule does not alternate back and forth between these two structures; it is a hybrid of these two forms. The ozone molecule, then, is more correctly shown with both Lewis structures, with the two-headed resonance arrow between them:. In contrast, the lone pairs on the oxygen in water are localized — i.

Resonance delocalization stabilizes a molecule by spreading out charges, and often occurs when lone pairs or positive charges are located next to double bonds. Resonance plays a large role in our understanding of structure and reactivity in organic chemistry. A more accurate picture of bonding in molecules like this is found in Molecular Orbital theory, but this theory is more advanced, and mathematically more complex topic, and will not be dealt with here.

Examples We can satisfy the octet rule and make the formal charges smaller by making a carbon-oxygen double bond. Once again, structure 1 is a resonance structure of 2 , 3 , and 4 , but it is a higher energy structure, and does not contribute as much to our picture of the molecule. Multi-Center Molecules Molecules with more than one central atoms are drawn similarly to the ones above. The octet rule and formal charges can be used as a guideline in many cases to decide in which order to connect atoms.

C 2 H 6 ethane C 2 H 4 ethylene The octet rule is not satisfied on the B, but the formal charges are all zero. In fact, trying to make a boron-fluorine double bond would put a positive formal charge on fluorine; since fluorine is highly electronegative, this is extremely unfavorable.

In this structure, the formal charges are all zero, but the octet rule is not satisfied on the N. Since there are an odd number of electrons, there is no way to satisfy the octet rule. Nitric oxide is a free radical, and is an extremely reactive compound. In the body, nitric oxide is a vasodilator, and is involved in the mechanism of action of various neurotransmitters, as well as some heart and blood pressure medications such as nitroglycerin and amyl nitrite.

Notice that the formal charge on the phosphorus atom is zero. Notice that the formal charge on the sulfur atom is zero. Notice that the formal charge on the xenon atom is zero. Structures 1 and 2 are resonance structures of each other, but structure 2 is the lower energy structure, even though it violates the octet rule. Sulfur can accommodate more than eight electrons, and the formal charges in structure 2 are all zero.

Lone pairs go in the equatorial positions, since they take up more room than covalent bonds. The Lewis structures of the previous examples can be used to predict the shapes around their central atoms:. Formula Lewis Structure Bonding Shape 1. CH 4 4 bonds. NH 3 3 bonds.

HCN 2 bonds. CO 2 2 bonds. CCl 4 4 bonds. COCl 2 3 bonds. O 3 2 bonds. CO 3 2- 3 bonds. O: bent BF 3 3 bonds. NO linear PCl 5 5 bonds. SF 6 6 bonds. SF 4 4 bonds. XeF 4 4 bonds. O: bent With Lewis structures involving resonance, it is irrelevant which structure is used to determine the shape, since they are all energetically equivalent.

Polar and Nonpolar Covalent Bonds Electronegativity is a measure of the ability of an atom in a molecule to attract shared electrons in a covalent bond. Electronegativity is a periodic property, and increases from bottom to top within a group and from left to right across a period:.

When two atoms of the same electronegativity share electrons, the electrons are shared equally , and the bond is a nonpolar covalent bond — there is a symmetrical distribution of electrons between the bonded atoms.

Water has four electron groups so it falls under tetrahedral for the electron-group geometry. The four electron groups are the 2 single bonds to Hydrogen and the 2 lone pairs of Oxygen. Since water has two lone pairs it's molecular shape is bent. According to the VSEPR theory, the electrons want to minimize repulsion, so as a result, the lone pairs are adjacent from each other. Carbon dioxide has two electron groups and no lone pairs.

Carbon dioxide is therefore linear in electron-group geometry and in molecular geometry. The shape of CO 2 is linear because there are no lone pairs affecting the orientation of the molecule. Therefore, the linear orientation minimizes the repulsion forces.

The VSEPR theory not only applies to one central atom, but it applies to molecules with more than one central atom. We take in account the geometric distribution of the terminal atoms around each central atom. For the final description, we combine the separate description of each atom. In other words, we take long chain molecules and break it down into pieces.

Each piece will form a particular shape. Follow the example provided below:. Butane is C 4 H C-C-C-C is the simplified structural formula where the Hydrogens not shown are implied to have single bonds to Carbon. You can view a better structural formula of butane at en.

Let's start with the leftmost side. We see that C has three single bonds to 2 Hydrogens and one single bond to Carbon. That means that we have 4 electron groups. By checking the geometry of molecules chart above, we have a tetrahedral shape. Now, we move on to the next Carbon. This Carbon has 2 single bonds to 2 Carbons and 2 single bonds to 2 Hydrogens.

Again, we have 4 electron groups which result in a tetrahedral. Continuing this trend, we have another tetrahedral with single bonds attached to Hydrogen and Carbon atoms.

As for the rightmost Carbon, we also have a tetrahedral where Carbon binds with one Carbon and 3 Hydrogens. Let me recap. We took a look at butane provided by the wonderful Wikipedia link. We, then, broke the molecule into parts. We did this by looking at a particular central atom. In this case, we have 4 central atoms, all Carbon.

By breaking the molecule into 4 parts each part looks at 1 of the 4 Carbons , we determine how many electron groups there are and find out the shapes. We aren't done, yet! We need to determine if there are any lone pairs because we only looked at bonds.

Remember that electron groups include lone pairs! Butane doesn't have any lone pairs. Hence, we have 4 tetrahedrals. Now, what are we going to do with 4 tetrahedrals?

Well, we want to optimize the bond angle of each central atom attached to each other. This is due to the electrons that are shared are more likely to repel each other.

With 4 tetrahedrals, the shape of the molecule looks like this: en. That means that if we look back at every individual tetrahedral, we match the central Carbon with the Carbon it's bonded to. Bond angles also contribute to the shape of a molecule. Bond angles are the angles between adjacent lines representing bonds. The bond angle can help differentiate between linear, trigonal planar, tetraheral, trigonal-bipyramidal, and octahedral.