Chem summary week 6 - Chemistry: Structure and Properties PDF

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Summary of chapter 6 with images to explain how to work practice problems. Focus of chapter is on formal charge....

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Chem Summary Week 6 5.4 Resonance and Formal Charge Resonance structure is when one of two or more Lewis structures that have the same skeletal formula but different electron arrangements. The atoms stay in the same location and the only thing that changes is the placement of the electrons. The actual structure of the molecule is intermediate between the two or more resonance structures and is called a residence hybrid. In the Lewis model, electrons are localized either on one atom or between atoms. However, in nature, electrons and molecules are often delocalized over several atoms or bonds. The resulting stabilization of the elements is sometimes called resonance stabilization. It makes an important contribution to the stability of many molecules. Example 5.5 writing resonance structures Write the Lewis structure for the NO3- ion. Include resonance structures.

Formal charge is the fictitious charge assigned to each atom in a Lewis structure that helps us to distinguish among competing structures. The formal charge of an atom in a Louis structure is the charge it would have if all bonding electrons were equally shared between the bonded atoms. The ideal formal charge is 0. The formula is Formal charge = number of valence electrons - (number of nonbonding electrons + 1/2 number of bonding electrons) 4 rules apply 1. 2. 3. 4.

The sum of all the formal charges in a neutral molecule must be zero The sum of all formal charges an ion must equal the charge of the ion Small or zero formal charges on individual atoms are better than large ones when the formal charge cannot be avoided, negative formal charge should reside on the most electronegative atom

Example 5.6 assigning formal charges Assign formal charges to each atom in the residence forms of the cyanate ion (OCN-). Which resonance form is likely to contribute most to the correct structure of OCN-?

For Practice 5.6 assign formal charges to each atom in the resonance forms of N2O. Which resonance form is likely to contribute most to the correct structure of N2O? Example 5.7 drawing resonance structures and assigning formal charge for organic compounds. Draw the Lewis structure (including resonance structures) for nitromethane (CH3NO2). For each resonance structure, assign formal charges to all atoms that have formal charge.

5.5 exceptions to the octet rule: Odd-Electron Species, Incomplete Octets, and Expanded Octets Exceptions to the octet rule include: (1) odd-electron species, molecules or ions with an odd number of electrons; (2) incomplete Octets, molecules or ions with fewer than eight electrons around an atom; and (3) expanded Octets, molecules or ions with more than eight electrons around an atom Molecules and ions with an odd number of electrons in their Lewis structure are free radicals. example : nitrogen monoxide has only 11 electrons so the top of nitrogen only gets one dot, not a pair 5.3 Conceptual Connection Which molecule would you expect to be a free radical? ClO Incomplete octets do not have enough electrons to satisfy the octet rule. Example : BH3. There are not enough electrons for boron to have 8 and there are no addition electrons to move into the bonding region. For BF3, the Lewis structure has octets for all atoms but when we assign formal charges, we get a negative formal charge on B and a positive formal charge on F. This predicts that Bf3 might react in ways that would complete its octet. Elements in the third row of the periodic table and beyond often exhibit expanded octets of up to 12 or 14 electrons. In AsF5 arsenic has an explained octet of 10 electrons. This is because the d orbital in these elements are energetically accessible. Example 5.8 writing Lewis structures for compounds having expanded octets Write the Lewis structure for XeF2.

5.4 Conceptual Connection Which molecule could have an expanded octet? H3PO4 5.6 bond energies and bond length The bond energy of a chemical bond is the energy required to break 1 molecule of the bond in the gas phase. Bond energies are positive because energy must be put into a molecule to break a bond. The process is endothermic which absorbs heat and carries a positive sign. Cl2 (g) -> 2 Cl (g) bond energy = 243 kJ A single bond is weaker than a double bond is weaker than a triple bond. The average bond length represents the average length of a bond between two particular atoms in a large number of compounds. It takes more energy to the bond of a triple bond. A triple bond (the shortest) is the hardest one to break and the single bond (the longest) is the easiest. 5.7 VSEPR Theory: The Five Basic Shapes Valence shell electron pair repulsion (VSEPR) theory is based on the simple idea that electron groups- which we define as lone pairs, single bonds, multiple bonds, and even single electrons- repel one another through coulombic forces. The electron groups are also attracted to the nucleus but VSEPR theory focuses on the repulsions. According to the theory, the repulsions between electron groups on interiors atoms (or the central atom) of a molecule determine the geometry of the molecule. BeCl2 has two electron groups (two single bonds) about the central atom. According to the theory, the geometry of BeCl2 is determined by the repulsion between these two electron groups, which maximize their desperation by assuming a 180 degree

bond or angle or a linear geometry. Molecules that form only two single bonds, with no lone pairs, are rare because they do not follow the octet rule. According to the theory, double bonds repel each other resulting in linear geometry as well. When three electron groups maximize their separation by assuming 120 degree bond angles in a plane it is called a trigonal planar geometry. Most bonds are around the ideal 120 degrees but different types of electron groups exert slightly different repulsions. Conceptual Connection 5.5 In determining electron geometry, why do we consider only the electron groups on the central atom? Why don’t we consider the electron groups on terminal atoms? Molecular geometry is completely based on the terminal atoms and the electron groups on the terminal atoms don’t affect it at all For molecules with four or more electron groups around the central atom, the geometries are three dimensional and are more difficult to imagine and draw. A common way to visualize them is like a bunch of balloons tied together. The balloons represent electron groups. If you tie four balloons together, they assume a 3D tetrahedral geometry with 109.5 degree angles between the balloons. Conceptual Connection 5.6 What is the geometry of the HCN molecule? It’s Lewis structure is H-C triple bonded to N: linear Five electron groups around a central atom assumes a trigonal bipyramidal geometry, like 5 balloons tied together. The angles in the trigonal bipyramidal structure are not all the same. The angles between the equatorial positions are 120 degrees while the angle between the axial positions of the trigonal plane is 90 degrees. Six electron groups around a central atom assume an octahedral geometry, like 6 balloons tied together. Four groups lie in a single plane with a fifth groups above the plane and one below it. All angles are 90 degrees. Example 5.9 VSEPR Theory and the Basic Shapes Determine the molecules geometry of NO3-

5.8 VSEPR Theory: The Effect of Lone Pairs In the Lewis structure of ammonia, we find that the electron geometry- the geometrical arrangement of the electron groups- is still tetrahedral. However, the molecular geometry- the geometrical arrangement of the atoms- is trigonal pyramidal. The electron geometry is relevant to the molecular geometry. Since water has four electron groups, water’s electron geometry is tetrahedral but it’s molecular geometry is bent. In general, electron group repulsions vary as follows: lone pair- lone pair > Lone pair- bonding pair> bonding pair-bonding pair. The bond angle gets progressively smaller. In an equatorial position, the lone pair has only two 90 degree interactions. The lone pair occupies an equatorial positions and the resulting geometry is called seesaw because it resembles a seesaw. When two of the five electron groups around the central atom are lone pairs, the lone pairs occupy two of the three equatorial positions resulting in a T-shaped molecular geometry. When the electron geometry has six electron groups, the molecule geometry is square pyramidal. When two of the six groups around the central atom are lone pairs, the molecular geometry is square planar. Conceptual Connection 5.7 Suppose that a molecule with six electron groups were confined to two dimensions and therefore had a hexagonal planar electron geometry. If two of the six groups were lone pairs, where would they be located in the figure shown here? Since 1 and 4 have a larger distance, they minimize lone pair- lone pair repulsions. Conceptual Connection 5.8

What statement is always true according to VSEPR theory? D 5.9 VSEPR Theory: predicting molecule geometries Table 5.5 electron and molecular geometries Example 5.10 predicting molecular geometries Predict the geometry and bond angles of PCl3

1. Draw the Lewis structure for the molecule 2. Determine the total number of electrons groups around the central atom 3. Determine the number of bonding groups in the number of lone pairs around the central atom 4. Refer to Table 5.5 to determine the electron geometry and molecular geometry Example 5.11 Predict the geometry and bond angles of ICl4-

Representing molecular geometries on paper straight line - Bond in plane of paper Hatched wedge. Bond going into the page Solid wedge bond coming out of the page Larger molecules may have two or mole interior atoms. We predict the shapes by determining the number of electron groups, number of lone pairs, and molecular geometry. Example 5.12 predicting the shape of larger molecules Predict the geometry about each interior atom in methanol (CH3OH) and sketch the molecule

5.10 molecular shape and polarity If a diatomic molecule has a polar bond, the molecule as a whole is polar. In a polyatomic molecule, the presence of polar bonds may or may not result in a polar molecule, depending on the molecule geometry.

Steps of determining molecular shape and polarity 1. Draw the Lewis structure for the molecule and determine its molecular geometry 2. Determine if the molecule contains polar bonds 3. Determine if the polar bonds add together to form a nat dipole moment One dimension vectors Vectors pointing in the same direction have positive magnitudes. Vectors pointing in the opposite direction have negative magnitudes. Two or more dimensions To add two vectors, draw a parallelogram in which the two vectors form two adjacent sides. Draw the other two sides of the parallelogram parallel to and the same length as the two original vectors. To add three or more, add two of them together than add the third vector to the result. Example 5.13 Determining if a molecule is polar Determine if NH3 is polar....