Brad Gietman's List: 2nd 21Sept

• No chemical bond can be 100 percent ionic, and, except for those between identical atoms, 100 percent covalent bonds do not exist either
• Electron density in covalent bonds shifts toward the more electronegative atom, producing partial charges on each atom and hence a dipole. In a polyatomic molecule
• In the case of symmetric molecules, the effects of individual bond dipoles cancel and a nonpolar molecule results.
• In addition to deficiency of electrons and expansion of the valence shell, the octet rule is violated by species which have one or more unpaired electrons
• Rearranging electrons (but not atomic nuclei) results in several structures which are referred to collectively as a resonance hybrid. Like an sp hybrid, a resonance hybrid is a combination of the contributing structures and has properties intermediate between them.

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• Instead the electron cloud is distorted by the attraction of the Li⁺ ion so that some of the H^– 1s² electron density is pulled into the bonding region between the Li and H nuclei. This contributes partial covalent character to the bond.
• Distortion of an electron cloud, as described in the previous paragraph, is called polarization.
• The tendency of an electron cloud to be distorted from its normal shape is referred to as its polarizability

• An ion pair like LiH has a negative end (H^–) and a positive end (Li⁺). That is, it has two electrical “poles,” like the north and south magnetic poles of a magnet. The ion pair is therefore an electrical dipole (literally “two poles“), and a quantity known as its dipole moment may be determined from experimental measurements. The dipole moment μ is proportional to the size of the separated electrical charges Q and to the distance r between them:

  
  

  μ = Qr       (1)
• If we increase the degree of polarization of an ionic bond, a bond which is more covalent than ionic is eventually obtained.

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• The Greek letter δ (delta) is used here to indicate that electron transfer is not complete and that some sharing takes place.
• The dipole moment of LiH shows that in effect only 77 percent of a full electronic charge has been transferred to H, and so δ = 0.77. If the transfer had been complete, δ would have been 1.0. Because the Li—H bond is only partially negative at the one end and partially positive at the other, we often say that the bond is polar or polar covalent, rather than 100 percent ionic

• The ability of an atom in a molecule to attract a shared electron pair to itself, forming a polar covalent bond, is called its electronegativity. The negative side of a polar covalent bond corresponds to the more electronegative element. Furthermore the more polar a bond, the larger the difference in electronegativity of the two atoms forming it.
• As can be seen from this table, elements with electronegativities of 2.5 or more are all nonmetals in the top right-hand comer of the periodic table.
• By contrast, elements with negativities of l.3 or less are all metals on the lower left of the table
• These elements have been coded in dark gray. They are often referred to as the most electropositive elements, and they are the metals which invariably form binary ionic compounds.
• In most of the typical ionic compounds discussed in the previous chapter, the difference is greater than l.5, although it is dangerous to attach too much significance to this figure since electronegativity is only a semiquantitative concept.
• An important example of an almost completely covalent bond between two different atoms is that between carbon (2.5) and hydrogen (2.l).

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• hese laws were originally based on the movement or transfer (dynamics) of heat (thermo), and the law of conservation of energy is therefore referred to as the first law of thermodynamics.
• We assign the symbol ΔH and the name enthalpy change to the quantity of heat absorbed by a chemical or physical change under conditions of constant pressure. Y
• The first law of thermodynamics (the law of conservation of energy) states that when heat energy is supplied to a substance, that energy cannot disappear-it must still be present in the atoms or molecules of the substance. Some of the added energy makes the atoms or molecules move faster. This is called translational energy. In the case of molecules, which can rotate and vibrate, some of the added energy increases the rotational and vibrational energies of the molecules. Finally, any atom or molecule will have a certain electronic energy which depends on how close its electron clouds are to positively charged nuclei.
• The total of translational, rotational, vibrational, and electronic energies is the internal energy of an atom or molecule.
• When chemical reactions occur, the internal energy of the products is usually different from that of the reactants, and the difference appears as heat energy in the surroundings. If the reaction is carried out in a closed container ( bomb calorimeter, for example), the increase in internal energy of the atoms and molecules is exactly equal to the heat energy absorbed from the surroundings
• When a chemical reaction occurs at constant pressure, as in a coffee-cup calorimeter, there is a change in potential energy of the atmosphere (given by P ΔV) as well as a change in heat energy of the surroundings. Because the heat energy absorbed can be measured more easily than P ΔV, it is convenient to define the enthalpy as the internal energy plus the increased potential energy of the atmosphere. Thus the enthalpy increase equals the heat absorbed at constant pressure.
• Enthalpy changes for a variety of reactions may be calculated from standard enthalpies of formation.
• They may also be estimated by summing the bond enthalpies of all bonds broken and subtracting the bond enthalpies of all bonds formed

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• In other sections we point out that when a chemical bond forms, negative charges move closer to positive charges than before, and so there is a lowering of the energy of the molecule relative to the atoms from which it was made. This means that energy is required to break a molecule into its constituent atoms. The bond enthalpy D_X–Y of a diatomic molecule X—Y is the enthalpy change for the (usually hypothetical) process:

  
  XY(g) → X(g) + Y(g)      ΔH°(298 K) = D_X―Y      (1)
• . Because of this fact, we must expect to obtain only approximate results, accurate only to about ± 50 kJ mol^–1, from the use of bond enthalpies.

• f less energy is needed to break up the reactant molecules into their constituent atoms than is released when these atoms are reconstituted into product molecules, then the reaction will be exothermic
• In summary, there are two factors which determine whether a gaseous reaction will be exothermic or not: (1) the relative strengths of the bonds as measured by the bond enthalpies, and (2) the relative number of bonds broken and formed. An exothermic reaction corresponds to the formation of more bonds, stronger bonds, or both.
• A second generalization is that the strengths of bonds usually increase with polarity. The bond enthalpies of the hydrogen halides, for instance, increase in the order HI < HBr < HCl < HF, and a similar order can be noted for bonds between carbon and halogens.
• A third factor affecting the strength of bonds is the size of the atoms. For the most part smaller atoms form stronger bonds.

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• : (1) Some stable molecules simply do not have enough electrons to achieve octets around all atoms. This usually occurs in compounds containing Be or B. (2) Elements in the third period and below can accommodate more than an octet of electrons.
• An atom like phosphorus or sulfur which has more than an octet is said to have expanded its valence shell.
• For example, in the case of phosphorus, the valence shell has a principal quantum number n = 3. An octet would be 3s²3p⁶. However, the 3d subshell is also available, and some of the 3d orbitals may also be involved in bonding. This permits the extra pair of electrons to occupy the valence (n = 3) shell of phosphorus in PF₅.

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• The majority of molecules discussed in general chemistry courses are demonstrated to have pairs of electrons. However, there are a few stable molecules which contain an odd number of electrons
• Free radicals are usually more reactive than the average molecule in which all electrons are paired. In particular they tend to combine with other molecules so that their unpaired electron finds a partner of opposite spin. Since most molecules have all electrons paired, such reactions usually produce a new free radical.

• For these molecules it is possible to draw more than one Lewis structure which obeys the octet rule but which is unsatisfactory in other ways. A simple example of such a molecule is ozone, an unusual form of oxygen, whose molecular formula is O₃.
• In other words the structure of O₂ is somehow intermediate in character between the two structures shown.
• On a mathematical level we can satisfactorily account for the properties of ozone by regarding its structure as a hybrid of the two structures shown above, the term hybrid having exactly the same sense as for hybrids. We then obtain an electron probability distribution in which both bonds receive equal treatment and are intermediate in character between double and single bonds. Such a structure is called a resonance hybrid and is indicated in one of the following ways:
• This hexagonal ring of six carbon  atoms is called a benzene ring.
• Any bonds which are intermediate in character can be treated as though they were single bonds, though perhaps a bit fatte

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• the electron-pair bonds are arranged so that they avoid each other in space to the maximum possible extent. This may be understood in terms of the repulsion between electron clouds due to their like charges
• BeH₂ central Be atom has only two electron pairs in its valence shell. These are arranged on opposite sides of the Be atom in a straight line, and they bond the two atoms to the Be atom. Thus the three nuclei are all in a straight line, and the H―Be-H angle is 180°

• A molecule such as BeH₂, whose atoms all lie on the same straight line, is said to be linear. In BCl₃ the three valence electron pairs, and hence the three Cl nuclei, are arranged in an equilateral triangle around the B atom. Each Cl―B―Cl angle is 120° and all four nuclei (B included) lie in the same plane. The three Cl atoms are said to be trigonally arranged around B.

  In CH₄ the four H nuclei are at the four corners of a geometric figure called a tetrahedron. A tetrahedron has six equal edges, four equilateral triangular faces, and four identical corners (apices). The C nucleus lies in the exact center of the tetrahedron, equidistant from each corner. All the H—C—H angles are the same, namely, 109.5°. This important angle is called the tetrahedral angle. The four H atoms are said to be tetrahedrally arranged around the C atom. This tetrahedral arrangement is the most important of those described in Table 1.In PF₅ the five F nuclei are arranged at the corners of a trigonal bi-pyramid. As drawn in the figure, one F atom lies directly above the P atom and one directly below. The remaining three F atoms are arranged in a triangle around the middle of the P. Some of the F―P―F angles are 90°, while others are 120°.

  In SF₆ the six F atoms are arranged at the corners of an octahedron. An octahedron has twelve edges, eight equilateral triangular faces, and six identical corners. The name octahedron is derived from the eight faces, but it is usually the six corners of this figure which are of interest to chemists. Thus you will have to remember that an octahedral arrangement involves six atoms, not the eight that the name seems to imply.In SF₆ the six F atoms are octahedrally arranged around the S. All the F―S―F angles are 90°. Octahedral arrangements are quite common in chemistry. In crystals of LiH and NaCl, for instance, six anions are arranged octahedrally around each cation while six cations are arranged octahedrally around each anion.

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• because a lone pair of electrons is always “fatter” than a bonding pair.

• In the sections on chemical bonding we showed that a covalent bond results from an overlap of atomic orbitals—usually one orbital from each of two bonded atoms. Maximum bond strength is achieved when maximum overlap occurs
• . Each of the two orbitals whose electron densities are shown in Fig. 1b is called an sp hybrid. The word hybrid indicates that each orbital is derived from two or more of the atomic orbitals discussed in sections on the Electronic Structure of Atoms, and the designation sp indicates that a single s and a single p orbital contributed to each sp hybrid.
• three electron pairs in a trigonal arrangement. If we place these three electron pairs in the valence shell of boron, one 2s and two 2p orbitals (2p_x and 2p_x, for example) will be filled, giving the electron configuration 2s²2_x²2_y².
• All three electron clouds are identical in shape, and they are oriented at angles of 120° with respect to each other. Again we have a set of hybrid orbitals, but this time they are derived from one s orbital and two p orbitals. Accordingly these orbitals are known as sp² hybrids
• In addition to the two hybrids just considered, a third combination of s and p orbitals, called sp³ hybrids, is possible. As the name suggests, sp³ hybrids are obtained by combining an s orbital with three p orbitals (p_x, p_y and p_z).

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• In a double bond, two electron pairs are shared between a pair of atomic nuclei
• three electron pairs in a triple bond behave as an “extra-fat” bond.

• A reaction is said to be unimolecular if, on the microscopic level, rearrangement of the structure of a single molecule produces the appropriate product molecules.
• The minimum quantity of energy required to surmount an energy barrier during a chemical reaction is called the activation energy, and the molecular species at the top of the barrier is called the activated complex or the transition state
• Quantities associated with this activated complex are usually denoted by a double dagger (‡)

• What we have just said applies to any unimolecular process. The reaction rate must always be directly proportional to the concentration of the reacting species. That is, for a general unimolecular process, A → products, the rate equation must be first order in A:

  
  

  Rate = kc_A

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