Chapter 5

helps determine activity Ex: morphine (molecular imposter) shaped like endorphine Analogy: key fits in a lock (plug in socket)

why does shape matter?

determining molecular shape and polarity

• Draw the Lewis structure for the molecule and determine its molecular geometry. • Determine if the molecule contains polar bonds. A bond is polar if the two bonding atoms have sufficiently different electronegativities (see Figure 5.4). If the molecule contains polar bonds, superimpose a vector, pointing toward the more electronegative atom, on each bond. Make the length of the vector proportional to the electronegativity difference between the bonding atoms. • Determine if the polar bonds add together to form a net dipole moment. Sum the vectors corre- sponding to the polar bonds together. If the vectors sum to zero, the molecule is nonpolar. If the vectors sum to a net vector, the molecule is polar.

electrostatic map

red= electron rich blue- electron poor

4 EGs w/ lone pairs

-Ammonia has The central nitrogen atom has four electron groups (one lone pair and three bonding pairs) that repel each other. -If we do not distinguish between bonding electron groups and lone pairs, we find that the electron geometry—the geometrical arrangement of the electron groups—is still tetrahedral, as we expect for four electron groups -However the molecular geometry—the geometrical arrangement of the atoms—is TRIGONAL PYRAMIDAL -although the electronic and molecular geometry are different, the electron geometry is relevant to the molecular geometry. The lone pairs exerts its influence on the bonding pairs. -Lone pair electrons generally exert slightly greater repulsions than bonding electrons. -if all 4 electron groups in NH3 exerted equal repulsions on one another, the bond angles in the molecule would all be the ideal tetrahedral angle, 109.5°. However, the actual angle between N¬H bonds in ammonia is slightly smaller, 107°. -A lone electron pair occupies more angular space than a bonding pair. exerting a greater repulsive force on neighboring electrons and compressing the N-H bond angles -MG=Trigonal pyramidal, EG= Tetrahedral

formal charge

-An electron bookkeeping system that allows us to discriminate between alternative Lewis structures. -The formal charge of an atom in a Lewis structure is the charge it would have if all bonding electrons were shared equally between the bonded atoms -formal charge is the calculated charge for an atom in a molecule if we completely ignore the effects of electronegativity. Ex: we know HF are dipole but if we ignore EN the formal charge is 0 for both -formal charge= # of VEs- (# of nonbonding electrons- half # shared electrons) -VEs (no bond+no unshared electrons), must add up to overall avg. -Best LS has the FC closer to 0 Rules that apply: 1. The sum of all formal charges in a neutral molecule must be zero. 2. The sum of all formal charges in an ion must equal the charge of the ion 3. Small (or zero) formal charges on individual atoms are better than large ones. 4. When formal charge cannot be avoided, negative formal charge should reside on the most electro- negative atom. -the best skeletal structure usually has the least electronegative atom in the central position and the negative formal charge on the more electronegative atom

incomplete octets

-Another significant exception to the octet rule involves those elements that tend to form incomplete octets. -The most important of these is boron, which forms compounds with only six electrons, rather than eight. -For example, BF3 and BH3 lack an octet for B -In BF3 we could add a double bond but this would create a + FC on F so its better to just keep it incomplete octet as the better LS

increase in repulsion

Electron pair repel each other, but some electron group want more spare than others (bonds need less space than lone pairs) Repulsion: this helps us decide bond angles B-B< B-LP < LP-LP

expanded octets

-Elements in the third row of the periodic table and beyond often exhibit expanded octets of up to 12 (and occasionally 14) electrons. -Examples: arsenic pentafluoride and sulfur hexafluoride -only central atom has expanded octet either 10 or 12 -AsF5 arsenic has an expanded octet of 10 electrons, and in SF6 sulfur has an expanded octet of 12 elec- trons. Both of these compounds exist and are stable. -d orbitals in these elements are energetically accessible (they are not much higher in energy than the orbitals occupied by the valence electrons) and can accommodate the extra electrons. period 3 -Expanded octets never occur in second-period elements because they do not have energetically accessible d orbitals and therefore never exhibit ex- panded octets. -However, we should never expand the octets of second-row elements.

polarity in diatomic molecules

-If a diatomic molecule has a polar bond, the molecule as a whole is polar. -an electrostatic potential map of HCl. In these maps, yellow/red areas indicate electron-rich regions in the molecule and the blue areas indicate electron-poor regions. -Yellow indicates moderate electron density. -If the bond in a diatomic molecule is nonpolar, the molecule as a whole will be nonpolar.

nonpolar

-If two atoms with identical electronegativities form a covalent bond, they share the electrons equally, and the bond is purely covalent or nonpolar. (a.k.a)

larger molecules

-Larger molecules may have two or more interior atoms. When predicting the shapes of these molecules, we apply the principles we just covered to each interior atom.

odd electron species

-Molecules and ions with an odd number of electrons in their Lewis structures are FREE RADICALS (or simply radicals) -Example, nitrogen monoxide cant achieve octets for both atoms yer it still exists -For free radicals, such as NO and NO2, we simply write the best Lewis structure that we can. -unusual

5 EGs w/ lone pairs

-SF4. The central sulfur atom has five electron groups (one lone pair and four bonding pairs). The electron geometry, due to the five electron groups, is trigonal bipyramidal. -the lone pair could occupy either an equatorial position or an axial position -the lone pair occupies the position that minimizes its interaction with the bonding pairs. -If the lone pair were in an axial position, it would have three 90° interactions with bonding pairs. In an equatorial position, however, the lone pair has only two 90° interactions. Consequently, the lone pair occupies an equatorial position. The resulting molecular geometry is called seesaw because it resembles a SEESAW (or teeter-totter/ irregular tetrahedron). -axial lone pair: does not occur -equatorial: lone pair -molecular geometry: seesaw, EG= TBP

6 EGs w/ lone pairs

-The Lewis structure of BrF5 is shown here. The central bromine atom has six electron groups (one lone pair and five bonding pairs) -EG= octahedral, MG= square pyramidal -Since all six positions in the octahedral geometry are equivalent, the lone pair can be situated in any one of these positions. The resulting molecular geometry is SQUARE PYRAMIDAL.

electronegativity

-The ability of an atom to attract electrons to itself in a chemical bond (which results in polar and ionic bonds) is electronegativity. -The ability of an atom to attract electrons density to itself when bonded to another atom (different element) -Increasing from left to right -increasing from bottom to top -most EN is at the top right corner of PT -least EN is at the bottom left corner of PT -electronegativity is inversely related to atomic size—the larger the atom, the less ability it has to attract electrons to itself in a chemical bond.

bond length

-The average bond length represents the average length of a bond between two particular atoms in a large number of compounds. -bond lengths depend not only on the kind of atoms involved in the bond, but also on the type of bond: single, double, or triple. -triple bonds are shorter than double bonds, which are in turn shorter than single bonds. -as the bond get longer it also become weaker= single=not a smooth trend really -single, double, triple= distance is shorter as we move down the like in amperes

bond energy

-The bond energy of a chemical bond is the energy required to break 1 mole of the bond in the gas phase. -bond energies are + because energy must be put into a molecule to break a bond (the process is endothermic. absorbs heat and carries a + sign) -the more energy it take to break a bond the more stronger it is -stronger bonds then to be more chemically stable, and less chemically reactive than compounds with weaker bonds. -In general, for a given pair of atoms, triple bonds are stronger than double bonds, which are, in turn, stronger than single bonds. -single, double, triple= increase down the line in bond energy/ bond strength

VSEPR theory

-The geometry of a molecule is determined by the number of electron groups on the central atom (or on each interior atom, if there is more than one). -The number of electron groups is determined from the Lewis structure of the molecule. If the Lewis structure contains resonance structures, we can use any one of the resonance structures to determine the number of electron groups. -Each of the following counts as a single electron group: a lone pair, a single bond, a double bond, a triple bond, or a single electron. -The geometry of the electron groups is determined by their repulsions as summarized in Table 5.5 on the next page. In general, electron group repulsions vary in relative ordering of repulsions as follows: (Lone pair ¬ lone pair 7 Lone pair ¬ bonding pair 7 Bonding pair ¬ bonding pair) -Bond angles can vary from the idealized angles because double and triple bonds occupy more space than single bonds (they are bulkier even though they are shorter), and lone pairs occupy more space than bonding groups. The presence of lone pairs usually makes bond angles smaller than the ideal angle for the particular geometry.

percent ionic character

-The percent ionic character is the ratio of a bond's actual dipole moment to the dipole moment it would have if the electron were completely transferred from one atom to the other, multiplied by 100%. -the percent ionic character generally increases as the electronegativity difference increases -However, as we can see, no bond is 100% ionic. In general, bonds with greater than 50% ionic character are referred to as ionic bonds.

one dimension

-To add two vectors that lie on the same line, assign one direction as positive. Vectors pointing in that direction have positive magnitudes. Vectors pointing in the opposite direction have negative magnitudes. Sum the vectors (always remembering to include their signs) -+5,+5= +10 --5,+10= +5 --5, +5= 0

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. Draw the resultant vector beginning at the origin and extending to the far corner of the parallelogram -To add three or more vectors, add two of them together first, and then add the third vector to the result

dipole moment

-We quantify the polarity of a bond by the size of its dipole moment -A dipole moment (m) occurs anytime there is a sepa- ration of positive and negative charge -(mu)= qr -q=charge, r=distance b/n charged particles -The debye (D) is the unit commonly used for reporting dipole moments (1 D = 3.34 * 10^30 C*m). -The smaller the magnitude of the charge separation, and the smaller the distance between the charges, the smaller the dipole moment.

5 EGs w/ lone pairs

-When three of the five electron groups around the central atom are lone pairs, as in XeF2, the lone pairs occupy all three of the equatorial positions and the resulting molecular geometry is LINEAR. -EG=TBP, MG=linear

5 EGs w/ lone pairs

-When two of the five electron groups around the central atom are lone pairs, as in BrF3, the lone pairs occupy two of the three equatorial positions—again minimizing 90° interactions with bonding pairs and also avoiding a lone pair-lone pair 90° repulsion -The resulting molecular geometry is T-SHAPED. -EG= trigonal bipyramidal, and MG= T Shaped

6 EGs w/ lone pairs

-When two of the six electron groups around the central atom are lone pairs, as in XeF4, the lone pairs occupy positions across from one another (to minimize lone pair-lone pair repulsions). The resulting molecular geometry is SQUARE PLANAR. -EG=octahedral. MG= square planar

VSEPR theory

-electron groups— which we define as lone pairs, single bonds, multiple bonds, and even single electrons—repel one another through coulombic forces -VSEPR theory focuses on the repulsions. -the repulsions between electron groups on interior atoms (or the central atom) of a molecule determine the geometry of the molecule -the preferred geometry is one that has the maximum separation and therefore the minimum energy possible. -molecular geometry depends on (a) the number of electron groups around the central atom and (b) how many of those electron groups are bonding groups and how many are lone pairs. -Valence shell electron pair repulsion (theory that explain molecular structure) -electron pains want to be as far away from each other as possible Using VSEPR theory: 1. Do the lewis dot structure 2. Count EG around the central atom 3. Decide on the Electronic geometry (one of 5) 4. Find the molecular geometry (shape of the bonded atoms ONLY) (ignore lone pairs) 5. Adjust the bond angles for different amounts of repulsion (if needed!)

exceptions to the octet rule

-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 8 electrons around an atom 3) expanded octets, molecules or ions with more than 8 electrons around an atom

ionic bond

-in an ionic bond the electron is essentially transferred from one atom to another. -If there is a large electronegativity difference between the two atoms in a bond, such as normally occurs between a metal and a nonmetal, the electron from the metal is almost completely transferred to the nonmetal, and the bond is ionic -Large (2.0+), NaCl -Electrons transferred -the more EN element wins the tug of war for electron density and steals an electron (s) electron transfer -cation/anion -solid at RT when melted they conducted electricity

polar covalent bond

-is intermediate in nature b/n a pure covalent bond and an ionic bond. -most covalent bonds b.n dissimilar atoms are actually polar covalent, somewhere b/n the two extremes -If there is an intermediate electronegativity difference between the two atoms, such as between two different nonmetals, then the bond is polar covalent. For example, HCl has a polar covalent bond. -Intermediate (0.4-2.0) -Electrons shared unequally -the more EN element hogs the electron density -various properties, is in between ionic and nonpolar bonds

visualization

-just look around central atom -0 LP= nonpolar, unless different outer atoms (like BF2cl) -1 LP= always polar -2 LP= It depends. Are they symmetrically disturbed? if yes then NP, if no then P -3 LP= NP

lewis structure for polyatomic ions

-pay special attention to the charge of the ion when calculating the number of electrons for the Lewis structure. -Add one electron for each negative charge and subtract one electron for each positive charge. -the Lewis structure for a polyatomic ion within brackets with the charge of the ion in the upper right-hand corner, outside the bracket.

endorphins

-short for endogenous morphine -bodies natural painkillers -Morphine binds to opioid receptors because it fits into a special pocket (called the active site) on the opioid receptor protein (just as a key fits into a lock) that normally binds endorphins.

lewis structure

-show dot-dash representation -show all VEs and S=N-A if the molecule obeys octet rule -s and p block elements like to have 8 VEs -N=needed electrons to complete valance shell with 8 electron unless H which is only 2 -A=avaiable electrons. add # of VEs of each atom. in ion add or subtract for the charge -S= N-A= shared electrons. this is how many electrons MUST be shared to complete the outer shells -lone pairs= (A-S)/2 -oxygen atoms do not bond to each other except O2, O3, O2-, and O2-2 -ternery acids (acids with H,O and one other element) the H will bond to an oxygen -if there are two or more choices: 1) nature like symmetry. and 2) look at formal charge -1 bond=2 shared electrons -C has 4 bonds, O has 2 bonds, Halogens (grp. 17) like 1 bond

4 EGs w/ lone pairs

-similar effect in water. The Lewis structure of water has two bonding pairs and two lone pairs -Since it has four electron groups, water's electron geometry is tetrahedral (like that of ammonia), but its molecular geometry is BENT. -the bond angles in H2O are smaller (104.5°) than the ideal tetrahedral bond angles because of the greater repulsion exerted by the lone pair electrons. -H2O has two lone pairs of electrons on the central oxygen atom so it is even smaller than that in NH3 -MG=Bent, EG= tetrahedral

geometries on paper

-straight line= bond in plane of paper -hatched wedge (multiple lines)= Bond going into the page -solid wedge (black solid)= bud coming out of the page

resonance

-the concept resonance, used when two or more valid lewis structures can be down fro the same compound -when we can write two or more valid structures for the same molecule in nature the molecule is an average of the the 2 LS. -put a double headed arrow b/n them -RESONANCE STRUCTURE is one of the two or more lewis structures that have the same skeletal formula (the atoms are in the same locations) but different electron arrangements -the actual structure of the molecule is intermediate b/n the 2 or more resonance structures and is called a RESONANCE HYBRID. this is the only structure that actually exists. NOT a lewis rep. bonds are (1 and 1/3) for CO3-2 -in LS electrons are localized= either on one atom (lone pair) or b/n atoms (bonding pari) -in nature delocalized= electrons in molecules are often delocalized over several atoms or bonds. this lowers energy and stableizes them

polarity in polyatomic molecules

-the presence of polar bonds may or may not result in a polar molecule, depending on the molecular geometry -if the molecular geometry is such that the dipole moments of individual polar bonds sum together to a net dipole moment, then the molecule is polar. -if the molecular geom- etry is such that the dipole moments of the individual polar bonds cancel each other (that is, sum to zero), then the molecule is nonpolar -Dipole moments cancel each other because they are vector quantities; they have both a magnitude and a direction. -water is a polar molecule (bent), CO2 is nonpolar (linear)

pauling

-was quantified by Linus Pauling, he used a heteronuclear diatomic molecule, HF. Bond energies was more than expected so he suggested it was due to the ionic character of the bond -F=EN of 4.0 (the most EN element on the PT)

vector addition

-we can determine whether a molecule is polar by summing the vectors associated with the dipole moments of all the polar bonds in the molecule -If the vectors sum to zero, the mol- ecule will be nonpolar. -If they sum to a net vector, the molecule will be polar. Here, we demonstrate how to add vectors together in one dimension and in two or more dimensions. -Oil and water do not mix because water molecules are polar and the molecules that compose oil are nonpolar. -A mixture of polar and nonpolar molecules is similar to a mixture of small mag- netic particles and nonmagnetic ones. The magnetic particles (which are like polar molecules) clump together, excluding the nonmagnetic particles (which are like nonpolar molecules) and separating into distinct regions. -Opposite partial charges on molecules attract one another. -if the arrows are going away from one another they cancel and are non polar (no net dipole) -if one arrow is pointing one way and the other is not there then it is polar (dipole)

electron groups

-we define as lone pairs, single bonds, multiple bonds, and even single electrons—repel one another through coulombic forces -electron groups are regions of electron density RED Theres two types... 1) lone pairs (unshared electron pairs, unbounded pairs) 2) bonds (single, double, or triple) (all count as 1 Electron group)

lewis structure for molecular compounds

1. Write the correct skeletal structure for the molecule -First, hydrogen atoms are always terminal. hydrogen, which has only a single valence electron to share and requires only a duet, can form just one bond and central atom needs at least 2 -Second, put the more electronegative elements in terminal positions and the less electronegative elements (other than hydrogen) in the central position. 2. Calculate the total number of electrons for the Lewis structure by summing the valence electrons of each atom in the molecule. -valence electrons for any main-group element is equal to its group number in the periodic table. -for a polyatomic ion, you must consider the charge of the ion when calculating the total number of electrons. - +1 e for each - charge, -1 e for each + charge 3. Distribute the electrons among the atoms, giving octets (or duets in the case of hydrogen) to as many atoms as possible. -Begin by placing two electrons between every two atoms. this represents the min. # of boding e's. -distribute the remaining electrons as lone pairs, first to terminal atoms and then to the central atom, giving octets (or duets for hydrogen) to as many atoms as possible. 4. If any atoms lack an octet, form double or triple bonds as necessary to give them octets -by moving lone electron pairs from terminal atoms into the bonding region with the central atom. 5. Check.

EG repulsions

In general, electron group repulsions vary as follows: (lone pair-lone pair) > (Lone pair-bonding pair) > (Bonding pair-bonding pair) Most repulsive to Least repulsive -bond angles get smaller as there is an increase in lone pairs

adding dipole moments

Linear= The dipole moments of two identical polar bonds pointing in opposite directions will cancel. The molecule is nonpolar. Bent= The dipole moments of two polar bonds with an angle of less than 180° between them will not cancel. The resulting dipole moment vector is shown in red. The molecule is polar. Trigonal Planar= The dipole moments of three identical polar bonds at 120° from each other will cancel. The molecule is nonpolar. Tetrahedral= The dipole moments of four identical polar bonds in a tetrahedral arrangement (109.5° from each other) will cancel. The molecule is nonpolar. Trigonal pyramidal= The dipole moments of three polar bonds in a trigonal pyramidal arrangement (109.5° from each other) will not cancel. The resulting dipole moment vector is shown in red. The molecule is polar.

morphine

a drug, usually a painkiller. -Morphine acts by binding to receptors (called opioid receptors) that exist within nerve cells. When morphine binds to an opioid receptor, the transmission of nerve signals is altered, resulting in less pain -morphine is a molecular impostor, mimicking the action of endorphins because of similarities in shape. -theres shape specific interactions b/n molecules and proteins

trigonal bipyramidal geometry

five EGs: -Five electron groups around a central atom assume a trigonal bipyramidal geometry, like five balloons tied together. -three of the groups lie in a single plane, as in the trigonal planar configuration, while the other two are positioned above and below this plane -The angles between the equatorial positions (the three bonds in the trigonal plane) are 120°, while the angle between the axial positions (the two bonds on either side of the trigonal plane) and the trigonal plane is 90°. -an example: PCl5 The three equatorial chlorine atoms are separated by 120° bond angles, and the two axial chlorine atoms are separated from the equatorial atoms by 90° bond angles. (ax=axial)(eq=equitorial) -angles include: 90, 180, 120

main group elements

for main group elements motile the following trend in EN -Electronegativity generally increases across a period in the periodic table. -Electronegativity generally decreases down a column in the periodic table. -Fluorine is the most electronegative element. -Francium is the least electronegative element (sometimes called the most electropositive)

tetrahedral geometry

four EGs: -geometries are three-dimensional -analogy of balloons tied together, In this analogy, each electron group around a central atom is like a balloon tied to a central point. the ballots represent EGs not atoms -tie four balloons together, however, they assume a three-dimensional tetrahedral geometry with 109.5° angles between the balloons. -balloons point toward the vertices of a tetrahedron—a geometrical shape with four identical faces, each an equilateral triangle, as shown here -Methane is an example of a molecule with four electron groups around the central atom -the tetrahedron is the three-dimensional shape that allows the maximum separation among the groups. -name of the electronic geometry is tetrahedral

Lewis dot

limitations: -One limitation of representing electrons as dots, and covalent bonds as two dots shared between two atoms, is that the shared electrons always ap- pear to be equally shared. Such is not the case. H has a + charge and F has a - charge . d+ (delta plus), d- (delta minus). The deltas sometimes triangles mean partial charge (fractional) -also not an ionic bond -HF is unequally shared. electron density is greater on F than on H. This is a polar bond. -When the +--> this is a dipole.

endogenous

mens produced within the organism -proteins are among the most important biological molecules serve many function in living organisms

covalent bond

shared electrons -Small (0-0.4), Cl2 -Pure (nonpolar) covalent bond -Electrons shared equally -The bond between C and H lies at the border between covalent and polar covalent; however, this bond—which is very important in organic chemistry—is normally considered covalent (nonpolar). -gases or liquids at RT -nonconductors

octahedral geometry

six EGs: -Six electron groups around a central atom assume an octahedral geometry, like six balloons tied together. -four of the groups lie in a single plane, with a fifth group above the plane and another below it. -The angles in this geometry are all 90°. As an example of a molecule with six electron groups around the central atom, consider SF6 The structure of this molecule is highly symmetrical; all six bonds are equivalent -angles include 90, 180

trigonal planar geometry

three EGs: -has three electron groups around the central atom -These three electron groups maximize their separation by assuming 120° bond angles in a plane— a trigonal planar geometry -formaldehyde: HCO bond angles are 121.9° and that the HCH bond angle is 116.2°. The HCO bond angles are slightly greater than the HCH bond angle because the double bond contains more electron density than the single bond and therefore exerts a slightly greater repulsion on the single bonds. -different types of electron groups exert slightly different repulsions—the resulting bond angles reflect these differences. -name of electronic geometry is trigonal planar -two- dimensional

linear geometry

two EGs: -beryllium often forms incomplete octets, as it does in this structure. (BeCl2). This has 2 EGs (two single bonds) -the repulsion between these two electron groups, which maximize their separation by assuming a 180° bond 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. -same geometry is observed in all molecules that have two electron groups (and no lone pairs) -a double bond counts as one EG shape: Eg-A-EG, where A=the central atom, name of the electronic geometry is linear -two- dimensional