The electrons donated from a Lewis base to a Lewis acid in a Lewis acid-base reaction are donated and accepted at particular atomic centers. For instance, the formation of an adduct between ammonia and BF 3 involves the donation of the lone pair on the ammonia nitrogen atom to the Lewis acid site on BF 3 . Because Lewis acid-base reactions involve electron donation and acceptance at particular sites, substituent groups which alter the electron density at a site through inductively donating or withdrawing electron density will affect the Lewis acid-base properties of that site. For instance, the BF 3 affinities of 4-substituted pyridines increase slightly as the substituent on the aromatic ring is changed from electron donating Me to electron withdrawing CF 3 . Substituent inductive effects are much as one might predict. Since Lewis bases donate electron pairs and Lewis acids accept them: electron withdrawing substituents tend to decrease the Lewis basicity of basic sites while electron donating substituents increase site Lewis basicity by making them more electron rich. electron withdrawing substituents increase the Lewis acidity of acidic sites by making those sites more electron deficient while electron donating substituents tend to decrease Lewis acidity by making sites less electron deficient. Nevertheless it can be difficult to predict substituent-based trends in Lewis acidity and basicity by inductive effects alone. This is because inductive effects are modest and often exists in competition with other substituent effects, such as Steric effects, discussed in section 6.4.7 Hardness effects, discussed in section 6.6 \(\pi\)-donation and acceptance effects, which can increase or decrease electron density at a given site as well as create an energy barrier for any structural distortions that might occur upon adduct formation. Avoid misconceptions about the impact of\(\pi\) donation and acceptance on Lewis acid-base affinity As with \(\sigma\)-based induction effects, \(\pi\)-donation tends to increase Lewis basicity and decrease Lewis acidity while \(\pi\)-withdrawal tends to decrease Lewis basicity and increase Lewis acidity. However, care needs to be taken in assessing the effect of \(\pi\)-donation effects on Lewis acidity and basicity. For example, some textbooks claim that the Lewis acidity of boron trihalides is dominated by the reduction of boron acidity through \(\pi\)-donation from the halide substituents: According to this explanation, the extent of this \(\pi\)-donation decreases down the halogen group as the boron-halogen bond distance decreases. This is consistent with the observed trend in Lewis acidity of the boron trihalides towards most bases, which runs counter to that suggested by inductive effects alone: BI 3 > BBr 3 > BCl 3 >> BF 3 However, computational work suggests that this explanation is incorrect, since atomic size effects are important mainly for substituents in which the connected atom is in row 3+ or higher; atomic size effects mainly involve changes in the extent of \(\sigma\)-overlap. In other words, the larger halogens are less able to reduce the electron deficiency at the boron center through \(\sigma\) interactions while \(\pi\)-interactions play little or no role. References: 1. Plumley, J. A.; Evanseck, J. D., Periodic Trends and Index of Boron Lewis Acidity. The Journal of Physical Chemistry A 2009, 113 (20), 5985-5992. 2. Jupp, A. R.; Johnstone, T. C.; Stephan, D. W., Improving the Global Electrophilicity Index (GEI) as a Measure of Lewis Acidity. Inorganic Chemistry 2018, 57 (23), 14764-14771. The conjugate bases of many Brønsted superacids have electron-withdrawing substituents that make them such poor Lewis bases that they are useful as noncoordinating anions. The nonreactivity of Brønsted superacids' conjugate bases towards hydrogen ions is often mirrored in nonreactivity towards other Lewis Acids/electrophiles, most notably metals. This makes these substances useful as inert or noncoordinating ions, although since all are reactive towards suitably electrophilic centers, they are perhaps better understood as weakly coordinating. A number of noncoordinating anions are commonly used in synthetic and other applications. The conjugate base of perchloric acid, perchlorate, was a common noncoordinating inert anion in classical coordination chemistry and continues to be used widely in electrochemistry. In contrast, the conjugate bases of triflic acid, hexafluoroboric acid, and tetrafluoroboric acid are now more commonly used as counterions for isolating reactive cations. Even less reactive noncoordination anions include derivatives of tetraphenylborate, particularly those with electron withdrawing substituents. Other classes of noncoordinating ions include fluoroantimonate clusters, derivatives of the carborane anion (\(CB_{11}H_{11}^-\)), and fluorinated aluminum tetraalkoxides. References Engesser, T. A.; Lichtenthaler, M. R.; Schleep, M.; Krossing, I., Reactive p-block cations stabilized by weakly coordinating anions. Chemical Society reviews 2016 , 45 (4 ), 789-899.