Specific heat of UPdâ Siâ

The specific heat of single-crystal U Pd2 Si2 has been studied using both the step heating and continious heating methods for the temperature range 2 to 250 K. Successive phase transitions at Tl = 136I< and T2 = 108I< are reported, which are consistent with current publications. The transition at 40K, which was previously reported, has not been detected. Recent published elastic neutron scattering data, magnetic susceptibility and resistivity results suggest that U Pd2 Si2 may be a heavy fermion compound, however, the electronic specific heat coefficient I (= 18.97 ;~), obtained from the specific heat Cv measurements, is smaller than that of the conventional heavy fermion system. The Debye temperature of U Pd2Si2 is found to be 116.55K. The possibility is discussed that the maximum in CIT in the low-temperature range 2 to 4K corresponds to Schottky anomaly induced by localized magnetic impurities .

Heats of formation of some hydrates /|nA. Phillipson. -- 260 St. Catharines, Ont. : [s. n.],

A great deal of data on the heats of formation of various hydrates has been compiled i n the J.A.N.A.F. and other tables such as the National Bureau of Standards circulars. Comparison of the heat of f ormation of a hydrate with that of the corresponding anhydrate exposes anomalies i n a surprising number of cases. Some of the results are so discordant that i t is apparent that one or the other value is seriously mistaken. No attempt has been made i n this work to determine which value may be correct, but measurements have been made of the difference between these two values. The procedure adopted has been to dissolve the hydrate and the anhydrate, to achieve the same final concentration of the compound in solution, and so to measure the difference in heats of solution .. Measurements were made at OOC in a modified Bunsen ice calorimeter, well insulated and surrounded by an icewater mixture . The observed differences in heats of solut ion were corrected t o 25°0 by using appropriate heat capacity data. These differences offer a direct measure of the enthalpy involved in binding a mole of water into the crystal structure and so should shed light on the nature of binding involved. The following hydrates were studied : MgS04.nH20 (n = 1,4,7), MnC12.nH20 (n = 1, 2), LiI. nH20 (n = 1,3), MnS04. nH20 (n = 1,4), CaC12. nH20 (n = 2,6) , K2C03.1~H20, LiCl.H20, LiBr.2H20, CdC12.2t H2o, and N2H4eH20.

NMR studies of the exchange reactions of CH3CN.BX3 with excess CH3CN /|nJoseph Fogelman. -- 260 St. Catharines, Ont. : [s. n.],

Exchange reactions between molecular complexes and excess acid or base are well known and have been extensively surveyed in the literature(l). Since the exchange mechanism will, in some way involve the breaking of the labile donor-acceptor bond, it follows that a discussion of the factors relating to bonding in molecular complexes will be relevant. In general, a strong Lewis base and a strong Lewis acid form a stable adduct provided that certain stereochemical requirements are met. A strong Lewis base has the following characteristics (1),(2) (i) high electron density at the donor site. (ii) a non-bonded electron pair which has a low ionization potential (iii) electron donating substituents at the donor atom site. (iv) facile approach of the site of the Lewis base to the acceptor site as dictated by the steric hindrance of the substituents. Examples of typical Lewis bases are ethers, nitriles, ketones, alcohols, amines and phosphines. For a strong Lewis acid, the following properties are important:( i) low electron density at the acceptor site. (ii) electron withdrawing substituents. (iii) substituents which do not interfere with the close approach of the Lewis base. (iv) availability of a vacant orbital capable of accepting the lone electron pair of the donor atom. Examples of Lewis acids are the group III and IV halides such (M=B, AI, Ga, In) and MX4 - (M=Si, Ge, Sn, Pb). The relative bond strengths of molecular complexes have been investigated by:- (i) (ii) (iii) (iv) (v] (vi) dipole moment measurements (3). shifts of the carbonyl peaks in the IIIR. (4) ,(5), (6) .. NMR chemical shift data (4),(7),(8),(9). D.V. and visible spectrophotometric shifts (10),(11). equilibrium constant data (12), (13). heats of dissociation and heats of reactions (l~), (16), (17), (18), (19). Many experiments have bben carried out on boron trihalides in order to determine their relative acid strengths. Using pyridine, nitrobenzene, acetonitrile and trimethylamine as reference Lewis bases, it was found that the acid strength varied in order:RBx3 > BC1 3 >BF 3 • For the acetonitrile-boron trihalide and trimethylamine boron trihalide complexes in nitrobenzene, an-NMR study (7) showed that the shift to lower field was. greatest for the BB~3 adduct ~n~ smallest for the BF 3 which is in agreement with the acid strengths. If electronegativities of the substituents were the only important effect, and since c~ Br ,one would expect the electron density at the boron nucleus to vary as BF3of back-bonding varies inversely as the bo~on halogen distance and one would therefore expect the B-F bond to exhibit greater back-bonding character than the B-Cl or B-Br bonds. Since back-bonding transfers electron density from substituent to the boron atom site, this process would be expected to weaken the Lewis acid strength. This explains the Lewis acid strength increasing in the order BF 3 BC1 3 BBr 3 . When the acetonitrile boron trihalide complex is formed, the boron atom undergoes ~_cbange of hybridization from sp2 to sp3. From a linear relationship between the heat of formation of ethyl acetate adducts and the shift in the carbonyl I.R. stretch, Drago (22) et al have proposed that the angular di~tortion of the X-B-X bonds from sp2 (12 ) to sp3 (10 hybridization is proportional to the amount of charge transferred, i.e. to the nature of the base, and they have rejected the earlier concept of reorganization energy in explaining the formation of the adduct bond (19).