Empirical formulas experimental determination

Divide the experimentally determined molar mass of succinic acid by the mass of the empirical formula to determine n. molar mass of succinic acid... 

The word empirical means derived from experiment. As we will see in Chapter 3, empirical formulas are determined experimentally. 

For calculations based on Eq. (X.72), we must determine the duration of the impact. For this purpose we will make use of experimental results reported by Malyshev [289], who proposed an empirical formula to determine the impact time when particles come into contact with a surface ... 

The size of the droplets formed in an aerosol has been examined for a range of conditions important in ICP/MS and can be predicted from an experimentally determined empirical formula (Figure 19.6). Of the two terms in the formula, the first is most important, except at very low relative flow rates. At low relative velocity of liquid and gas, simple droplet formation is observed, but as the relative velocity increases, the stream of liquid begins to flutter and to break apart into long thinner streamlets, which then break into droplets. At even higher relative velocity, the liquid surface is stripped off, and the thin films so-formed are broken down into... 

Experimentally it has been shown that for air-water systems the value of Tj /Zc c, the psychrometric ratio, is approximately equal to 1. Under these conditions the wet-bulb temperatures and adiabatic-saturation temperatures are substantially equal and can be used interchangeably. The difference between adiabatic-saturation temperature and wet-bulb temperature increases with increasing humidity, but this effect is unimportant for most engineering calculations. An empirical formula for wet-bulb temperature determination of moist air at atmospheric pressure is presented by Liley [Jnt. J. of Mechanical Engineering Education, vol. 21, No. 2 (1993)]. 

The empirical formula is C3H2N02, with a formula mass of 84.0 g/mol. This is one-half the experimentally determined molar mass. Thus, the molecular formula is C6H4N204. 

A reasonable empirical formula is C4H4S, which has an empirical mass of 84.1 g/mol. Since this is the same as the experimentally determined molar mass, the molecular formula of thiophene if C4H4S. 

In the next section of Addnl Ref O, Evaluation of Parameters , it is stated that the parameters, y, a, and B/Q, which appear in the LSZK equation of state, must be evaluated by using experimental data and it is explained how this is done on pp 9 10. As an example, compressed TNT of various densities was investigated. Detonation velocities determined by LSZK expression (34) proved to be in good agreement with those detd by the empirical formula ... 

Other early experimenters obtained picric acid by nitrating various organic substances such as silk, natural resins, etc. The correct empirical formula for picric acid was determined by Laurent in 1841 who prepared the acid by reacting phenol with nitric acid and isolated dinit-rophenol which was formed in an intermediate stage of the reaction. 

Use the following experimental information to determine the empirical formula of an oxide of silicon. 

Column (5) is determined in this problem, as in the previous problem, by dividing both numbers by the smallest (0.0559) which preserves the mole ratio of 0.0559 0.391 and leads to the whole-number ratio required to write a chemical formula. Column (5) contains 6.99, a number that is so close to a whole number that the difference can be taken for experimental error. The mole ratio of Na2SC>4 to H2O is 1 to 7, providing us with an empirical formula of Na2S04- 7H20. 

The empirical formula of a compound can be determined in a laboratory experiment by finding the ratio between the number of moles of the elements in the compound. The number of moles of each element can be calculated from the experimental values of the weights in which the elements combine by dividing by their corresponding atomic weights. If the molecular weight and the empirical formula of the compound are known, then the molecular formula of the compound can be determined. 

From the experimental data, determine the empirical formula of copper(II) chloride, and the error in determining the percent of copper. 

At this point, you need to make a decision. If the outcome of your calculation produces all integers within a small degree of experimental error, you are ready to use these numbers as the subscripts in the empirical formula. If they are not, then you need to determine what you can multiply all of the numbers by to produce all integers. In most cases, it is fairly obvious. For instance, in our example, you can see that by doubling 3.50, it becomes 7.00—an integer. 

At this point, you need to determine integers that are multiples of 1 and 2.66. Some quick calculations determine that 3 X 2.66 = 7.98, which is well within experimental error of 8. Therefore, the empirical formula must be C3H8. To determine the molecular formula, we can use the information in the problem. One mole of the combustion of unknown C requires 5 moles of Oz. The balanced equation for the complete combustion of C3H8 requires 5 moles of 02. The next possible formula, C6H16, requires 10 moles. The number of moles of oxygen will continue to increase, so the molecular formula must be C3H8. 

The determination of the empirical formula of a compound can be made experimentally, by determining the percentage amounts of elements present in the substance using the methods of quantitative chemical analysis. At the same time the relative molecular mass of the compound has to be measured as well. From these data the empirical formula can be determined by a simple calculation. If, for some reason, it is impossible to determine the relative molecular mass the simplest (assumed) formula only can be calculated from the results of chemical analysis the true formula might contain multiples of the atoms given in the assumed formula. 

The partial pressure of water is determined with the empirical formula from Ref. 21 and is shown in Eq. 14. This relationship enables the partial pressure of water to be calculated from experimental data as the temperature of the DI water into the stack anode varies,... 

The formula of a compound gives the relative number of atoms of the different elements present. It also gives the relative number of moles of the different elements present. As was shown in Sec. 7.5, the percent by mass of each element in a compound may be computed from its formula. Conversely, if the formula is not known, it may be deduced from the experimentally determined composition. This procedure is possible because once the relative masses of the elements are found, the relative numbers of moles of each may be determined. Formulas derived in this manner are called empirical formulas or simplest formulas. In solving a problem in which percent composition is given, any size sample may be considered, since the percentage of each element does not depend on the size of the sample. The most convenient size to consider is 100 g, for with that size sample, the percentage of each element is equal to the same number of grams. 

Formulas describe the composition of compounds. Empirical formulas give the mole ratio of the various elements. However, sometimes different compounds have the same ratio of moles of atoms of the same elements. For example, acetylene, C2H2, and benzene, CeHe, each have 1 1 ratios of moles of carbon atoms to moles of hydrogen atoms. That is, each has an empirical formula CH. Such compounds have the same percent compositions. However, they do not have the same number of atoms in each molecule. The molecular formula is a formula that gives all the information that the empirical formula gives (the mole ratios of the various elements) plus the information of how many atoms are in each molecule. In order to deduce molecular formulas from experimental data, the percent composition and the molar mass are usually determined. The molar mass may be determined experimentally in several ways, one of which will be described in Chap. 12. 

More information is required to construct a molecular formula than is required to obtain the empirical formula of a substance. The empirical formula can be obtained from the elemental analysis of a substance. To obtain the molecular formula, the total molecular mass must be determined experimentally. The molecular for-... 

The molecular formula of a compound is always an integer multiple (e.g., 1, 2,3,...) of the empirical formula. If the empirical formula of a compound is known, the molecular formula can be determined by the experimental determination of the molecular weight of the compound... 

The first step is to divide the experimentally determined molecular weight of the compound by the molecular weight of the empirical formula in order to determine the integer multiple that represents the number of empirical formula units in the molecular formula, hi the second step, the molecular formula is obtained by multiplying the subscripts of the empirical formula by the integral multiple of empirical formula units. 

The empirical formula for a compound is P2O5. Its experimental molar mass is 284 g/mol. Determine the molecular formula of the compound. 

The molecular formula is determined from the empirical formula and the experimentally determined molar mass. 

Summarize briefly the process of using empirical formula and the value for experimental molar mass to determine the molecular formula. 

Your team needs to develop a procedure to experimentally determine the correct empirical formulas for both hydrates of this anhydrous compound. You will use gypsum samples from the mine and samples of the plaster of Paris product. 

To determine the molecular formula for a compound, the molar mass of the compound must be determined through experimentation and compared with the mass represented by the empirical formula. For example, the molar mass of acetylene is 26.04 g/mol and the mass of the empirical formula, CH, isl3.02 g/mol. Dividing the actual molar mass by the mass of the empirical formula indicates that the molar mass of acetylene is two times the mass of the empirical formula. 

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