Elements calculating moles

Mass % composition of compound — Assume 100 g — Grams of each element — X 1/molar mass — Moles of each element Calculate mole ratios Empirical formula... 

Strategy First (1), convert the masses of the three elements to moles. Knowing the number of moles (n) of K, Cr, and O, you can then (2) calculate the mole ratios. Finally (3), equate the mole ratio to the atom ratio, which gives you the simplest formula. 

From the density of each element, calculate the volume occupied by one mole of its atoms in... 

Knowing a compound s percent composition makes it possible to calculate the compound s chemical formula. As shown in Figure 3.8, the strategy is to find the relative number of moles of each element in the compound and then use the numbers to establish the mole ratios of the elements. The mole ratios, in turn, give the subscripts in the chemical formula. 

The mass of elements in a compound is crucial when solving problems with the Empirical Formula = Simplest Formula empirical formula. First, the number of moles in atoms are found by dividing their masses by their molar masses. Next, the number of moles are either divided by the smallest mole number or, if necessary to calculate mole ratio of elements, multiplied with certain multipliers to get whole numbers for each element. These numbers for different elements provide the empirical formula. 

The product gas from a solid fuel combustion reaction has the following dry-basis molar composition. 72.0% CO2,2.57% CO, 0.0592% SO , and 25.4% O2. Pure oxygen is fed to the furnace in 20% excess of that required to burn the fuel completely. There is no oxygen in the fuel. Calculate the elemental composition (mole% of the various elements) of the fuel, stating any assumptions you have to make to arrive at your answer. 

Step 1 Assume 100 g sample calculate moles of each element... 

If we can determine the number of moles of each element in any amount of a substance, we can calculate the molar ratio of these elements in the compound, which can then be simplified to the positive integers that represent the simplest molar ratio. We found in Sections 9.2 and 9.3 that we can calculate moles of a substance from grams. Thus one path to determining the empirical formula for a compound is ... 

Next, calculate the simplest ratio of moles of elements by dividing the moles of each element by the smallest value in the calculated mole ratio. 

From the electronegativities of the elements calculate a value for the standard enthalpy of formation of SiC(c). The experimental value is—111 kJ mole . (Answer —96 kJ mole . )... 

The empirical formula of a compound shows Ihe ratios of numbras of atoms in the compound. You can find this formula from the composition of the compound by converting masses of Ihe elements to moles. The next two examples show these calculations in detail. 

According to Figure 3.7, the mass of an element in grams can be calculated by multiplying the amount of the element in moles by the element s molar mass., grams Cu moles Cu x, = grams Cu moles Cu... 

Avogadro s number can be used to find the number of atoms of an element from the amount in moles or to find the amount of an element in moles from the number of atoms. While these types of problems are less common in chemistry than converting between amoimt in moles and mass in grams, they are usefiil in demonstrating the meaning of Avogadro s number. Note that in these calculations, Avogadro s number is expressed in units of atoms per mole. 

Strategy From the product masses, determine the mass of C and the mass of H in the 5.50-g sample of benzene. Sum the masses of C and H the difference between this sum and the original sample mass is the mass of O in the sample (if O is in feet present in benzene). Convert the mass of each element to moles, and use the results as subscripts in a chemical formula. Convert the subscripts to whole numbers by dividing each by the smallest subscript. This gives the empirical formula. To calculate the molecular formula, first divide the molar mass given in the problem statement by the empirical-formula mass. Then, multiply the subscripts in the empirical formula by the resulting number to obtain the subscripts in the molecular formula. 

Answer Since a molecular formula represents the number of moles of each element per mole of compound, first calculate the mass of each element in 1 mole of compound from its molecular weight and composition, (Recall that mass percentage of an element is its mass in 100 g of the compound.)... 

Without doing any calculations, determine which of the samples contains the greatest amount of the element in moles. Which contains the greatest mass of the element ... 

A more useful quantity for comparison with experiment is the heat of formation, which is defined as the enthalpy change when one mole of a compound is formed from its constituent elements in their standard states. The heat of formation can thus be calculated by subtracting the heats of atomisation of the elements and the atomic ionisation energies from the total energy. Unfortunately, ab initio calculations that do not include electron correlation (which we will discuss in Chapter 3) provide uniformly poor estimates of heats of formation w ith errors in bond dissociation energies of 25-40 kcal/mol, even at the Hartree-Fock limit for diatomic molecules. 

Plot the calculated first IPs as a function of the atomic number Z for the elements from H to Ne in the atomic table. The plot has a characteristic shape that should be familiar from earlier courses. These plots are frequently given in the experimental units of electron volts (eV hartrees x 27.21 = eV) or kilojoules per mole (kJ mol hartrees x 2625 = kJmol ). Write a paragraph or two in your project report explaining why the graph of IP vs. Z appears as it does. 

These data can be used to obtain the value of the equilibrium constant at any temperature and this in turn can be used to calculate the degree of dissociation through the equation for the conceiiuation dependence of the constant on the two species for a single element, die monomer and the dimer, which coexist. Considering one mole of the diatomic species which dissociates to produce 2x moles of the monatomic gas, leaving (1 — jc) moles of the diatomic gas and producing a resultant total number of moles of (1 +jc) at a total pressure of P atmos, the equation for the equilibrium constant in terms of these conceiiU ations is... 

The heat capacity, the heat required to raise 1 g-mole drrough one kelvin, can be calculated at temperatures generally above 300 K by two simple empirical rules. The first of these, Dulong and Petit s rule, was discovered in the course of calorimeU ic smdies of the heat capacities of the elements and shows drat the heat capacity has a value,... 

Knowing the formula of a compound, Fe203> you can readily calculate the mass percents of its constituent elements. It is convenient to start with one mole of compound (Example 3.4a). The formula of a compound can also be used in a straightforward way to find the mass of an element in a known mass of the compound (Example 3.4b). 

The strategy used to calculate the simplest formula depends to some extent on which of these types of information is given. The basic objective in each case is to find the number of moles of each element, then the simplest mole ratio, and finally the simplest formula. 

The Gibbs-Helmholtz equation can be used to calculate the standard free energy of formation of a compound. This quantity, AGf, is analogous to the enthalpy of formation, AH . It is defined as the free energy change per mole when a compound is formed from the elements in their stable states at 1 atm. 

The volume per mole of atoms of some fourth-row elements (in the solid state) are as follows K, 45.3 Ca, 25.9 Sc, 18.0 Br, 23.5 and Kr, 32.2 ml/mole of atoms. Calculate the atomic volumes (volume per mole of atoms) for each of the fourth-row transition metals. Plot these atomic volumes and those of the elements given above against atomic numbers. 

---