Molecules and Formula Units

Formulas for ionic compounds represent ont formula unit. However, molecules, too, hive formulas and thus formula units. Even uncombined atoms of an element have formulas. Thus, formula units may refer to uncombined atoms, molecules, or atoms combined in ionic compounds one formula unit may be 

One atom of uncombined element, for example, Pt One molecule of a covalently bonded compound, for example, H2O One simple unit of an ionic compound, for example, NaCl or (NH4)2S04 

EXAMPLE 7.1. What are the formula units of the element copper, the compound dinitrogen trioxidi and the compound sodium bromide  

The formula unit of copper is Cu, one atom of copper. The formula unit of dinitrogen trioxide is f 03, one molecule 

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The periodic table can also be used to calculate the molar mass of molecules and formula units as well. If you can add up the mass of all of the atoms in a molecule to find the molecular mass in atomic mass units, the molar mass of the same molecular compound will have the same value with the unit, grams (g). Following are some examples ... 

But we won t be dealing with pennies here. We start with atoms of the elements, and then move onto molecules and formula units of compounds. As the discussion progresses, the definition of important terms is repeated to help you stay on course. 

Whether you are dealing with elements or compounds, the molar mass of a species is the mass in grams of 1 mole (6.022 x 1023) of that species 1 mole of atoms, 1 mole of molecules, or 1 mole of formula units. With compounds, you re dealing with molecules and formula units, so it is necessary to calculate the molecular or formula mass of each compound to get its molar mass. It s easy to get confused by the language ... 

Atoms, most molecules, and formula units of ionic compounds are extremely tiny. Their formula masses are measured in atomic mass units, which are useful for comparison purposes only, in order to get weighable quantities of matter, a huge collection of formula units is required. The mole is defined as the number of atoms in exactly 12 grams of A millimole is 0.001 mol, and is useful for calculations with small quantities of substances. The mole is abbreviated mol, not m or M, which are used for related quantities, and millimole is abbreviated mmol. 

From the quantities above, it can be seen that the equation can be viewed in terms as small as the smallest number of molecules and formula units. In this case, that involves simply 1 formula unit of CaCOg (s), 1 formula unit of CaO is), and 1 molecule of CO2. This would involve a total of 1 Ca atom, 1 C atom, and 3 O atoms. From such a small scale, it is possible to expand to moles by scaling up by 6.022 X 10 (Avogadro s number), giving 100.1 g of CaCOg, 56.Ig of CaO, and 44.0 g of CO2. Actually, these quantitative relationships are applicable to any amount of matter, and they enable the calculation of the amounts of material reacting and produced in a chemical reaction. Next, it is shown how these kinds of calculations are performed. 

Balance this equation. Using the balanced equation, describe the amounts of each reactant and product in terms of atoms, molecules, and formula units (for any ionic compound). Then describe the amounts of each reactant and product in terms of moles. Finally, describe the amounts of each reactant and product in terms of mass, in grams. How did you determine the numbers of moles Are the masses in the same proportions as the numbers of moles Compare the mass of the mercury(II) oxide with the total mass of mercury and oxygen. 

Chemical formulas, which we defined in Section 1.3.2 as symbolic designations for compounds, are used for both molecular compounds and ionic compounds. Thus we can tell what elements are chemically combined to make up a compound by looking at its formula. Examples we presented in Section 1.3.2 included H2O (water), a combination of hydrogen (H) and oxygen (O), and NaCl (sodium chloride, or salt), a combination of sodium (Na) and chlorine (Cl). Water is an exanple of a molecular compound, while sodium chloride is an example of an ionic compound. Since compounds are composed of elements in chemical combination, it follows that molecules and formula units are composed of the atoms of the elements in chemical combination. Thus, molecules of water are composed of atoms of hydrogen in chemical combination with atoms of oxygen, and formula units of salt are composed of sodium atoms (actually ions) and chlorine atoms (ions) in chemical combination. It is the nature of this chemical combination that defines whether a compound is molecular or ionic. This will be detailed further as our study proceeds. 

The number of atoms of each element that make up a molecule or formula unit can be seen in the formulas. The numbers written as subscripts in the formulas, such as the 2 in the formula for water, H2O, and the 2 and the 4 in the formula for potassium chromate, K2Cr04, represent the number of atoms of the elements immediately to the left of each subscript in the formula. If there is no subscript immediately to the right of an element symbol, then there is only one atom of that element in the molecule or formula unit. Thus, in a molecule of water, H2O, there are 2 atoms of hydrogen and one atom of oxygen that are chemically combined. In a formula unit of salt, NaCl, there is one atom (ion) of each element, sodium and chlorine. In a formula unit of potassium chromate (also an ionic compound), K2Cr04, there are two atoms of potassium, one atom of chromium, and four atoms of oxygen. Thus chemical formulas are reasonable representations of the molecules and formula units of which compounds are composed. 

We previously indicated that chemists think of atoms as tiny spheres of various sizes. Chemists think of molecules and formula units as entities that have a number of these spheres linked together or stuck to each other. The forces that cause them to stick differs, depending on whether the compound is ionic or molecular. For example, a chemist s view of the molecules of water (H2O) is shown in Figure 1.4. The sizes of molecules and formula units, then, are most often on the same scale as the incredibly small atoms of which they are composed. 

While we have dealt with the atomic weights of elements here and in Chapter 1, we have not dealt with the weights of molecules and formula units, which are the fundamental particles of compounds. First, by way of review, it is important to distinguish between molecular (or covalent) compounds and ionic compounds. Molecular compounds are those that consist of distinct individual particles called molecules. Molecules interact with one another in a sample of a molecular compound, but only in a rather mild way. Ionic compounds are those that consist of indistinct particles, which are most appropriately referred to as formula units, because a formula represents the simplest ratio between the cations and anions that constitute the compound (and not a distinct particle). Ionic compounds, on the atomic scale, consist of a three-dimensional array of cations and anions, all of which strongly interact with each other. Thus there is no such particle as a molecule in these cases. Figure 7.1 depicts the difference between these two kinds of compounds. 

Besides identifying what substances are the reactants and products of a chemical reaction and how many atoms, molecules, and formula units are involved, chemical equations can provide additional information if the writer so chooses. Symbols may be used to provide some qualitative information about the substances involved or about the reaction in general. However, such symbols are usually optional. Symbols giving specific information about the reactants or products are placed in parentheses immediately to the right of the chemical symbol or formula in the equation. Symbols giving information about the reaction in general are usually placed above the arrow. 

State with words the following equation in terms of atoms, molecules, and formula units and then restate the equation in terms of moles. 

According to international rules the mass of molecules and formula units is given as relative molecular mass Mr, defined as follows ... 

The molar mass of a compound, the mass per mole of its molecules or formula units, is used to convert between the mass of a sample and the amount of molecules or formula units that it contains. 

STRATEGY We convert from the given volume of gas into moles of molecules (by using the molar volume), then into moles of reactant molecules or formula units (by using a mole ratio), and then into the mass of reactant (by using its molar mass). If the molar volume at the stated conditions is not available, then use the ideal gas law to calculate the amount of gas molecules. 

Avogadro s number is the number of particles (atoms, molecules,. or formula units) that are in a mole of a substance. In this lab, you will relate a common object to the concept of Avogadro s number by finding the mass and volume of one mole of the object. 

Ans. There are 6.02 X 10 i atoms in 1.00 mol Na (Avogadro s number). There is 23.0 g of Na in LOO mol Na (equal to the atomic weight in grams). This problem requires use of two of the most important conversion factors involving moles. Note which one is used with masses and which one is used with numbers of atoms (or molecules of formula units). With numbers of atoms, molecules, or formula units, use Avogadro s number with mass or weight use the formula weight. 

One of the most important electrolytic processes is the extraction of aluminum from an ore called bauxite. This ore is mainly composed of hydrated aluminum oxide, AI2O3 XH2O. (The x in the formula indicates that the number of water molecules per formula unit is variable.) In industry, the scale of production of metals is huge. The electrolytic production of aluminum is over two million tonnes per year in Canada alone. As you know from Faraday s law, the amount of a metal produced by electrolysis is directly proportional to the quantity of electricity used. Therefore, the industrial extraction of aluminum and other metals by electrolysis requires vast quantities of electricity. The availability and cost of electricity greatly influence the location of industrial plants. 

K2La2Ti3Oio is a typical oxide in the Ruddlesden-Popper series exhibiting interlayer reactivity44) It forms a hydrate with about 1-2 H20 molecules per formula unit under ambient conditions.44,74) Fig. 16.9 shows schematic structures of anhydrous and hydrated oxides. Variation of the number of waters of hydration with humidity and exposure time has been reported for the analogous compound K2Nd2Ti3O10.70) A band gap of 3.5 eV has been estimated for K2La2Ti3Oi0 from reflectance measurements.102)... 

How do we use molecular masses Because the mass ratio of one HC1 molecule to one ethylene molecule is 36.5 28.0, the mass ratio of any given number of HC1 molecules to the same number of ethylene molecules is always 36.5 28.0. In other words, a 36.5 28.0 mass ratio of HC1 and ethylene corresponds to a 1 1 number ratio. Equal numbers of different molecules (or formula units) always have a mass ratio equal to their molecular (or formula) mass ratio (Figure 3.1). 

When referring to the enormous numbers of molecules or ions that take part in a visible chemical reaction, it s convenient to use a special unit called a mole, abbreviated mol. One mole of any substance is the amount whose mass—its molar mass—is equal to the molecular or formula mass of the substance in grams. One mole of ethylene has a mass of 28.0 g, one mole of HC1 has a mass of 36.5 g, one mole of NaCl has a mass of 58.5 g, and so on. (To be more precise, one mole is formally defined as the amount of a substance that contains the same number of molecules or formula units as there are atoms in exactly 12 g of carbon-12.)... 

When substances—the reactants—react to form new substances—the products—we say that a chemical reaction has occurred. A chemical equation is a statement of such an event in which the formulas of the reactants are on the left, followed by a right-pointing arrow and then the formulas of the products. In a balanced equation, coefficients specify the number of molecules (or formula units) of each species involved. The coefficients must satisfy Dalton s requirement that atoms are not created or destroyed in a chemical reaction. There is no fixed procedure for balancing an equation. Although a systematic algebraic approach is in principle possible, a trial-and-error approach often works. 

To indicate that a binary compound is made up of two non-metals, a prefix is usually added to both non-metals in the compound. This prefix indicates the number of atoms of each element in one molecule or formula unit of the compound. For example, P205 is named diphosphorus pen-toxide. Alternatively, the Stock System may be used, and P205 can be named phosphorus (V) oxide. AsBr3 is named phosphorus tribromide. 

To begin to balance an equation, you can add numbers in front of the appropriate formulas. The numbers that are placed in front of chemical formulas are called coefficients. They represent how many of each atom, molecule, or formula unit take part in each reaction. For example, if you add a coefficient of 2 to NaCl in the equation in Figure 4.2, you indicate that two formula units of NaCl are produced in the reaction. Is the equation balanced now As you can see by examining Figure 4.3, it is not. The chlorine atoms are balanced, but now there is one sodium atom on the left side of the equation and two sodium atoms on the right side. 

The mole is literally the chemist s dozen. Just as egg farmers and grocers use the dozen (a unit of 12) to count eggs, chemists use the mole (a much larger number) to count atoms, molecules, or formula units. When farmers think of two dozen eggs, they are also thinking of 24 eggs. 

In section 5.2, you explored the relationship between the number of atoms or particles and the number of moles in a sample. Now you are ready to relate the number of moles to the mass, in grams. Then you will be able to determine the number of atoms, molecules, or formula units in a sample by finding the mass of the sample. 

Using the mole concept and the periodic table, you can determine the mass of one mole of a compound. You know, however, that one mole represents 6.02 x 1023 particles. Therefore you can use a balance to count atoms, molecules, or formula units ... 

As part of his atomic theory, John Dalton stated that atoms combine with one another in simple whole number ratios to form compounds. For example, the molecular formula of benzene, C6H6, indicates that one molecule of benzene contains 6 carbon atoms and 6 hydrogen atoms. The empirical formula (also known as the simplest formula) of a compound shows the lowest whole number ratio of the elements in the compound. The molecular formula (also known as the actual formula) describes the number of atoms of each element that make up a molecule or formula unit. Benzene, with a molecular formula of C6H6, has an empirical formula of CH. Table 6.1 shows the molecular formulas of several compounds, along with their empirical formulas. 

The mole is the amount of substance of a system that contains as many elementary entities as there are atoms in 0.012 kg of carbon-12. Although it is defined in terms of the number of entities, in practice, 1 mol of atoms, molecules, or specific formula units of a substance is measured by weighing M x (1 mol) of the substance, where M is the molar mass, the mass per unit amount of substance. Molar mass is synonymous with the terms atomic weight, for atoms, and molecular weight, for molecules or formula units, respectively, and is reported in grams per mole (g mol ). 

Hydrated salts contain one or more HiO molecules per formula unit in addition to the cations and anions. A dot is used in the formula of these salts. 

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