Empirical and Molecular Formula
The empirical formula represents the simplest whole-number ratio of atoms in a compound, while the molecular formula gives the actual number of each type of atom in a molecule. The empirical formula can be determined from the molecular formula by dividing the subscripts by the greatest common factor. Both formulas provide important information about the composition of a compound.
7 Key excerpts on "Empirical and Molecular Formula"
eBook - ePub- Jeffrey Gaffney, Nancy Marley(Authors)
- 2017(Publication Date)
- Elsevier(Publisher)
...As such, the empirical formula does not necessarily represent the actual numbers of atoms present in a single molecule of the compound. For example, the molecular formula for sulfur monoxide is SO and the molecular formula for disulfur dioxide is S 2 O 2. However, both compounds have the same empirical formula, which is SO. So, the molecular formula can either be the same as the empirical formula (as with sulfur monoxide) or it can be a multiple of the empirical formula (as with disulfur dioxide). Since there is no molecular formula for an ionic compound, the chemical formula and the empirical formula are always the same. An empirical formula is important because it is determined from elemental composition data obtained on a compound directly by experiment. This composition data is either given as the weight percentage of each of the elements present in the compound or as the number of grams of each element present in the sample. In either case, in order to determine the empirical formula from this experimental data, the results must first be converted to moles by using the molar mass of each element. After converting the composition values to moles, the empirical formula is then derived by determining the smallest positive whole number ratio of the moles present in the compound. This is done by dividing the number of moles of each element by the smallest molar value. If this does not result in a whole number, fractional values that are close to a whole number can be rounded to the nearest whole number. For example, 2.03 can be rounded off to 2.00 and 6.98 can be rounded up to 7.00. However, a fractional value that is greater than 0.1 or less than 0.9 must be multiplied by a small integer to obtain a whole number as shown in Table 4.1...
eBook - ePub- Adrian Dingle, Derrick C. Wood(Authors)
- 2014(Publication Date)
- Research & Education Association(Publisher)
...A sample that does not match the percentage by mass of a pure sample is not pure. 4. The empirical formula of a compound is the simplest wholenumber (integer) ratio of the atoms of each element in that compound. 5. The empirical formula can be calculated from percentage by mass data. i. Take the percentage of each element and assume a sample of 100 g. (This assumption converts percentages to masses.) ii. Convert the masses (per 100 g) of each element to the moles (see II. Moles below) of each element by dividing each element’s mass (per 100 g) by the corresponding average atomic mass taken from the periodic table. iii. Find the ratio of the moles calculated in ii above by dividing each of the moles by the smallest number of moles. iv. The results from iii above will be in a convenient ratio and gives the empirical formula. It may be that the ratio includes a decimal (fraction) such as 0.500, 0.333, 0.250, and so on. Since empirical formulae are the simplest whole -number ratios, you must multiply all of the numbers in the ratio by 2, 3, or 4 as appropriate in order to remove the decimal (fraction). C. Molecular Formula 1. Unlike an empirical formula (which shows the simplest wholenumber ratio), the molecular formula of a compound shows the exact whole-number ratio of the different elements in a compound. 2. Like the empirical formula, the numbers of each element are recorded using a subscript to the right of the element’s symbol. When only 1 atom is present, the subscript 1 is assumed (understood) and not written. 3. The following describes the relationship between the molecular formula and the empirical formula. i. The molecular formula will be some simple, integer-multiple of the empirical formula...
eBook - ePubChemistry
Concepts and Problems, A Self-Teaching Guide
- Richard Post, Chad Snyder, Clifford C. Houk(Authors)
- 2020(Publication Date)
- Jossey-Bass(Publisher)
...The simplest possible chemical formula that represents the ratios of the atoms within an unknown molecule is called an empirical formula. The simple formula (CH) is called a(n) ______________. Answer: empirical formula An unknown compound with an empirical formula of CH could be any of several compounds. For example, both acetylene (C 2 H 2) and benzene (C 6 H 6) have an empirical formula of CH. The formulas for acetylene and benzene (C 2 H 2 and C 6 H 6) are called molecular formulas because they describe the actual number of atoms contained in each molecule. A formula describing the simplest ratio between atoms in an unknown molecule is called a(n) _____________. A formula describing the actual number of atoms contained in each molecule is called the _____________. Answer: (a) empirical formula; (b) molecular formula An empirical formula is sometimes the same as an actual molecular formula. A formula determined by percentage weight analysis is considered to be an empirical formula since it only represents the ratio of one atom to another. By weight analysis, a compound is found to consist of carbon and oxygen in a 1:1 ratio. A formula of CO is assigned to the compound and you suspect that the compound could be carbon monoxide but have no proof. The formula CO is in this case a(n) ___________ (empirical, molecular) formula. Answer: empirical (If by simple weight analysis, the compound has a carbon and oxygen ratio of 1 : 1, then the assigned formula CO is empirical. If it is determined by additional testing that the compound is indeed carbon monoxide, then the formula CO would be a molecular formula.) To determine an empirical formula if given percentage composition, follow these three steps. Drop the percentage signs and replace them with amu. The result is the weight composition of a 100 amu sample. Multiply each of the results of step 1 by the appropriate atoms per atomic weight conversion factor (as shown in frames 26 and 27)...
eBook - ePub- Kevin R. Reel(Author)
- 2013(Publication Date)
- Research & Education Association(Publisher)
...However, you must go through the stoichiometric calculation with both sets of given information. • The limiting reactant is the reactant that runs out first as the chemical reaction proceeds. Therefore, it is the reactant that produces the least amount of product. Example: In the following reaction, a 12.8 L sample of CO at 2.0 atm and 300 K is combined with 6.33 g of Fe 2 O 3. How many liters of carbon dioxide are formed at the same temperature, and which is the limiting reactant? Solution: Therefore, iron oxide is the limiting reactant and 1.46 L of carbon dioxide is formed. EMPIRICAL FORMULAS AND PERCENT COMPOSITION • The empirical formula is the simplest ratio of atoms in a molecule. The empirical formula may or may not be the same as the molecular formula. The molar mass is the information that is needed to convert from the empirical formula to the molecular formula. • Mass-mole conversions are necessary in finding empirical formulas from experimental data using the following steps: 1. Change the “%” to grams by assuming that you have 100 grams of material. 2. Convert from mass into grams for all atoms involved. 3. Divide the moles of each atom by the value that represents the smallest number of moles of atoms. Example: (a) What is the empirical formula of the hydrocarbon that contains 85.7% carbon? (b) Vapor pressure calculations determine the molar mass of the compound to be 28.0 g/mol. What is the molecular formula of the compound? Solution: Empirical formula: C 7.135 H 14.186, or CH 2 (b) Since the molar mass is twice the mass of the empirical formula, then the molecular formula is twice the empirical formula; C 2 H 4....
eBook - ePubThe Science For Conservators Series
Volume 1: An Introduction to Materials
- The Conservation Unit Museums and Galleries Commission(Author)
- 2008(Publication Date)
- Routledge(Publisher)
...These are much more useful than the ones we have just seen for discussing chemical reactions. Consequently they are the ones most frequently found in chemistry books and in conservation texts. B1 Molecular formulae molecular formula The first of these more symbolic models is the molecular formula. It tells concisely how many atoms of which elements are contained in each molecule of a compound. subscript The molecular formula for methane is CH 4. Comparing this combination of two letters and a number with the pictures in figures 3.2 and 3.3 shows you immediately that it describes the molecule as having one carbon atom and four hydrogen atoms. The elements present are identified by their symbols (see page 33 in Chapter 2), and there is another convention to show how many of each sort of atom is present. Although the information might have been written CHHHH, the convention is to represent any number of atoms greater than one by a little numeral at the bottom right-hand corner of the element’s symbol (a subscript). Thus in CH 4 the letter C stands for one atom of carbon and H 4 for four atoms of hydrogen. (The symbols must always be written carefully; Co is the symbol for cobalt but CO is the molecular formula for carbon monoxide.) Exercises In order to help you become familiar with writing molecular formulae it may be useful for you to do the following: 1 Write in the appropriate numbers in the blank spaces in the following examples; a the molecular formula for carbon dioxide is CO 2. Each molecule contains _____ atom(s) of carbon and _____ atom(s) of oxygen. b the molecular formula for sulphuric acid is H 2 SO 4. Each molecule contains _____ atom(s) of hydrogen, _____ atom(s) of sulphur and _____ atom(s) of oxygen. c the molecular formula of acetone is C 3 H 6 O...
eBook - ePub- Atef Korchef(Author)
- 2022(Publication Date)
- CRC Press(Publisher)
...the empirical formula. Since the mole ratios are integers, so the empirical formula is C 2 H 6 O. 4.12 Balancing Chemical Equations and Stoichiometry 4.12.1 Law of Conservation of Mass The law of conservation of mass states that, in an isolated system (closed to all transfers of matter and energy), the mass of the system must remain constant over time, as the system’s mass is neither created nor destroyed by chemical reactions or physical transformations. According to the law of conservation of mass, the mass of products (substances produced) in a chemical reaction must equal the mass of the reactants (substances consumed). 4.12.2 Balancing Chemical Equations A balanced chemical equation is an equation that represents the correct amount of reactants and products in a chemical reaction. Balancing a chemical equation is determining the coefficients that provide equal numbers of each type of atom on each side of the equation. To balance a chemical equation, follow the following steps: Step 1: Write the correct formulas for the reactants (on the left side of the equation) and the products (on the right side of the equation). Do not change the subscripts that indicate the actual numbers of atoms in one molecule). Example: Butane reacts with oxygen to form carbon dioxide and water. The correct formula of butane is C 4 H 10. Do not write for example 2(C 2 H 5) or (C 2 H 5) 2 The correct formula of oxygen is O 2 and the correct formula of water is H 2 O C 4 H 10 + O 2 → CO 2 + H 2 O Step 2: Change the coefficients (numbers in front of the molecular formulas) to make the number of atoms of each element the same on both sides of the equation. Start by balancing those elements that appear in only one reactant and one product. In the equation: C 4 H 10 + O 2 → CO 2 + H 2 O carbon (C) and hydrogen (H) appear in only one reactant. However, oxygen (O) appears in two products, which are carbon dioxide (CO 2) and water (H 2 O). Therefore, start by balancing C or H and not O...
eBook - ePubFoundations for Teaching Chemistry
Chemical Knowledge for Teaching
- Keith S. Taber(Author)
- 2019(Publication Date)
- Routledge(Publisher)
...To spot periodic patterns it is necessary to know the order of the elements. To a chemist today the elements are inherently ordered by atomic number, which is equated to proton number: that is, an aspect of the model of the structure of matter at submicroscopic scales. The first element is the one that has only one proton in the nucleus of its atoms. This is another abstract idea (our students cannot directly see for themselves that hydrogen is composed of atoms that have only one proton in their nuclei), but a very definitive one that we can teach to students. Yet this model was not available when the periodic table was first conceived, and the ordering of elements was in terms of best available measures of atomic weight (what is now called relative atomic mass) at a time when the atomic nature of matter was still very much conjectural and modern models of atomic structure were yet to be proposed. A key aspect of periodicity relates to valency, which is about the combining power of atoms, and so the formulae of compounds formed. Historically, these formulae were found from empirical investigations. These were complicated by the need to know accurate atomic masses – so, as tends to be the case, chemistry proceeded iteratively (Chang, 2004). Molecular masses and empirical formulae needed to be found by experiments involving careful weighing out of reagents. Yet determining either depended upon knowledge of the other – at least until a completely independent means of finding atomic masses was available, for example by exploring positive ray spectra in vacuum tubes (Thomson, 1922), which led to mass spectroscopy. Today we can see clear patterns in valency in the main groups of the (sic) periodic table, e.g., LiCl; BeCl 2 ; BCl 3 ; CCl 4 ; NCl 3 ; OCl 2 ; ClF; [Ne- no chloride formed] The chemical investigations that provided such information were often difficult at the time and are not always suitable for repeating in school laboratories...
