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Chemical Formulas and Composition Stoichiometry

Chemical Formulas and Composition Stoichiometry. Chapter 2. Stoichiometry. Describes the ___________ relationships among elements in compounds (composition stoichiometry) and among substances as they undergo chemical changes (reaction stoichiometry). Dalton’s Atomic Theory.

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Chemical Formulas and Composition Stoichiometry

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  1. Chemical Formulas and Composition Stoichiometry Chapter 2

  2. Stoichiometry • Describes the ___________ relationships among elements in compounds (composition stoichiometry) and among substances as they undergo chemical changes (reaction stoichiometry).

  3. Dalton’s Atomic Theory • An element is composed of extremely small, indivisible particles called _____. • All atoms of a given element have ______ properties that are ______. • Atoms cannot be _____, ______, or ______ into atoms of another element. • ______ are formed when atoms of different elements combine with one another in small whole-number ratios. • The relative _____ and ____ of atoms are constant in a given compound.

  4. Atoms and Molecules • We know what atoms are and the fundamental particles that compose them. • _______ is the smallest particle of an element or compound that can have a stable, ______ existence. • This means that if a molecule contains only one type of atom it is also an element.

  5. Atoms and Molecules • Monatomic – contains only ___ atom • Na and He For these elements, an atom and a molecule are the same. • Diatomic – molecules that contain only __ atoms • H2, O2, and all halogens • Polyatomic – molecules than contain ______ atoms • P4, S8, and C60 Allotropes – different _____ (geometry or number of atoms) of the same element

  6. Compounds • Compounds are substances that contain two or more elements in fixed proportions • Compounds containing only covalent bonds • The atoms in a compound _________ electrons. • These compounds are also called molecules since they exist as individual particles. • Compounds containing ionic bonds • These compounds are held together by electrostatic interactions. • These are not molecules since they do not exist as individual particles.

  7. Covalent Compounds • The chemical formula represents the ____ present and the ______ in which the atoms of the elements occur. • Table 2.2 Show a few models to illustrate that all atoms are held together by covalent bonds. • H2O and C4H10 (models) • Notice that the formula is not necessarily the simplest ratio • _________ compounds are largely composed of C-C,C-H, and C-O bonds. • Ethanol and benzene

  8. Covalent Compounds • Inorganic compounds largely do not contain _____ and/or ____ bonds. • H2O and NH3 • Law of Definite Proportions – different pure samples of a compound always contain the same elements in the same proportion by mass • _______ of atoms can also affect the properties of compounds even if they have the same chemical formula. Demo: John Cullen’s demonstration

  9. Compounds Containing Ionic Bonds • Held together by __________ between ions • They do not exist as individual units • Collection of a large number of ions • The formula represents the ______ of ions in the compound • Each positively-charged species is surrounded by negatively-charged species (and vice-versa) • DEMO: NaCl structure

  10. Writing Formulas and Names for Ionic Compounds (Table 2-3) • Positive and negative ions will combine in such a way to make the ionic compound _____. • Binary Ionic Compounds • You already know how to write the formulas. • Write the formula for calcium bromide and lithium sulfide • Naming a binary compound from the formula • Cation – name of the metal • Anion – drop the last portion of the name and add ‘ide’ • Name AlBr3 and BeI2

  11. Common Ions

  12. More Polyatomic Ions

  13. Writing Formulas and Names for Ionic Compounds (Table 2-3) • Ionic compounds containing a cation that can have multiple charges • The charge on the cation has to be specified • What is the charge on FeCl3 and FeCl2? • In most cases, the charge can be determined by the negatively-charged ion • Write the names for these compounds? • Write the formulas for copper(II) sulfide and iron(III) oxide.

  14. Writing Formulas and Names for Ionic Compounds (Table 2-3) • Ionic compounds containing polyatomic ions • The atoms in polyatomic ions are held together by _________. • Write the Lewis structure for SO32- • The polyatomic ions and the oppositely-charged ions are held together in the solid by _________

  15. Writing Formulas and Names for Ionic Compounds (Table 2-3) • Write the names for (NH4)2S and K2(SO4). • Notice that the charges are fixed and given on Table 2-3 for the polyatomic ions. • Write the formulas for magnesium nitrate and lithium carbonate. • If more than one polyatomic ion is needed parentheses are placed around it

  16. Atomic Weight • The atomic weight is based on the atomic mass unit. • One amu is _____ the mass of a carbon-12 atom • The atomic weight, which is given on the periodic table, is a _______ of all the naturally occurring isotopes of that element

  17. The Mole • A mole describes a ________ of objects, particles, or atoms. This is similar to other measurements that describe a particular quantity. • 1 dozen = 12 items • Gross = 144 items • 1 mole = _____________ (much larger) • Avogardo’s number

  18. The Mole • A mole is the amount of substance that contains as many entities as there are atoms in 12.0 grams of carbon-12. • One mole of H2O would contain 6.02  1023 molecules of H2O. • The mass of one mole of atoms of an element is equal to its atomic weight (periodic table) in grams. • This is called the molar mass (units ________).

  19. The Mole • Iron (Fe) • These conversion factors will be used extensively!! • Calculate the mass of a single Mg atom. • Calculate the number of atoms in 1.40  10-15 moles of Cu. • Calculate the number of moles in 85.3 grams of Al.

  20. Molar Mass • The ________ of a substance is the sum of the atomic weights of all the atoms in the formula. Expressed in amu. • _________ can be used if the compound is composed of molecules. • Molar mass is _________ to the formula weight of a compound. Expressed in grams • Indicates how many grams in ______ of the compound

  21. Molar Mass • This molar mass of a substance will contain 6.02  1023 formula units • What is the molar mass of Ca(NO3)2? • Calculate the number of C3H8 molecules in 75.4 grams of C3H8 • Calculate the number of millimoles in 0.234 g of oxalic acid (CHOO)2

  22. More Problems • The chemical formula indicates the relative number of ______ and _____ in a compound. • Shown by the ________ C3H8 • How many formula units in 2.5 moles of CaCl2? How many Ca2+ and Cl- ions in 2.5 moles of CaCl2? • Calculate the number of oxygen atoms in 54.2 grams of sucrose, C6H12O6. • Calculate the number of moles of hydrogen in 54.2 grams of sucrose.

  23. Percent Composition • The mass percent of each element present in a compound is simply determined by dividing its mass by the molar mass of the compound multiplied by 100%. • Let’s do H2O. What is the percent composition of oxygen and hydrogen? • What is the percent composition of carbon and hydrogen in C3H8?

  24. Percent Composition • The percentages must add up to _____. Any pure compound will contain the same percentages of each element (____ __ _____ ________). • One more – What is the percent composition of Fe in Fe2(SO4)3?

  25. Determining Empirical Formulas from Percent Composition • Empirical formula (or simplest formula) is the __________ ratio of elements in a compound. • Molecular formula is the ________ of atoms present in a compound containing covalent bonds. The empirical formula and molecular formula are not necessarily the same. • Benzene (show structure) Molecular vs. empirical • Hydrogen peroxide (show structure) Molecular vs. empirical • Sucrose

  26. Determining Empirical Formulas from Percent Composition • A compound contains 24.74% K, 34.76% Mn, and 40.50% O. What is the empirical formula of this compound? • Step 1: Consider how many grams of each element would be present in 100.0 grams. • Step 2: Convert to moles for each element. • Step 3: Obtain smallest whole-number ratio by dividing by the smallest mole number. • Step 4: If necessary, convert fraction to whole numbers by multiplying an integer.

  27. Determining Empirical Formulas from Percent Composition • A compound contains 7.537% H, 26.185% N, and 66.278% Cl. What is the empirical formula? • A sample of a compound contains 6.541 grams of Co and 2.369 grams of O. Find the simplest formula.

  28. Combustions Train – Analysis of Hydrocarbons • Combustion trains are used to perform quantitative carbon-hydrogen analysis. • A known mass is burned in a stream of oxygen. • Carbon is converted to CO2 • Hydrogen is converted to H2 Weight increase in the absorber region is measured.

  29. Combustions Train – Analysis of Hydrocarbons • Determine the mass of carbon in the compound • Use the conversion factor • Determine the mass of oxygen in the compound • Use the conversion factor Why do we use these conversion factors? Which units will cancel?

  30. Combustions Train – Analysis of Hydrocarbons • A 0.2815-gram sample of pure hydrocarbon was burned in a combustion train to produce 0.4362 grams of H2O and 0.8527 grams of CO2. Determine the masses of C and H in the sample and the percentages of these elements in this hydrocarbon. I’ll give you one on your chapter problems.

  31. Combustions Train – Analysis of Hydrocarbons • Compounds that contain C, H, and O. • A 0.3615-gram compound containing on C, H, and O was burned in a combustion train to produce 0.8219 grams of CO2 and 0.3364 grams of H2O. Determine the masses of C, H, and O in the sample and the percentages of these elements in this hydrocarbon.

  32. Determining Molecular Formulas • Percent composition only yields empirical formulas. • Determine the empirical formulas for the last two problems. • In order to determine the molecular formula, both the _______ and the _________ must be known.

  33. Determining Molecular Formulas • The molecular formula is an integer multiple of the simplest formula. • C6H6 and C6H12O6 • Molecular formula = n simplest formula • It follows that; • Molecular weight = n simplest formula weight (rearrange to solve for n) • The molecular weight of the compound in the first problem was determined to be 58.123 amu. What is the molecular formula?

  34. Determining Molecular Formulas • The molecular weight of the compound in the second problem was determined to be 174.240 amu. What is the molecular formula? Have fun!!!

  35. More Fun with Conversions • What mass of ammonium phosphate, (NH4)3PO4 would contain exactly 15.00 grams of N? • What mass of sucrose, C6H12O6, would contain the same mass of oxygen as is contained in 25.1 grams of ammonium phosphate?

  36. Solid Hydrates • Solid hydrates are crystalline solids that contain _______ which can be generally removed by heating. • DEMO: CuSO45H2O The water can be removed by heating. • Every mole of CuSO45H2O contains _____ of water. • An unidentified hydrate of magnesium sulfate is heated, MgSO4xH2O. The mass before heating was 78.30 grams. The new mass after heating is 38.24 grams. What is the formula of the hydrate?

  37. Purity of Samples • Samples are never 100% pure. Percent purity is used to specify the purity of a particular sample. • In many cases, 100 grams of sample can be used for convenience. • A bottle of Na3PO4 is 98.3% pure. What are the masses of Na3PO4 and impurity in 251 grams of sample? • Calculate the number of moles of C2H4Cl2 in 58.2 g of 92.1% pure C2H4Cl2.

  38. Solving Problems

  39. Let’s Do Some More Problems

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