Week 4 and 5_Molecules and Compounds CHEM1010 PDF

CHEM 1010 General Chemistry 1 Week 4 and 5: Molecules and Compounds  Molecules and Compounds Objectives: 1. Write chemical formulas and names for ionic compounds 2. Write chemical formulas and names for hydrated ionic compounds 3. Write chemical formulas and nam...

CHEM 1010 General Chemistry 1 Week 4 and 5: Molecules and Compounds  Molecules and Compounds Objectives: 1. Write chemical formulas and names for ionic compounds 2. Write chemical formulas and names for hydrated ionic compounds 3. Write chemical formulas and names for molecular compounds 4. Write chemical formulas and names for acids and oxyacids 5. Calculate molar masses of compounds and use Avogadro’s number to conduct mass to mole conversions. 6. Calculate the mass percent of an atom in a given compounds. 2  3.1 Compounds  Remember: COMPOUNDS are chemicals made of 2 or more elements  There are different types of compounds o Ionic & molecular compounds vs. Salt Sugar 3  3.1 Hydrogen, Oxygen, and Water When two or more elements combine to form a compound, an entirely new substance results 4 Definite Proportion (recall the law of definite proportion) 5  3.2 Chemical Bonds  Compounds are composed of atoms held together by chemical bonds  Chemical bonds may result from the attractions between two charged species (cation and anion)  Two broad classes of bonds – ionic and covalent  So, there are 2 types of compounds: Salt o IONIC compounds made of a metal and a nonmetal o COVALENT compounds Also called molecular compounds Sugar Made of only nonmetals 6  Ionic Bonds Ionic bonds involve the complete transfer of electrons from the metal atom to the nonmetal atom The metal atom becomes a cation while the nonmetal atom becomes an anion These oppositely charged ions attract one another by electrostatic forces and form an ionic bond 7  8 The Formation of Ionic Compounds In the solid phase, the ionic compound is composed of a lattice—a regular three-dimensional array of alternating cations and anions. 8  9 Covalent Bonds Covalent bonds occur between two or more nonmetals The two atoms share electrons between them, forming a molecule Covalently bonded compounds are also called molecular compounds 9  10 What type of bond—ionic or covalent—forms between nitrogen and oxygen? a. Ionic b. Covalent 10 3.3 Representing Compounds: Chemical 11 Formulas and Molecular Models A compound’s chemical formula indicates the elements present in the compound and the relative number of atoms or ions of each For example: Water is represented as H2O Sodium Chloride is represented as NaCl Carbon dioxide is represented as CO2 Carbon tetrachloride is represented as CCl4 11  12 Types of Chemical Formulas Chemical formulas can generally be categorized into three different types: empirical formulas molecular formulas structural formulas Ball-and-stick model 12  13 Types of Chemical Formulas An empirical formula gives the relative number of atoms of each element in a compound A molecular formula gives the actual number of atoms of each element in the molecule of a compound For example: a. For H2O2, the greatest common factor is 2. The empirical formula is therefore HO. b. For B2H6, the greatest common factor is 2. The empirical formula is therefore BH3. c. For CCl4, the only common factor is 1, so the empirical formula and the molecular formula are the same 13  14 Types of Chemical Formulas Molecular formula shows the atoms in the molecule, but they do not show how they are attached. For example: C3H6 H2O (Propene) (Water) 14  15 Types of Chemical Formulas A structural formula uses lines to represent covalent bonds and shows how atoms in a molecule are connected or bonded to each other It can also show the molecule’s geometry The structural formula for H2O2 can be shown as either of the following 15  Types of Chemical Formulas 16 Models show you the attachment pattern but also give you an idea about the molecule's shape. 16  17 Types of Chemical Formulas The type of formula we use depends on how much we know about the compound and how much we want to communicate A structural formula communicates the most information. An empirical formula communicates the least Example: Select the structural formula for water (H2O). a. H — O b. H — H c. H — O — H d. H2O 17  18 3.4 An Atomic-Level View of Elements and Compounds Elements may be either atomic or molecular Compounds may be either molecular or ionic 18  19 View of Elements and Compounds Atomic elements exist with single atoms as their basic units Most elements fall into this category Examples include Na, Ne, K, Mg, etc. Molecular elements do not normally exist in nature with single atoms as their basic units; they exist as molecules (2 or more atoms of the element bonded together) There are only seven diatomic elements (H2, N2, O2, F2, Cl2, Br2, and I2) Polyatomic elements: Molecular phosphorus (P4), Molecular sulfur (S8). 19  Conceptual Connection 3.4 20 Classify the substance represented by the molecular view shown here. a. Atomic element b. Molecular element c. Molecular compound d. Ionic compound 20  21 Ionic Compounds Ionic compounds are composed of cations (usually a metal) and anions (usually one or more nonmetals) bound together by ionic bonds. The basic unit of an ionic compound is the formula unit, the smallest, electrically neutral collection of ions. Table salt is an ionic compound with the formula unit NaCl, which is composed of Na+ and Cl– ions in a one-to-one ratio. 21  22 Molecular Compounds Molecular compounds are usually composed of two or more covalently bonded nonmetals The basic units of molecular compounds are molecules composed of the constituent atoms Water is composed of H2O molecules Dry ice is composed of CO2 molecules Propane (often used as a fuel for grills) is composed of C3H8 molecules 22  23 Which statement best summarizes the difference between ionic and molecular compounds? a. Molecular compounds contain highly directional covalent bonds, which result in the formation of molecules. Ionic compounds contain nondirectional ionic bonds, which result (in the solid state) in the formation of ionic lattices. b. Molecular compounds and ionic compounds both contain molecules as their smallest identifiable unit, but in ionic compounds the molecules are smaller. c. A molecular compound is composed of covalently bonded molecules. An ionic compound is composed of ionically bonded molecules (in the solid phase). 23  24 3.5 Ionic Compounds: Formulas and Names Ionic Compound Formulas  Ionic compounds always contain positive and negative ions  In a chemical formula, the sum of the charges of the cations must equal the sum of the charges of the anions  The formula of an ionic compound reflects the smallest whole- number ratio of ions 24  Ionic Compounds The charges of the main group elements can be predicted from their group numbers. − The representative elements form only one type of charge − Note: If the anion is a polyatomic ion use the table of common polyatomic ion in the Chemistry data book to predict the anion 25  26 3.5 Ionic Compounds: Formulas and Names Ionic Compound Formulas  Transition metals tend to form different charges (see the periodic table in the chemistry databook, e.g. Fe2+, Fe3+; Mn4+, Mn7+) Only 1 charge listed. More than 1 charge listed. We must use math to find the charge of Fe in the compound No charge listed – Alkali metals always have a charge of +1 26  27 Naming Ionic Compounds Ionic compounds can be categorized into two types, depending on the metal in the compound 27 Metals Whose Charge Is Invariant from 28 One Compound to Another 28  29 Which metal has the same charge in all of its compounds? a. Fe b. Mo c. Pb d. Sr 29  30 Naming Binary Ionic Compounds of Type I Cations Binary compounds contain only two different elements The names of binary ionic compounds take the following form: 1- Cations are named first, then anions. 2- Cation has same name of its element without any change. 3- Use –ide at the end of the anion name. 30  31 Examples: Type I Binary Ionic Compounds 1- Cations are named first, then anions. 2- Cation has same name of its element Ions without any change. Compound Name present 3- Use –ide at the end of the anion name. potassium KCl K+, Cl- chlroride calcium CaS Ca2+, S2- sulfide aluminum AlP Al3+, P3- phosphide barium BaI2 Ba2+, I- iodide Ag2O Ag+, O2- silver oxide lithium Li3N Li+, N3- nitride aluminum Al2O3 Al3+, O2- oxide 31  32 Naming Type II Ionic Compounds These contain metals that can form more than one kind of cation So, the metal’s charge must be specified in the name The proportion of metal cation to nonmetal anion helps us determine the charge on the metal ion 32  33 Naming Type II Ionic Compounds Iron, for instance, forms 2+ and 3+ cations Metals of this type are often transition metals FeS: here, iron is +2 cation (Fe2+) Fe2S3: here, iron is +3 cation (Fe3+) Cu2O: here, copper is +1 cation (Cu+) CuO: here, copper is +2 cation (Cu2+) Some main group metals, such as Pb and Sn, form more than one type of cation 33  34 Naming Type II Ionic Compounds The full name of compounds containing metals that form more than one kind of cation have the following form: The charge of the metal cation can be determined by inference from the sum of the charges of the nonmetal 34  35 Naming Type II Ionic Compounds 1- Cations are named first (without any change) 2- Indicate the charge in Roman numerals in parenthesis 3- Use –ide at the end of the anion name.  Examples: CuCl Copper(I) chloride CuCl2 Copper(II) chloride CoCl3 Cobalt(III) chloride HgI2 Mercury(II) iodide PbO2 Lead(IV) oxide PbCl4 Lead(IV) chloride 35  36 Naming Type II Ionic Compounds Example: To name CrBr3,determine the charge on the chromium. Total charge on cation + total anion charge = 0. Cr charge + 3(Br− charge) = 0. Since each Br has a −1 charge, then: Cr charge + 3(−1) = 0 Cr charge −3 = 0 Cr = +3 Hence, the cation Cr3+ is called chromium(III), while Br− is called bromide Therefore, CrBr3 is chromium(III) bromide. 36  37 Polyatomic Ions Many common ionic compounds contain polyatomic ions These are composed of a group of covalently bonded atoms with an overall charge (see chemistry data booklet) Examples: NaNO3 contains Na+ and NO3− CaCO3 contains Ca2+ and CO32− Mg(ClO3)2 contains Mg2+ and ClO3− Named similarly, except we use the name of polyatomic ion when necessary NaNO3 (sodium nitrate) CaCO3 (calcium carbonate) Mg(ClO3)2 (magnesium chlorate) 37  38 38  39 Oxyanions Most polyatomic ions are oxyanions, anions containing oxygen plus another element When a series of oxyanions contains different numbers of oxygen atoms, they are named according to the number of oxygen atoms in the ion If there are two ions in the series: the one with more oxygen atoms has the ending -ate the one with fewer has the ending -ite Examples NO3– is nitrate SO42– is sulfate NO2– is nitrite SO32– is sulfite 39  40 Oxyanions If there are more than two ions in the series, then the prefixes hypo-, meaning less than, and per-, meaning more than, are used ClO− hypochlorite BrO– hypobromite ClO2– chlorite BrO2– bromite ClO3– chlorate BrO3– bromate ClO4– perchlorate BrO4– perbromate 40  41 Identify the polyatomic ion and its charge in each compound: KNO2, CaSO4, Mg(NO3)2. a. NO2−, SO42−, and NO3− b. K+, Ca2+, and Mg2+ c. K+, Ca2+, Mg2+, NO2−, SO42−, and NO3− d. NO22−, SO4−, and NO32− 41  Learning Check Write the formulae for the following ionic compounds: a) potassium thiocyanate KSCN b) silver sulfate Ag2SO4 c) copper (II) chloride CuCl2 d) barium hydroxide Ba(OH)2 42  Learning Check Write the names of the following compounds: a) NaF sodium fluoride b) K2O potassium oxide c) CoO cobalt (II) oxide d) CrF2 chromium (II) fluoride (NH4)2S ammonium sulfide e) KBrO3 potassium bromate f) Fe(ClO2)3 iron (III) chlorite g) Li2Cr2O7 lithium dichromate 43  44 Hydrated Ionic Compounds Hydrates are ionic compounds containing a specific number of water molecules associated with each formula unit For example, the formula for epsom salts is MgSO4 ⋅ 7H2O. Its systematic name is magnesium sulfate heptahydrate CoCl2 ⋅ 6H2O is cobalt (II) chloride hexahydrate 44  45 Common Hydrate Prefixes Common hydrate prefixes Hemi = 1/2 tetra = 4 hepta = 7 mono = 1 penta = 5 octa = 8 di = 2 hexa = 6 nona = 9 tri = 3 Deca = 10 Other common hydrated ionic compounds and their names are as follows: CaSO4 ⋅ ½ H2O is called calcium sulfate hemihydrate BaCl2 ⋅ 6H2O is called barium chloride hexahydrate CuSO4 ⋅ 6H2O is called copper sulfate hexahydrate 45  Learning Check  Write the name for the following hydrates: o MgSO4. 5 H2O magnesium sulfate pentahydrate o CoCl2. 6 H2O cobalt (II) chloride hexahydrate  Write the formula for the following hydrates: o sodium carbonate decahydrate Na2CO3. 10 H2O o iron (III) phosphate tetrahydrate FePO4. 4 H2O 46  47 3.6 Molecular Compounds: Formulas and Names Unlike ionic compounds, the formula for a molecular compound cannot readily be determined from its constituent elements The same combination of elements may form many different molecular compounds, each with a different formula Nitrogen and oxygen form all of the following unique molecular compounds: NO, NO2, N2O, N2O3, N2O4, and N2O5 47  48 Molecular Compounds: Formulas and Names Molecular compounds are composed of two or more nonmetals Generally, write the name of the element with the smallest group number first If the two elements lie in the same group, then write the element with the greatest row number first − The prefixes given to each element indicate the number of atoms present 48  49 Molecular Elements Phosphorus (P4) Sulfur (S8) 49  50 Binary Molecular Compounds These prefixes are the same as those used in naming hydrates: mono = 1 hexa = 6 di = 2 hepta = 7 tri = 3 octa = 8 tetra = 4 nona = 9 penta = 5 deca = 10 If there is only one atom of the first element in the formula, the prefix mono- is normally omitted. 50  51 Binary Molecular Compounds Examples  HCl Hydrogen chloride  SiC Silicon carbide  NO Nitrogen monoxide  N2O Dinitrogen monoxide Only used  N2O5 Dinitrogen pentoxide for the  SO2 Sulfur dioxide second element  SO3 Sulfur trioxide  CBr4 Carbon tetrabromide  PCl5 Phosphorus pentachloride  SF6 Sulfur hexafluoride 51  52 The compound NCl3 is nitrogen trichloride, but AlCl3 is simply aluminum chloride. Why? a. The name forms differ because NCl3 is an ionic compound and AlCl3 is a molecular compound. Prefixes such as mono-, di-, and tri- are used for ionic compounds but not for molecular compounds. b. The name forms differ because NCl3 is a molecular compound and AlC l3 is an ionic compound. Prefixes such as mono-, di-, and tri- are used for molecular compounds but not for ionic compounds. 52  Some compounds have TRIVIAL names 53 that must be memorized  These names are used more often than the systematic names o H2O – water o CH4 – methane o NH3 – ammonia o C2H6 – ethane o PH3 – phosphine o CH3OH – methanol o SiH4 – silane o C2H5OH – ethanol o O3 – ozone o C6H12O6 – glucose o NO – nitric oxide o C12H22O11 – sucrose (sugar) o N2O – nitrous oxide o H2O2 – hydrogen peroxide 53  Learning Check Complete the table below. Compound Name Compound Formula diphosphorus tetroxide P 2O4 nitrogen trichloride NCl3 nitric oxide NO methanol CH3OH ammonia NH3 54  Acids  Acids are molecular compounds that release hydrogen ions (H+) when dissolved in water o Must be aqueous (aq) Examples: HCl (aq)  H+ + Cl- H2SO4 (aq)  H+ + HSO4-  2 H+ + SO42- 55  56 Acids Sour taste Dissolve many metals − such as Zn, Fe, and Mg; but not Au, Ag, or Pt Acids are composed of hydrogen, usually written first in their formulas, and one or more nonmetals, written second 56  57 Acids Binary acids have H+ cation and nonmetal anion [e.g., HCl(aq)] Oxyacids have H+ cation and polyatomic anion [e.g. , HNO3(aq) contains the nitrate (NO3- )] 57  58 Naming Binary Acids Write a hydro- prefix. Follow with the nonmetal base name. Add -ic. Write the word acid at the end of the name.  Example: o HCl (aq) hydrochloric acid 58  59 Naming Oxyacids If the polyatomic ion name ends in -ate, change ending to -ic. If the polyatomic ion name ends in -ite, change ending to -ous. Write word acid at the end of all names. oxyanions ending with -ate  Example: H3PO4 (aq) phosphoric acid oxyanions ending with -ite  Example: H2SO3 (aq) sulfurous acid 59  60 Practice: Name the Acid 1. H2S hydrosulfuric acid 2. HClO3 chloric acid 3. HC2H3O2 acetic acid 60  Learning Check  Write the names these acids: a) HIO2 (aq) iodous acid b) HCH3COO (aq) acetic acid c) H2S (aq) hydrosulfuric acid  Write the formulas of these acids: a) hydrobromic acid HBr (aq) b) borous acid H3BO2 (aq) c) thiosulfuric acid H2S2O3 (aq) 61  62 Inorganic Nomenclature Flowchart 62  Atomic masses: The carbon-12 scale The mass of an atom is a small quantity. – Hydrogen atom 1.67 x 10-24 g – Oxygen atom 2.66 x 10-23 g – Carbon atom 2.00 x 10-23 g – Iron atom 9.27 x 10-23 g 63  Atomic masses: The carbon-12 scale Because it is inconvenient to use these masses, a system of relative masses of atoms has been devised. The scale is based on an arbitrarily assigned value of exactly 12 atomic mass unit (amu) for the carbon-12 atom. This mean that an atom twice as heavy as a C-12 atom would have a mass of 24 amu, an atom half as heavy as C- 12 atom would weight 6 amu. 1 amu = 1.66054 x 10-24 g 64  Atomic Mass ATOMIC MASS is the mass of 1 atom –Because atoms are so small, we measure the mass in amu (atomic mass unit) You can find an element’s atomic mass on the periodic table: The atomic mass of carbon is 12.0107 amu 65  Atomic Mass In real life, we do not work with individual atoms The unit “gram” (g) is much more helpful 1 g = 6.022 x 1023 amu 66  Molar Mass MOLAR MASS is the mass (in g) for 1 mole of a chemical We can also find this on the periodic table The molar mass of carbon is 12.0107 g/mol 67  Expressing Mass The atomic mass of Na is 22.989769 amu. The atomic mass (AM) of Na is 22.989769 g/mol. 68  69 Formula Mass The mass of an individual molecule or formula unit also known as molecular mass (MM) Sum of the masses of the atoms in a single molecule or formula unit whole = sum of the parts! Mass of 1 molecule of H2O = (2 atoms H) (1.01 amu / H atom ) + (1 atom O )(16.00 amu / atom O ) = 18.02 amu 69  70 Molar Mass (MM) of Compounds The molar mass of a compound—the mass, in grams, of 1 mol of its molecules or formula units—is numerically equivalent to its formula mass with units of gram/mol. 70  71 Molar Mass of Compounds The relative masses of molecules can be calculated from atomic masses: formula mass = 1 molecule of H2O = 2 (1.01 amu H) + 16.00 amu O = 18.02 amu 1 mole of H2O contains 2 moles of H and 1 mole of O: molar mass = 1 mole H2O 2 mol (1.01 g / 1 mol H) + 1 mol (16.00 g / 1 mol O ) = 18.02 g / 1 mol H2O So the molar mass of H2O is 18.02 g /mole. ram Molar mass = formula mass (in g /mole) ram 71  Concept Check  Example: Find the molar mass of Li2O. MM = 2 Li + 1 O = 2 (6.941 g/mol) + 15.9994 g/mol = 29.8814 g/mol  Example: Find the molar mass of Mg(NO3)2 MM = Mg + 2 N + 6 O = 24.3050 + 2(14.0067) + 6(15.9994) = 148.3148 g/mol 72  How many atoms are in exactly 12.0107 g of carbon? MOLE: 1 mole = 6.022 x 1023 particles AVOGADRO’S CONSTANT (NA) NOTE: PARTICLES can mean atoms, ions, formula units, molecules, or any other small particle. 73  Mole From the definition of a mole, we can calculate the number of particles in any sample: 𝐍𝐍 = 𝐧𝐧 × 𝐍𝐍𝐀𝐀 Number of Avogadro’s particles Number of moles Number 74  Mole Example: How many atoms are in 3.74 mol of Fe? N = n x NA N = (3.74 mol)(6.022 x 1023 atoms/mol) N = 2.252228 x 1024 atoms 75  Mole Example: How many moles are in 7.17 x 1024 molecules of CO2? N = n x NA n = (7.17 x 1024 molecules)/(6.022 x 1023 molecule/mol) n = 11.90634341 mol 76  Moles Because molar mass is measured in g/mol, we can calculate the moles of any mass with this formula: 𝐦𝐦 = 𝐧𝐧 × 𝐌𝐌𝐌𝐌 Mass (g) Molar mass (g/mol) Number of moles (mol) 77  Moles Example: 0.150 mol of O2 (g) (MMO2 = 31.9988 g/mol) will have what mass? 𝐦𝐦 = 𝐧𝐧 × 𝐌𝐌𝐌𝐌 𝐦𝐦 = 𝟎𝟎. 𝟏𝟏𝟏𝟏𝟏𝟏 𝐦𝐦𝐦𝐦𝐦𝐦 (𝟑𝟑𝟑𝟑. 𝟗𝟗𝟗𝟗𝟗𝟗𝟗𝟗 𝐠𝐠/𝐦𝐦𝐦𝐦𝐦𝐦) 𝐦𝐦 = 𝟒𝟒. 𝟕𝟕𝟕𝟕𝟗𝟗𝟗𝟗𝟗𝟗 𝐠𝐠 78  Moles We now know 2 formulas with moles. 𝐍𝐍 = 𝐧𝐧 × 𝐍𝐍𝐀𝐀 𝐦𝐦 = 𝐧𝐧 × 𝐌𝐌𝐌𝐌 We can now do two-step problems. Particles Atom, molecule, formula unit Mass MOLE 79  Concept Check  Find the mass of each sample: a) 0.2500 mol of CuCl2 (MM = 134.452 g/mol) 33.613 g = 33.61 g b) 1.53 mol of Au 301.3588521 g = 301 g 80  Concept Check  Find the number of moles in each sample: a) 150 g of Co(NO3)2 (MM = 182.9430 g/mol) 0.819927518 mol = 0.820 mol b) 25 g of ZnO MM = 81.4084 g/mol n = 0.307093916 mol = 0.30 mol 81  82 Composition of Compounds A chemical formula, in combination with the molar masses of its constituent elements, indicates the relative quantities of each element in a compound. Mass percentage of each element in a compound can be determined from the formula of the compound (number of atoms element X)(atomic mass of element X) Mass % Element X = x 100% (MM of the compound) 82  83 Learning Check Calculate the mass percent composition of Cl in CCl4? Molar mass of CCl4 = 12.0107 g/mol + 4(35.453 g/mol) = 153.8227 g/mol 4 ∗ molar mass Cl 𝐌𝐌𝐌𝐌𝐌𝐌𝐌𝐌 % 𝐨𝐨𝐨𝐨 𝐂𝐂𝐂𝐂 = ∗ 𝟏𝟏𝟏𝟏𝟏𝟏 % molar mass CCl4 4 ∗ 35.453 g /mol 𝐌𝐌𝐌𝐌𝐌𝐌𝐌𝐌 % 𝐨𝐨𝐨𝐨 𝐂𝐂𝐂𝐂 = ∗ 𝟏𝟏𝟏𝟏𝟏𝟏 % 153.8227 g /mol = 92.19185468 % 83  84 Learning Check Calculate the percentage of carbon in ethane (C2H6)? (2)(12.0 amu) %C = x 100 (30.0 amu) 24.0 amu = x 100 30.0 amu %C = 80.0% 84  85 Concept Check Without doing any calculations, list the elements in C6H6O in order of decreasing mass percent composition. a. C>O>H b. O>C>H c. H>O>C d. C>H>O 85  86 Conversion Factors from Chemical Formulas Chemical formulas show the relationship between numbers of atoms and molecules. − Or moles of atoms and molecules 58.64 g Cl : 100 g CCl2F2 1 mol CCl2F2 : 2 mol Cl These relationships can be used to determine the amounts of constituent elements and molecules. − Like percent composition 86  87 Concept Check The molecular formula for water is H2O. Which ratio can be correctly derived from this formula? Explain a. 2 g H : 1 g H2O b. 2 mL H : 1 mL H2O c. 2 mol H : 1 mol H2O 87  Pearson Platform and eBook Readings/Exercises: Week 4 and 5. Molecules and Compounds Readings: Chapter 3.1, 3.2, 3.5, 3.6, 3.8, 3.9 Pearson Platform Tutorial 3 (Practice Questions) Note: Tutorial questions will be available soon on the Pearson Platform 88