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Chapter 7 Chemical Formulas & Chemical Compounds DHMO Chapter 7 Chemical Formulas and Chemical Compounds Section 1 Chemical Names and Formulas Section 2 Oxidation Numbers Section 3 Using Chemical Formulas Section 4 Determining Chemical Formulas Chapter 7 Section 1 Chemical Names and Formulas Objectives • Explain the significance of a chemical formula. • Determine the formula of an ionic compound formed between two given ions. • Name an ionic compound given its formula. • Using prefixes, name a binary molecular compound from its formula. • Write the formula of a binary molecular compound given its name. • Chemical Formula – shorthand that uses symbols to tell the kind and number of atoms in a molecule or formula unit. Al2(SO4)3 Subscript 2 refers to 2 aluminum atoms Subscript 4 refers to 4 oxygen atoms in the sulfate ion Totals for each atom: Al – 2 atoms S – 3 atoms O – 12 atoms Subscript 3 refers to everything in the parentheses Section 1 Chemical Names and Formulas Chapter 7 Naming Monatomic Ions • Monatomic cations are identified simply by the element’s name. – examples: • K+ is called the potassium cation • Mg2+ is called the magnesium cation • For monatomic anions, the ending of the element’s name is dropped, and the ending ide is added to the root name. – examples: • F– is called the fluoride anion • N3– is called the nitride anion Monatomic Ions – Table 1 pg.221 Chapter 7 Section 1 Chemical Names and Formulas Common Monatomic Ions Chapter 7 Section 1 Chemical Names and Formulas Common Monatomic Ions Why no Roman Numeral For silver or zinc? Note: these are all cations What about anions? Cu2O CuO Cuprous oxide Cupric oxide Naming Monatomic Ions Writing and Naming Binary Ionic Compounds • Mg+2 and Br-1 MgBr2 Magnesium Bromide • Al+3 and O-2 Al2O3 Aluminum Oxide • Fe+2 and O-2 FeO Iron(II) Oxide Ferrous Oxide • Fe+3 and O-2 Fe2O3 Iron(III) Oxide Ferric Oxide Polyatomic Ions – Table 2 pg.226 Compounds with Polyatomic Ions • NH4+1 and Cl-1 NH4Cl Ammonium Chloride • Ba+2 and NO3-1 Ba(NO3)2 Barium Nitrate • Fe+3 and CrO4-2 Fe2(CrO4)3 Iron(III) Chromate Chapter 7 Section 1 Chemical Names and Formulas Naming Binary Ionic Compounds, Compounds Containing Polyatomic Ions, continued Sample Problem Write the formula for tin(IV) sulfate. Chapter 7 Section 1 Chemical Names and Formulas Naming Binary Ionic Compounds, Compounds Containing Polyatomic Ions, continued Sample Problem Solution Write the symbols for the ions side by side. Write the cation first. 4+ Sn 2 4 SO Cross over the charges to give subscripts. Add parentheses around the polyatomic ion if necessary. 4+ 2 2 4 4 Sn (SO ) Chapter 7 Section 1 Chemical Names and Formulas Naming Binary Ionic Compounds, Compounds Containing Polyatomic Ions, continued Sample Problem Solution, continued The total positive charge is 2 4+ = 8+. The total negative charge is 4 2 = 8. The largest common factor of the subscripts is 2, so the smallest whole-number ratio of ions in the compound is 1:2. The correct formula is therefore Sn(SO4)2. Chapter 7 Polyatomic Ions Polyatomic Ions with Multiple Oxygens • • • • ClOClO2ClO3ClO4- Hypochlorite Chlorite Chlorate Perchlorate 2 less oxygens 1 less oxygen Root ion 1more oxygen Root Polyatomic Ions CO3-2 NO3-1 SiO4-4 PO4-3 SO4-2 ClO3-1 AsO4-3 SeO4-2 BrO3-1 TeO4-2 IO3-1 Molecular compounds are… made of just nonmetals smallest piece is a molecule can’t be held together by opposite charge attraction can’t use charges to figure out how many of each atom (there are no charges present) Molecular compounds are easier! Ionic compounds use charges to determine how many of each. –You have to figure out charges. –May need to criss-cross numbers. compounds: the name tells you the number of atoms. Molecular – Uses prefixes to tell you the exact number of each element present! Prefixes for Naming Covalent Compounds Naming Binary Molecular Compounds 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. Mono Di Tri Tetra Penta Hexa Hepta Octa Nona Deca • • • • • • N2O NO NO2 N2O3 N2O4 N2O5 dinitrogen monoxide nitrogen monoxide nitrogen dioxide dinitrogen trioxide dinitrogen tetraoxide dinitrogen pentaoxide Acids are… Compounds that give off hydrogen ions (H1+) when dissolved in water (the Arrhenius definition) Will start the formula with H. There will always be some Hydrogen next to an anion. The anion determines the name. Acids and Salts • An acid is a certain type of molecular compound. Most acids used in the laboratory are either binary acids or oxyacids. – Binary acids are acids that consist of two elements, usually hydrogen and a halogen. – Oxyacids are acids that contain hydrogen, oxygen, and a third element (usually a nonmetal). Acids and Salts, continued • In the laboratory, the term acid usually refers to a solution in water of an acid compound rather than the acid itself. • Many polyatomic ions are produced by the loss of hydrogen ions from oxyacids. sulfuric acid nitric acid phosphoric acid H2SO4 HNO3 H3PO4 sulfate nitrate phosphate 2 4 SO 3 NO 3 4 PO Naming Acids • Binary Acids – HF – HCl – HBr – HI Hydrofluoric Acid Hydrochloric Acid Hydrobromic Acid Hydroiodic Acid Oxyacids • H2SO4 • HNO3 • H3PO4 – HClO4 – HClO3 – HClO2 – HClO Sulfuric Acid Nitric Acid Phosphoric Acid Perchloric Acid Chloric Acid Chlorous Acid Hypochlorous Acid Rules for Naming acids: Name it as a normal compound first 1) If the anion attached to hydrogen ends in -ide, put the prefix hydro- and change -ide to -ic acid • HCl - hydrogen ion and chloride ion = hydrochloric acid • H2S hydrogen ion and sulfide ion = hydrosulfuric acid Naming Acids • If the anion has oxygen in it, then it ends in -ate or -ite 2) change the suffix -ate to -ic acid (use no prefix) • Example: HNO3 Hydrogen and nitrate ions = Nitric acid 3) change the suffix -ite to -ous acid (use no prefix) • Example: HNO2 Hydrogen and nitrite ions = Nitrous acid 2 additional rules 4) If the acid has 1 more oxygen than the –ic acid, add the prefix pera. HClO3 (Hydrogen Chlorate) is chloric acid b. HClO4 would be perchloric acid 5) If there is 1 less oxygen than the -ous acid, add the prefix hypo• HClO2 (Hydrogen Chlorite) is chlorous acid, then HClO would be hypochlorous acid Practice by naming these: HF H3 P H2SO4 H2SO3 Writing Acid Formulas – in reverse! • Hydrogen will be listed first • The name will tell you the anion • Be sure the charges cancel out. • Starts with prefix hydro?- there is no oxygen, -ide ending for anion • no prefix hydro? 1) -ate anion comes from –ic ending 2) -ite anion comes from –ous ending Write formulas for these: hydroiodic acid acetic acid carbonic acid hydrobromic acid Visual Concepts Salt Click below to watch the Visual Concept. Visual Concept Diatomic Elements –elements that are always bonded with some other atom even alone they will bond with themselves H O N Cl Br I F O N Cl Br I F Double Salts – compound that contains two different metal ions with a nonmetal or negative polyatomic ion NHCO3 – Sodium Hydrogen Carbonate KNaSe – Potassium Sodium Selenide Chapter 7 Section 2 Oxidation Numbers Objectives • List the rules for assigning oxidation numbers. • Give the oxidation number for each element in the formula of a chemical compound. • Name binary molecular compounds using oxidation numbers and the Stock system. Oxidation Numbers • The charges on the ions in an ionic compound reflect the electron distribution of the compound. • In order to indicate the general distribution of electrons among the bonded atoms in a molecular compound or a polyatomic ion, oxidation numbers are assigned to the atoms composing the compound or ion. • Unlike ionic charges, oxidation numbers do not have an exact physical meaning: rather, they serve as useful “bookkeeping” devices to help keep track of electrons. Text’s Rules for Assigning Oxidation Numbers • In general when assigning oxidation numbers, shared electrons are assumed to “belong” to the more electronegative atom in each bond. • More-specific rules are provided by the following guidelines. 1. The atoms in a pure element have an oxidation number of zero. examples: all atoms in sodium, Na, oxygen, O2, phosphorus, P4, and sulfur, S8, have oxidation numbers of zero. 2. The more-electronegative element in a binary compound is assigned a negative number equal to the charge it would have as an anion. Likewise for the less-electronegative element. 3. Fluorine has an oxidation number of –1 in all of its compounds because it is the most electronegative element. 4. Oxygen usually has an oxidation number of –2. Exceptions: • In peroxides, such as H2O2, oxygen’s oxidation number is –1. • In compounds with fluorine, such as OF2, oxygen’s oxidation number is +2. 5. Hydrogen has an oxidation number of +1 in all compounds containing elements that are more electronegative than it; it has an oxidation number of –1 with metals. 6. The algebraic sum of the oxidation numbers of all atoms in an neutral compound is equal to zero. 7. The algebraic sum of the oxidation numbers of all atoms in a polyatomic ion is equal to the charge of the ion. 8. Although rules 1 through 7 apply to covalently bonded atoms, oxidation numbers can also be applied to atoms in ionic compounds similarly. Rules for Oxidation Numbers 1. The oxidation number of an element in its elemental form is zero. Examples of this are N2 (g), O2 (g), Na (s), Cl2 (g), etc. 2. The oxidation number of a monatomic ion is exactly the same as its charge. So, Group IA ions will all have an oxidation number of +1, since they all lose one electron. Group IIA ions will all have an oxidation number of +2. Aluminum ions only exist as Al+3 and will have an oxidation number of +3. 3. The oxidation state of oxygen, in most compounds is -2, i.e., it tends to pull 2 shared electrons toward itself. The exceptions are H2O2, hydrogen peroxide and O2-2, peroxide, when it is -1. In O2 it is zero (see rule 1.). 4. The oxidation state of hydrogen is almost always +1. The exceptions are H2 (oxidation number = zero , rule 1.) and when hydrogen is bonded to metals in binary compounds, like LiH, when it is -1. 5. Fluorine is always -1. The other halogens are -1, except when bonded to oxygen (rule 3. gives oxygen an oxidation number of -2, making halogens bonded to oxygen positive). 6. Oxidation number of the atoms in a compound must add up to the total charge on that molecule or ion. Rules for Assigning Oxidation Numbers Click below to watch the Visual Concept. Visual Concept NaNO3 Na = +1 O = -2 N=? Chapter 7 Section 2 Oxidation Numbers Assigning Oxidation Numbers, continued Sample Problem E Assign oxidation numbers to each atom in the following compounds or ions: a. UF6 b. H2SO4 c. ClO3 Chapter 7 Section 2 Oxidation Numbers Assigning Oxidation Numbers, continued Sample Problem E Solution a. Place known oxidation numbers above –1 the appropriate elements. UF6 Multiply known oxidation numbers by the appropriate number of atoms and place the totals underneath the –1 corresponding elements. UF6 –6 Chapter 7 Section 2 Oxidation Numbers Assigning Oxidation Numbers, continued Sample Problem E Solution, continued The compound UF6 is molecular. The sum of the oxidation numbers must equal zero; therefore, the total of positive oxidation numbers is +6. –1 UF6 +6 – 6 Divide the total calculated oxidation number by the appropriate number of atoms. There is only one uranium atom in the molecule, so it must have an oxidation number of +6. +6 – 1 UF6 +6 – 6 Chapter 7 Section 2 Oxidation Numbers Assigning Oxidation Numbers, continued Sample Problem E Solution, continued b. Hydrogen has an oxidation number of +1. Oxygen has an oxidation number of 2. The sum of the oxidation numbers must equal zero, and there is only one sulfur atom in each molecule of H2SO4. Because (+2) + (8) = 6, the oxidation number of each sulfur atom must be +6. 1 6 2 H2 SO4 2 6 8 Chapter 7 Section 2 Oxidation Numbers Assigning Oxidation Numbers, continued Sample Problem E Solution, continued c. The total of the oxidation numbers should equal the overall charge of the anion, 1. The oxidation number of a single oxygen atom in the ion is 2. The total oxidation number due to the three oxygen atoms is 6. For the chlorate ion to have a 1 charge, chlorine must be assigned an oxidation number of +5. +5 2 ClO°V 3 +5 6 Chapter 7 Section 2 Oxidation Numbers Using Oxidation Numbers for Formulas and Names • As shown in the table in the next slide, many nonmetals can have more than one oxidation number. • These numbers can sometimes be used in the same manner as ionic charges to determine formulas. – example: What is the formula of a binary compound formed between sulfur and oxygen? From the common +4 and +6 oxidation states of sulfur, you could predict that sulfur might form SO2 or SO3. Both are known compounds. Chapter 7 Section 2 Oxidation Numbers Common Oxidation States of Nonmetals Chapter 7 Section 2 Oxidation Numbers Using Oxidation Numbers for Formulas and Names, continued • Using oxidation numbers, the Stock system, introduced in the previous section for naming ionic compounds, can be used as an alternative to the prefix system for naming binary molecular compounds. Homework Pages 251-254 Numbers for first half/ first quiz 4,6,7,10,11,14,15,16,18,21,23,24,25 Numbers for second half/ second quiz 28,31,32,36,38,39,40,44,50