Naming and Writing Formulas for Ionic Compounds

Chemistry

CH 8 Tutorial

Naming and Writing Formulas for Ionic Compounds I. Binary Ionic Compounds A. Binary Compounds 1. Compounds composed of two different elements B. Naming Binary Ionic Compounds from their formulas 1. Name the cation 2. Name the anion C. The Stock System of Nomenclature 1. Roman numerals are used to denote the charge of metals that can form two or more cations. 2. The numeral is enclosed in parentheses and placed immediately after the metal name a. Iron (II) and Iron (III), pronounced “iron two” and “iron three” 3. Roman numerals are never used: a. For anions b. For metals that form only one ion D. Writing Formulas for Binary Ionic Compounds 1. Write the symbols for the ions side by side. ALWAYS write the cation first! 2. Cross over the charges by using the absolute value of each ion’s charge as the subscript for the other ion 3. Check that the subscripts are in smallest whole number ratio E. The Stock System of Nomenclature 1. Roman numerals are used to denote the charge of metals that can form two or more cations. 2. The numeral is enclosed in parentheses and placed immediately after the metal name a. Iron( II) and Iron (III), pronounced “iron two” and “iron three” 3. Roman numerals are never used: a. For anions b. For metals that form only one ion F. Naming compounds containing polyatomic ions a. Same as for monatomic ions G. Writing formulas including polyatomic ions a. Use parentheses when you need MORE THAN one of a polyatomic ion b. Parentheses are NEVER used for monatomic ions, regardless of how many are in the formula

Naming and Writing Formulas for Molecular Compounds I. Naming Binary Molecular Compounds A. Binary Molecular Compounds 1. Covalently bonded molecules containing only two elements, both nonmetals B. Naming 1. Least electronegative element is named first 2. First element gets a prefix if there is more than 1 atom of that element 3. Second element ALWAYS gets a prefix, and an “-ide” ending Examples: N2O3 = dinitrogen trioxide CO = carbon monoxide, not monocarbon monoxide Numerical Prefixes Number 1 2 3 4 5 6 7 8 9 10 Prefix mono di tri tetra penta hexa hepta octa nona deca Naming and Writing Formulas for Acids and Bases I. Naming Acids A. Binary Acids 1. Acids that consist of two elements, usually hydrogen and one of the halogens HCl Use Hydro the element –ic Hydrochloric B. Oxyacids 1. Acids that contain hydrogen, oxygen and a third element (usually a nonmetal) C. Naming Acids 1. Poly ends in ate, change to- ic; poly ends in ite, change to –ous

HNO3 Nitrate Nitric Acid HNO2 Nitrite Nitrous Acid Describing Chemical Reactions I. Introduction A. Reactants 1. Original substances entering into a chemical rxn B. Products 1. The resulting substances from a chemical rxn Reactants -> Products C. Chemical Equation 1. Represents with symbols and formulas, the identities and relative amounts of the reactants and products in a chemical rxn II. Writing Chemical Equations A. The equation must represent known facts 1. This can be done with a word equation: "hydrogen reacts with oxygen to form water" Hydrogen + Oxygen  Water B. The equation must contain the correct formulas for reactants and products 1. This is done with a formula equation

H2 + O2  H2O C. The law of conservation of atoms (mass) must be satisfied (mass of reactants = mass of products) 1. Balancing is done with coefficients - small whole numbers that appear in front of a formula

2H2 + O2  2H2O (Also represent MOLES) D. Additional symbols used in chemical equations (balanced equation) Symbols Used in Chemical Equations Symbol Explanation + Used to separate two reactants or products  "Yields," separates reactants from products <-> Used in place of a single arrow to indicate a reversible reactions (s) Reactant or product in the solid state. Also a precipitate (ppt) (l) Reactant or product in the liquid state. (aq) Reactant or product in an aqueous solution (dissolved in water) (g) Reactant or product in the gaseous state Δ Reactants are heated -Pt---> A formula written above or below the yield sign indicates its use as a catalyst (in this case, platinum) -2atm---> Pressure at which rxn is carried out exceeds normal atmospheric pressure --25 C--> Temperature at which the rxn is carried out, in this case 25 °C --MnO2--->Formula of catalyst, in this case manganese dioxide, used to alter the rate of the reaction III. Balancing Chemical Equations A. Identify the names of reactants and products, and write a word equation B. Write a formula equation by substituting correct formulas for the names of the reactants and the products C. Balance the formula equation according to the law of conservation of atoms D. Count atoms to be sure that the equation is balanced CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g) Carbon tetra hydride gas reacts with 2 molecules of oxygen gas to produce carbon dioxide gas and 2 molecules of steam.

IV. Evidence that a chemical reaction has occurred. A. Change in temperature B. Production of light C. Production of gas D. Formation of a precipitate (ppt)- a solid appears (turns cloudy) E. Color change

Types of Chemical Reactions I. Combination Reactions (Synthesis Rxns) A. Two or more substances combine to form a more complex substance A + X  AX (ONE PRODUCT) B. Types of Synthesis Rxns 1. Metals react with oxygen to form oxides

4Al(s) + 3O2(g) 2Al2O3(s) 2. Nonmetals react with oxygen to form oxides

C(s) + O2(g)  CO2(g) 3. Metals react with halogens to form salts (halogen means "salt maker")

2Na(s) + Cl2(g)  2NaCl(s) 4. Active metal oxides react with water to form metallic hydroxides

MgO(s) + H2O(l)  Mg(OH)2(s) 5. Nonmetal oxides react with water to form oxyacids (acid rain)

SO2(g) + H2O  H2SO3(aq)

II. Decomposition Reactions A. Decomposition Rxns 1. One substance breaks down to form two or more simpler substances AX  A + X (ONE REACTANT) B Kinds of Decomposition Rxns MEMORIZE 1. Metallic carbonates, when heated, form metallic oxides and carbon dioxide

CaCO3(s)  CaO(s) + CO2(g) 2. Metallic hydroxides , when heated, decompose into metallic oxides and water

Ca(OH)2(s)  CaO(s) + H2O(g) 3. Metallic chlorates, when heated, decompose into metallic chlorides and oxygen

2KClO3(s)  2KCl(s) + 3O2(g) 4. Some acids, when heated, decompose into nonmetallic oxides and water

H2SO4(aq)  H2O(l) + SO3 (g) 5. Binary compounds

2NaCl(s)  2Na(s) + Cl2 (g)

III. Single-Replacement Reactions A. Single-Replacement Rxns (ELEMENT + COMPOUND) 1. One substance is replaced in its compound by another substance A + BX  AX + B (cation replaces cation) Y + BX  BY + X (anion replaces anion)

Zn(s) + CuSO4(aq) ZnSO4(aq) + Cu(s) + - Ca(s) + 2H2O(l)  Ca(OH)2(aq) + H2 (g) (H2O is H and OH )

Cl2(g) + 2KBr(aq)  2KCl(aq) + Br2 (g) B. Activity Series 1. A list of elements organized according to the ease with which the elements undergo certain chemical rxns 2. Each element in the list displaces from a compound any of the elements below it. The larger the interval between elements in a series, the more vigorous the replacement rxn. 3. Metals may replace other metals 4. Halogens may replace other halogens IV. Double-Replacement Reactions A. Double-Replacement Rxn 1. The ions of two compounds exchange places in an aqueous solution to form two new compounds AB + CD  AD + CB (COMPOUND + COMPOUND) B. Types of Double-Replacement Rxns

BaCl2(aq) + Na2SO4(aq)  2NaCl(aq) + BaSO4(s)

FeS(aq) + H2SO4(aq)  FeSO4(aq) + H2S(g)

NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(l) C. Use Solubility Table to see if a reaction will occur, will a precipitate form. I. Predicting the Formation of a Precipitate A. Solubility Rules 1. No compound is completely insoluble 2. Compounds of very low solubility can be considered insoluble 3. Dissociation equations cannot be written for insoluble compounds General Solubility Guidelines 1. Most sodium, potassium, and ammonium compounds are soluble in water (will not form ppt) 2. Most nitrates, acetates, and chlorates are soluble (will not form ppt) 3. Most chlorides are soluble, except those of silver, mercury(I), and lead. Lead(II) chloride is soluble in hot water 4. Most sulfates are soluble, except those of barium, strontium, and lead 5. Most carbonates, phosphates, and silicates, are insoluble, except those of sodium, potassium, and ammonium 6. Most sulfides are insoluble, except those of calcium, strontium, sodium, potassium, and ammonium D. Precipitation Reactions 1. A reaction between two soluble compounds in solution, resulting in at least one insoluble product a. Identify the insoluble product if there is one

Chemical Equation Na2SO4(aq) + BaCl2(aq)  2NaCl(aq) + BaSO4(s) + -2 +2 - + - Ionic Equation 2Na (aq) + 2SO4 (aq) + Ba (aq) + 2Cl (aq)  2Na (aq) + 2Cl (aq) + BaSO4(s) +2 -2 Net Ionic Ba (aq) + 2SO4 (aq)  BaSO4(s) Spectator ions Na+(aq) ,Cl-(aq) Net Ionic Equations A. Definition 1. An equation that includes only those compounds and ions that undergo a chemical change in a reaction in an aqueous solution B. Writing a Net Ionic Equation Spectator ions are those ions that do not take part in a chemical rxn and are found in solution both before and after the rxn:

V. Combustion Reactions A. Combustion Rxns B. Hydrocarbon combustion always produces carbon dioxide and water

2C2H6(g) + 7O2(g)  4CO2(g) + 6H2O(g)

VI Oxidation/ Reduction Reactions A. Rules 1. The oxidation number for an atom in its elemental form is always zero. . S8: The oxidation number of S = 0 . Fe: The oxidation number of Fe = 0 2. The oxidation number of a monoatomic ion equals charge of the monatomic ion. o Examples: . Oxidation number of S2- is -2. . Oxidation number of Al3+ is +3. 3. The oxidation number of all Group 1A metals = +1 (unless elemental). 4. The oxidation number of all Group 2A metals = +2 (unless elemental). 5. Hydrogen (H) has two possible oxidation numbers: o +1 when bonded to a nonmetal HCl o -1 when bonded to a metal NaH 6. Oxygen (O) has two possible oxidation numbers: 2- o -1 in peroxides (O2 )....pretty uncommon o -2 in all other compounds...most common 7. The oxidation number of fluorine (F) is always -1. 8. The sum of the oxidation numbers of all atoms (or ions) in a neutral compound =0. 9. The sum of the oxidation numbers of all atoms in a polyatomic ion equals the charge on the polyatomic ion. When assigning oxidation numbers to the elements in a substance, take a systematic approach. Ask yourself the following questions: 1. Is the substance elemental? 2. Is the substance ionic? 3. If the substance is ionic, are there any monoatomic ions present? 4. Which elements have specific rules? 5. Which element(s) do(es) not have rules? 6. Use rule 8 or 9 from above to calculate these.

Example: Determine the oxidation number of each element in Na2SO4. 1. Is the substance elemental? o No, there are three elements present. 2. Is the substance ionic? o Yes, because metal + non-metal = ionic. 3. If the substance is ionic, are there any monoatomic ions present? o Yes, the sodium ion (Na+) is monoatomic. o Therefore, the oxidation number of Na is +1. 4. Which elements have specific rules? o Oxygen has a rule (-2 in most compounds). o Oxidation number of O = -2. 5. Which element(s) do(es) not have rules? o S does not have a rule. o Use rule 8 (and a little simple algebra) to find the oxidation number of S. . Let S = oxidation number of sulfur. . According to rule 8: . (# Na) (Oxid. # of Na) + (# S)(Oxid. # S) + (# Oxygens) (Oxid. # of O) = 0 . So: 2(+1) + S + 4(-2) = 0 . Solve for S using algebra. S = +6. . Therefore, the oxidation number of S = +6

Example: Determine the oxidation number of each element in K2C2O4. 1. Is the substance elemental? 1. No, there are three elements present. 2. Is the substance ionic? 1. Yes, because metal + non-metal = ionic. 2. If the substance is ionic, are there any monoatomic ions present? 1. Yes, the potassium ion (K+) is monoatomic. 2. Therefore, the oxidation number of K is +1. 3. Which elements have specific rules? 1. Oxygen has a rule (-2 in most compounds). 2. Oxidation number of O = -2. 4. Which element(s) do(es) not have rules? 1. C does not have a rule. 2. Use rule 8 (and a little simple algebra) to find the oxidation number of C. . Let C = oxidation number of carbon. . According to rule 8: . (# K) (Oxid. # of K) + (# C)(Oxid. # C) + (# Oxygens) (Oxid. # of O) = 0 . So: 2(+1) + 2C + 4(-2) = 0 . Solve for C using algebra. C = +3. . Therefore, the oxidation number of C = +3 B. Oxidation and Reduction Processes 1.. Oxidation A. Loss of Electrons B. Rxns in which atoms or ions attain a more positive or less negative oxidation state C. Na  Na+1 + e- 2. Reduction A. Gaining Electrons B. Rxns in which atoms or ions attain a more negative or less positive oxidation state C. Cl + e-  Cl-1 3. Redox rxns A. Any chemical process in which elements undergo a change in oxidation number B. Ox. cannot occur without red. occuring simultaneously C. number of electrons lost = number of electrons gained D. If none of the atoms in a reaction change oxidation state, then it is NOT a redox reaction E. Half-reactions C. The part of the reaction involving oxidation or reduction alone can be written as a half-reaction:

Balancing Redox Equations I. Electron Transfer Method A. Write the skeleton equation for the rxn

NH3 + CuO  Cu + H2O + N2 B. Assign oxidation #’s to all elements and determine what is oxidized and what is Reduced

C Write the electronic eqns for both the oxidation and reduction processes

D. Adjust the coefficients in both electronic eqns so that the number of electrons lost equals the number of electrons gained

E. Place these coefficients in the skeleton eqn.

F. Supply the proper coefficients for the rest of the eqn to satisfy conservation of atoms Oxidizing and Reducing Agents I. Oxidizing and Reducing Agents A. Oxidizing agents 1. A substance that has the potential to cause another substance to be oxidized 2. The substance that is reduced in a redox rxn 3. The halogens and oxygen are active oxidizing agents B. Reducing agents 1. A substance that has the potential to cause another substance to be reduced 2. The substance that is oxidized in a redox rxn 3. Group I and II metals are active reducing agents