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Chemical Names, Formulas, Reactions, Equations, and Stoichiometry Chemical Names and Formulas Do Now • Take out Reference Tables and open to Table E – QUIZ ON POLYATOMIC IONS ON WEDNESDAY • Turn to Table S – Where is the oxidation number noted? • How many aluminum atoms combine with how many sulfur atoms? How do we figure this out? OXIDATION NUMBER • Oxidation number = the charge an atom would acquire if all its bonds were treated as ionic bonds. • To determine how many atoms combine with one another in a compound we must determine each element’s OXIDATION NUMBER. Predicting Ionic Charges Group 1: Lose 1 electron to form 1+ ions H+ Li+ Na+ K+ Predicting Ionic Charges Group 2: Loses 2 electrons to form 2+ ions Be2+ Mg2+ Ca2+ Sr2+ Ba2+ Predicting Ionic Charges B3+ Al3+ Ga3+ Group 13: Loses 3 electrons to form 3+ ions Predicting Ionic Charges Group 14: Lose of 4 electrons or gain of 4 electrons? Neither! Group 14 elements rarely form ions. Predicting Ionic Charges N3- Nitride P3- Phosphide As3- Arsenide Group 15: Gains 3 electrons to form 3- ions Predicting Ionic Charges O2- Oxide S2- Sulfide Se2- Selenide Group 16: Gains 2 electrons to form 2- ions Predicting Ionic Charges F1- Fluoride Br1- Bromide Cl1-Chloride I1- Iodide Group 17: Gains 1 electron to form 1- ions Predicting Ionic Charges Group 18: Stable Noble gases do not form ions! Predicting Ionic Charges Groups 3 - 11: Many transition elements have multiple oxidation states. Iron(II) = Fe2+ Iron(III) = Fe3+ Predicting Ionic Charges Groups 3 - 11: Some metals have only one possible oxidation state. Silver = Ag+ Zinc = Zn2+ Rules for assigning Oxidation Numbers 1. The atoms in a pure element have an oxidation number of zero. 2. Alkali metals always have an oxidation number of +1; alkaline earth metals always have an oxidation number of +2. 3. Fluorine always has an oxidation number of -1. 4. Oxygen has an oxidation number of -2 in almost all compounds. Exceptions are in compounds with a halogen, when it has an oxidation number of +2, and in peroxides (H2O2), when it has an oxidation number of -1. 5. Hydrogen has an oxidation number of +1 in almost all compounds except when combined with a metal when it has an oxidation number of -1. 6. The sum of all the oxidation numbers in a neutral compound is zero. 7. The sum of all the oxidation numbers in a polyatomic ion is equal to the charge of the ion. ASSIGNING OXIDATION NUMBERS • A compound has a total charge of ZERO so set your equation equal to ZERO. • Assign the variable X to your unknown oxidation number. • Given a compound, find the oxidation number of every element you know for certain. Then solve for others using algebra. • KMnO4 • CaCO3 Do you notice a pattern with the elements in these compounds? H2+1 O1-2 Mg1+2F2-1 Al2+3S3-2 Cu3+1P1-3 Ca1+2Cl2-1 Fe2+3O3-2 CRISS-CROSS METHOD to determine the chemical formula 1. Write the symbols for the elements side by side. 2. Write the oxidation states of each element to the top right of the symbol. When the nonmetal is combines with a metal, the oxidation state will always be the first number (the negative one) in the list of oxidation states. 3. Criss cross the charges DOWN and use the absolute values (-2 becomes 2). 4. Check to make sure the subscripts are the lowest ratio. (Empirical formula) Practice Criss-Cross Method 1. Na and S 2. K and P 3. Al and S 4. Mg and Br 5. Al and O What about Polyatomic Ions? • Write down the only cations on Table E: (positively charges ions) • Write down the only polyatomic ions that end in –ide. Polyatomic Ions Chart (Table E) Find the charge of the polyatomic ion using Table E Put parenthesis around the polyatomic ion. 1. PO4 2. CO3 3. SO3 4. NH4 5. ClO 6. ClO2 7. HSO4 8. ClO4 9. CN 10. OH 11. S2O3 12. SCN Polyatomic Ions • If you see a group of atoms together with a charge it is a polyatomic ion from Table E. • Put the polyatomic ion within Parenthesis. • Find the charge of the polyatomic ion. • Use the criss-cross method to determine subscripts. Na +1 CO3-2 How many atoms are present in the compound? NH4 +1 S-2 How many atoms are present in the compound? Rules for writing the formula of the compound composed of ions 1. Place all Polyatomic Ions in Parenthesis (Table E) 2. Determine all oxidation numbers of elements and polyatomic ions 3. Use Criss-Cross Method 4. Reduce to Empirical Form (simplest ratio) Practice 1. NH4 S 2. Na NO3 3. Cu Br 4. Al SO4 5. Fe CO3 6. Pb PO4 7. Ag ClO 8. Ca F 9. NH4 SO3 10. Cu OH 11. Ni I 12. Zn SO4 13. Pb ClO2 14. H I 15. Fe HSO4 16. Cu CO3 17. NH4 O 18. Ag S 19. Al ClO4 Do Now • What is the name of H2O? • What is the name of NaCl? • What is the name of NH4? Naming Compounds • Many chemical compounds have common names. • H2O = water • NaCl = table salt • NH4 = ammonia Review • Chemical formula = the relative number of atoms of each kind in a chemical compound • Empirical formula = the elements appear in the smallest whole number ratio. • Molecular formula = the number and kind of atoms in a molecule (not necessarily the smallest whole number ratio) • Structural formula = shows how the atoms in a molecule are bonded to one another NAMING IONIC COMPOUNDS • An Ionic compound can quickly be determined if a METAL (or cation) is bonded to a NONMETAL (or anion). • Naming Ionic compounds with metals that have only ONE oxidation state is fairly simple. Of the following metals listed below, check off all the elements that have more than one oxidation state: a. Mn___ b. Zn ___ c. K ____ d. Pb___ e. Mg___ f. Au ___ g. Ag ___ h. Ga__ i. Sn __ j. Li ____ k. U ____ l. Cr ___ Naming Ionic Compounds that contain metals with 1 oxidation state • Binary Ionic Compounds = ionic compounds with only 2 different elements. • Name the metal and end the nonmetal in – ide. • For example: CaBr2 = _______________ Naming Binary Ionic Compounds *using the IUPAC system • Naming salts is very easy, because they are binary ionic compounds (made up of two elements). – The cation is named by borrowing the name of the element. – The anion named by combining the name of the element with an –ide ending. • The name of compound is made up of both the cation and anion name – Ex: NaCl = sodium chloride – Ex: ZnS = zinc sulfide – Ex: K2O = potassium oxide – Ex: Mg3N2 = magnesium nitride – Ex: Al2S3 = aluminum sulfide Name the following Binary compounds: 1. MgO= ________________________ 2. CaCl2 = _______________________ 3. AlBr3 = _______________________ 4. Ag3N = _______________________ 5. Al2O3 = ______________________ 6. LiI = _________________________ 7. BaF2 = _______________________ 8. Zn2C = _______________________ 9. Ba3N2 = ______________________ 10.CdO = _______________________ 11. Ga2S3 = _______________________ 12. K3N = ________________________ 13. SrO = ________________________ What about Polyatomic ions? (Table E) • These are ions consisting of more than one atom. The names of polyatomic ions end in – ate or –ite, except for two (ammonium and cyanide) Polyatomic ion chart (Table E) Ternary Ionic Compounds • • • • Ionic Compounds with 3 different elements. They usually contain Polyatomic Ions (Table E) Name the following Polyatomic Ions: a. NO3- _________ d. SO42- ___________ • b. ClO2- _________ e. SO32- __________ • c. CO32- _________ f. SCN- ____________ Naming Ternary Compounds • Name the metal and then name the polyatomic ion (if it has a negative oxidation number). • For example: KNO3 = ________________ Naming Ternary Compounds • Name the polyatomic ion (if it has a positive oxidation number) and then name the nonmetal. • For example: NH4Cl = ________________ Naming Ternary Compounds • If there are two polyatomic ions, name the positive polyatomic ion first and then name the negative polyatomic ion. • For example: NH4NO3 = ______________ • DO NOT CHANGE THE ENDINGS OF POLYATOMIC IONS! Name the following Ternary compounds: 1. NaC2H3O2 _____________________ 2. AgHCO3 _______________________ 3. LiNO2 _________________________ 4. Ga2(S2O3)3 ____________________ 5. Ca3(PO4)2 _____________________ 6. ZnSO3 __________________________ 7. KClO3 ___________________________ 8. Al(OH)3 _________________________ 9. RbSCN __________________________ 10. SrCO3 _________________________ Naming Ionic Compounds with Metals with Multiple oxidation states (multiple charges): Using the STOCK SYSTEM 1. Determine the oxidation state of the metal in the compound. 2. Name the metal, put the oxidation state in ROMAN NUMERALS in parenthesis and end the nonmetal in –ide. Review: Find the Formula Criss-Cross Pb+4 O-2 Cu+2 (SO4) -2 Sn+2(CO3) -2 Find Empirical formula Work Backwards: Start with the Empirical Formula to determine the Oxidation state of a Metal with Multiple Oxidation States Empirical Formula Fe1O1 Write in the oxidation number for the nonmetal or polyatomic ion you are sure of and criss-cross. Fe O Fe1(SO4)1 Fe (SO4) Cu1(SO4)1 Cu (SO4) Non-reduced form with Oxidation States Work Backwards: Start with the Empirical Formula to determine the Oxidation state of a Metal with Multiple Oxidation States Empirical Formula Write in the oxidation number for the nonmetal or polyatomic ion you are sure of and criss-cross. Sn1(SO3)1 Sn (SO3) Mn1(SO4)2 Mn (SO4) Cr1(PO4)2 Cr (PO4) Non-reduced form with Oxidation States Name the following compound using the Stock System: 1. Fe O ____________________________ 2. Fe Cl2 ___________________________ 3. Cu SO4 __________________________ 4. Pb Cl2 ___________________________ 5. Pb O2 ___________________________ 6. Cu3(PO4)2 ________________________ 7. Cu2 S ___________________________ 8. Fe2(CrO4)3 ______________________ 19. Sn CO3 _________________________ 10. Sn F4 ___________________________ Name each of the following compounds, use Roman Numerals only when necessary. Put a check next to every compound that begins with a metal with more than 1 oxidation state. Put parenthesis around all the polyatomic ions. 1. NH4 Cl _________________________ 2. Pb SO4 _________________________ 3. Co Cl3 __________________________ 4. Ba (NO3)2 ______________________ 5. Co2 (SO3)3 _____________________ 6. KH ____________________________ 7. NH4 F _________________________ 8. K2Cr2O7 _______________________ 9. Cu S __________________________ 10. Cu ClO2 ______________________ 11. Ag NO3 _______________________ 12. Fe Cl3 ________________________ 13. Cr F2 _________________________ 14. Na Cl _________________________ 15. Fe PO4 ______________________ 16. Li F _________________________ 17. Fe F3 _______________________ 18. Al (OH)3 _____________________ 19. Mg I2 ________________________ 20. Fe Cl3 _______________________ Another Naming System • For metallic elements that can form two positive ions, the suffixes –ous and –ic may be attached to the Latin name of the element: – Ex: FeCl2 = ferrous chloride – Ex: FeCl3 = ferric chloride – Ex: Cu2O = cuprous oxide – Ex: CuO = cupric oxide – Ex: Hg2Br2 = mercurous bromide – Ex: HgBr2 = mercuric bromide Do Now • What types of bonds are in molecular compounds? _____________________ • How can we tell if a formula has a covalent bond? ___________________________ Review Molecular Compounds • Molecule = a neutral compound held together by covalent bonds • Molecules may consist of identical atoms bonded together (O2) or different atoms bonded together (H2O) Naming Covalent Compounds *using the IUPAC system • Covalent compounds are named in a similar way to ionic compounds – The first element in the formula is usually written first in the name – The second element has an –ide ending • Ex: SO2 = sulfur oxide – However, this is not completely correct… Naming Covalent Compounds *using the IUPAC system • Since multiple covalent compounds can be made from the same elements, the name must distinguish them as different. – Prefixes are used to indicate the number of atoms of each element in the molecule. • Ex: SO2 = sulfur dioxide • Ex: SO3 = sulfur trioxide Naming Covalent Compounds *using the STOCK SYSTEM • Most nonmetals have more than 1 oxidation state, therefore you can use the Stock System (Roman Numerals) or the Prefix System. • The Prefix System includes: MonoDiTriTetraPenta- HexaHectaOctaNonaDeca- Naming Molecular Compounds • According to the number of atoms of each element, state the prefix for the number of each atom before the name of the element and end the nonmetal in –ide. • For example: P2O5 = ______________ CO2 = ______________ Naming Molecular Compounds • Name the first element. Use a prefix ONLY if there is more than one. • Name the second element. ALWAYS use a prefix. Change the ending to –ide. • The prefix mono only needs to be used for the second half of the compound NOT the first element. Name the following covalent compounds: 1. CI4 ________________________________ 2. PCl5 ______________________________ 3. SI6 _______________________________ 4. P2S6 ______________________________ 5. N3O4 ______________________________ 6. SO2 ____________________________ 7. N2O4 ___________________________ 8. CO _____________________________ 9. NF3 ____________________________ 10. ICl5 ___________________________ 11. H2S3 __________________________ 12. N2O3 _________________________ 13. ClF7 __________________________ 14. SO3 ___________________________ 15. NI5 ___________________________ 16. BN2 __________________________ 17. P2O5 _________________________ 18. IF7 ___________________________ Do Now • What is the chemical formula for dihydrogen monoxide? • What is the chemical formula for acetate? Writing Formulas • The chemical formula of a compound can be determined from the chemical name! General Steps to Determining Chemical Formulas: • 1. Determine what elements and/or polyatomic ions are in the compound. • 2. Write the symbols for each substance with oxidation states of each substance to the top right of the symbol. – If there is a metal with more than one oxidation state, it will be indicated as roman numerals in parenthesis after the element. – If there is a nonmetal with more than one oxidation state, use the negative number. General Steps to Determining Chemical Formulas (continued): • 3. All oxidation states in a neutral compound add up to ZERO, so figure out how many of each substance you need to make all oxidation states add up to ZERO. • 4. Put these numbers as subscripts and write the chemical formula. • 5. If you have more than one polyatomic ion, make sure you put the symbols in parenthesis. • 6. Check to make sure the subscripts are the lowest ratio. Polyatomic Ions (Table E) Writing Formulas for Ionic Compounds • All ionic compounds must be in empirical form. (Reduced Form) • If the compound ends in –ide, most likely it is a binary compound. Except for cyanide and hydroxide. • If the substance ends in –ate or –ite it contains a polyatomic ion. Put the polyatomic ion in parenthesis with oxidation numbers indicated before you crisscross. Put the oxidation states on top and criss-cross the numbers. • For example: Aluminum Sulfide = Write the formulas for each of the following compounds: (Make sure you reduce it to empirical form) 1. Aluminum Chloride ___________________ 2. Silver Phosphate _____________________ 3. Lithium Hydride ______________________ 4. Magnesium Acetate __________________ 5. Potassium Sulfite _____________________ 6. Zinc Thiosulfate ____________________ 7. Strontium Nitride ____________________ 8. Calcium Oxide _______________________ 9. Gallium Oxalate ______________________ 10. Ammonium Hydroxide ________________ Writing formulas for compounds with metals with more than one oxidations state: (Roman Numerals will be given) 1. If there are Roman Numerals with the name, the Roman Numeral is the charge of the metals. For example: Iron (II) Oxide= Fe+2O-2=Fe2O2=FeO 2. If the compound ends in –ate or –ite, most likely you should look on the polyatomic ion chart. 3. Write the metal, then look up the polyatomic ion and place it in parenthesis. Put the charges on top and criss-cross. Reduce if necessary. For example: Zinc Carbonate= Gold (III) Thiocyanate = For each compound listed below, write the correct formula using the stock system. 1. Iron (II) Chloride _____________________ 2. Lead (IV) Phosphide __________________ 3. Tin (II) Oxide ________________________ 4. Copper (I) Iodide _____________________ 5. Nickel (III) Sulfide ____________________ 6. Cobalt (II) Thiocyanate ______________ 7. Manganese (IV) Oxide _________________ 8. Titanium (IV) Chromate _________________ 9. Iron (III) Sulfate _______________________ 10. Lead (II) Nitrate ______________________ 11. Tin (IV) Carbonate ______________ 12. Copper (II) Acetate _____________ Writing Formulas for Covalent Compounds • Sometimes the name will have prefixes (only when the two elements in the compound are nonmetals). Simply use the prefixes to figure out the formula! For each compound listed below, write the correct formula using the prefix system. 1. Diphosphorous pentoxide _____________________ 2. Silicon tetrafluoride ________________________ 3. Dihydrogen monoxide ________________________ 4. Tetraphosphorous trisulfide __________________ Naming Ionic Compounds using the Formula Determining subscripts of elements or polyatomic in ionic compounds Use Criss-cross method Example: Naming ionic compounds with metals with only 1 oxidation state Binary Ionic Compounds Ionic compounds with only 2 different elements. For example: CaBr2 = Calcium Bromide Name the metal and end the nonmetal in –ide. Ternary Ionic Compounds Ionic Compounds with 3 different elements. •Name the metal and then name the polyatomic ion. Naming ionic compounds with metals with multiple oxidation state 1. Determine the oxidation state of the metal in the compound. 2. Name the metal, put the oxidation state in ROMAN NUMERALS in parenthesis and end the nonmetal in –ide. For example: KNO3 = Potassium Nitrate Example: FeO= Iron (II) Oxide Writing the Formula of Ionic Formulas using the name of the Ionic Compound Formula Writing for Ionic Compounds Writing formulas for compounds with metals with more than one oxidations state: (Roman Numerals will be given) All ionic compounds must be in empirical form. (Reduced Form) If the compound ends in –ide, most likely it is a binary compound. Except for cyanide and hydroxide. If the substance ends in –ate or –ite it contains a polyatomic ion. Put the polyatomic ion in parenthesis before you criss-cross. Put the oxidation states on top and crisscross the numbers. If there are Roman Numerals with the name, the Roman Numeral is the charge of the metals. 2. If the compound ends in –ate or –ite, most likely you should look on the polyatomic ion chart. 3. Write the metal, then look up the polyatomic ion and place it in parenthesis. Put the charges on top and criss cross. Reduce if necessary. For example: Zinc Carbonate= Zn+2(CO3) -2 = Zn2(CO3)2 =Zn(CO3) Gold (III) Thiocyanate = Au+3 (SCN) Au(SCN)3 -1 = For example: Iron (II) Oxide= Fe+2O-2=Fe2O2=FeO Quiz • Short Quiz on Oxidation Numbers and Naming CHEMICAL REACTIONS Do Now • What are some signs that a chemical change may have taken place? • Where are the reactants and products in a reaction? What do they represent? • Describe the law of conservation of mass. Chemical Reaction • A process in which one or more substances are converted into new substances with different chemical and physical properties. Chemical Reaction • In chemical reactions, existing bonds are broken, atoms are rearranged, and new bonds are formed. • Reactant: substances that enter the reaction (start) • Product: substances that come out of reaction (end) Evidence • How do we know that a reaction has taken place? – Production or absorption of energy (heat and light) – Production of a gas – Formation of a precipitate (a solid that forms as a result of mixing 2 solutions). – Color change Chemical Equations • Represent reactions with symbols and formulas, tells us the identities and relative amounts of reactants and products in a chemical reaction. Word Equations • Give names of reactants and products + = “reacts with,” “add,” or “ added to” → = “yields” or “produces” • Describe the following as a word equation: Calcium + Oxygen → Calcium Oxide Formula Equations • Give symbols or formulas of reactants and products • When going from word to formula, be careful to use the correct formulas. (Remember the diatomic molecules Br2, I2, N2, Cl2, H2, O2, F2) • Write the formula for the previous word equation: • Ca + O2 → CaO Additional symbols in equations: • (s) or ↓ = Solid • (l) = Liquid • (g) or ↑ = Gas • (aq) = Aqueous (dissolved in water) • Write the word equation for the reaction that occurs when solid sodium oxide is added to water and forms sodium hydroxide (dissolved in water). • Na2O(s) + H2O → NaOH (aq) • Write the word equation and formula for methane reacting with oxygen to produce carbon dioxide and water. Methane + oxygen carbon dioxide + water CH4 + O2 CO2 + H2O LAW OF CONSERVATION OF MATTER (MASS) • Formula equations must be written in accordance with the LAW OF CONSERVATION OF MATTER (MASS) = atoms cannot be created or destroyed in a chemical reaction. They can only be rearranged! *If you have 50g of reactants, you must end up with 50g of products. Law of Conservation of Mass Mass is neither created nor destroyed during chemical or physical reactions. Total mass of reactants = Total mass of products Formula equations must be written in accordance with the LAW OF CONSERVATION OF MATTER (MASS) Antoine Lavoisier Law of Conservation of Mass • The same number and kind of atoms must appear on both sides of the arrow!!! • COEFFICIENTS= small whole numbers that appear in front of a formula in a chemical equation. • The coefficient multiplies the number of atoms of each element in the formula that follows. – Ex: 2 H2O = 4 H and 2 O Trial and Error is the best way to balance equations. Steps to Writing Balanced Equations: 1. Read the word equation that describes a reaction: zinc + hydrochloric acid → zinc chloride + hydrogen 2. Replace the words with symbols. Zn +HCl → ZnCl2 + H2 • 3. Count the number of atoms on each side of the arrow. (Keep polyatomic ions that exist on both sides of the equation as one!) Zn +HCl → ZnCl2 + H2 1 Zn 1 Zn 1H 2H 1 Cl 2 Cl • 4. Balance using trial and error. Zn +2HCl → ZnCl2 + H2 • 5. Check 1 Zn 2H 2 Cl 1 Zn 2H 2 Cl 1mol of Zn will react with 2mol HCl to produce 1mol ZnCl2 and 1mol H2 Practice • Na + H2O NaOH + H2 1 Na 1 Na 2H 3H 1O 1O • 2Na + 2H2O 2NaOH + H2 2 Na 2 Na 4H 4H 2O 2O Practice • Iron (III) oxide and hydrogen are the reactants. Iron and Water are the products. Find the balanced equation. Fe2O3 + H2 Fe + H2O 2 Fe 1 Fe 3O 1O 2H 2H Practice Fe2O3 + H2 Fe + H2O 2 Fe 1 Fe 3O 1O 2H 2H Fe2O3 + 3H2 2Fe + 3H2O 2 Fe 2 Fe 3O 3O 6H 6H Balanced Particle Diagrams • What are the chemical formulas for the reactants and products? What did you learn today? What did you learn today? • In all chemical reactions there is a conservation of mass, energy, and charge. • A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction. Do Now • What do you notice about the arrangement of elements in the reactants and products? • 2Al + Fe2O3 2Fe + Al2O3 • HCl + NaOH H2O + NaCl TYPES OF CHEMICAL REACTIONS • Thousands of chemical reactions occur in nature. Reactions can be classified into one of several categories. Classification is important for predicting products of reactions. – 1. SYNTHESIS REACTIONS – 2. DECOMPOSITION REACTIONS – 3. SINGLE REPLACEMENT REACTIONS – 4. DOUBLE REPLACEMENT REACTIONS – 5. COMBUSTION 1. SYNTHESIS REACTIONS • Two or more substances combine to form a new compound (composition or direct combination). • Examples: • A + X → AX • 2 Mg +O2 → 2 MgO • 2 CO + O2 → 2 CO2 (car exhaust) • CO2 + H2O H2CO3 (carbonic acid) 2. DECOMPOSITION REACTIONS • A single compound undergoes a reaction that produces two or more simpler substances (breakdown). • Examples: • AX → A + X • 2 H2O2(l) → 2 H2O(l) +O2(g) • 2 H2O(l) → 2 H2(g) +O2(g) – Electrolysis- the decomposition of a substance by electric current. Electrolysis of Water 3. SINGLE REPLACEMENT REACTIONS • One element replaces a similar element in a compound (displacement). • A + BX → B + AX • Mg(s) + 2HCl(aq) → MgCl2(aq) + H2 ↑ • 2Al(s)+3Pb(NO3)2(aq) → 2Al(NO3)3(aq)+3Pb(s) • This type of reaction occurs only under certain conditions! What does “certain conditions” mean? • A + BX → B + AX • A (metal or nonmetal) • B (metal) • X (nonmetal) Using the Activity Series (Table J) • Some elements are considered to be very reactive, some are considered unreactive. The ability of an element to react is called the element’s “activity.” • Higher metals are more reactive than lower metals. A metal will only replace another metal in a compound if it is HIGHER on Table J. Using the Activity Series • Metals: the higher the activity, the greater its tendency is to lose electrons and therefore more reactive. • Nonmetals: the higher the reactivity, the greater its tendency to gain electrons and therefore more reactive. • The activity series is used to predict whether or not a chemical reaction will occur (single replacement reactions). • If it is lower no reaction (NR) will occur! Using the Activity Series • Single Replacement reactions only occur when element A is more active than the element being replaced. • Element A must be more active on the Activity Series (Table J) than the element being replaced!! • Metals replace metals, nonmetals replace nonmetals. METALS: Single-replacement • If element A lies further up on the activity chart (Table J) than element B, element A is more active, therefore the reaction will occur. • For example: • Al + CuCl2 → AlCl3 + Cu • A BX AX B • Aluminum is more active than copper, it lies further up on the chart, therefore the reaction will occur! In each of the following, circle which metal is more active: • 1. Li or Al • 2. Au or Mg • 3. Na or Pb *notice precious metals at the bottom = less reactive and more valuable In each of the following, predict whether there will be a reaction, then balance the reactions: • Mg + CuSO4 → • Al + LiCl → • Sn + Ba3PO4 → NONMETALS: Single-replacement • If element A is a stronger nonmetal than element X (nonmetals replace nonmetals) then Element A is more active and the reaction will occur! • For example: • F2 + 2LiBr → Br2 + 2LiF • A BX X BA • Fluorine is more active than Bromine, therefore the reaction will occur!! In each of the following, circle which nonmetal is more active: • 4. F2 or I2 • 5. Br2 or H2 • 6. Cl2 or I2 In each of the following, predict whether there will be a reaction, then balance the reactions: • I2 + NaF → • F2 + MgBr2 → • Cl2 + BaH2 → Rules for completing and writing single replacement reactions. 1. Determine whether the reaction will occur or not Li + CuSO4 → Yes a reaction will occur (Lithium is higher up on the chart) 2. Determine the products of the reaction by switching the appropriate elements. Li +CuSO4 → LiSO4 + Cu 3. Place the charges on the elements and balance each of the new compounds formed. Li + CuSO4 → Li2(SO4) + Cu 4. Balance the entire equation. 2 Li + CuSO4 → Li2(SO4) + Cu 4. DOUBLE REPLACEMENT REACTIONS • The ions of 2 compounds exchange places in an AQUEOUS solution to form 2 new compounds. • Examples: • AB + CD → AD + CB • Pb(NO3)2(aq) +KI(aq) → KNO3(aq) + PbI2(s) • Double replacement reactions do not always occur. • We use the term No Reaction (NR) when no significant change occurs in a reaction. Conditions when Double Replacement Reactions occur • 1. If the product is insoluble. (A precipitate is formed) ↓ – Insoluble: cannot be dissolved in water – Precipitate: an insoluble product (solid) • 2. If water is a product • 3. If the product decomposes into a gas. ↑ – H2CO3 → – H2SO3 → – NH4OH → H2O + CO2 (g) H2O + SO2 (g) H2O + NH3 (g) PRODUCT IS INSOLUBLE • Use the Solubility Chart (Table F) on your reference table to figure out the solubility for the problems on the next slide: • Soluble = dissolvable in water, therefore it is aqueous (aq) • Insoluble = cannot dissolve in water, therefore it forms a precipitate (s) Solubility Chart (Table F) Determine whether the compound is soluble or insoluble a. NH4OH b. PbCl2 c. BaCO3 d. Al(OH)3 e. Na2S f. K2SO4 g. CaSO4 h. Mg3(PO4)2 i. Zn(OH)2 Now complete each of the following double replacement equations. For each of the products use Table F to assign a notation of (s) or (aq) below each formula. 1. BaCl2 (aq) + Na2CO3 (aq) → 2. Zn(NO3)2 (aq) + KOH (aq) → 3. CaI2 (aq) + Na3PO4 (aq) → 4. CuSO4 (aq) + (NH4)2CO3 (aq) → 5. COMBUSTION REACTIONS • A substance combines with oxygen (O2) to release large amounts of energy in the form of light and heat. • CH4 +O2 → CO2 +2H2O http://www.chem.uiuc.edu/clcwebsite/meth.html Type of reaction Definition Example Synthesis Two or more substances combine to form a new compound. 2Mg + O2→ 2MgO Decomposition A single compound undergoes a reaction that produces two or more simpler substances (breakdown). H2O2(l) → 2H2O(l) + O2 (g) Single Replacement One element replaces a similar element in a compound. Mg (s) + HCl (aq) → MgCl2 (aq) + H2 (g) Double Replacement The ions of 2 compounds exchange places in an AQUEOUS solution to form 2 new compounds. The metals in the compounds switch places. 2 Li + CuSO4 → Li2(SO4) + Cu A substance combines with oxygen (O2) to release large amounts of energy in the form of light and heat. CH4 +O2 → CO2 +2H2O Combustion Reaction When will the reaction take place •Single Replacement reactions only take occur when one element is more active than the element being replaced. •Use Table J Activity Series •Metals replace metals in a compound. •Nonmetals replace nonmetals in a compound. 1. If the product is insoluble. (A precipitate is formed) ↓ 2. If water is a product 3. If the product decomposes into a gas. ↑ Lab • Balancing Reactions web lab Lab • Single Replacement lab Lab • Double Replacement lab Quiz • Quiz on Balancing and Types of Reactions Chemical formulas and equations Do Now • What is a mole in chemistry? • What was Avogadro's number? What did it represent? • What is molar mass? Do Now • What is a mole in chemistry? • Remember: Mole (mol) = the SI unit for amount • What was Avogadro's number? What did it represent? • Remember: Avogadro's number = the number of particles in a mole = 6.02 x 1023 • What is molar mass? • Remember: Molar mass = mass in grams of one mole of an element or compound Remember • You can convert from moles to particles, or moles to mass, and vice versa. • Ex: How many particles do you have if you have 2.5 mol of sulfur? • Ex: How many grams of carbon do you have if you have 2.44 x 1022 atoms? * You must use conversion factors to convert from number of atoms to moles and then to grams. Formula Mass • The molar mass is numerically equal to the atomic mass of monatomic elements and the formula mass of compounds and diatomic elements. • Ex: Find the formula mass of KBr • Ex: Find the formula mass of H2O Chemical Formula • Chemical formulas indicate the relative number of atoms of each kind in a chemical compound. • Identify the number of atoms in each of the following compounds: 1. KCl= _____________________ 2. C6H12O6=__________________ 3. NH3=__________________ Chemical Formulas • Formulas for covalent compounds show both the elements and the number of atoms of each element in a molecule. • Formulas for ionic compounds do not show numbers of atoms, but show the simplest ratio of cations and anions. Chemical Formulas • The meaning of formulas do not change when polyatomic ions are involved. • Polyatomic ions = a group of covalently bond elements that behave as a single ion Finding the number of atoms in a compound with polyatomic ions Compound 1. MgCO3 2. Al(ClO4)3 What does it look like expanded? Numbers of atoms Compound 3. Zn3(PO4)2 4. (NH4)2S What does it look like expanded? Numbers of atoms Molar Mass • Formulas can be used to calculate molar mass • Ex: ZnCl2 • Ex: ZnSO4 Do Now • Do you see a pattern in the following formulas? NH4NO2 NH2O Law of Definite Proportions • Law of Definite Proportions states that every pure substance always contains the same elements combined in the same proportions by weight. • For example: H2O, will always have the same percent by weight (11.2% H and 88.8% O) Percentage Composition • Percentage composition = the percentage by mass of each element in a compound • Percentage composition helps verify a substance’s identity and can be used to compare the ratio of masses contributed by the elements in two substances • Ex: Fe2O3 FeO 69.9% Fe 77.7% Fe 30.1% O 22.3% O Review Formulas • Molecular Formula = represents the number and kind of atoms in a molecule (not necessarily the smallest whole number ratio). • Structural Formula = indicates twodimensional arrangement of the bonds and lone pairs of electrons in a molecule. Empirical Formulas • Empirical formula = a chemical formula that shows the composition of a compound in terms of the relative numbers and kinds of atoms in the simplest ratio. Empirical Formula for Molecules (have covalent bonds) • The chemical formula that represents the simplest (lowest) atomic ratio in which elements can combine. • Formulas for molecular compounds are NOT NECESSARILY the empirical formula Molecular Formula Empirical Formula C2H6 C6H12O6 C4H8 C3H8 C6H10 C4H6 C2H4 C5H10 Empirical Formula for Ionic Compounds (Always written in empirical form) • The chemical formula that represents the simplest (lowest) atomic ratio in which elements can combine. • Formulas for ionic compounds are the empirical formula because they are ALWAYS written in the expressed as the lowest possible ratios. Empirical Form Ca2O2 Fe2O2 Pb2O4 Mg2O2 Empirical Formula for Ionic Compounds with Polyatomic Ions Empirical Form Fe2(CO3)2 Fe2(HSO4)2 Zn2(SO4)2 Pb2(SO3)2 Empirical or Molecular Formula C6H12O6 LiNO3 H2O2 C11H22O11 If written in molecular formula what is the empirical formula? Type of Bond Finding Empirical Formulas • You can find the empirical formula from percentage composition! – Step 1: Convert to mass in grams (assume you have 100g of the given substance) – Step 2: Convert from grams to moles (using the molar mass conversion factor) – Step 3: Reduce the molar ratio to the simplest wholenumber ratio by dividing by the smaller amount – Step 4: Round to whole numbers and insert subscripts Finding Empirical Formulas • Ex: A given liquid has 60.0% C, 13.4% H, and 26.6% O by mass. Calculate the empirical formula. – Step 1: Convert to mass in grams (assume you have 100g of the given substance) 60.0% C x 100g = 60.0g C 13.4% H x 100g = 13.4g H 26.6% O x 100g = 26.6g O Finding Empirical Formulas – Step 2: Convert from grams to moles (using the molar mass conversion factor) 60.0g C x 1mol = 5.00mol C 12.01g C 13.4g H x 1mol = 13.3mol H 1.01g H 26.6g O x 1mol = 1.66mol O 16.00g O C5H13.3O1.66 ? Finding Empirical Formulas – Step 3: Reduce the molar ratio to the simplest wholenumber ratio by dividing by the smaller amount C5H13.3O1.66 ? 5.00mol C / 1.66 = 3.01mol C 13.3mol H / 1.66 = 8.01mol H 1.66mol O / 1.66 = 1.00mol O – Step 4: Round to whole numbers and insert subscripts C3H8O Practice • What is the empirical formula of a compound that is 78.6% B and 21.4% H? Practice • What is the empirical formula of a compound containing 32.38% Na, 22.65% S, and 44.99% O? Practice • What is the empirical formula of a compound containing 26.56% potassium, 35.41% chromium, and the remainder is oxygen? Practice • In a 10.150g sample, 4.433g are phosphorous and the rest is oxygen. What is the empirical formula for this compound? Practice • A compound contains 0.606g nitrogen and 1.390g oxygen. What is the empirical formula of the compound? What did you learn today? What did you learn today? • A compound is a substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by a specific chemical formula. • Types of chemical formulas include empirical, molecular, and structural. • The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that compound. Do Now • What was the difference between the empirical formula and molecular formula? Do Now • What was the difference between the empirical formula and molecular formula? • Empirical formula = smallest possible whole number ratio of elements • Molecular formula = the actual formula of a molecule Molecular Formulas • For ionic compounds, the molecular formula is the same as the empirical formula. • For molecular compounds, the molecular formula is a whole number multiple of the empirical formula. *Both formulas are just different ways of representing the composition of the same molecule. Molar Mass • In order to determine the molecular formula, you must know the molecular mass! • The molar mass of a compound is equal to the molar mass of the empirical formula time a whole number, n. n(empirical formula) = molecular formula Molecular Formula Examples • Formaldehyde, acetic acid, and glucose each have the same empirical formula, CH2O – For formaldehyde, n = 1 – For acetic acid, n = 2 – For glucose, n = 6 Determining Molecular Formula • Step 1: Find the molar mass of the empirical formula using the molar masses of the elements from the periodic table • Step 2: Solve for n, the factor multiplying the empirical formula to get the molecular formula n = experimental molar mass of compound molar mass of the empirical formula • Step 3: Multiply the empirical formula by this factor to get the molecular formula Determining Molecular Formula • The empirical formula of a compound is P2O5. The experimental molar mass is 284 g/mol. Determine the molecular formula. – Step 1: Find the molar mass of the empirical formula using the molar masses of the elements from the periodic table 2 x molar mass of P = 2(30.97) = 61.94g/mol 5 x molar mass of O = 5(16.00) = 80.00g/mol molar mass of P2O5 = 141.94g/mol Determining Molecular Formula • The empirical formula of a compound is P2O5. The experimental molar mass is 284 g/mol. Determine the molecular formula. – Step 2: Solve for n, the factor multiplying the empirical formula to get the molecular formula n = experimental molar mass molar mass of empirical formula n = 284 g/mol 141.94 g/mol n=2 Determining Molecular Formula • The empirical formula of a compound is P2O5. The experimental molar mass is 284 g/mol. Determine the molecular formula. – Step 3: Multiply the empirical formula by this factor to get the molecular formula n(empirical formula) = 2(P2O5) = P4O10 Determining Molecular Formula • You can verify your answer by finding the molar mass of the molecular formula and compare it to the experimental molar mass 4 x molar mass of P = 4(30.97g/mol) = 123.88g/mol 10 x molar mass of O = 10(16.00g/mol) = 160.0g/mol molar mass of P4O10 = 283.88g/mol Practice • What is the molecular formula of a compound with the empirical formula BH3 and molecular mass of 28g? Practice • What is the molecular formula of a compound with a molecular mass of 34g that consists of 0.44g H and 6.92g O? Practice • What is the empirical formula of a compound that contains 65.5% carbon, 5.5% hydrogen, and 29.0% oxygen? What is the molecular formula if the molecular mass is 110g? Do Now • What does percentage mean? • What does a percentage represent? Review • Chemical formulas allow scientists to calculate a number of characteristics values for a given compound. • Chemical formulas represent the number and kind of atoms in a molecule. • If you know the chemical formula, then you can calculate the percentage composition. Review • Formula mass = the sum of the average atomic masses of all the atoms represented in the formula • Formula mass is numerically equal to the molar mass or gram formula mass • Ex: Find the formula mass of KClO3 • Ex: What is the molar mass of Ba(NO3)2? Percent Composition • Percentage composition = the percentage by mass of each element in a compound Finding Percent Composition • From the subscripts, you can determine the mass contributed by each element and add these to get the molar mass. • Divide the mass of each element by the molar mass. • Multiply by 100 to find the percentage composition of that element Finding Percent Composition • CO2 1mol x 12.01 g/mol = 12.01 g C +2mol x 16.00 g/mol = 32.00 g O mass of 1mol CO2 = 44.01 g Finding Percent Composition % C = 12.01 g C x 100 = 27.29% 44.01 g CO2 % O = 32.00 g O x 100 = 72.71% 44.01 g CO2 Finding Percent Composition • CO 1mol x 12.01 g/mol = 12.01 g C +1mol x 16.00 g/mol = 16.00 g O mass of 1mol CO = 28.01 g Finding Percent Composition % C = 12.01 g C x 100 = 42.88% 28.01 g CO % O = 16.00 g O x 100 = 57.71% 28.01 g CO Percent Composition • Percentage composition helps verify a substance’s identity and can be used to compare the ratio of masses contributed by the elements in two substances • Ex: Fe2O3 FeO 69.9% Fe 77.7% Fe 30.1% O 22.3% O Practice • Calculate the percentage composition of Cu2S, a copper ore called chalcocite. Practice • Calculate the percent of both elements in sulfur dioxide Practice • Calculate the percentage composition of ammonium nitrate, NH4NO3 Hydrates • Hydrates = salts that have crystallized from water solution • In the process of crystallization, the water molecules bind to the salt to form hydrates. • Hydrates are represented as follows: Na2CO3 ● 10H2O (sodium carbonate decahydrate) Practice • Calculate the molar mass of Na2CO3 ● 10H2O Practice • Calculate the percent composition of water in Na2CO3 ● 10H2O What did you learn today? What did you learn today? • The formula mass of a substance is the sum of the atomic masses of its atoms. The molar mass (gram formula mass) of a substance equals one mole of that substance. • The percent composition by mass of each element in a compound can be calculated mathematically. Lab • Percent Composition of Sugar in Gum Lab • Hydrate lab Quiz • Quiz on Formulas and Percent Composition Stoichiometry Do Now • A recipe calls for one cup of milk and three eggs per serving. You quadruple the recipe because you are expecting guests. How much milk and eggs do you need? Stoichiometry • Stoichiometry = the proportional relationship between two or more substances during a chemical reaction. • Reaction stoichiometry = problems involving the amount of products in relation to the amount of reactants – In order to correctly solve stoichiometry problems, you MUST be able to correctly balance equations! • Stoichiometry problems are solved by using ratios from the balanced equation. Balanced Equations • Balanced equations show proportions • A balanced chemical equation is very similar to a recipe in that the coefficients show the proportions of the reactants and products involved in the reaction. – For example: 2 H2 + O2 2 H2O So, 2 moles of hydrogen react with 1 mole of oxygen to produce 2 moles of water. Mole Ratio • You can use mole ratios to determine how much of a reactant is needed to produce a quantity of product, and vice versa. • Since the coefficients in a balanced equation show the relative numbers of moles of the substances in a reaction, ALWAYS convert between amounts in moles! – For example: Al2O3 (l) Al (s) + O2 (g) Moles to Moles • Steps to converting between amounts in moles: 1. Identify the amount in moles that you know from the problem. 2. Using coefficients from the balanced equation, set up the mole ratio with the known substance on the bottom and the unknown substance on top. 3. Multiply the original amount by the mole ratio. Moles to Moles Setup Amount of known moles x unknown moles = unknown moles known moles Moles to Moles Example • How many moles of hydrogen are needed to prepare 312 moles of ammonia? N2 + 3H2 2NH3 Moles to Moles Example • How many moles of hydrogen are needed to prepare 312 moles of ammonia? N2 + 3H2 2NH3 • Amount of NH3 = 312 mol • Amount of H2 = unknown • From the equation: 3 mol H2 = 2 mol NH3 Moles to Moles Example • How many moles of hydrogen are needed to prepare 312 moles of ammonia? N2 + 3H2 2NH3 • 312 mol NH3 x 3 mol H2 = ? 2 mol NH3 = 468 mol H2 Mole to Mole Problems Al2O3 (l) Al (s) + O2 (g) • How many moles of Al(s) will be produced from the decomposition of 13.0 mol of Al2O3(l)? Mole to Mole Problems Al2O3 (l) Al (s) + O2 (g) • How many moles of O2(g) will be produced when 36 mol of Al(s) are produced from the decomposition of Al2O3(l)? Mole to Mole Problems Al2O3 (l) Al (s) + O2 (g) • How many moles of O2(g) are produced when 15 mol of Al2O3(l) react? Problems • Complete the additional problems in your notes packet… Do now • How would you get from Van Cortlandt park to Flushing Meadows park? Do Now • Take out your Reference Table and turn to the back page! • List two conversion factors that relate to the mole. Remember to Use Table T Mass Calculations • Substances are usually measured by mass, so before using the mole ratio, you will need to convert from mass to moles! • The conversion factor for converting between mass and moles is the molar mass of the substance. • Molar mass = the sum of the atomic masses of the elements using the chemical formula The Mole Map Mole Triangle Multiply by atomic/molar mass from periodic table Atoms or Molecules A Divide by 6.02 X 1023 Multiply by 6.02 X 1023 Divide by atomic/molar mass from periodic table 1 mole A ↔ Use ratio of moles 1 mole B Multiply by 6.02 X 1023 Atoms or Molecules B Divide by 6.02 X 1023 Mass Grams A Divide by atomic/molar mass from periodic table Multiply by atomic/molar mass from periodic table Mass (grams) B Mole-Mass Conversions • Most of the time in chemistry, the amounts are given in grams instead of moles • We still go through moles and use the mole ratio, but now we also use molar mass to get to grams • Example: How many grams of chlorine are required to react completely with 5.00 moles of sodium to produce sodium chloride? 2 Na + Cl2 2 NaCl 5.00 moles Na mol Cl2 mol Na g Cl2 1 mol Cl2 = g Cl2 Practice • Calculate the mass in grams of Iodine required to react completely with 0.50 moles of aluminum. Mass-Mole • We can also start with mass and convert to moles of product or another reactant • We use molar mass and the mole ratio to get to moles of the compound of interest – Calculate the number of moles of ethane (C2H6) needed to produce 10.0 g of water – 2 C2H6 + 7 O2 4 CO2 + 6 H20 10.0 g H2O mol H2O g H2O mol C2H6 = mol H20 mol C2H6 Practice • Calculate how many moles of oxygen are required to make 10.0 g of aluminum oxide Mass to Mass • Steps to converting between amounts in mass: 1. Convert you given amount from mass to moles 2. Use the mole ratio to determine moles of the unknown substance (from the balanced equation) 3. Convert the moles calculated into the units required to solve the problem (grams) Mass to Mass Setup mass of known x 1 mol = mol of known molar mass mol of known x mol of unknown = mol of unknown mol of known mol of unknown x molar mass = grams of unknown 1 mol Mass to Mass Example • N2 + 3H2 2NH3 • What mass of NH3 can be made from 21 g H2 and excess N2? Mass to Mass Example • N2 + 3H2 2NH3 • What mass of NH3 can be made from 21 g H2 and excess N2? • Given: mass of H2 = 21 g molar mass of H2 = 2.02 g/mol mass of NH3 = unkwown molar mass of NH3 = 17.04 g/mol mole ratio: 3 mol H2 = 2 mol NH3 Mass to Mass Example • N2 + 3H2 2NH3 • What mass of NH3 can be made from 21 g H2 and excess N2? 21 g H2 x 1mol H2 x 2mol NH3 x 17.04 g NH3 = 2.02 g H2 3mol H2 1 mol NH3 = 118 g NH3 Mass-Mass Conversions • Most often we are given a starting mass and want to find out the mass of a product we will get (called theoretical yield) or how much of another reactant we need to completely react with it (no leftover ingredients!) • Now we must go from grams to moles, mole ratio, and back to grams of compound we are interested in Mass-Mass Conversion • Ex. Calculate how many grams of ammonia are produced when you react 2.00g of nitrogen with excess hydrogen. N2 + 3 H2 2 NH3 g N2 mol N2 mol NH3 g NH3 g N2 mol N2 mol NH3 = g NH3 N2 + 3H2 → 2NH3 1. Convert grams to moles. 2. Solve as a mole-mole problem to get # moles of NH3. 3. Convert answer back to grams NH3. Practice • How many grams of calcium nitride are produced when 2.00 g of calcium reacts with an excess of nitrogen? Mass to Mass Practice • Fe2O3 + Al Fe + Al2O3 • How many grams of Al are needed to completely react with 135 g Fe2O3? Mass to Mass Practice • Fe2O3 + Al Fe + Al2O3 • How many grams of Al2O3 can form when 23.6 g Al react with excess Fe2O3? Mass to Mass Practice • Fe2O3 + Al Fe + Al2O3 • How many grams of Fe2O3 react with excess Al to make 475 g Fe? Mass to Mass Practice • Fe2O3 + Al Fe + Al2O3 • How many grams of Fe will form when 97.6 g Al2O3 form? Conversions liters to moles 22.4 Liters of gas = 1 mole of gas 22.4 Liters gas 1 mole of gas = 1 mole of gas 22.4 Liters gas ↔ 22.4 Liters of gas 6.022x1023 (CO2) or formula units (NaCl) atoms (Ne) or molecules ↔ 1 mole Grams compound or element ↔ (Found on Periodic Table) How many liters of O2 are present if you have 15 moles O2? ↔ 22.4 Liters of gas 6.022x1023 (CO2) or formula units (NaCl) atoms (Ne) or molecules ↔ 1 mole Grams compound or element ↔ (Found on Periodic Table) How many moles are there in 30 liters of O2 gas? ↔ 22.4 Liters of gas 6.022x1023 (CO2) or formula units (NaCl) atoms (Ne) or molecules ↔ 1 mole Grams compound or element ↔ (Found on Periodic Table) How many liters of H2O are present in 1.5 moles of H2O? Do Now (Honors) • A bicycle mechanic has 10 frames and 16 wheels in the shop. How many complete bicycles can he assemble using these parts? Limiting Reactants • If you have the following available ingredients, how many peanut butter and jelly sandwiches can you make? – 4 slices of bread – 1 jar of peanut butter – 1/2 jar of jelly Definitions • Limiting Reactant • Excess Reactant Limiting Reactants 1. Write a balanced equation. 2. For each reactant, calculate the amount of product formed. 3. Smaller answer indicates the limiting reactant Limiting Reactants • 79.1 g of zinc react with 0.90 L of 2.5M HCl. Identify the limiting and excess reactants. How many liters of hydrogen are formed at STP? Zn + 2HCl ZnCl2 + H2 Limiting Reactants How many liters of hydrogen are formed if we are given 79.1g of zinc? Zn + 2HCl 79.1 g 0.90 L 2.5M g Zn ZnCl2 + H2 ?L mol Zn mol H2 L H2 g Zn mol Zn mol H2 = H2 L Limiting Reactants How many liters of hydrogen are formed when we are given 0.90 Liters of HCl? Zn + 2HCl 79.1 g 0.90 L 2.5M L HCl ZnCl2 + H2 ?L mol HCl mol H2 1L mol HCl L H2 = mol H2 L H2 Limiting Reactant is the one that produces the least amount of Hydrogen Zn:____L H2 HCl:_____L H2 Limiting reactant: Excess reactant: Product Formed: L H2 left over ____ • Solid silicon dioxide (quartz) is usually quite unreactive but reacts readily with hydrogen fluoride gas to form silicon tetrafluoride gas and liquid water. Write the balanced chemical reaction for this. If 2.0mol of HF are exposed to 4.5mol of SiO2, which is the limiting reactant? Limiting Reactants • Limiting reactants are not present in excess and they are entirely consumed in a reaction. • First, you need to determine which reactant is present in a limiting quantity by comparing moles. • Then, you can calculate the mole ratio. • Lastly, you can determine the amount of product produced. Practice • N2 + 3H2 2NH3 • If you have 2 moles of N2 and 2 moles of H2, which of the reactants is the limiting agent? • Mole ratio from problem: 2mol N2 = 1 2mol H2 • Mole ratio from equation: 1mol N2 = 0.33 3mol H2 Practice Larger • Mole ratio from problem: 2mol N2 = 1 Limiting reactant 2mol H2 • Mole ratio from equation: 1mol N2 = 0.33 3mol H2 • If the mole ratio from the problem is larger, then the reactant in the numerator is present in excess. • If the mole ratio from the problem is smaller, then the reactant in the denominator is present in excess. • The reactant not present in excess is the limiting reactant. Practice • N2 + 3H2 2NH3 • If H2 is the limiting reactant, then all of it will be consumed in the reaction. • So, how much N2 will react with 2mol of H2? 2mol H2 x 1mol N2 = 0.67mol N2 3mol H2 Practice Problem • C3H8 + 5O2 3CO2 + 4H20 • Determine the limiting reactant and the amount of excess reactant that reacts if you have 1.5mol C3H8 and 15mol O2. Practice Problem • C3H8 + 5O2 3CO2 + 4H20 • Mole ratio from problem: 1.5mol C3H8 = 0.1 Smaller 15mol O2 • Mole ratio from equation: 1mol C3H8 = 0.2 5mol O2 • O2 is present in excess • C3H8 is the limiting reactant Practice Problem • C3H8 + 5O2 3CO2 + 4H20 1.5mol C3H8 x 5mol O2 = 7.5mol O2 1mol C3H8 Practice Problem • 2H2 + O2 2H2O • How much water is produced if 100g H2 and 160g O2 react? • You must convert mass moles moles mass Practice Problem • 2H2 + O2 2H2O • 100g H2 x 1mol H2 2g H2 = 50mol H2 • 160g O2 x 1mol O2 32g O2 = 5mol O2 • Mole ratio from problem: 50mol H2 =10 5mol O2 • Mole ratio from equation: 2mol H2 = 2 1mol O2 Limiting reactant Practice Problem • 2H2 + O2 2H2O 5mol O2 x 2mol H2O x 18g H20 = 1mol O2 1mol H2O = 180g H2O Check • 2H2 + O2 2H2O • 10H2 + 5O2 10H2O • Conservation of mass: 10H2 x 2g H2 = 20g H2 1mol H2 Sum of reactants 5O2 x 32g O2 = 160g O2 = Sum of products 1mol O2 10H2O x 18g H2O = 180g H2O 1mol H2O Percent Yield measured in lab actual yield (g) % yield 100 theoretical yield(g) calculated on paper Percent Yield • When 45.8 g of K2CO3 react with excess HCl, 46.3 g of KCl are formed. Calculate the theoretical and % yields of KCl. K2CO3 + 2HCl 2KCl + H2O + CO2 45.8 g ?g actual: 46.3 g Percent Yield K2CO3 + 2HCl 2KCl + H2O + CO2 45.8 g ?g actual: 46.3 g Theoretical Yield: g mol K2CO3 K2CO3 mol KCl g KCl g mol K2CO3 K2CO3 = mol g KCl KCl Percent Yield K2CO3 + 2HCl 2KCl + H2O + CO2 45.8 g actual: 46.3 g Theoretical Yield = % Yield = 46.3 g g g KCl 100 = % • Methanol (CH3OH) can be produced through the reaction of CO and H2 in the presence of a catalyst. CO(g) + 2H2(g) catalyst > CH3OH(l) If 75.0g of CO reacts to produce 68.4g CH3OH, what is the percent yield of methanol? Quiz • Quiz on Double Replacement Reactions and Mole Calculations (Stoichiometry) Lab • Hydrate lab Lab • Stoichiometry lab Lab • Mole lab Test • Test on Chemical Formulas, Naming, Reactions, and Moles/Stoichiometry