Unit 5

This unit will introduce the following tools that scientists use to “understand reactions” Section 1 Chemical Reactions & Types Section 2 Balancing Equations Section 3 Predicting Products of Synthesis & Decomposition Reactions Section 4 Predicting Products of Single, Double Replacement & Combustion Reactions

Section 1 What is a Chemical Reaction?

What is a Chemical Reaction? Bonds are broken and atoms are rearranged to form new compounds. 2 H2 + O2  2 H2O O H H Watch as 2 H2 and 1 O2 undergo a chemical reaction

Chemical Reactions Bonds and atoms are rearranged to form new compounds. 2 H2 + O2  2 H2O H O O H O H The compounds in the end are different from those in the beginning H Bonds are broken and formed between different atoms

What is a Chemical Reaction? Also known as a chemical change

Observations (Indicators) of a Chemical Reaction(change) Gas production Odor change Color change Precipitate formation: insoluble solid that is formed when 2 aqueous solutions are reacted Energy change(change in temperature) Exothermic Reactions: reactants release heat & temp. of the reaction increases Endothermic Reactions: reactants absorb heat & temp. of the reaction decreases

What are Chemical Equations? The “sentence” of chemistry Uses chemical formulas & other symbols instead of words to show a chemical reaction Reactants Products Starting Materials New substances formed NaCl(aq) + AgNO3(aq)  AgCl(s) + NaNO3(aq) Arrow : Read as: Yields, Produces, Forms Physical State s = solid l = liquid g = gas aq = aqueous (dissolved in water)

Examples of Chemical Equations Word equation: Magnesium metal is reacted with aqueous hydrochloric acid to produce aqueous magnesium chloride and hydrogen gas  Formula equation: Mg (s) + 2 HCl (aq) MgCl2 (aq) + H2 (g) Visualization

Writing Chemical Equations You must write the correct chemical formula for each reactant & product “and”, “is mixed with” or “reacts with” = + “yield”, “produces” and “forms” =  Remember “BrINClHOF”: to represent the 7 diatomic elements that have a subscript of 2 in their chemical formula Br2 I2 N2 Cl2 H2 O2 F2

Write the word equation into symbol form Let’s Practice #1 Potassium metal is reacted with calcium bromide to form potassium bromide and calcium Example: Write the word equation into symbol form

Write the word equation into symbol form Let’s Practice #1 Potassium metal is reacted with calcium bromide to form potassium bromide and calcium Example: Write the word equation into symbol form K + CaBr2  KBr + Ca

Write the word equation into symbol form Let’s Practice #2 Copper (II) nitrate and sodium hydroxide form copper (II) hydroxide and sodium nitrate Example: Write the word equation into symbol form

Write the word equation into symbol form Let’s Practice #2 Copper (II) nitrate and sodium hydroxide form copper (II) hydroxide and sodium nitrate Example: Write the word equation into symbol form Cu(NO3)2 + NaOH  Cu(OH)2 + NaNO3

Synthesis(Combination): when 2 or more substances combine to form a SINGLE substance Element + Element  Compound Cu + S  CuS Compound + Compound  Compound Li2O + H2O  LiOH

Decomposition: when 1 substance breaks apart into 2 or more substances; requires Energy to initiate Compound  Element + Element CaS  Ca + S Compound  Compound + Compound H2O2  H2O + O2

Single Replacement: when a more reactive element replaces a less reactive element in a compound Compound + Element  Element + Compound AgI + K  Ag + KI

Double Replacement: when the positive ions of each aqueous compound replace each other + +  Compound + Compound  Compound + Compound KI + AgNO3  AgI + KNO3

Combustion: reaction of an element or compound with OXYGEN to form an oxide & produce heat hydrocarbons combust to always produce carbon dioxide,water & heat Mg + O2  MgO CH4 + O2  CO2 + H2O

Section 2—Balancing Equations We need to finish writing those equations we started!

Law of Conservation of Mass Matter cannot be created nor destroyed just changed in form. Thus mass must remain constant through changes.

Conservation of Mass CaCl2 + Na2SO4  CaSO4 + 2NaCl Mass before(reactants) = mass after(products) # atoms before (reactants) = # atoms after(products)

How Does the Law Lead to Balancing? Matter cannot be created nor destroyed during a chemical or physical change. The number of atoms of each element must be the same on both sides To ensure this we must balance the equation .

How Do We Begin? Subscripts tell you how many atoms Start by Counting the Atoms of each Substance Subscripts tell you how many atoms Coefficients are always multiplied by subscripts With a parentheses, inner & outer subscripts are always multiplied together

Try These! Count the atoms in each compound (a) 2 (NH4)3PO4 N= ___ H= ___ P= ___ O= ___ (b) 4 KC2H3O2 K= ___ C= ___ H= ___ O= ___

How Do We Balance Equations? Practice Problems: How many atoms of each type are indicated in the following compounds? (a) 2 (NH4)3PO4 N= ___ H= ___ P= ___ O= ___ (b) 4 KC2H3O2 K= ___ C= ___ H= ___ O= ___ (c) 3 Ca(NO3)2 Ca= ___ N= ___ O= ___ 6 24 2 8 4 8 12 8 3 6 18

How to Balance Chemical Equations: 1 Make a table of elements & count the atoms of each substance _____ C H4 + _____ O2  _____ H2 O + _____ C O2 Reactants Products C H O

How to Balance By Inspection: 2 Count the number of each element or ion on the reactants and products side. Don’t forget to add all the atoms of the same element together—even if it appears in more than one compound! _____ C H4 + _____ O2  _____ H2 O + _____ C O2 Reactants Products C 1 1 H 4 2 O 2 3

How to Balance By Inspection: 2 Use coefficients to balance the numbers Each time you add a coefficient, update your table with the new quantities of each atom. _____ C H4 + _____ 2 O2  _____ 2 H2 O + _____ C O2 Reactants Products C 1 1 H 4 2 4 O 2 4 3 4

How to Balance By Inspection: OPTIONAL Place a “1” in any empty coefficient location 3 Filling each coefficient location lets you and the grader know that you finished the problem rather than you left some blank because you weren’t done! _____ 1 C H4 + _____ 2 O2  _____ 2 H2 O + _____ 1 C O2 Reactants Products C 1 1 H 4 2 4 O 2 4 3 4

Coefficients are added to change the number of atoms in a substance! Total: 4 H 2 O Total: 4 H 2 O The equation is balanced.

Do You Remember Polyatomic Ions? Polyatomic ion – Group of atoms that together has a net charge e.g. Nitrate NO31- Carbonate CO32- A polyatomic ion is a group of atoms that together have a charge.

Balancing with Polyatomic Ions: 1 Make a table of elements Hint: Chunk the polyatomic ions—IF they appear on both sides OH is a polyatomic ion that is sometimes “hidden” in H2O. Re-write H2O as HOH to “see” the OH polyatomic ion. H OH _____ H3 PO4 + _____ Ca (OH)2  _____ Ca3 (PO4)2 + _____ H2 O Reactants Products H PO4 Ca OH

Balancing with Polyatomic Ions: Count the number of each element or ion on the reactants and products side. 2 H OH _____ H3 PO4 + _____ Ca (OH)2  _____ Ca3 (PO4)2 + _____ H2 O Reactants Products H 3 1 PO4 1 2 Ca 1 3 OH 2 1

Balancing with Polyatomic Ions: 3 Add coefficients to balance the numbers H OH _____ 2 H3 PO4 + _____ 3 Ca (OH)2  _____ Ca3 (PO4)2 + _____ 6 H2 O Reactants Products H 3 6 1 6 PO4 1 2 2 Ca 1 3 3 OH 2 6 1 6

Balancing with Polyatomic Ions: 5 Place a “1” in any empty coefficient location H OH _____ 2 H3 PO4 + _____ 3 Ca (OH)2  _____ 1 Ca3 (PO4)2 + _____ 6 H2 O Reactants Products H 3 6 1 6 PO4 1 2 2 Ca 1 3 3 OH 2 6 1 6

Let’s Practice : without a chart Example: Balance the following equation __ HCl + __ Ca(OH)2  __ CaCl2 + __ H2O

Balance the following equation Example: Balance the following equation Did you see the “OH” polyatomic ion & change H2O to HOH? 2 1 1 2 __ HCl + __ Ca(OH)2  __ CaCl2 + __ H2O HOH

Balance the following equation SELF CHECK Example: Balance the following equation _____ Li + _____ H2 O  ____ LiOH + _____ H2

Balance the following equation Example: Balance the following equation 2 2 __ Li + __ H2O  __ LiOH + ___ H2 2 1

Balance the following equation SELF CHECK Example: Balance the following equation __ Fe + __ O2  ___ Fe2O3

Balance the following equation Example: Balance the following equation 4 3 2 __ Fe + __ O2  ___ Fe2O3

BALANCING HINTS Never start balancing by placing a 1 as a coefficient. Do that last. Leave H & O for last when balancing. Chuck the polyatomic Ions Change H2O to HOH if the hydroxide ion is present on both sides If there is an odd number of 1 type of atom on one side and an even number of the same atom on the other side, make the odd number even by multiplying by 2. Make sure the set of coefficients are in the simplest ratio When finished balancing, always re-check that all of your atoms are equal on both sides. In a combustion reaction, always start balancing with the C and H of the hydrocarbon

1. 2 5 4 2 combustion 2. 1 1 1 1 SR 3. 1 2 2 1 SR 4. 2 2 1 decomp Practice Problems 1. 2 5 4 2 combustion 2. 1 1 1 1 SR 3. 1 2 2 1 SR 4. 2 2 1 decomp 5. 1 1 1 synthesis 6. 16 1 8 synthesis 7. 1 1 1 2 DR 8. 4 3 4 3 SR

Continuation of Answers 9. 1 2 1 2 DR 10. 3 1 1 synthesis 11. 4 1 2 synthesis 12. 1 3 1 decomp 13. 1 2 1 1 SR 14. 2 3 2 3 DR 15. 1 1 1 2 DR

Predicting Products of a Chemical Reaction Section 3 To be successful in this section, you will need to use the reaction models found on the Reference Sheet

Steps to Predicting Products of a Chemical Reaction Identify the type of reaction. Use the reference sheet to determine which model to use. Create products for your reaction following the model as a guide. As you form new ionic compounds, don’t forget to check charges. Make sure to recognize the elements of BrINClHOF & format them correctly. Balance the chemical equation.

Synthesis Reaction: Model 1A Formation of a Binary Ionic Compound What do you do? Bring symbols together and check charges Example: K + Br2 

Synthesis Reaction: Model 1B Metal oxide & water forms a base MOH What do you do? Combine metal symbol with hydroxide (OH) & check charges Example: CaO + H2O 

Synthesis Reaction: Model 1C Nonmetal oxide & water forms an acid What do you do? Add up all atoms starting with H, then the nonmetal, ending with O. Example: CO2 + H2O 

Decomposition Reaction: Model 2A Break down of a Binary Compound What do you do? Break binary compound apart into the 2 elements that make it up Look for diatomic elements! Example: CaCl2 

Decomposition Reaction: Model 2B Metal Carbonates What do you do? Place metal next to oxygen & check charges Then add carbon dioxide, CO2 Example: K2CO3 

Decomposition Reaction: Model 2C Metal Hydrogen Carbonates (bicarbonates) What do you do? Place metal next to oxygen & check charges Then add water, H2O & then carbon dioxide, CO2 Example: NaHCO3 

Decomposition Reaction: Model 2D Metal Hydroxides What do you do? Place metal next to oxygen & check charges Then add water, H2O Example: LiOH 

Decomposition Reaction: Model 2E Metal Chlorates What do you do? Place metal next to chlorine & check charges Then add oxygen gas, O2 Example: AgClO3 

Decomposition Reaction: Model 2F OxyAcids What do you do? Remove water (H2O) from the oxyacid to form the nonmetal oxide Example: H2SO3 

Using the Activity Series: The Activity Series is a chart that shows the reactivity of metals found on your reference sheet Elements on top of the chart are more reactive than elements on the bottom More reactive elements can only replace less reactive elements within a compound

Self Check Using the activity series, determine if the following replacements can happen: K + CaCl2  Cu + Zn(OH)2  Na + H2O  Ni + H2O  Mg + HCl 

Single Replacement Reaction: Model 3A More Reactive Metal replacing Less Reactive Metal What do you do? Check Activity Series: Is the single metal above the metal in the compound? If YES, replace it & check charges of new compound. Be on alert for diatomic elements! If NO, write NR for “no reaction”.

A + BC  AC + B Example: Zn + Cu(NO3)2  metal + ionic comp  ionic comp. + metal Example: Zn + Cu(NO3)2 

Single Replacement Reaction: Model 3B More Reactive Metal replacing Hydrogen from Water What do you do? Check Activity Series: Is the single metal above the Hydrogen? Look at the specific details about temperature of water. If YES, replace it & check charges of new compound. Then add hydrogen gas, H2. If NO, write NR for “no reaction”.

Examples: Na + H2O  Al + H2O  Al + H2O (g) M + H2O  MOH + H2 metal + water  metal hydroxide + hydrogen gas Examples: Na + H2O  Al + H2O  Al + H2O (g)

Single Replacement Reaction: Model 3C More Reactive Metal replacing Hydrogen from an Acid What do you do? Check Activity Series: Is the single metal above the Hydrogen? If YES, replace it & check charges of new compound. Then add hydrogen gas, H2. If NO, write NR for “no reaction”.

M + HX  MX + H2 Example: Mg + HCl  metal + acid  ionic comp. + hydrogen Example: Mg + HCl 

Single Replacement Reaction: Model 3D More Reactive Halogen replacing Less Reactive Halogen What do you do? Check Activity Series: Is the single halogen above the halogen in the chart? If YES, replace it & check charges of new compound. Be on alert for diatomic elements! If NO, write NR for “no reaction”.

D + BC  BD + C Examples: Cl2 + KBr  I2 + NaCl  Element + ionic compound  ionic compound + element (halogen) (halogen) Examples: Cl2 + KBr  I2 + NaCl 

Solubility Rules Chart

Solubility Rules Chart: The Solubility Rules Chart classifies substances as soluble or insoluble found on your reference sheet Categorized by Anion SOLUBLE- means that the substance will dissolve in water INSOLUBLE- means that the substance will not dissolve in water and so it is considered a PRECIPITATE when on the product side (PP) Insoluble = Precipitate

Let’s Practice #1: decide whether each is soluble or insoluble NaNO3 Fe(C2H3O2)2 CaBr2 Ba(OH)2 Cu(OH)2

NaNO3 Fe(C2H3O2)2 CaBr2 Ba(OH)2 Cu(OH)2 Soluble Not Soluble

Self Check Using the Solubility Rules, determine if the following substances are soluble or insoluble(precipitate): 1. CaCl2 4. Zn(OH)2 2. K3PO4 5. NiSO4 3. MgCO3 6. AgNO3

Double Replacement Reaction: Model 4A Formation of a Precipitate What do you do? Swap the positive ions with each other & check the charges on the new compounds Using the Solubility Rules chart, identify the product that is a precipitate (INSOLUBLE)

Precipitation Reactions A precipitation reaction is when 2 soluble substances are mixed together and they form an insoluble substance. This is called a precipitate. Reactants 2 soluble chemicals: NaOH and Cu(NO3)2 NaOH Cu(NO3)2

Precipitation Reactions(DR Rxns) NaOH + CuNO3  NaNO3 + Cu(OH)2 Precipitation Reactions(DR Rxns) Na+1 OH-1 Cu+2 NO3 -1 Cu(OH)2(S) Products: 1 soluble chemical: NaNO3 1 insoluble chemical (the precipitate): Cu(OH)2 Na+1 NO3 -1

AB + CD  AC + BD Example: NaCl + AgNO3 metal + ionic comp  ionic comp. + metal Example: NaCl + AgNO3

Double Replacement Reaction: Model 4B Acid Base Reaction What do you do? Identify the reactants as an acid & base. No need to use solubility rules chart! Swap the positive ions with each other & check the charges on the new compounds which will always be a salt & water.

AB + CD  AC + BD Example: KOH + H2SO4  Acid + Base  ionic comp. + water Example: KOH + H2SO4 

CombustionReaction: Model 5 CxHy + O2  CO2 + H2O Hydrocarbon + oxygen  carbon dioxide + water What do you do? Identify the hydrocarbon and O2 as reactants Products will always be CO2 and H2O no matter what the formula of the hydrocarbon Examples: C3H8 + O2  CH3OH + O2 

Section 3: Part C Net Ionic Reactions – Precipitation Shows the details of aqueous reactions that involve ions in aqueous solution Molecular Equation: the typical equation you are use to writing keeping all molecules together Complete Ionic Equation: shows all the particles in a solution as they really exist, as IONS or MOLECULES. Anything aqueous needs to be split apart into the cation and anion Anything solid stays intact Coefficients need to be multiplied by subscripts to determine the exact amount of each cation and anion.

Spectator ions: ions that do not participate in a reaction; they are identical on both sides of the equation & are crossed out! Net Ionic Equation: the final equation showing the major players. All spectator ions have been removed.

NET IONIC REACTIONS for Precipitation Reactions Molecular equation: KI(aq) + AgNO3(aq)  AgI(s) + KNO3(aq) Complete Ionic equation: K+1 + I-1 + Ag+1+ NO3-1  AgI + K+1 + NO3-1 Spectator ions: ions that do not participate in a reaction; they are identical on both sides of the equation & are crossed out! Net Ionic equation: I-1 + Ag+1  AgI

NET IONIC REACTIONS for Precipitation Reactions Molecular equation: 2 NaOH(aq) + CuCl2(aq)  2 NaCl(aq) + Cu(OH)2(s) Complete Ionic equation: 2 Na+1 + 2 OH-1 + Cu+2 + 2 Cl-1  2 Na+1 + 2 Cl-1 + Cu(OH)2 Net Ionic equation: 2 OH-1 + Cu+2  Cu(OH)2

Take Home Practice: Predict products and balance Iron (III) chloride reacts with sodium hydroxide Molecular equation: 1 FeCl3(aq) + 3 NaOH(aq)  1 Fe(OH)3(s) +3 NaCl(aq) Complete Ionic equation: 3 Na+1 + 3 OH-1 + Fe+3 + 3 Cl-1  3 Na+1 + 3 Cl-1 + Fe(OH)3 Net Ionic equation: 3 OH-1 + Fe+3  Fe(OH)3