Chemistry Unit 7 - Chemical Reactions Notes

Chemistry Unit 7 - Chemical Reactions Notes A chemical reaction is represented by a chemical equation such as:

1. Na + Cl2 → NaCl

read as: sodium and chlorine yield sodium chloride (→ is yields sign) sodium reacts with chlorine to produce sodium chloride reactants: left of yield sign (react) products: right of yield sign (are produced)

Balancing Chemical Equations:  First , count and list how many elements you have on each side  start balancing (usually metals first, then non-metals) WITH COEFFICIENTS  balance polyatomic ions as a group if they are on both sides  leave anything that is on one side in more than one place until last (usually H, O)

 the seven diatomic elements: Br I N Cl H O F, when alone have a subscript of 2 2. examples:

A. MgBr2 + Cl2 → MgCl2 + Br2 already balanced (SR)

B. Cl2 + 2 NaI → 2 NaCl + I2 (SR)

C. 2 KNO3 → 2 KNO2 + O2 (D)

D. Zn + 2 HCl → ZnCl2 + H2 (SR)

E. CaO + 2 HCl → CaCl2 + H2O (DR) Homework: Write on another sheet of paper if you need more room!

F. Al + 3 Pb(NO3)2 → 2 Al(NO3)3 + 3 Pb (SR)

G. Cu + 2 AgNO3 → Cu(NO3)2 + 2 Ag(SR)

H. 2 K + 2 H2O → 2 KOH + H2 (SR)

I. Cl2 + 2 LiI → 2 LiCl + I2 (SR)

J. Ca(OH)2 + 2 HCl → CaCl2 + 2 H2O (DR)

K. 3 KOH + H3PO4 → 3 H2O + K3PO4 (DR)

L. 2 Al(NO3)3 + 3 H2SO4 → Al2(SO4)3 + 6 HNO3 (DR)

M. Na2SO3 + 2 HCl → 2 NaCl + H2O + SO2 no type (redox)

N. 2 KNO3 → 2 KNO2 + O2 (D)

O. 2 PbO2 → 2 PbO + O2 (D)

P. 2 NaOH → Na2O + H2O (D)

Q. MgCO3 → MgO + CO2 already balanced (D)

R. 2 Na + Cl2 → 2 NaCl (S)

S. Br2 + 2 H2O + SO2 → 2 HBr + H2SO4 no type (redox)

T. CaO + H2O → Ca(OH)2 already balanced (S) U. P2O5 + 3 BaO → Ba3(PO4)2 (S)

Types of Reactions:  synthesis: A + B → AB

 decomposition: AB → A + B

 combustion

fuel (C, H, O) + O2 → CO2 + H2O

 Single Replacement ( + activity series):

A + BC → AC + B

 Double Replacement: ( + solubility rules):

AB + CD → AD + CB Symbols: subscripts: (s) solid or ↓ precipitate formed (l) liquid

(g) gas or ↑ gas formed

(aq) aqueous

“yield” sign: ↔ or reversible (at equilibrium) ∆ → heat added Pt/Pd → catalyst added What does a catalyst do to a chemical reaction? speeds up the reaction

Take a look at this picture for a better idea of what a catalyst does:

If the products have lower energy than the reagents, the reaction is said to be exothermic. Exothermic reactions give off heat, causing the things that were formed to have less energy than they did before they reacted.

It may seem weird to say that something that gives off heat would have less energy than something that doesn't, but let's consider yet another analogy. Let's say that for some reason or another, you've decided to go lie on a block of ice in your swimsuit. From the perspective of the block of ice, you're awfully warm. However, from your perspective, you're losing energy like crazy to the block of ice. When the ice has melted all the way, you'll have a lot less energy than you did before. That's how something that feels hot loses energy.

If the products have greater energy than the reagents, the reaction is endothermic. Endothermic reactions absorb heat, making them feel cold.

Again, let's say that you're lying on this block of ice. From your perspective, the block of ice feels very cold. Why is this? It's because the ice is absorbing all of your energy from you. The water actually has more energy than it did before you sat on it in ice form, because it absorbed all that energy from your body. Generally, endothermic reactions take place more slowly than exothermic reactions, because their activation energies are higher.

Heat of reaction, or H, is just a measure of how exothermic or endothermic a reaction is. A reaction that's very exothermic would have a very negative H value, while a reaction that's only slightly exothermic would have a slightly negative H value. Heats of reaction aren't affected at all by the addition of a catalyst to the reaction, because no matter how it happens, the result of the reaction is the same.

After all, if your friend buys a candy bar for you, in the end, you still have the same candy bar as you would have had if you'd argued with your enemy to get it.

Potential Energy Diagram 1. Which of the letters a–f in the diagram represents the potential energy of the products? e

2. Which letter indicates the potential energy of the activated complex? c

3. Which letter indicates the potential energy of the reactants? a

4. Which letter indicates the activation energy? b

5. Which letter indicates the heat of reaction (ΔH)? f

6. Is the reaction exothermic or endothermic? endothermic (+ΔH)

7. Which letter indicates the activation energy of the reverse reaction? d

8. Which letter indicates the heat of reaction of the reverse reaction (ΔH)? f

9. Is the reverse reaction exothermic or endothermic? exothermic (-ΔH)

10. DRAW on the diagram to show how the curve would be different if there was a catalyst. Write the chemical reactions and balance.  balance charges for all ionic compounds before balancing equation with coefficients  positive ion always comes first in an ionic compound  diatomic molecules Br I N Cl H O F, subscript of 2 when alone  Identify the type of reaction - synthesis, decomposition, combustion, single replacement or double replacement  Include state of matter subscripts when possible (s, l, g, aq)

1. Aqueous magnesium bromide reacts with gaseous chlorine to make aqueous magnesium chloride and bromine gas.

MgBr2 (aq) + Cl2 (g) → MgCl2 (aq) + Br2 (g)

already balanced

2. Aqueous silver nitrate reacts with aqueous sodium chloride to make solid silver chloride and aqueous sodium nitrate.

AgNO3 (aq) + NaCl (aq) → AgCl( s) + NaNO3 (aq)

already balanced

3. Gaseous potassium nitrate decomposes into aqueous potassium nitrite and oxygen gas.

2 KNO3 (g) → 2 KNO2 (aq) + O2 (g)

4. Zinc metal reacts with aqueous hydrochloric acid to make aqueous zinc chloride and hydrogen gas.

Zn (s) + 2 HCl (aq) → ZnCl2 (aq) + H2 (g)

5. Solid calcium oxide reacts with aqueous hydrochloric acid to make aqueous calcium chloride and water.

CaO( s) + 2 HCl (aq) → CaCl2 (aq) + H2O (l)

6. Sodium (s) reacts with water to produced aqueous sodium hydroxide and hydrogen gas.

2 Na (s) + 2 H2O (l) → 2 NaOH (aq) + H2 (g) 7. Mercury (II) oxide decomposes to give mercury and oxygen. 2 HgO (s) → 2 Hg (l) + O2 (g)

8. Magnesium reacts with aluminum chloride and makes magnesium chloride and aluminum.

3 Mg (s) + 2 AlCl3 (aq) → 3 MgCl2 (aq) + 2 Al (s)

9. Iron and iodine produce iron (II) iodide.

Fe (s) + I2 (s) → FeI2 (aq)

already balanced

10. Chlorine reacts with potassium iodide to make iodine and potassium chloride.

Cl2 (g) + 2 KI (aq) → I2 (s) + 2 KCl(aq)

Chemistry Name:______Single Replacement Date:______Single Replacement Reactions: 1. Decide which ion the element would replace 2. Compare them on the activity series 3. If the element is higher on the list, the reaction will occur 4. If the element is lower on the list, the reaction will not occur, write NR. 5. If the reaction occurs, write the products *balance charges! 6. The displaced element is neutral * (diatomics) Br I N Cl H O F 7. Balance the reaction using coefficients.

1. Ca(NO3)2 + 2 Li → 2 LiNO3 + Ca

2. 2 K + FeCO3 → K2CO3 + Fe

3. Mg + 2 HCl → MgCl2 + H2

4. 2 NaI + Cl2 → 2 NaCl + I2

5. Fe2O3 + 2 Al → Al2O3 + 2 Fe

6. Na + AuClO3 → NaClO3 + Au

7. I2 + NiF2 → NR

Redox Reactions Oxidation: loss of electrons (example: Ca0 → Ca2+)

Reduction: gain of electrons (example: Ca2+ → Ca0)

OXIDATION # IS REDUCED

LEO the lion says GER

OIL RIG LEO GER

Chemistry Name:______Double Replacement Date:______Double Replacement Reactions: 1. The elements exchange their positive and negative ions 2. Write the products (the ions will keep the same charges) *balance charges! 3. Balance the reactions using coefficients. Assume both reactant are (aq) 4. Label the products with an (aq) or (s) as a subscript using solubility rules

8. H2SO4 + Sr(OH)2 → 2 HOH(l) + SrSO4 (s) (neutralization reaction)

9. Al2(SO4)3 + 3 Ba(ClO3)2 → 2 Al(ClO3)3 (aq) + 3 BaSO4 (s)

10. (NH4)2SO4 + Ca(OH)2 → NH4OH(aq) + CaSO4 (aq) (no precipitate)

11. 3 AgC2H3O2 + Na3PO4 → Ag3PO4 + 3 NaC2H3O2

12. 2 H3PO4 + 3 Mg(OH)2 → 6 HOH (l) + Mg3(PO4)2 (s) (neutralization reaction)

______Not in solubility rules, use rule that’s closest: 2- 2- 2- 3- 2- - - SO4 for SO3 and S2O3 , PO4 for H2PO4 , ClO3 for BrO3

13. Mg(H2PO4)2 + CuSO3 → MgSO3 + Cu(H2PO4)2

14. Sc2(S2O3)3 + 6 CuBrO3 → 2 Sc(BrO3)3 + 3 Cu2S2O3

Activity Series Metals Non-Metals

K Cl2 Ba Br2

Sr I2 Ca Na Mg Al Mn Zn Fe Cd Co Ni Sn Pb H Cu *Diatomic elements: subscript of 2 Ag Hg Br I N Cl H O F Au *Solubility Rules for ionic compounds: Compounds containing the following ions are generally soluble in water. + + + + 1) alkali metal ions and ammonium ions (Li , Na , K , NH4 ) - 2) acetate ions (C2H3O2 ) - 3) nitrate ions (NO3 ) - 4) chlorate ions (ClO3 ) 5) halide ions (Cl-, Br-, I-) (except for silver, mercury (I) and lead (II), which are insoluble) -2 6) sulfate ion (SO4 ) except for strontium, barium and lead (II), which are insoluble)

Compounds containing the following ions are generally INsoluble in water. -2 7) carbonate ions (CO3 ) (see rule 1 for exceptions that are soluble) -2 8) chromate ions (CrO4 ) (see rule 1 for exceptions that are soluble) -3 9) phosphate ions (PO4 ) (see rule 1 for exceptions that are soluble) 10) sulfide ions (S-2) (calcium, barium, lead II and rule 1 exceptions are soluble) 11) hydroxide ions (OH-) (calcium, barium, strontium and rule 1 exceptions are soluble)

Net Ionic Equations

Double Replacement Reaction (general form): AB + CD → AD + CB (see example below!) 1. Predict products 2. Balance the reaction 3. Label products as aq (soluble) or s (insoluble) 4. Write complete ionic reaction: all ions – precipitate will NOT DISSOCIATE because it does not dissolve 5. Cancel spectator ions to have the net ionic reaction

Vocab: spectator ions: “watch” reaction take place, do not participate, not included in net ionic + - (K and NO3 in following example) Reaction:

AgNO3 (aq) + KCl(aq) → AgCl (aq) + KNO3 (aq)

Complete Ionic:

+ - + - + - Ag (aq) + NO3 (aq) + K (aq) + Cl (aq) → AgCl(s) + K (aq) + NO3 (aq)

Net Ionic: cancel spectators (shown above)

+ - Ag (aq) + Cl (aq) → AgCl(s) Practice:

1. Al2(SO4)3 (aq) + 3 Sr(NO3)2 (aq) → 2 Al(NO3)3 (aq) + 3 SrSO4 (s)

2+ 2- Sr (aq) + SO4 (aq) → SrSO4 (s)

2. Pb(NO3)2 (aq) + 2 NH4I(aq) → 2 NH4NO3 (aq) + PbI2 (s)

2+ - Pb (aq) + 2 I (aq) → PbI2 (s)

3. Li2CrO4 (aq) + CaBr2(aq) → 2 LiBr (aq) + CaCrO4 (s)

2+ 2- Ca (aq) + CrO4 (aq) → CaCrO4 (s)

4. NaNO3 (aq) + KCl(aq) → KNO3 (aq) + NaCl (aq) no net ionic Chemistry Name:______Double Replacement Reactions Worksheet Date:______For each of the precipitation reactions, do the following: a. Predict the products b. Balance the equation to get the molecular (un-ionized) equation. c. Label products as (aq) or (s) d. Determine (and balance) the net ionic equation – also using (aq) and (s). Remember**  The metal ions will keep the same charge they have as a reactant and product  Assume that all the reactants are dissolved in solution  Zn always form a Zn2+ ion  Ag always forms a Ag+ ion  HCl is a strong acid meaning that it completely ionizes (dissolves) in solution

+ - 1. NaI + AgNO3 → NaNO3 (aq) + AgI(s) Net: Ag (aq) + I (aq) → AgI(s)

2+ 3- 2. 2Na3PO4 + 3BaCl2 → 6NaCl (aq) + Ba3(PO4)2(s)Net:3Ba (aq) + 2PO4 (aq) → Ba3(PO4)2(s)

2+ 2- 3. Na2SO4 + BaCl2 → 2NaCl (aq) + BaSO4(s)Net: Ba (aq) + SO4 (aq) → BaSO4(s)

+ - 4. NaCl + AgNO3 → NaNO3 (aq) + AgCl(s) Net: Ag (aq) + Cl (aq) → AgCl(s) 2+ 2- 5. BaCl2 + H2SO4 → 2HCl(aq) + BaSO4 (aq) Net: Ba (aq) + SO4 (aq) → BaSO4(s)

3+ - 6. FeCl3 + 3NH4OH → 3NH4Cl (aq) + Fe(OH)3(s)Net: Fe (aq) + 3OH (aq) →Fe(OH)3(s)

3+ 2- 7. 2BiCl3 + 3H2S → 6HCl (aq) + Bi2S3(s) Net: 2Bi (aq) + 3S (aq) → Bi2S3(s)

2+ 2- 8. K2CO3 + NiBr2 → 2KBr (aq) + NiCO3(s)Net: Ni (aq) CO3 (aq) → NiCO3(s)

2+ 2- 9. ZnCl2 + (NH4)2S → 2NH4NO3 (aq) + ZnS(s)Net: Zn (aq) + S (aq) → ZnS(s)

10. aluminum nitrate + sodium phosphate

Al(NO3)3 + Na3PO4 → 3NaNO3 (aq) + AlPO4 (s) 3+ 3- Net: Al (aq) PO4 (aq) → AlPO4 (s)

11. sodium carbonate + silver nitrate

Na2CO3 + 2AgNO3 → 2NaNO3 (aq) + Ag2CO3 (s) 2+ 2- Net: Ag (aq) + CO3 (aq) → Ag2CO3 (s)

12. iron (II) nitrate + sodium phosphate

3Fe(NO3)2 + 2Na3PO4 → 6NaNO3 (aq) + Fe3(PO4)2 (s) 3+ Net: 3Fe(aq) + 2PO4 (aq) → Fe3(PO4)2 (s)

For each of the neutralization reactions, assume that the only net ionic is simply the formation of + - water: H (aq) + OH (aq) → HOH(l) Do the following: a. Predict the products b. Balance the equation c. Label the acid, base and salt

13. HBr + KOH → HOH(l) + KBr(aq)

14. 2HClO4 + Ca(OH)2 → 2HOH(l) + Ca(CrO4)2 (s)

15. Fe(OH)2 + H2S → 2HOH(l) + FeS(s)

16. hydrochloric acid + barium hydroxide 2HCl + Ba(OH)2 → 2HOH(l) + BaCl2 (aq)

17. magnesium hydroxide + nitric acid

Mg(OH)2 + 2HNO3 → 2HOH(l) + Mg(NO3)2 (aq)

18. sulfuric acid + sodium hydroxide H2SO4 + 2NaOH → 2HOH(l) + Na2SO4 (aq)

Chemistry Name:______

Unit 6 Review Sheet Test Date: ______

Vocabulary:  reactant  product  yields  precipitate  soluble  insoluble  aqueous  spectator ions  catalyst  balance  coefficient  subscript  diatomic elements  exothermic  endothermic  activation energy  synthesis  decomposition  single replacement  double replacement  combustion

Symbols (know what they mean and how to use them):  →  (s)  (l)  (g)  (aq)  ∆ above →

 Pt, Pd, MnO2 above →  ↔ or  know the seven diatomic elements (Br I N Cl H O F) Do:  balance a chemical equation (from formulas or words)  classify a reaction as synthesis, decomposition, single replacement, double replacement or combustion  identify the two products of a complete combustion reaction (between a hydrocarbon and oxygen)  use the activity series to determine if a single replacement reaction will occur, write the products and balance with coefficients, identify the species oxidized and reduced  use solubility rules to determine aqueous solutions and precipitates in double replacement reactions, write the products, balance with coefficients, and write the net ionic equation, if there is one write “no net ionic” (remember that in a neutralization reaction the formation of water is a net ionic)

Explain:  how the “law of conservation of mass” applies to balanced chemical equations  explain how a catalyst works

Bonus:  More difficult predicting products and balancing

Chemistry Name:______balancing and net ionic equations Date:______

Write the chemical reactions and balance.  balance charges for all ionic compounds before balancing equation with coefficients  positive ion always comes first in an ionic compound  diatomic molecules Br I N Cl H O F, subscript of 2 when alone  Identify the type of reaction - synthesis, decomposition, combustion, single replacement or double replacement  Include state of matter subscripts when possible (s, l, g, aq) 1. Carbon and oxygen combine to produce carbon monoxide.

2. Hydrogen and oxygen combine to produce hydrogen peroxide. 3. Zinc is replaced in zinc bromide by manganese. Manganese (IV) bromide is produced in this single displacement reaction.

4. Potassium bromide combines with barium selenide in a standard double displacement reaction.

5. Barium and fluorine react. There is only one possible product.

6. Strontium chloride will react with tellurium. There is only one possible set of products in this single replacement reaction.

7. Cesium sulfate reacts with cobalt to form cobalt (III) sulfate.

8. Zirconium (III) nitrate decomposes, producing zirconium (II) nitride, nitrogen and oxygen.

9. Heptane (C7H16) reacts in a classical combustion reaction.

10. Nitric acid combines with barium hydroxide to produce barium nitrate and water. Write the full (molecular) equation AND the net ionic equation for the reaction between each of the following pairs of compounds. If no reaction can occur, explain why. 11. sodium phosphate and magnesium iodide

12. sodium carbonate and hydrosulfuric acid

13. iron (III) acetate and sodium hydroxide

14. lead (II) acetate and hydrobromic acid

15. A barium nitrate solution is added to a sodium sulfate solution.