Ch. 9 Chemical Reactions & Equations SAVE PAPER AND INK!!! When you print out the notes on PowerPoint, print "Handouts" instead of "Slides" in the print setup. Also, turn off the backgrounds (Tools>Options>Print>UNcheck "Background Printing")! Ch. 9 Chemical Reactions & Equations Zn + I2 Reactants Zn I2 Product

Introduction Chemical reactions occur when bonds between atoms are formed or broken Chemical reactions involve changes in matter, the making of new materials with new properties, and energy changes. Symbols represent elements, Formulas represent compounds, Chemical equations represent chemical reactions

Combustion of Methane CH4 + 2O2  CO2 + 2H2O Atoms are rearranged!

Chemical Equations Yields or produces! Their Job: Depict the kind of reactants and products and their relative amounts in a reaction. 4 Al (s) + 3 O2 (g) > 2 Al2O3 (s) The numbers in the front of formulas are called COEFFICIENTS The letters (s), (g), and (l) are the physical states of compounds. Starting Materials What’s created

The charcoal used in a grill is basically carbon The charcoal used in a grill is basically carbon. The carbon reacts with oxygen to yield carbon dioxide.

This is the “Word Equation” for that reaction Word Equations: show the names of the reactants and the products Lavoisier, 1788 This is the “Word Equation” for that reaction carbon + oxygen carbon dioxide A skeleton equation does NOT indicate the relative amounts of reactants and products (no coefficients!) The skeleton equation for that reaction is: C + O2  CO2 Chemical Equations must be balanced in order to conform to the Law of Conservation of Mass - same # & type of atoms on each side of the yield arrow.

Skeleton equation: CH4 + O2 CO2 + H2O Word Equation: Methane + Oxygen gas carbon dioxide + water Skeleton equation: CH4 + O2 CO2 + H2O

Symbols Used in Equations Solid (s) Liquid (l) Gas (g) Aqueous solution (aq) (dissolved in water) Catalyst H2SO4 Escaping gas () Change of temperature () Precipitate ( )

Balancing Equations When balancing a chemical reaction you may add coefficients in front of the compounds to balance the reaction, but you may not change the subscripts. Changing the subscripts changes the compound. Subscripts are determined by the valence electrons (charges for ionic or sharing for covalent) Never put a coefficient in the middle of a formula 2 NaCl is ok Na2Cl is not.

Subscripts vs. Coefficients The subscripts tell you how many atoms of a particular element are in a compound. The coefficient tells you about the quantity, or number, of molecules of the compound.

There are a number of ways to interpret balanced equations 2 H2(s) + O2(g) ---> 2 H2O(s) This equation means: 2 molecules H2 + 1 molecules O2 ---->2 molecules H2O(g)

Steps to Balancing Equations There are four basic steps to balancing a chemical equation. Write the correct formula for the reactants and the products. DO NOT TRY TO BALANCE IT YET! You must write the correct formulas first. And most importantly, once you write them correctly DO NOT CHANGE THE FORMULAS! Find the number of atoms for each element on the left side. Compare those against the number of the atoms of the same element on the right side. Determine where to place coefficients in front of formulas so that the left side has the same number of atoms as the right side for EACH element in order to balance the equation. Check your answer to see if: The numbers of atoms on both sides of the equation are equal. The coefficients are in the lowest possible whole number ratios. (reduced)

Some Suggestions to Help You Helpful Hints for balancing equations: Take one element at a time. Save pure elements for last IF everything balances except for that last pure element, and there is no way to balance it with a whole number, double all the coefficients and try again. (Shortcut) Polyatomic ions that appear on both sides of the equation should be balanced as independent units That is, don’t separate them into individual atoms!

Example Make a table to keep track of atoms H2 + O2  H2O R P 2 H O 1 Need another O on the product side

Example H2 + O2  2H2O Place a coefficient of 2 in front of H2O R P 2 1 Changes the O

Example H2 + O2  2H2O R P 2 H O 1 2 This also changes the H

Example H2 + O2  2H2O R P 2 H O 1 4 2 Now we need twice as much H in the reactant

Example 2H2 + O2  2H2O R P 2 H O 1 4 2 Add a coefficient of 2 in front of H2

Example 2H2 + O2  2H2O Your answer R P 2 H O 1 4 4 2 Recount to check

Balancing Equations Na3PO4 + Fe2O3 ---> Na2O + FePO4 Sodium phosphate + iron (III) oxide  sodium oxide + iron (III) phosphate Na3PO4 + Fe2O3 ---> Na2O + FePO4 3 2 2 R L 2 1 3 1 2 6 2 3 Na PO4 Fe O 6 2

4 10 2 11 __B4H10 + ___O2 ---->___ B2O3 + ___H2O __C3H8 + __O2 ----> __CO2 + __ H2O 5 3 4 3 Balancing Equations

Types of Reactions There are millions of chemical reactions. The only way to be sure what your products will be is to carry them out in the lab! Not very practical – or cost effective, however… There are five types of chemical reactions we will talk about: Synthesis reactions Decomposition reactions Single Replacement reactions Double Repalcement reactions Combustion reactions You need to be able to identify the type of reaction and predict the product(s)

1. Synthesis reactions Synthesis reactions occur when two or more substances (generally elements) combine to form a single compound. (Sometimes these are called combination reactions.) reactant + reactant  1 product Basically: A + B  AB Example: 2H2 + O2  2H2O Example: C + O2  CO2

Synthesis Reactions Here is another example of a synthesis reaction

Balancing Equations 2 3 ___ Al(s) + ___ Br2(l) ---> Al2Br6(s)

Practice Predict the products. Write and balance the following synthesis reaction equations. Sodium metal reacts with chlorine gas Na(s) + Cl2(g)  Solid Magnesium reacts with fluorine gas Mg(s) + F2(g)  Aluminum metal reacts with fluorine gas Al(s) + F2(g) 

2. Decomposition Reactions Decomposition reactions occur when a single compound breaks up into one or more elements or simpler compounds 1 Reactant  Product + Product In general: AB  A + B Example: 2 H2O  2H2 + O2 Example: 2 HgO  2Hg + O2

Decomposition Reactions Another view of a decomposition reaction:

Example: Decomposition of hydrogen peroxide in the presence of a catalyst. Word equation: hydrogen peroxide  water + oxygen

Example: Decomposition of hydrogen peroxide in the presence of a catalyst. Skeleton equation: H2O2  H2O + O2

Example: Decomposition of hydrogen peroxide in the presence of a catalyst. Balanced Equation: 2H2O2  2H2O + O2

Example: Decomposition of hydrogen peroxide in the presence of a catalyst. Balanced equation showing the catalyst (MnO2) the state of the reactants and products: MnO2 2H2O2(l)  2H2O (l) + O2 (g)

Decomposition Exceptions Carbonates and chlorates are special case decomposition reactions that do not go to the elements. Carbonates (CO32-) decompose to carbon dioxide and a metal oxide Example: CaCO3  CO2 + CaO Chlorates (ClO3-) decompose to oxygen gas and a metal chloride Example: 2 Al(ClO3)3  2 AlCl3 + 9 O2 There are other special cases, but we will not explore those in Chemistry I

Practice N2(g) + O2(g)  BaCO3(s)  Co(s)+ S(s)  Nitrogen monoxide Identify the type of reaction for each of the following synthesis or decomposition reactions, and write the balanced equation: N2(g) + O2(g)  BaCO3(s)  Co(s)+ S(s)  NH3(g) + H2CO3(aq)  NI3(s)  Nitrogen monoxide (make Co be +3)

3. Single Replacement Reactions Single Replacement Reactions occur when one element replaces another in a compound. A metal can replace a metal (+) OR a nonmetal can replace a nonmetal (-). element + compound element + product A + BC  AC + B (if A is a metal) OR A + BC  BA + C (if A is a nonmetal) (remember the cation always goes first!) Use Activity Series of Metals to see if reaction works.

We have looked at several reactions: Fe + CuSO4  Cu + Fe2(SO4)3 Li + H2O  LiOH + H2 Such experiments reveal trends. The activity series ranks the relative reactivity of metals. It allows us to predict if certain chemicals will undergo single displacement reactions when mixed: metals near the top are most reactive and will displacing metals near the bottom. Q: Which of these will react? Fe + CuSO4  Ni + NaCl  Li + ZnCO3  Al + CuCl2  K Na Li Ca Mg Al Zn Fe Ni Sn Pb H Cu Hg Ag Au No, Ni is below Na Yes, Li is above Zn Yes, Al is above Cu Yes, Fe is above Cu Cu + Fe2(SO4)3 NR (no reaction) Zn + Li2CO3 Cu + AlCl3

A: No Mg + H2O  No Reaction H is the only nonmetal listed. H2 may be displaced from acids or can be given off when a metal reacts with H2O (producing H2 + metal hydroxide). The reaction with H2O depends on metal reactivity. Q: will Mg react with H2O? K Na Li Ca Mg Al Zn Fe Ni Sn Pb H Cu Hg Ag Au A: No Mg + H2O  No Reaction Complete these reactions: Al + HCl Cu + Sn2S  Ca + Cu2SO4  Na + PbO  No Reaction NR 2 Cu + CaSO4 2 Pb + Na2O

Single Replacement Reactions Another view:

2 Al + Fe2O3 2 Fe + Al2O3 Here is a mixture of aluminum powder and iron (III) oxide powder. Once ignited the aluminum rips the oxygen off of the iron oxide resulting in molten iron, which falls down to the pot of sand below. One time someone used wet sand, which was a mistake because the molten iron turned the water to steam instantly. This sprayed molten iron all over the front two rows of students. Write the balanced equation for this reaction.

Single Replacement Reactions Write and balance the following single replacement reaction equation: Zinc metal reacts with aqueous hydrochloric acid Zn(s) + HCl(aq)  ZnCl2 + H2(g) Note: Zinc replaces the hydrogen ion in the reaction 2

Single Replacement Reactions Sodium chloride solid reacts with fluorine gas NaCl(s) + F2(g)  NaF(s) + Cl2(g) Note that fluorine replaces chlorine in the compound Aluminum metal reacts with aqueous copper (II) nitrate Al(s)+ Cu(NO3)2(aq) 2 2

4. Double Replacement Reactions Double Replacement Reactions occur when a metal in one compound replaces a metal in another compound. compound + compound  compound + compound AB + CD  AD + CB

Double Replacement Reactions Think about it like “foil”ing in algebra, first and last ions go together + inside ions go together Example: AgNO3(aq) + NaCl(s)  AgCl(s) + NaNO3(aq) Another example: K2SO4(aq) + Ba(NO3)2(aq)  KNO3(aq) + BaSO4(s) 2

Practice Predict the products. Then, write and balance the following decomposition reaction equations: Solid Lead (IV) oxide decomposes PbO2(s)  Aluminum nitride decomposes AlN(s) 

MnO2 + CO --> Mn2O3 + CO2 Carbon monoxide is commonly used to strip off oxygen atoms from metals. This reaction shows the first step. If more CO is present, eventually all oxygen atoms will be grabbed by CO and manganese (Mn) metal will be left. This is how metal ores get converted to metals. 2 MnO2 + CO --> Mn2O3 + CO2 Balance this equation!

Practice Predict the products. Balance the equation HCl(aq) + AgNO3(aq)  CaCl2(aq) + Na3PO4(aq)  Pb(NO3)2(aq) + BaCl2(aq)  FeCl3(aq) + NaOH(aq)  H2SO4(aq) + NaOH(aq)  KOH(aq) + CuSO4(aq) 

5. Combustion Reactions Combustion reactions occur whenever something reacts with oxygen gas. This is also called burning!!! In order to burn something you need the 3 things in the “fire triangle”: 1) A Fuel (hydrocarbon) 2) Oxygen to burn it with 3) Something to ignite the reaction (spark)

Complete Combustion of Hydrocarbons In general: CxHy + O2  CO2 + H2O Products are ALWAYS carbon dioxide and water. (although incomplete burning does cause some by-products like carbon monoxide) Combustion is used to heat homes and run automobiles (octane, as in gasoline, is C8H18)

Combustion Example C5H12 + O2  CO2 + H2O Write the products and balance the following combustion reaction: C10H22 + O2  5 8 6

Combustion Reactions Edgar Allen Poe’s drooping eyes and mouth are potential signs of CO poisoning.

Mixed Practice State the type, predict the products, and balance the following reactions: BaCl2 + H2SO4  C6H12 + O2  Zn + CuSO4  Cs + Br2  FeCO3 

Some steps for Writing Reactions Identify the type of reaction Predict the product(s) using the type of reaction as a model Remember! If you write a formula – You MUST balance it! 3. Balance the equation (using coefficients) Don’t forget about the diatomic elements! (BrINClHOF or HON 7) For example, Oxygen is O2 as an element. In a compound, it can’t be a diatomic element because it’s not an element anymore, it’s part of a compound!