CHEMICAL REACTIONS Reactants: Zn + I2 Product: Zn I2 1 Introduction – Chemical reactions occur when bonds (between the electrons of atoms) are formed or broken – Chemical reactions involve • changes in the chemical composition of matter (the making of new materials with new properties) • energy changes – Symbols represent elements – Formulas describe compounds – Chemical equations describe a chemical reaction 2 Chemical Equations (eqns.) What is a chemical equation? How do you balance a chemical equation? How do you identify the type of chemical equation? Chemical Equations A chemical equation is written as an expression similar to a mathematic equation that can be compared to a recipe that a chemist follows in order to produce desired results. 5 Chemical Equations 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 are called stoichiometric coefficients The letters (s), (g), and (l) are the physical states of compounds. Chemical Equations Because of the principle of the conservation of matter (matter can not be created or destroyed) must be balanced. an equation It must have the same number of atoms of the same kind on both sides. Law of Conservation of Energy MUST ALSO BE FOLLOWED! • Energy changes are written in (endo-/ exothermic reactions) Lavoisier, 1788 6 Chemical Equations All chemical equations have reactants and products. We express a chemical equation as follows: Reactants Products The arrow is equivalent to an “=“ math. When we describe the equation we use the word “yields” or “produces” instead of equals Example C + O2 CO2 This reads “carbon plus oxygen react to yield carbon dioxide” Balancing a Chemical Equation A chemical equation is balanced when the ions or atoms found on the reactant side of the equation equals that found on the product side. The arrow can be considered the balance point. 9 Symbols Used in Equations • Solid (s) • Liquid (l) • Gas (g) • Aqueous solution (aq) H2SO4 or Pt • Catalyst • Escaping gas () • Depositing solid (↓) • Change of temperature/ heat energy ( or + 3kJ or – 3kJ) 10 Chemical Reactions and Chemical Equations Reactants Products Signs of Driving forces: a) Color change b) Formation of a solid/precipitate c) Evolution of a gas d) Evolution or absorption of heat In all cases, a more energetically stable (lower energy product) is created 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) – Think back to naming compounds/ determining formulas. NaCl exists, because Na is + and Cl is -, but NaCl2 does NOT exist since you would not have a neutral compound! You can’t just add a number to a formula to balance an equation. 11 Subscripts vs. Coefficients 12 • 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. Chemical Equations 4 Al(s) + 3 O2(g) ---> 2 Al2O3(s) This equation means 4 Al atoms + 3 O2 molecules ---produces---> 2 molecules of Al2O3 AND/OR 4 moles of Al + 3 moles of O2 ---produces---> 2 moles of Al2O3 13 14 Steps to Balancing Equations 15 There are four basic steps to balancing a chemical equation. 1. 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! 2. 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. 3. 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. 4. Check your answer to see if: – The numbers of atoms on both sides of the equation are now balanced. – The coefficients are in the lowest possible whole number ratios. (reduced) Some Suggestions to Help You Some helpful hints for balancing equations: • Take one element at a time, working left to right except for H and O. Metals, then nonmetals are a good way, too. Save H for next to last, and O until last. • IF everything balances except for O, and there is no way to balance O with a whole number, double all the coefficients and try again. (Because O is diatomic as an element) • (Shortcut) Polyatomic ions that appear on both sides of the equation should be balanced as independent units 16 Balancing Equations ___ 2 H2(g) + ___ O2(g) ---> ___ 2 H2O(l) What Happened to the Other Oxygen Atom????? This equation is not balanced! Two hydrogen atoms from a hydrogen molecule (H2) combine with one of the oxygen atoms from an oxygen molecule (O2) to form H2O. Then, the remaining oxygen atom combines with two more hydrogen atoms (from another H2 molecule) to make a second H2O molecule. 17 18 Balance this equation! Na + Cl2 Na- 1 Cl- 2 NaCl Na-1 Cl-1 **note that the number of sodiums balance but the chlorine does not. We will have to use coefficients in order to balance this equation. 19 Inserting subscripts Na + Cl2 Na- 1 Cl- 2 2 NaCl Na- 1 2 Cl- 1 2 ** Now the chlorine balances but the sodium does not! So we go back and balance the sodium. 20 Finally balanced! 2Na + Cl2 Na- 1 2 Cl- 2 2 NaCl Na-1 2 Cl-1 2 **Since the number of each element on the reactant side and the product side of the equation are equal, the equation is balanced. Balancing Equations 2 Al(s) + ___ 3 Br2(l) ---> ___ Al2Br6(s) ___ 21 Balancing Equations Sodium phosphate + iron (III) oxide sodium oxide + iron (III) phosphate 2 Na3PO4 + 3 Na2O + Fe2O3 ----> 2 FePO4 22 23 Balancing Equation Practice 1. CuCl3 + Li2S Cu2S3 + LiCl 2. NiNO3 + KCl NiCl + KNO3 3. FeCl3 + Na2O Fe2O3 + NaCl 24 Answers: • 1. 2CuCl3 + 3Li2S Cu2S3 + 6LiCl • 2. NiNO3 + KCl NiCl + KNO3 (already balanced) • 3. 2FeCl3 + 3Na2O Fe2O3 + 6NaCl Types of Chemical Reactions (rxns.) Types of Reactions • • Reactions are classified by their products. There are five types of chemical reactions we will talk about: 1. 2. 3. 4. 5. • Synthesis reactions Decomposition reactions Single displacement reactions Double displacement reactions Combustion reactions You need to be able to identify the type of reaction and predict the product(s) Steps to Writing Reactions • Some steps for doing reactions 1. 2. 3. Identify the type of reaction Predict the product(s) using the type of reaction as a model Balance it Don’t forget about the diatomic elements! (BrINClHOF) 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 a compound! Synthesis or Combination reactions Synthesis (meaning to make) or combination reactions are typified by their single product. If you have a reaction in which at least 2 elements or compounds are reacted and produce a single product, the reaction is a synthesis reaction. 1. Synthesis reactions • • Synthesis reactions are sometimes called combination or addition reactions. reactant + reactant 1 product Basically: A + B AB • • • Example: 2H2 + O2 2H2O Example: C + O2 CO2 Note: Single Product! This is your clue that this is a synthesis or combination reaction. Synthesis Reactions • Here is another example of a synthesis reaction Practice • • • • Predict the products. Write and balance the following synthesis reaction equations. Sodium metal reacts with chlorine gas 2 Na(s) + Cl2(g) 2 NaCl(s) Solid Magnesium reacts with fluorine gas Mg(s) + F2(g) MgF2(s) Balanced Aluminum metal reacts with fluorine gas 2 Al(s) + 3 F2(g) 2 AlF3(s) 2. Decomposition Reactions Decomposition reactions are really just the opposite of a synthesis reaction. Remember, if you can make a substance, you should be able to break it back apart into its components. A good way to remember decomposition reactions to to remember what happens when something decomposes. It falls apart! Decomposition Reactions • • • • • • Decomposition reactions occur when a compound breaks up into the elements or in a few to simpler compounds 1 Reactant Product + Product Basically: AB A + B Example: 2 H2O 2H2 + O2 Example: 2 HgO 2Hg + O2 Note: Single Reactant! The single reactant is your clue that this is a decomposition reaction. Decomposition Reactions • Another view of a decomposition reaction: 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 this class Practice • • • Predict the products. Then, write and balance the following decomposition reaction equations: Solid Lead (IV) oxide decomposes PbO2(s) Pb(s) + O2(g) Aluminum nitride decomposes 2 AlN(s) 2 Al(s) + N2(g) Practice Identify the type of reaction for each of the following synthesis or decomposition reactions, and write the balanced equation: N2(g) + O2(g) 2 NO (g) BaCO3(s) BaO(s) + CO2 (g) Co(s)+ S(s) Co2S3 (s) NH3(g) + H2CO3(aq) (NH4)2CO3(s) I2 (s) NI3(s) N2 (g) + 3. Single Replacement Reactions Single replacement reactions occur when one chemical takes the place of another in a reaction. In the typical single replacement reaction, an element trades places with one of the ions in a compound. 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 product + 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!) When H2O splits into ions, it splits into H+ and OH- (not H+ and O-2 !!) Single Replacement Reactions • Another view: The Activity Series ► Not all single replacement reactions will occur. ► This depends upon the location of the elements present in the activity series ► Elements above MAY replace elements below; elements below MAY NOT replace elements above them on the series You will be given a copy of this!!!! Single Replacement Reactions Write and balance the following single replacement reaction equation: • Zinc metal reacts with aqueous hydrochloric acid Zn(s) + 2 HCl(aq) ZnCl2 + H2(g) Note: Zinc replaces the hydrogen ion in the reaction • [If ZnCl2 + H2(g) Zn(s) + HCl(aq) the reaction WOULD NOT OCCUR because Hydrogen is below zinc on the activity series] Single Replacement Reactions • Sodium chloride solid reacts with fluorine gas 2 NaCl(s) + F2(g) 2 NaF(s) + Cl2(g) Note that fluorine replaces chlorine in the compound • Aluminum metal reacts with aqueous copper (II) nitrate 2 Al(s)+ 3 Cu(NO3)2(aq)3 Cu(s) +2 Al(NO3)3(aq) 4. Double Replacement Reactions Double replacement reactions are identified by two ions trading places and forming new compounds. Double Replacement Reactions • • • • Double Replacement Reactions occur when a metal replaces a metal in a compound and a nonmetal replaces a nonmetal in a compound Compound + compound product + product AB + CD AD + CB Notice that one ion from compound AB replaces one ion from compound CD. 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) 2 KNO3(aq) + BaSO4(s) Solubility For a double replacement reaction to have occurred, a solid (precipitate) MUST be formed OR water must be formed in an acid-base reaction. There are rules to determine which of the materials formed is the solid If no solid is formed, there is said to be no reaction. Figure 8.4: The forming of solid AgCl. Solubility Tables Solubility tables help determine which materials are soluble in water and which are not In general, Solubility Rules can be summarized as follows 1. All compounds containing alkali metal cations and the ammonium ion 2. 3. 4. 5. 6. are soluble. All compounds containing NO3-, ClO4-, ClO3-, and C2H3O2- anions are soluble. All chlorides, bromides, and iodides are soluble except those containing Ag+, Pb2+, or Hg22+. All sulfates are soluble except those containing Hg22+, Pb2+, Sr2+, Ca2+, or Ba2+. All hydroxides are insoluble except compounds of the alkali metals, Ca2+, Sr2+, and Ba2+. All compounds containing PO43-, S2-, CO32-, and SO32- ions are insoluble except those that also contain alkali metals or NH4+. You will be given a copy of this!!!! Practice • Predict the products. Balance the equation 1. HCl(aq) + AgNO3(aq) 2. Pb(NO3)2(aq) + BaCl2(aq) 3. FeCl3(aq) + NaOH(aq) 4. H2SO4(aq) + NaOH(aq) HNO3(aq) + AgCl(s) PbCl2(s) + Ba(NO3)2(aq) Fe(OH)3(s) + NaCl(aq) H2O(l) + Na2SO4(aq) Combustion Reactions Combustion reactions are the ones that burn (or explode!). There are two types of combustion reactions—complete or incomplete reactions. These reactions are identified by their products. They either produce carbon monoxide and water or carbon dioxide and water. Complete Combustion Reactions These reactions burn “efficiently” which means they produce carbon dioxide and water. These reactions typically burn cleanly and leave very little residue behind. 5. Combustion Reactions • Combustion reactions occur when a hydrocarbon 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) Combustion Reactions • In general: CxHy + O2 CO2 + H2O • Products in combustion 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) Complete Combustion Reactions CH4 + O2 CO2 + H2O They may also be written: CH4 CO2 + H2O (O2 is usually written above the arrow.) Clue: CO2 (carbon dioxide) in the product along with water Combustion Reactions Edgar Allen Poe’s drooping eyes and mouth are potential signs of CO poisoning. Combustion • Example • C5H12 + 8 O2 5 CO2 + 6 H2O • Write the products and balance the following combustion reaction: 10 CO2 + 22 11 H2O • 2 C10H22 + 31O2 20 Incomplete Combustion Reactions Incomplete combustion reactions occur when something does not burn efficiently. This can cause a lot of harm if the gases produced cannot escape. Carbon monoxide,an odorless and colorless gas, is dangerous. People poisoned by this gas usually become sleepy and can die due to exposure. Incomplete Combustion Reactions CH4 + O2 CO + H2O These reactions may also be written by: CH4 CO + H2O (the O2 is usually written over the arrow.) Clue: CO (Carbon monoxide as a product.) Mixed Practice • 1. 2. 3. 4. 5. State the type, predict the products, and balance the following reactions: BaCl2 + H2SO4 BaSO4 + HCl C6H12 + O2 CO2 + H2O Zn + CuSO4 ZnSO4 + Cu Cs + Br2 CsBr FeCO3 FeO + CO2 Predicting Products of Reactions Completing reactions requires knowledge of the different reaction types (sometimes called mechanisms). You must first identify the reaction type by the reactants. The only type of reaction that cannot be predicted this way is the combustion reaction since the products are very similar. First Step: Identify reaction type Example: Al + O2 Clue: 2 elements – Synthesis or combination reaction Second Step: Write the net ionic equation for the reactants Al + O2 becomes Al3+ + O2- Step 3 Using clues, complete reaction taking care to write each formula correctly by checking charges and “criss-crossing” if necessary. Al + O2 Al3+O2Al + O2 Al2O3 Predicting Products of Reactions (cont.) For Single Replacement reactions, check activity series to make sure the reaction goes. Once you write the molecular equation, you should check for reactants and products that are soluble or insoluble. (Double Replacement only) Reactions in Aqueous Solutions a.k.a. Net Ionic Equations Molecular Equations: shows complete formulas for reactants and products – Does not show what happens on the molecular level Total (or Complete) Ionic Equations: All substances that are strong electrolytes (are soluble and dissociate) are written as their ions. – Some ions participate in the reaction – Some ions do NOT participate in the reaction-called spectator ions. Net Ionic Equations: show only the ions that participate in the reaction Writing Total Ionic Equations Once you write the molecular equation (synthesis, decomposition, etc.), you should check for reactants and products that are soluble or insoluble. We usually assume the reaction is in water We can use a solubility table to tell us what compounds dissolve in water. If the compound is soluble (does dissolve in water), then splits the compound into its component ions If the compound is insoluble (does NOT dissolve in water), then it remains as a compound Writing Total Ionic Equations Molecular Equation: K2CrO4 + Pb(NO3)2 PbCrO4 + 2 KNO3 Soluble Insoluble Soluble Soluble Total Ionic Equation: 2 K+ + CrO4 -2 + Pb+2 + 2 NO3- PbCrO4 (s) + 2 K+ + 2 NO3- Net Ionic Equations These are the same as total ionic equations, but you should cancel out ions that appear on BOTH sides of the equation Total Ionic Equation: 2 K+ + CrO4 -2 + Pb+2 + 2 NO3- PbCrO4 (s) + 2 K+ + 2 NO3 (Spectator ions) Net Ionic Equation: CrO4 -2 + Pb+2 PbCrO4 (s) Net Ionic Equations Try this one! Write the molecular, total ionic, and net ionic equations for this reaction: Silver nitrate reacts with Lead (II) Chloride in hot water. AgNO3 + PbCl2 Molecular: 2 AgNO3 + PbCl2 2 AgCl + Pb(NO3)2 Total Ionic: 2 Ag+ + 2 NO3- + Pb+2 + 2 Cl- 2 AgCl (s) + Pb+2 + 2 NO3Net Ionic: Ag+ + Cl- AgCl (s) Acid-Base Reactions Acid: – produces hydrogen ions (H+) in solution (Arrhenius) – proton donor (Lewis) Base: – produces hydroxide ions (OH-) in solution (Arrhenius) – proton acceptor (Lewis) The reaction ALWAYS forms water and an ionic compound (sometimes aqueous). Acid-Base Reactions Example HNO3 (aq) + KOH (aq) Molecular: HNO3 (aq) + KOH (aq) H2O (l) + KNO3 (aq) Total Ionic: H+ (aq) + NO3- (aq) + K+ (aq) + OH- (aq) H2O (l) + K+ (aq) + NO3- (aq) Net Ionic: H+ (aq) + OH- (aq) H2O (l) Oxidation-Reduction Reactions a.k.a. Redox Equations Between a metal and a nonmetal forming an ionic compound Electron transfer occurs Oxidation numbers: assigning an excess or deficiency in electrons for each element (the charge based on the compound). Rules for Oxidation Number (ox. #) Determination 1. The sum of the oxidation numbers add up to the charge a. all elements have an ox. # of 0 b. ions of elements, ox. # is the charge (Cl-) c. the sum of the ox. # of a complex ion equals the charge (CO3-2 ) 2. H is 1+ when combined with a nonmetal and 1- with a metal H3PO4 H= 1+ PO4 = 3- CaH2 Ca = 2+ H= 1- Rules for Oxidation Numbers (cont.) 1. 2. 3. F is always 1-; Cl, Br, I are 1- except when combined with each other or O O is 2- except when combined with F (F2O) Group I is 1+ and Group II is 2+ in their compounds Recognizing Redox Rxns. oxidation 1+ 2 HCl (aq) + Mg (s) MgCl2 (aq) + H20 (q) 0 Net: 2+ reduction 2 H+ (aq) + Mg2+ (aq) Mg 2+ (aq) + H2 0 (q) Loss of electron = oxidation Gain of electron = reduction “LEO the lion goes GER" Half Reactions Separate the individual oxidation and reduction reactions. Look at electron movement Half rxn.: Mg0 (aq) Mg 2+ + 2 e2e- + 2 H+ H20 Net: 2 H+ (aq) + Mg0 (aq) Mg 2+ (aq) + H2 0 (q) Oxidizing agent: the one reduced (H+) Reducing agent: the one oxidized (Mg0) Recognition of Redox rxns. Oxidation # changes Reactions with oxygen Reaction of any element (forms a new compound) Balancing Balance by mass Balance by charge Balance net ionic equation Example Problem: Fe (s) + Cl2 (aq) Fe3+ (aq) + Cl- (aq) 1. Balance by mass Fe (s) + Cl2 (aq) Fe3+ (aq) + 2 Cl- (aq) 2. Write half reaction Fe0 (s) Fe 3+ + 3 e2e- + Cl2 2 Cl- 3. Balance by charge (want # of e- to cancel) (Fe0 (s) Fe 3+ + 3e-) 2 + (2e- + Cl2 2 Cl ) 3 2 Fe0 (s) + 6e- + 3Cl2 2 Fe 3+ + 6e- + 6 ClFinal eqn.: 2 Fe (s) + 3Cl2 2 Fe 3+ + 6 Cl- Again on types of reactions/ summary/ new ways of looking at it The Myth! MYTH: Little Mikey (“Mikey Likes It”) from the LIFE cereal commercials in the late 70s died when he ate pop rocks then drank a coke, causing the pop rocks to explode inside his stomach. The Myth! Not surprisingly, this one is completely false. The Evidence: – Mikey is alive and well. His real name is John Gilchrist and he’s an advertising manager for a New York Radio Station. – There isn’t enough carbonation in pop rocks to release more than a tiny amount of CO2 – much less than in a coke. If the myth were true, coke alone should be able to explode your stomach. Types of Chemical Reactions We can divide chemical reactions into five basic types. First Type: – Synthesis Reaction: – Looks like this: A+B AB – In this reaction, substance A and substance B “combine” to form a new compound. – Substances A and B can be elements in atomic form, elements in molecular form, or compounds. Examples of Synthesis Reactions 2Na + S Na2S – This one is an example of two elements in atomic form (Na and S) combining to form a compound (sodium sulfide). 2H2 + O2 2H2O – In this example, A and B are two elements in molecular form (hydrogen and oxygen molecules), and the product is water, which is simply the chemical combination of hydrogen and oxygen. Examples of Synthesis Reactions 2Fe + O2 2FeO – In this example, substance “A” is an element in atomic form (Fe), and substance “B” is an element in molecular form (O2). The result is a direct chemical combination of the two elements (FeO, iron oxide, which is “rust”). CuO + H2O Cu(OH)2 – This is an example where both substances going into the reaction are molecules. The result is what you get when you “add” all of the atoms in the reaction together. Chemical Reactions Second Type: – Decomposition Reaction: – Looks like: AB A+B – Basically, the opposite of a synthesis reaction. – A and B (i.e. the products of the reaction) can be elements (atomic or molecular form), or compounds. – Example: 2HgO 2Hg + O2 Chemical Reactions Third Type: – Single Replacement – Looks like this: AB + C AC + B – As usual, A and B can be elements or compounds, but C is typically an element. – Example: 2H2O + 2Na 2NaOH + H2 Here a sodium atom replaces a hydrogen atom in water. The reaction is a bit easier to see if we write water as hydrogen hydroxide and keep the freed hydrogen as an atom: HOH + Na NaOH + H Chemical Reactions Fourth Type: – Double Replacement: – Looks like: AB + CD AD + BC – A, B, C, and D can be elements or compounds (they are usually an element or a polyatomic ion). – Basically, the atoms (or polyatomic ions) in the reactions “trade partners.” Chemical Reactions Examples: – Most of our “practice” equations are examples of this type of reaction. – 2H3PO4 + 3Mg(OH)2 Mg3(PO4)2 + 6H2O – H2SO4 + Mg(OH)2 MgSO4 + 2H2O Chemical Reactions Important Type of Double Replacement: Acid-Base Neutralization Reactions An acid: lots of free hydrogen ions (H+). A base: lots of free hydroxide ions (OH-). When you combine the two, the hydrogen ions in the acid combine with the hydroxide ions in the base to make water. The substances that “donated” the hydrogen and hydroxide ions also chemically combine to make a special class of compounds called salts. Acid-Base Neutralization Chemically the reaction looks like this: Acid + Base Salt + Water A classic example: HCl + NaOH NaCl + H2O Hydrochloric Acid Water Sodium Hydroxide (Lye) Sodium Chloride (Table Salt) Acid-Base Neutralization Here’s the equation again: – HCl + NaOH NaCl + H2O – Chemically, this is a double replacement reaction: The H traded its Cl for an OH The Na traded its OH for a Cl. Chemical Reactions Hydrocarbon – A class of compounds formed by combining hydrogen and carbon. Often go by their common names. Examples: – CH4 is methane – C2H6 is ethane – C3H8 is propane – C4H10 is butane – C8H18 is octane Chemical Reactions Fifth Type: – Combustion Reaction – A hydrocarbon reaction with oxygen. – “Complete” combusion: the products are always CO2 and H2O (and a lot of energy in the form of heat and light). – “Incomplete” combusion: the products are CO2, H2O, CO, and a “simpler” hydrocarbon (one with fewer carbons). Combustion Reactions Complete reactions occur when there is plenty of oxygen available for the reaction. – Typically hard to ignite. – Burn very hot. Incomplete reactions occur when there isn’t enough oxygen available for a “complete” reaction. – Easier to ignite. – Cooler burning. Combustion Internal combustion engines (i.e. car engines) typically don’t mix enough oxygen in with the hydrocarbon (gasoline) to support a complete reaction. Result: Leftover hydrocarbons (toxic), and the production of carbon monoxide (very toxic). Both are major contributors to pollution. Combustion Would letting more “air into the mix” (for a more complete reaction) solve air pollution caused by cars? No! – “Hard to ignite” means the car is more prone to flooding and stalling. – “Hotter burning” has several detrimental effects: Hard on the engine’s oil. Creates various nitrogen-oxygen compounds that can form powerful acids (e.g. nitric acid), and other very toxic pollutants.