IIIIIIIVV I.Intro to Reactions (p. 282 – 285) Ch. 9 – Chemical Reactions

n Chemical change n Atoms of one or more substances (reactants) are rearranged into new substances (products) n Signs of a chemical reaction  Change of temperature or light  Formation of a gas (bubbles)  Formation of a precipitate  Odor  Color change  Change in volume Chemical Reaction

Law of Conservation of Mass n mass is neither created nor destroyed in a chemical reaction 4 H 2 O 4 H 2 O 4 g32 g 36 g n total mass stays the same n atoms can only rearrange

Chemical Equations A+B  C+D REACTANTSPRODUCTS Starting substances Substances formed

p. 283 Chemical Equations

Writing Equations n Identify the substances involved. n Use symbols to show: 2H 2 (g) + O 2 (g)  2H 2 O(g)  How many? - coefficient  Of what? - chemical formula  In what state? - physical state n Remember the diatomic elements!

Describing Equations n Describing Coefficients:  individual atom = “atom”  covalent substance = “molecule”  ionic substance = “unit” 3 molecules of carbon dioxide 2 atoms of magnesium 4 units of magnesium oxide 3CO 2  2Mg  4MgO 

Writing Equations Two atoms of solid aluminum react with three units of aqueous copper(II) chloride to produce three atoms of solid copper and two units of aqueous aluminum chloride. How many? Of what? In what state? 2 Al (s) + 3 CuCl 2 (aq)  3 Cu (s) + 2 AlCl 3 (aq)

Describing Equations How many? Of what? In what state? Zn(s) + 2HCl(aq)  ZnCl 2 (aq) + H 2 (g) One atom of solid zincreact with two molecules of aqueous hydrochloric acid to produce one unit of aqueous zinc chloride and one molecule of hydrogen gas.

IIIIIIIVV II. Balancing Equations Ch. 9 Chemical Reactions

Why is there a 2 after the oxygen? __________________ Oxygen is diatomic

Balancing Steps 1.Write the unbalanced equation. 2.Count atoms on each side. 3.Add coefficients to make #s equal. Coefficient  subscript = # of atoms 4.Reduce coefficients to lowest possible ratio, if necessary. 5.Double check atom balance!!!

Helpful Tips n Balance one element at a time. n Update ALL atom counts after adding a coefficient. n If an element appears more than once per side, balance it last. n Balance polyatomic ions as single units.  “1 SO 4 ” instead of “1 S” and “4 O”

Counting Atoms 1. H 3 PO KC 2 H 3 O 2 2. Mg 3 (PO 4 ) Al 2 (SO 4 ) 3 H – 3 PO 4 – 1 K – 2 C 2 H 3 O 2 – 2 Mg – 3 PO 4 – 2 Al – 6 SO 4 – 9

Practice Balancing! n All the atoms on the reactants side must equal all the atoms on the products side n First, you need to count all the atoms H 2 + O 2  H 2 O H O

How to Balance n Are they equal? n Add coefficients to change the number of atoms H 2 + O 2  H 2 O H O Balanced!! Good job Balanced!! Good job

Al + CuCl 2  Cu + AlCl 3 Al Cu Cl  3  6   Balancing Example Aluminum and copper(II) chloride react to form copper and aluminum chloride. 2  2  6 Good Job!

Practice… Sn +HF  SnF 2 + H 2 Cu+ AgNO 3  Cu(NO 3 ) 2 +Ag Sn H F Great! Cu Ag NO Good Job!

IIIIIIIVV Ch. 9 – Chemical Reactions III. Types of Chemical Reactions

Combustion CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(g) n The burning of any substance in O 2 to produce heat, usually a hydrocarbon n Products are always carbon dioxide and water vapor C X H Y + O 2  CO 2 + H 2 O

Synthesis (composition) n the combination of 2 or more elements or compounds to form a more complex compound n Basic Form: A + X  AX

Examples of Synthesis Reactions 1) Metal + oxygen → metal oxide n EX. 2 Mg (s) + O 2(g) → 2 MgO (s) 2) Nonmetal + oxygen → nonmetallic oxide n EX. C (s) + O 2(g) → CO 2(g) 3) Metal + nonmetal → salt n EX. 2 Na (s) + Cl 2(g) → 2 NaCl (s) 4) A few nonmetals combine with each other. n EX. 2 P (s) + 3 Cl 2(g) → 2 PCl 3(g)

Decomposition n A single compound breaks down into its component parts of simpler compounds n Basic form: AX  A + X

Examples of Decomposition Rxns 1) Some oxides, when heated, decompose. n EX. 2 HgO (s) → 2 Hg (l) + O 2(g) 2) Some are produced by electricity. n EX. 2 H 2 O (l) → 2 H 2(g) + O 2(g) EX. 2 NaCl (l) → 2 Na (s) + Cl 2(g)

Single Replacement n a more active element takes the place of another element in a compound and sets the less active one free n Basic Form: A + BX  AX + B or AX + Y  AY + X Examples of replacement reactions: 1) Replacement of a metal in a compound by a more active metal. n Fe (s) + CuSO 4(aq) → FeSO 4(aq) + Cu (s)

Examples of Single Replacement Rxns 2) Replacement of hydrogen in water by an active metal. n Note: it is helpful to think of water as HOH (H + OH - ) n 2 Na (s) + 2 H 2 O (l) → 2 NaOH (aq) + H 2(g) or 2 Na (s) + 2 HOH (l) → 2 NaOH (aq) + H 2(g) n Mg (s) + 2 H 2 O (g) → Mg(OH) 2 (s) + H 2(g) or Mg (s) + 2 HOH (g) → Mg(OH) 2 (s) + H 2(g)

Examples of Single Replacement Rxns 3) Replacement of hydrogen in acids by active metals. n EX. Zn (s) + 2 HCl (aq) → ZnCl 2(aq) + H 2(g) 4) Replacement of nonmetals by more active nonmetals. n EX. Cl 2(g) + 2 NaBr (aq) → 2 NaCl (aq) + Br 2(l) n If…… n EX. Br 2(g) + 2 NaCl (aq) → No Reaction!

Single Replacement n NOTE: Refer to the activity series for metals and nonmetals to predict products of replacement reactions.activity series for metals n If the free element is above the element to be replaced in the compound, then the reaction will occur. n If it is below, then no reaction (NR) occurs.

Double Replacement (Ionic) n ions in two compounds “change partners” ; cation of one compound combines with anion of the other n Occurs between ions in aqueous solution. A reaction will occur when a pair of ions come together to produce at least one of the following:  a precipitate (s)  water (l)  a gas (g) n Basic form: AX + BY → AY + BX n Complex form: ABX + CDY → ADY + CBX

Examples of Double Replacement Rxns 1) Formation of precipitate (s): n NaCl (aq) + AgNO 3(aq) → NaNO 3(aq) + AgCl (s) n BaCl 2(aq) + Na 2 SO 4(aq) → 2 NaCl (aq) + BaSO 4(s) 2) Formation of water (l). (If the reaction is between an acid and a base it is called a neutralization reaction.) n HCl (aq) + NaOH (aq) → NaCl (aq) + HOH (l) 3) Formation of a gas (g). n HCl (aq) + FeS (s) → FeCl 2(aq) + H 2 S (g)

4) Formation of a product which decomposes. n Ex: CaCO 3(aq) + 2 HCl (aq) → CaCl 2(aq) + H 2 CO 3 (aq) → CaCl 2 (aq) + CO 2 (g) + H 2 O (l) n Ex: 2 NaOH (aq) + (NH 4 ) 2 S (aq) → Na 2 S (aq) + 2 NH 4 OH (aq) → Na 2 S (aq) + 2 NH 3 (g) + H 2 O (l)

Examples of unstable products: (you need to know these three) A) H 2 CO 3(aq) → CO 2(g) + H 2 O (l) n Carbonic acid, as in soft drinks, spontaneously decomposes when it is formed. B) NH 4 OH (aq) → NH 3(g) + H 2 O (l) n Ammonium hydroxide produces the odor of ammonia gas because whenever it is formed it spontaneously decomposes into ammonia & water. C) H 2 SO 3(aq) → SO 2(g) + H 2 O (l) n Sulfurous acid, as a bleaching agent and disinfectant, spontaneously decomposes when it is formed.

Double Replacement n NOTE: Use the solubility rules to decide whether a product of an ionic reaction is insoluble in water and will thus form a precipitate. If a compound is soluble in water then it should be shown as being in aqueous solution, or left as separate ions. n If it products are both in solution (aqueous) then  No Net Ionic Reaction occurs!

Aqueous Rxns: Double Replacement n Ex: in the reaction involving the ionic compounds silver nitrate and potassium chloride we have: n AgNO 3 (aq) + KCl (aq)  AgCl (s) + KNO 3 (aq) n The driving force for double replacement reactions is the removal of ions from solution

Precipitation Reactions n Double replacement rxns that result in an insoluble precipitate n Solubility: amount of a substance that can be dissolved in a given quantity of water n The solubility of an ionic compound determines whether it will precipitate or not. n The reaction of KI and Pb(NO 3 ) 2 : 2KI (aq) +Pb(NO 3 ) 2(aq)  PbI 2(s) + 2KNO 3 (aq) n The PbI 2 is an insoluble ionic compound that will precipitate out of the solution

Precipitation Reactions n Can we predict whether an ionic compound will be soluble or not? n If an ionic compound is insoluble it means that neighboring ions have an attraction for each other that is greater than the attraction of water for the ions n Unfortunately, there are no clear rules for solubility based on physical properties of ions. n General behaviors of certain ions are observed by consulting A Solubility Table!

Neutral Molecular Compounds n Even though the neutral molecular compound may be soluble in aqueous solution, its formation is essentially irreversible. n Thus, ions are effectively removed from solution by this irreversible process n The neutralization reaction of HCl and NaOH: HCl (aq) + NaOH (aq)  H 2 O (l) + NaCl (aq) n The formation of the covalent compound (H 2 O) from the proton and hydroxide ions is essentially irreversible and drives the double replacement reaction (even though we would consider H 2 O to be "highly soluble" in H 2 O).

Gas Formation in Double Replacement n When a double replacement reaction involves the formation of a gas (& the gas is not soluble in H 2 O) the loss of the gas can drive the double replacement reaction (i.e. the gas is lost - therefore, it is an irreversible process) n Gasses that can form from ionic compounds include:  CO 2 (carbon dioxide)  H 2 S (hydrogen sulfide - smells like rotten eggs)  NH 3 (ammonia).

Formation of CO 2 from carbonic acid n The bicarbonate ion (HCO 3 - ) can combine with a proton to produce carbonic acid (H 2 CO 3 ): HCl (aq) + NaHCO 3 (aq)  NaCl (aq) + H 2 CO 3 (aq) n Carbonic acid in water is unstable & spontaneously decomposes to form water and carbon dioxide gas: H 2 CO 3 (aq)  H 2 O (l) + CO 2 (g) n The carbon dioxide gas is lost, & thus the formation of carbonic acid is irreversible and drives the double replacement reaction. The reaction is thus: HCl (aq) + NaHCO 3 (aq)  NaCl (aq) + H 2 O (l) + CO 2 (g)

Gases produced in Spontaneous Rxns n The three typical gasses produced take place with the following spontaneous reactions (only when they appear on the product side): KNOW THEM n H 2 CO 3(aq) → CO 2(g) + H 2 O (l) n NH 4 OH (aq) → NH 3(g) + H 2 O (l) n H 2 SO 3(aq) → SO 2(g) + H 2 O (l)

Net Ionic Reactions - DR n Complete ionic equation = equation that shows all dissolved ionic compounds as dissociated free ions n Net ionic equation = Equation for a reaction in solution that shows only those particles that are directly involved in the chemical change n Spectator Ion = an ion that appears on both sides of an equation and is not directly involved in the reaction

Neutralization of Nitric Acid & KOH n The molecular equation would be: n HNO 3 (aq) + KOH (aq)  KNO 3 (aq) + H 2 O (l) n The complete ionic equation would be: n H +1 (aq) + NO 3 -1 (aq) + K +1 (aq) + OH -1 (aq)  K +1 (aq) + NO 3 -1 (aq) + H 2 O (l) n The net ionic equation would therefore be: n H +1 (aq) + OH -1 (aq)  H 2 O (l) n Spectator Ions: NO 3 -1 (aq) + K +1 (aq)

Net Ionic Reactions! - DR AgNO 3 (aq) + KCl(aq)  AgCl(s) + KNO 3 (aq) Ag + (aq) + NO 3 – (aq) + K + (aq) + Cl – (aq)  AgCl(s) + K + (aq) +NO 3 – (aq) CIE: Molecular (or Full) Equation: NIE: Ag + (aq) + Cl – (aq)  AgCl(s) SI: K + (aq) and NO 3 – (aq)

Net Ionic Reactions! - DR Pb(NO 3 ) 2 (aq) +2LiCl(aq)  PbCl 2 (s) + 2LiNO 3 (aq) Pb +2 (aq) +2NO 3 – (aq) +2 Li + (aq) +2Cl – (aq)  PbCl 2 (s)+2Li + (aq) +2NO 3 – (aq) CIE: Molecular (or Full) Equation: NIE: Pb +2 (aq) + 2Cl – (aq)  PbCl 2 (s) SI: Li + (aq) and NO 3 – (aq)