1 Chapter 11 “Chemical Reactions”

Describing Chemical Reactions l OBJECTIVES: –Describe how to write a word equation –Describe how to write a skeleton equation –Describe the steps for writing a balanced chemical equation

3 Characteristics of Reactions l All chemical reactions involve a change in matter Reactants = the substances you start with, change into… Products = the substances you end up with Reactants  Products

4 - Page 321 Reactants Products

5 Describing Chemical Reactions l Atoms aren’t created or destroyed (according to the Law of Conservation of Mass) l A reaction can be described several ways: #1. In a sentence every chemical is a word Copper metal reacts with chlorine gas to form copper (II) chloride. #2. In a word equation some symbols used Copper + chlorine  copper (II) chloride

6 Symbols in Equations l (→) separates reactants from products (arrow points to → products) –Read as: … “yields” l The plus sign = “and” l (s) after formula = solid: Fe(s) l (g) after formula = gas: CO 2 (g) l (l) after formula = liquid: H 2 O(l) l (aq) after formula = dissolved in aqueous solution: KCl(aq) = ions in solution (K +, Cl – )

7 Symbols used in equations  used after a product indicates a gas has been produced: H 2 ↑  used after a product indicates a solid (precipitate) has been produced: PbI 2 ↓

8 Symbols used in equations ■ double arrow indicates a reversible reaction (more later) ■ shows that heat is supplied to the reaction ■ is used to indicate a catalyst is supplied (in this case, platinum is the catalyst)

9 What is a catalyst? l A substance that speeds up a reaction, without being changed or used up by the reaction. l Enzymes are biological catalysts in your body (proteins).

10 Describing Chemical Equations #3: The Skeleton Equation l Uses formulas and symbols to describe a reaction –but doesn’t indicate how many; this means they are NOT balanced l All chemical equations are a description of the reaction.

11 Write a skeleton equation for: 1. Solid iron (III) sulfide reacts with gaseous hydrogen chloride to form iron (III) chloride and dihydrogen sulfide gas. 2. Nitric acid dissolved in water reacts with solid sodium carbonate to form liquid water and carbon dioxide gas and sodium nitrate dissolved in water.

12 Now, read these equations: Fe (s) + O 2(g)  Fe 2 O 3(s) Cu (s) + AgNO 3(aq)  Ag (s) + Cu(NO 3 ) 2(aq) NO 2(g) N 2(g) + O 2(g)

13 Describing Chemical Equations #4: Balanced Chemical Equations l Atoms can’t be created or destroyed in an ordinary reaction: –All the atoms we start with we must end up with (meaning: balanced!) l A balanced equation has the same number of atoms of each element on both sides of the equation.

14 Rules for balancing: 1)Assemble the correct formulas for all the reactants and products, using “+” and “→” 2)Count the atoms of each type appearing on both sides 3)Treat polyatomic ions like an “element” if they are unchanged by the reaction 4)Balance the elements one at a time by adding coefficients (the numbers in front) where you need more - save balancing the pure elements until LAST! (things like H 2, O 2, N 2, Cu, Fe, etc.) 4)Double-Check to make sure it is balanced.  Count ‘em up on both sides!!

15 l I said NEVER change a subscript to balance an equation (can only change coefficients) –If you change the subscript (formula) you are describing a different chemical –H 2 O is a different compound than H 2 O 2 l Never put a coefficient in the middle of a formula; they must go only in the front 2 NaCl is okay, but Na 2 Cl is not.

16 Practice Balancing Examples _AgNO 3 + _Cu  _Cu(NO 3 ) 2 + _Ag _Mg + _N 2  _Mg 3 N 2 _P + _O 2  _P 4 O 10 _Na + _H 2 O  _H 2 + _NaOH _CH 4 + _O 2  _CO 2 + _H 2 O

Chemical Equation Practice l Write the balanced equation for the reaction between hydrogen and oxygen that produces water. l Write the balanced equation for the reaction that produces iron(II) sulfide from iron and sulfur. l Write the balanced equation representing the heating of magnesium carbonate to produce solid magnesium oxide and carbon dioxide gas. l Write a balanced equation for the production of HCl gas from its elements. l Write a balanced equation representing the formation of aqueous sulfuric acid from water and sulfur trioxide gas.

Types of Chemical Reactions l OBJECTIVES: –Describe the five general types of reactions –Predict the products of the five general types of reactions

19 Types of Reactions l There are millions of reactions. l We can’t remember them all, but most fall into one of several categories. l We will learn the 5 major types, and given the reactants, predict products for a given type l For some, we will be able to: c) predict whether or not they will happen at all. l How? Recognize them by their reactants

20 #1 - Combination Reactions l Combine = put together l 2 substances combine to make one compound (also called “synthesis”) Ca + O 2  CaO SO 3 + H 2 O  H 2 SO 4 l We can often predict products, especially if the reactants are both elements. Mg + N 2  Mg 3 N 2 (symbols, charges, cross)

21 Complete and balance: Ca + Cl 2  Fe + O 2  (assume iron (II) in product) Al + O 2  l Remember the first step is to write the correct formulas – you can still change subscripts, but not later while balancing! l Then balance by changing just coefficients only

22 #1 – Combination Reactions l Additional Important Notes: a) Some nonmetal oxides react with water to produce an acid: SO 2 + H 2 O  H 2 SO 3 b) Some metallic oxides react with water to produce a base: CaO + H 2 O  Ca(OH) 2 (This contributes to “acid rain”)

23 #2 - Decomposition Reactions l decompose = fall apart l one reactant breaks apart into two or more elements or compounds l NaCl Na + Cl 2 l CaCO 3 CaO + CO 2 l Note that energy (heat, sunlight, electricity, etc.) is usually required

24 #2 - Decomposition Reactions l We can predict the products if it is a binary compound (made up of only two elements) –It breaks apart into the elements: lH2OlH2O l HgO

25 #2 - Decomposition Reactions l If compound has more than two elements you must be given one of the products –other product will be from missing pieces l NiCO 3 CO 2 + ___ H 2 CO 3 (aq)  CO 2 + ___ heat

26 #3 - Single Replacement Reactions l One element replaces another l Reactants must be an element and a compound. l Products will be a different element and a different compound. Na + KCl  K + NaCl F 2 + LiCl  LiF + Cl 2 (Cations switched) (Anions switched)

27 #3 Single Replacement Reactions l Metals will replace other metals (and they can also replace hydrogen) K + AlN  Zn + HCl  l Think of water as: HOH –Metals replace the first H, and then combines with the hydroxide (OH). Na + HOH 

28 #3 Single Replacement Reactions l We can even tell whether or not a single replacement reaction will happen: –Because some chemicals are more “active” than others –More active replaces less active l There is a list on page called the Activity Series of Metals l Higher on the list replaces those lower.

29 The “Activity Series” of Metals Lithium Potassium Calcium Sodium Magnesium Aluminum Zinc Chromium Iron Nickel Lead Hydrogen Bismuth Copper Mercury Silver Platinum Gold 1)Pure metals can replace other metals in compounds, provided they are above the metal they are trying to replace (for example, zinc will replace lead) 2)Metals above hydrogen can replace hydrogen in acids. 3)Metals from sodium upward can replace hydrogen in water. Higher activity Lower activity

30 The “Activity Series” of Halogens Fluorine Chlorine Bromine Iodine Elemental halogens (F 2, Cl 2, I 2, etc.) will replace other halogens in compounds, IF they are above (on the periodic table) the halogen they are replacing. 2NaCl (s) + F 2(g)  2NaF (s) + Cl 2(g) MgCl 2(s) + Br 2(g)  No Reaction! Higher Activity Lower Activity

31 #3 Single Replacement Reactions Practice: Fe + CuSO 4  Pb + KCl  Al + HCl 

32 #4 - Double Replacement Reactions l Two things replace each other. –Reactants must be two ionic compounds, in aqueous solution NaOH + FeCl 3  –The positive ions change place. NaOH + FeCl 3  Fe +3 OH - + Na +1 Cl -1 = NaOH + FeCl 3  Fe(OH) 3 + NaCl

33 #4 - Double Replacement Reactions l O nly happen if one of the products: a) doesn’t dissolve in water and forms a solid (a “precipitate”), or b) If one of the products is a gas, or c) If one of the products is a molecular compound (usually water).

34 Complete and balance: l assume all of the following reactions actually take place: CaCl 2 + NaOH  CuCl 2 + K 2 S  KOH + Fe(NO 3 ) 3  (NH 4 ) 2 SO 4 + BaF 2 

35 How to recognize which type? l Look at the reactants: E + E =Combination C =Decomposition E + C =Single replacement C + C =Double replacement

36 Practice Examples: H 2 + O 2  H 2 O  Zn + H 2 SO 4  HgO  KBr + Cl 2  AgNO 3 + NaCl  Mg(OH) 2 + H 2 SO 3 

37 #5 – Combustion Reactions l Combustion means “add oxygen” l Normally, a compound composed of only C, H, (and maybe O) is reacted with oxygen – usually called “burning” l If the combustion is complete, the products will be CO 2 and H 2 O. l If the combustion is incomplete, the products will be CO (or possibly just C) and H 2 O.

38 Combustion Reaction Examples: C 4 H 10 + O 2  (complete) C 6 H 12 O 6 + O 2  (complete) C 8 H 8 + O 2  (incomplete)

39 SUMMARY: An equation... l Describes a reaction l Must be balanced in order to follow the Law of Conservation of Mass l Can only be balanced by changing the coefficients. l Has special symbols to indicate the physical state, if a catalyst or energy is required, etc.

40 Reactions l Come in 5 major types. l We can tell what type they are by looking at the reactants. l Single Replacement happens based on the Activity Series l Double Replacement happens if one product is: 1) a precipitate (an insoluble solid), 2) water (a molecular compound), or 3) a gas.

41 l I have lots more practice problems for balancing equations –Just ask if you think you need more practice!

Reactions in Aqueous Solution l OBJECTIVES: –Describe the information found in a net ionic equation –Predict the formation of a precipitate in a double replacement reaction

43 Net Ionic Equations l Many reactions occur in water - that is, in aqueous solution l When dissolved in water, many ionic compounds “dissociate”, or separate, into cations (+) and anions (–) l We’ve learned how to write balanced molecular equations; the last type we need to learn how to write is an ionic equation

Net Ionic Equations Example (usually a single or double replacement reaction) Example (usually a single or double replacement reaction) AgNO 3 (aq) + NaCl(aq)  AgCl(s) + NaNO 3 (aq) 1. this is the full balanced equation 2. write an ionic equation by splitting the compounds into their ions (this is really what’s in the solution!) Ag + + NO 3 – + Na + + Cl –  AgCl (s) + Na + + NO 3 – Note: AgCl is NOT soluble, it is a “precipitate” Simplify the equation by crossing out what is unchanged. Ag + + NO 3 – + Na + + Cl –  AgCl (s) + Na + + NO 3 – This leaves the NET IONIC EQUATION Ag + + Cl –  AgCl(s)

45 Predicting the Precipitate l Precipitates are always insoluble salts l Use the solubility rules to predict whether a particular pair of ions will be soluble or not

46 Let’s do some examples of net ionic equations, starting with these reactants: BaCl 2 + AgNO 3 → NaCl + Ba(NO 3 ) 2 → NH 4 +

47 More Practice l Use the solubility rules to determine IF a precipitation reaction occurs. For each that does occur, write a balanced, net ionic equation. AgNO 3 + NaSO 4  ?? NH 4 Cl + Ba(NO 3 ) 2  ?? CaCl 2 + K 2 SO 4  ?? Pb(NO 3 ) 2 + HCl  ?? NH 4 +