1 Chapter 21 “Chemical Reactions”

2 All chemical reactions… l have two parts: 1.Reactants = the substances you start with 2.Products = the substances you end up with l The reactants will turn into the products. Reactants  Products

3 Reactants Products

4 Chemical Reaction l Here’s a simple reaction of water, food coloring and mineral turpentine. The author uses light so you can see the reaction. om/watch?v=AanR rqZx2Gk&NR=1&fe ature=fvwp

5 Antoine Lavoisier’s Contributions 1770s - he experimented with mercury by placing a measured mass of solid mercury oxide into a sealed container. By adding heat, the red powder transferred into mercury metal, and a gas was produced. He found that no mass was lost: mercury (II) oxide oxygen plus mercury 10 g =.7 g g The total starting mass of all reactants equals the total final mass of all products.

6 The Father of Modern Chemistry Lavoisier realized in Elements of Chemistry that substances needed to be named based on their composition. The International Union of Pure and Applied Chemistry (IUPAC) was formed in This is the commission that coordinates guidelines for naming chemical compounds systematically. This is the commission that provides the name of an element before it receives a permanent name. Lavoisier also pioneered early experiments in biochemistry, medicine and sports science.

7 In a chemical reaction 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 item is a word Copper reacts with chlorine to form copper (II) chloride. #2. In a word equation some symbols used Copper + chlorine  copper (II) chloride

8 Symbols in equations? l the arrow (→) separates the reactants from the products (arrow points to products) –Read as: “reacts to form” or yields l The plus sign = “and” l (s) after the formula = solid: Fe (s) l (g) after the formula = gas: CO 2(g) l (l) after the formula = liquid: H 2 O (l)

9 Symbols used in equations l (aq) after the formula = dissolved in water, an aqueous solution: NaCl (aq) is a salt water solution  used after a product indicates a gas has been produced: H 2 ↑  used after a product indicates a solid has been produced: PbI 2 ↓

10 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)

11 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 are protein catalysts in your body.

12 Coefficients: Unit Managers Atoms are rearranged (never created or destroyed.) Numbers to the left of the formulas represent the number of units of each substance taking part in a reaction. Ex: 2NaOH is two units of sodium hydroxide. Read the analogy of making sandwiches on page 636 of your textbook.

13 Metals in the Atmosphere 1) When oxygen reacting with iron breaks down the metal bond (causing rust.) 2) When oxygen reacts with aluminum, the resulting aluminum oxide adheres to and protects the metal. 3)When copper reacts with oxygen, the corrosion creates a blue-green coating called patina (i.e., the Statue of Liberty.)

14 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.

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

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

17 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 each element on both sides of the equation.

18 Rules for balancing: 1)Assemble the correct formulas for all the reactants and products, using “+” and “→” 2)Count the number of atoms of each type appearing on both sides 3)Balance the elements one at a time by adding coefficients (the numbers in front) where you need more - save balancing the H and O until LAST! (hint: I prefer to save O until the very last) 4)Double-Check to make sure it is balanced.

19 l Never change a subscript to balance an equation (You 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.

20 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

21 Types of Reactions Chemists have identified 5 main categories of chemical reactions: 1) Combustion 2) Synthesis 3) Decomposition 4) Single Displacement 5) Double Displacement

22 Types of Reactions

23 #1-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 A combustion reaction occurs when a substance reacts with oxygen to produce energy in the form of heat and light. l Many combustion reactions can fit into other categories of reactions.

24 Combustion Reaction FHVjH2w

25 Synthesis Reactions l A very easy reaction in which 2 or more substances combine to form another substance: A + B -> AB Example: Water 2H 2 (g) + O 2 (g) -> 2H 2 O (g)

26 #2 - Decomposition Reactions l decompose = fall apart AB -> A + B l one reactant breaks apart into two or more elements or compounds. l NaCl Na + Cl l CaCO 3 CaO + CO 2 l Note that energy (heat, sunlight, electricity, etc.) is usually required

27 Single Displacement Reactions When one element replaces another element in a compound: A + BC -> AC + B Ex: Cu (s) + 2AgNO 3 (aq) -> Cu(NO 3 ) 2 (aq) + 2Ag (s) A B C -> A C + B

28 Single Displacement Reactions Single Displacement can cause problems (ex: cooking spinach in an aluminum pan.)

29 Single Displacement 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 

30 Single Displacement 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 Activity Series of Metals l Higher on the list replaces those lower.

31 The “Activity Series” of Metals Lithium Potassium Calcium Sodium Magnesium Aluminum Zinc Chromium Iron Nickel Lead Hydrogen Bismuth Copper Mercury Silver Platinum Gold 1)Metals can replace other metals, 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

32 Synthesis, Decomposition and Single Replacement WtsFinTM

33 Double Displacement Reactions l Two things replace each other AB + CD -> AD + CB –Reactants must be two ionic compounds, in aqueous solution Ba(NO 3 ) (aq) + K 2 SO 4 (aq) -> BaSO 4 (s) + 2KNO 3 (aq) 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

34 #4 - Double Displacement Reactions l Have certain “driving forces”, or reasons –Will only happen if one of the products: a) doesn’t dissolve in water and forms a solid (a “precipitate”), or b) is a gas that bubbles out, or c) is a molecular compound (which will usually be water).

35 Double Displacement Examples 3FLrPTa4

36 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 

37 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 

38 Oxidation-Reduction Reactions A common characteristic of many chemical reactions is the tendency of the substances to lose or gain electrons. Oxidation – lost of electrons Reduction – gain of electrons These reactions often involve oxygen (highly reactive) as they pull electrons from metallic elements. Oxidation and reduction always works in pairs (called redox.)

39 Chemical Reactions and Energy l All chemical reactions release or absorb energy (thermal, light, sound, electricity.) l Bonding takes energy. Bond formations release energy. Sometimes more energy is needed for one than the other.

40 Exergonic Reactions l In exergonic reactions, the chemical reactions release energy. Less energy is required to break the original bonds than is released when new bonds form. Example: glow sticks Heat is another example of exergonic reactions (ex: heat packs for sore muscles.)

41 Exothermic Reactions You should remember from studying states of matter that, when thermal energy is released, it is called an exothermic reaction. For example, burning wood (even rusting is exothermic, even though it is difficult to detect a temperature change.) It is exothermic reactions, through the burning of fossil fuels, that provide power used in most homes and industries. Unfortunately impurities (sulfur) can create pollution in our atmosphere.

42 Endergonic Reactions l An endergonic reaction occurs when the chemical reaction requires more energy to break bonds then to form new ones. The energy absorbed can be in the form of heat, light or electricity. l Electricity is often used to supply energy to endergonic reactions. The electroplating of metals is an example:

43 Electroplating 5wamQe0&feature=related

44 Endothermic Reaction When energy is needed in the form of thermal energy, the reaction is endothermic. This can be either chemical or physical (adding Epsom salts to water absorbs energy and lowers the temperature of the water.) Cold packs are an example of an endothermic reaction.

45 Endothermic Reaction l RhQM4DHV4c&feature=related&safet y_mode=true&persist_safety_mode= 1&safe=active RhQM4DHV4c&feature=related&safet y_mode=true&persist_safety_mode= 1&safe=active

46 Catalysts You have already learned that a catalyst can be used to speed up a chemical reaction without being permanently changed (the mass of the product is the same, just the reaction is faster.) Catalyst can sometimes be recovered and reused. Example: Polymers used to make plastics and fibers.

47 A Catalyst Halloween l 8OIFWYnto&safety_mode=true&persi st_safety_mode=1&safe=active 8OIFWYnto&safety_mode=true&persi st_safety_mode=1&safe=active

48 Inhibitors While catalysts speed up reactions, inhibitors are used to prevent certain reactions from occurring. Example: food preservatives are added to prevent the chemical breakdown (spoilage) of certain foods. Catalysts and Inhibitors can only increase or decrease the rate of reaction.

49 Rate of Reaction Factors Before adding a catalyst or inhibitor, scientists must take into account other factors that may affect reaction rate: -Concentration of reactants -Pressure -Temperature -Particle Size