# Chemical Reactions. Chemical equation

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1 Handout Kimia Dasar I 2013/2014 Chemical Reactions Chemical equation All chemical reactions take place according to a set of general principles that relate the amounts of materials consumed in a reaction to the amounts of products formed How much material is needed to make a desired amount of product? How efficient a chemical synthesis is? 1

2 Chemical equation A shorthand expression for a chemical change or reaction. Writing chemical equations. reactants products Balancing chemical equations Each kind of atom should contains the same number on each side of the equation. The ratio of the number of molecules is equal to the ratio of the number of moles. Example: H 2 + O 2 H 2 O 2H 2 + O 2 2H 2 O 2 molecules+ 1 molecule 2 molecules 2 moles + 1 mole 2 moles 4.04g g g g reactants g products 2

3 Sample Question 1 & 2 1. K (s) + H 2 O (l ) H 2(g) + KOH (aq) 2. Under appropriate conditions at 1000 C, ammonia gas reacts with oxygen gas to produce gaseous nitrogen monoxide (common name, nitric oxide) and gaseous water. Write the unbalanced and balanced equations for this reaction. Group Discussion Balance each of the following chemical equations. a. FeCl 3(aq) + KOH (aq) Fe(OH) 3(s) + KCl (aq) b. Pb(C 2 H 3 O 2 ) 2(aq) + KI (aq) PbI 2(s) + KC 2 H 3 O 2(aq) c. P 4 O 10(s) + H 2 O (l ) H 3 PO 4 (aq) d. Li 2 O (s) + H 2 O (l ) LiOH (aq) e. MnO 2(s) + C (s) Mn (s) + CO 2(g) f. Sb (s) + Cl 2(g) SbCl 3(s) g. CH 4 (g) + H 2 O (g) CO (g) + H 2(g) h. FeS (s) + HCl (aq) FeCl 2(aq) + H 2 S (g) 3

4 Reactions Why does a chemical reaction occur? The most common of the driving forces are: 1. Formation of a solid 2. Formation of water 3. Transfer of electrons 4. Formation of a gas Types of reactions 4

5 Types of reactions Precipitation reactions The formation of an insoluble product (precipitate). Usually involve ionic compounds. Ions in solution combine to form a solid salt. Example: The precipitation reaction that occurs when yellow potassium chromate, K 2 CrO 4(aq), is mixed with a colorless barium nitrate solution, Ba(NO 3 ) 2(aq). What Happens When an Ionic Compound Dissolves in Water? The designation Ba(NO 3 ) 2(aq) means that barium nitrate (a white solid) has been dissolved in water. Note from its formula that barium nitrate contains the Ba 2+ and NO 3 - ions. In virtually every case when a solid containing ions dissolves in water, the ions separate and move around independently. That is, Ba(NO 3 ) 2(aq) does not contain Ba(NO 3 ) 2 units. Rather, it contains separated Ba 2+ and NO 3 - ions. In the solution there are two NO 3 - ions for every Ba 2+ ion. 5

6 What Happens When an Ionic Compound Dissolves in Water? When each unit of a substance that dissolves in water produces separated ions, the substance is called a strong electrolyte. Barium nitrate is a strong electrolyte in water, because each Ba(NO 3 ) 2 unit produces the separated ions (Ba 2+, NO 3-, NO 3- ). Similarly, aqueous K 2 CrO 4 also behaves as a strong electrolyte. Potassium chromate contains the K + and CrO 4 2- ions, so an aqueous solution of potassium chromate (which is prepared by dissolving solid K 2 CrO 4 in water) contains these separated ions. 6

7 General Rules for Solubility of Ionic Compounds (Salts) in Water at 25 C 1. Most nitrate (NO 3- ) salts are soluble. 2. Most salts of Na +, K +, and NH 4+ are soluble. 3. Most chloride salts are soluble. Notable exceptions are AgCl, PbCl 2, and Hg 2 Cl Most sulfate salts are soluble. Notable exceptions are BaSO 4, PbSO 4, and CaSO Most hydroxide compounds are only slightly soluble.* The important exceptions are NaOH and KOH. Ba(OH) 2 and Ca(OH) 2 are only moderately soluble. 6. Most sulfide (S 2- ), carbonate (CO 3 2- ), and phosphate (PO 4 3- ) salts are only slightly soluble.* *The terms insoluble and slightly soluble really mean the same thing: such a tiny amount dissolves that it is not possible to detect it with the naked eye. 7

8 Sample Question 3-5 Predict what will happen when the following solutions are mixed. Write the balanced equation for any reaction that occurs. 3. KNO 3(aq) and BaCl 2(aq) 4. Na 2 SO 4(aq) and Pb(NO 3 ) 2(aq) 5. KOH (aq) and Fe(NO 3 ) 3(aq) Complete Net Ionic Equation Ba(NO 3 ) 2(aq) + K 2 CrO 4(aq) BaCrO 4(s) + 2KNO 3(aq) This is called the molecular equation for the reaction; it shows the complete formulas of all reactants and products. However, although this equation shows the reactants and products of the reaction, it does not give a very clear picture of what actually occurs in solution. The complete ionic equation, better represents the actual forms of the reactants and products in solution. 8

9 The complete ionic equation reveals that only some of the ions participate in the reaction. K + and NO 3- ions are present in solution both before and after the reaction. Ions such as these, which do not participate directly in a reaction in solution, are called spectator ions. The ions that participate in this reaction are the Ba 2+ and CrO 4 2- ions, which combine to form solid BaCrO4: Acid - base reactions - One of the most fundamental chemical reactions is the combination of a hydroxide ion and a hydronium ion to produce two molecules of water: OH - (aq) + H 3 O + (aq) 2H 2 O (l) - Acid as a proton donor and base as a proton acceptor. - Any reaction in which H + moves from one species to another. - Also called a neutralization reaction. acid + base salt + water 9

10 Strong acids: React with water to produce hydronium ions. Ex: HNO 3, HClO 4, H 2 SO 4, HCl, HBr and HI. Weak acids Substances that can t readily donate protons to water molecule. Ex: H 3 PO 4, CH 3 COOH Strong base Substances that are completely ionized into the metals ions and hydroxide ions. Ex: NaOH, Ba(OH) 2 Weak base Al(OH) 3, NH 3 Sample Question 6 6. Nitric acid is a strong acid. Write the molecular, complete ionic, and net ionic equations for the reaction of aqueous nitric acid and aqueous potassium hydroxide. 10

11 Besides water, which is always a product of the reaction of an acid with OH, the second product is an ionic compound, which might precipitate or remain dissolved, depending on its solubility. This ionic compound is called a salt. Base and acid oxides - Metals-oxides are bases. The oxide anion, O 2-, is a strong base that readily accepts a proton from a water molecule. Na 2 O (s) + H 2 O (l) 2Na + (aq) + 2OH - (aq) CaO (s) + H 2 O (l) Ca(OH) 2(s) 11

12 - Non metal oxides react with the water to produce acids. Non metal oxides whose molecules contain one non metal atom react in a 1:1 ratio with water: SO 2(g) + H 2 O (l) H 2 SO 3(aq) SO 3(g) + H 2 O (l) H 2 SO 4(aq) CO 2(g) + H 2 O (l) H 2 CO 3(aq) - Oxide whose molecules contain more than one non metal atom react to form more than one molecule of acid: N 2 O 5(g) + H 2 O (l) 2HNO 3(aq) Cl 3 O 7(g) + H 2 O (l) 2HClO 4(aq) P 4 O 10(s) + 6H 2 O (l) 4H 3 PO 4(aq) Oxidation reduction reactions - Oxidation reduction reactions occur when electrons from one chemical substance are transferred to another. - Example: 2Na (s) + Cl 2(g) 2NaCl (s) 2Mg (s) + O2 (g) 2MgO (s) 2Al (s) + Fe 2 O 3(s) 2Fe (s) + Al 2 O 3(s) 12

13 Sodium metal is composed of sodium atoms, each of which has a net charge of zero. (The positive charges of the 11 protons in its nucleus are exactly balanced by the negative charges on the 11 electrons.) Similarly, the chlorine molecule consists of 2 uncharged chlorine atoms (each has 17 protons and 17 electrons). However, in the product (sodium chloride), the sodium is present as Na and the chlorine as Cl. By what process do the neutral atoms become ions? After the electron transfer, each sodium has ten electrons and eleven protons (a net charge of 1), and each chlorine has eighteen electrons and seventeen protons (a net charge of 1). 13

14 Metal displacement - The reaction occurs when one metal in solution is displaced by another metal by means of a redox reaction. - Oxidation: Zn (s) Zn 2+ (aq) + 2e - - Reduction: Cu 2+ (aq) + 2e - Cu (s) - Redox: Zn (s) + Cu 2+ (aq) Zn 2+ (aq) + Cu (s) Oxidation of metals by H 3 O + and H 2 O - 2Li (s) + 2H 3 O + 2Li + (aq) + H 2(g) + 2H 2 O (l) - 2Fe (s) + 4H 3 O + 2Fe 2+ (aq) + 2H 2(g) + 4H 2 O (l) - 2Al (s) + 6H 3 O + 2Al 3+ (aq) + 3H 2(g) + 6H 2 O (l) - 2Na 2Na + + 2e - - 2H 2 O + 2e - H 2 + 2OH - 2Na + 2H 2 O H 2 + 2Na + + 2OH - Oxidation by molecular oxygen - Almost all elements combine with molecular oxygen to form binary oxides; the loss of electrons is called oxidation because elements lose electrons when they combine with oxygen. The more easily a metal is oxidized, the more readily it reacts with molecular oxygen. 14

15 The stoichiometry of chemical reactions - The knowledge on the relationship among atoms, moles, and masses including how much they are present in combination with the concept of a balanced chemical equation is required. Sample Question 7 7. How many grams of hydrogen do we need to produce 68 g of ammonia? N 2 + 3H 2 2NH 3 15

16 Sample Question 8 8. Poisonous hydrogen cyanide (HCN) is an important industrial chemical. It is produced from methane, ammonia, and molecular oxygen. The reaction also produces water. An industrial manufacturer wants to convert 175 kg of methane into HCN. How much molecular oxygen will be required for this synthesis? CH 4 + NH 3 + O 2 HCN + H 2 O Yields of chemical reactions - The amount of a product obtained from a reaction is often described in terms of the yield of the reaction. - The quantity of product predicted by stoichiometry the theoretical yield - the amount actually obtained the actual yield Percent yield = (actual yield) / (theoretical yield) (100%) 16

17 Under practical conditions, chemical reactions almost always produce smaller amounts of products than the amounts predicted by stoichiometric analysis. There are three major reasons for this: Many reactions stop before reaching completion. Other reactions do not go to completion because they reach dynamic equilibrium. While reactant molecules continue to form product molecules, product molecules also interact to re-form reactant molecules. 17

18 Competing reactions often consume some of the starting materials. When the product of a reaction is purified and isolated, some of it is inevitably lost during the collection process. Gases may escape while being pumped out of a reactor. Liquids adhere to glass surfaces, making it impossible to transfer every drop of a liquid product. Likewise, it is impossible to scrape every trace of a solid material from a reaction vessel. Sample Question 9 The Haber synthesis of ammonia stops when 13% of the starting materials have formed products. Knowing this, how much ammonia could an industrial producer expect to make from 2.0 metric tons of molecular hydrogen? N 2 + 3H 2 2NH 3 18

19 Sample Question 10 The industrial production of hydrogen cyanide is described in Example. If the yield of this synthesis is 97.5%, how many kilograms of methane should be used to produce 1.5 x 10 5 kg of HCN? The limiting reagent - often chemical reactions are run with an excess of one or more starting materials - One reactant will run out before the others. - The reactant that runs out is called the limiting reagent because it limits how much product can be made. - The other starting materials are said to be in excess. 19

20 N 2(g) + 3H 2(g) 2NH 3(g) Consider the following container of N 2(g) and H 2(g) : Each N 2 requires 3H 2 molecules to form 2NH 3 In this case, the mixture of N 2 and H 2 contained just the number of molecules needed to form NH 3 with nothing left over. That is, the ratio of the number of H 2 molecules to N 2 molecules was 15H 2 : 5N 2 = 3H 2 : 1H 2 This ratio exactly matches the numbers in the balanced equation: N 2(g) + 3H 2(g) 2NH 3(g) This type of mixture is called a stoichiometric mixture. 20

21 In this case H 2 is limiting. That is, the H 2 molecules are used up before all of the N 2 molecules are consumed. In this situation, the amount of hydrogen limits the amount of product (ammonia) that can form hydrogen is the limiting reactant. Some N 2 molecules are left over in this case because the reaction runs out of H 2 molecules first. The reactant that runs out first and thus limits the amounts of products that can form is called the limiting reactant (limiting reagent). 21

22 Sample Question 11 For the ammonia synthesis, if we start with 84.0 g of molecular nitrogen and 24.2 g of molecular hydrogen, what mass of ammonia can be prepared? N 2(g) + 3H 2(g) 2NH 3(g) Sample Question 12 Nitric acid is used in the production of fertilizers and explosives. One step in the industrial production of nitric acid is the reaction of ammonia with molecular oxygen to form nitrogen oxide: 4 NH O 2 4 NO + 6 H 2 O In a study of this reaction, a chemist mixed 125 g of ammonia with 256 g of oxygen and allowed them to react to completion. What masses of NO and were produced, and what mass of which reactant was left over? 22

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