Chemical Reactions and Quantities Chapter 7 Chemical Reactions occur Everywhere when fuel burns with oxygen in our cars to make the car move when we cook our food when we dye our hair in our bodies, chemical reactions convert food into molecules that build and move muscle in leaves of trees and plants, carbon dioxide and water are converted to carbohydrates
Some chemical reactions are simple, whereas others are quite complex However, they can all be written by chemical equations that chemists use to describe chemical reactions. In every chemical reaction, atoms are rearranged to give new substances. Just like following a recipe, certain ingredients are combined (and often heated) to form something new. Chapter Seven 7.1 Equations for Chemical Reactions 7.2 Types of Reactions 7.3 Oxidation-Reduction Reactions 7.4 The Mole 7.5 Molar Mass and Calculations 7.6 Mole Relationships in Chemical Equations 7.7 Mass Calculations for Reactions 7.8 Limiting Reactants and Percent Yield 7.9 Energy in Chemical Reactions
7.1 Equations for Chemical Reactions Write a balanced equation from formulas of the reactants and products for a reaction; determine the number of atoms in the reactants and products. Chemical Change A chemical change occurs when a substance is converted into one or more new substances that have different formulas and properties. For example, when silver tarnishes, the shiny, silver metal (Ag) reacts with sulfur (S) to become the dull, black substance we call tarnish (Ag 2 S).
Chemical Reaction A chemical reaction always involves a chemical change because atoms of the reacting substances form new combinations with new properties. For example, a chemical reaction (and chemical change) takes place when a piece of iron (Fe) combines with oxygen (O 2 ) in the air to produce a new substance, rust (Fe 2 O 3 ), which has a reddishbrown color. Evidence of a Chemical Reaction During a chemical change, new properties are often visible, which indicates that a chemical reaction took place.
Chemical Equation When you build a model airplane or prepare a new recipe, you follow a set of direction. These directions tell you what materials to use and the products you will obtain. In chemistry, a chemical reaction tells us the materials we need and the products that will form. Writing a chemical equation Suppose you work in a bicycle shop assembling wheels and frames into bicycles. You could represent this by a simple equation:
Writing a chemical equation Return to the silver example: When silver tarnishes, the shiny, silver metal (Ag) reacts with sulfur (S) to become the dull, black substance we call tarnish (Ag 2 S). * Unbalanced equation Writing a chemical equation Return to the iron example: A chemical reaction takes place when a piece of iron (Fe) combines with oxygen (O 2 ) in the air to produce a new substance, rust (Fe 2 O 3 ), which has a reddish-brown color. Generally, each formula is followed by an abbreviation, in parentheses, that gives the physical state of the substance. solid (s) liquid (l) gas (g) aqueous (aq) dissolved in water * Unbalanced equation
Writing a chemical equation Some reactions require heat to be added in order for the change to occur. For example, when you burn charcoal in a grill, the carbon (C) in the charcoal combines with the oxygen (O 2 ) to form carbon dioxide. * Unbalanced equation Chemical equation symbols Cu (s) + S (s) CuS (s) CaCO 3(s) CaO (s) + CO 2(g) Na 2 SO 4(aq) + BaCl 2(aq) BaSO 4(s) + 2NaCl (aq)
Identifying a balanced chemical equation When a chemical reaction takes place, the bonds between the atoms of the reactants are broken and new bonds are formed to give the products. Identifying a balanced chemical equation All atoms are conserved which means that atoms cannot be gained, lost, or changed into another type of element during the chemical reactions Every chemical reaction must be written as a balanced reaction, which shows the same number of atoms for each element for the reactants and in the products.
Balanced chemical equation Consider the following in which hydrogen (H 2 ) reacts with oxygen (O 2 ) to form water: H 2(g) + O 2(g) H 2 O (g) unbalanced there are the same number of hydrogens on both sides (2) but different numbers of oxygen atoms (2 and 1) so it is unbalanced. Balanced chemical equation H 2(g) + O 2(g) H 2 O (g) unbalanced We use whole numbers called coefficients in front of formulas. Coefficients indicate how many of each molecule are in an equation. In the balanced equation, there are two H 2 molecules and two H 2 O molecules: This illustrates the Law of Conservation of Matter which states that matter cannot be created or destroyed during a chemical reaction.
Practice counting atoms Indicate the number of each type of atom in the following balanced chemical equation: Fe 2 S 3(s) + 6HCl (aq) 2FeCl 3(aq) + 3H 2 S (g) Fe S H Cl reactants products Practice counting atoms Indicate the number of each type of atom in the following balanced chemical equation: 2C 2 H 6(g) + 7O 2(g) 4CO 2(g) + 6H 2 O (g) C H O reactants products
Balancing chemical equations When silver tarnishes, the shiny, silver metal (Ag) reacts with sulfur (S) to become the dull, black substance we call tarnish (Ag 2 S). What is the balanced chemical equations describing this reaction? Balancing chemical equations A chemical reaction takes place when a piece of iron (Fe) combines with oxygen (O 2 ) in the air to produce a new substance, rust (Fe 2 O 3 ), which has a reddish-brown color. What is the balanced chemical equation describing this reaction?
Balancing chemical equations Balance the following chemical reaction: Na 3 PO 4(aq) + MgCl 2(aq) Mg 3 (PO 4 ) 2(s) + NaCl (aq) Chapter Seven 7.1 Equations for Chemical Reactions 7.2 Types of Reactions 7.3 Oxidation-Reduction Reactions 7.4 The Mole 7.5 Molar Mass and Calculations 7.6 Mole Relationships in Chemical Equations 7.7 Mass Calculations for Reactions 7.8 Limiting Reactants and Percent Yield 7.9 Energy in Chemical Reactions
7.2 Types of Reactions Identify a reaction as a combination, decomposition, single replacement, double replacement, or combustion. A great number of reactions occur in nature, in biological systems, and in the laboratory. However, there are some general patterns among all reactions that help us classify reactions. * note some reactions may fit in more than one category.
Combination Reactions In a combination reaction, two or more elements or compounds bond to form one product. For example, sulfur and oxygen combine to form the product sulfur dioxide. Combination Reactions 2Mg (s) + O 2(g) 2MgO (s) N 2(g) + 3H 2(g) 2NH 3(g) Cu (s) + S (s) CuS (s) MgO (s) + CO 2(g) MgCO 3(s)
Decomposition Reactions In a decomposition reaction, a reactant splits into two or more simpler products. For example, when mercury (II) oxide is heated, the compound breaks apart into mercury atoms and oxygen. 2HgO (s) 2Hg (l) + O 2(g) Decomposition Reactions CaCO 3(s) CaO (s) + CO 2(g) Fe 2 S 3(s) 2Fe (s) + 3S (s)
Replacement Reactions In a replacement reaction, elements in a compound are replaced by other elements. In a single replacement reaction, one element switches places with another element in the reactants. *A and B switch places. Replacement Reactions Zn (s) + 2HCl (aq) H 2(g) + ZnCl 2(aq) Cl 2(g) + 2KBr (s) 2KCl (s) + Br 2(l)
Replacement Reactions In a double replacement reaction, the positive ions in the reacting compounds switch places. Replacement Reactions BaCl 2(aq) + Na 2 SO 4(aq) NaOH (aq) + HCl (aq)
Combustion Reactions The burning of a candle and the burning of fuel in the engine of a car are examples of combustion reactions. In a combustion reaction, a carbon-containing compound burns in oxygen (O 2 ) to produce carbon dioxide (CO 2 ) and water (H 2 O) and energy in the form of heat or flame. Fuel + O 2 CO 2 + H 2 O *unbalanced Fuel often: methane (CH 4 ), propane (C 3 H 8 ), or similar. Sometimes also has oxygen atoms in the formula: C 3 H 7 O Combustion Reactions Fuel + O 2 CO 2 + H 2 O unbalanced Methane gas (CH 4 ) aka natural gas undergoes combustion when used to cook food on a gas stove: CH 4(g) + 2O 2(g) CO 2(g) + 2H 2 O (g) + energy Combustion of propane (C 3 H 8 ): C 3 H 8(g) + 5O 2(g) 3CO 2(g) + 4H 2 O (g) + energy
Summary of Reaction Types Practice Classify as combination, decomposition, single replacement, double replacement, or combustion reaction. 2Fe 2 O 3(s) + 3C (s) 3CO 2(g) + 4Fe (s)
Practice Classify as combination, decomposition, single replacement, double replacement, or combustion reaction. 2KClO 3(s) 2KCl (s) + 3O 2(g) Practice Classify as combination, decomposition, single replacement, double replacement, or combustion reaction. C 2 H 4(g) + 3O 2(g) 2CO 2(g) + 2H 2 O (g) + energy
Chapter Seven 7.1 Equations for Chemical Reactions 7.2 Types of Reactions 7.3 Oxidation-Reduction Reactions 7.4 The Mole 7.5 Molar Mass and Calculations 7.6 Mole Relationships in Chemical Equations 7.7 Mass Calculations for Reactions 7.8 Limiting Reactants and Percent Yield 7.9 Energy in Chemical Reactions 7.3 Oxidation-Reduction Reactions Define the terms oxidation and reduction, identify the reactants as oxidized and reduced.
Another type of reaction is the: oxidation-reduction reaction This is a continuation of the previous section Types of Reactions. The oxidation-reduction reaction is arguably one of the most important and undeniably the most complicated of the types we will discuss in this chapter. Because of this, oxidation-reduction reactions has it s own section in this chapter. Perhaps you have never heard of an oxidation and reduction reaction. However, this type of reaction has many important applications in your everyday life. Rusty nail Tarnish on a silver spoon Corrosion on metal When you turn on the lights in your car, an oxidation-reduction reaction within the car battery provides the electricity. When you light a campfire, as the wood burns, oxygen combines with carbon and hydrogen to produce carbon dioxide, water, and heat. - In section 7.2 we called this a combustion reaction. It is. It s also a oxidation-reduction reaction. When you eat starchy food, the starches break down to give glucose which is oxidized in our cells to give you energy.
Oxidation-Reduction Reactions In an oxidation-reduction reaction (redox), electrons are transferred from one substance to another. Oxidation-Reduction Reactions If one substance loses an electron, another substance must gain the electron. The substance that lost electrons was oxidized and we define oxidation as the loss of electrons The substance that gained electrons was reduced and we define reduction as the gain of electrons
Oxidation-Reduction Reactions The substance that lost electrons was oxidized and we define oxidation as the loss of electrons The substance that gained electrons was reduced and we define reduction as the gain of electrons Redox Example 1 The green color that appears on copper surfaces is due partly to: 2Cu (s) + O 2(g) 2CuO (s) This is a redox reaction, and we can break the equation up to show this.
Redox Example 1 OIL RIG Break the equation into oxidation and reduction. 2Cu (s) + O 2(g) 2CuO (s) Oxidation Going from charge 0 to +2, Cu lost two electrons and Cu was oxidized. Redox Example 1 OIL RIG Break the equation into oxidation and reduction. 2Cu (s) + O 2(g) 2CuO (s) Reduction Going from charge 0 to -2, O 2 gained four electrons and O 2 was reduced. (Each O went from 0 to -2 and gained 2 electrons. Because there are two oxygen atoms in O 2, that s a total of four electrons.)
Redox Example 1 To summarize: 2Cu (s) 2Cu 2+ (s) + 4e - oxidation O 2(g) + 4e - 2O 2- (s) reduction In a redox reaction, there s always two pieces that occur. One substance is oxidized (loses electrons) and another substance is reduced (gains the lost electrons). Redox Example 1 Original equation: 2Cu (s) + O 2(g) 2CuO (s) The two half reactions can be added back together to give the original equation. 2Cu (s) 2Cu 2+ (s) + 4e - oxidation O 2(g) + 4e - 2O 2- (s) reduction
Redox Example 2 OIL RIG Zn (s) + CuSO 4(aq) ZnSO 4(aq) + Cu (s) Oxidation half-reaction: Reduction half-reaction: was oxidized was reduced Redox Example 3 OIL RIG Zn (s) + Cl 2(g) ZnCl 2(s) Oxidation half-reaction: Reduction half-reaction: was oxidized was reduced
Chapter Seven 7.1 Equations for Chemical Reactions 7.2 Types of Reactions 7.3 Oxidation-Reduction Reactions 7.4 The Mole 7.5 Molar Mass and Calculations 7.6 Mole Relationships in Chemical Equations 7.7 Mass Calculations for Reactions 7.8 Limiting Reactants and Percent Yield 7.9 Energy in Chemical Reactions 7.4 The Mole Use Avogadro s number to determine the number of particles in a given number of moles.
At the grocery store, you buy eggs by the dozen, or soda by the case. At Staples, pencils are ordered by the gross and paper by the ream. The terms dozen, case, gross, and ream are used to count the number of items present. Chemists do the same thing, and they call it the mole. The mole In chemistry, particles (such as atoms, molecules, and ions) are counted by the mole which contains 6.02 x 10 23 items.
Avogadro s number 6.02 x 10 23 is known as Avogadro s number. 602 000 000 000 000 000 000 000 = 6.02 x 10 23 It is a very large number and we use it because atoms are so small that it takes an extremely large number of atoms to provide a usable amount to use in chemical reactions in a laboratory. The mole and Avogadro s number 1 mole of any element always contains Avogadro s number of atoms. 1 mole of carbon atoms contains 6.02 x 10 23 carbon atoms. 1 mole of aluminum atoms contains 6.02 x 10 23 aluminum atoms. 1 mole of any element = 6.02 x 10 23 atoms of that element. 1 mole of any compound = 6.02 x 10 23 of that compound 1 mole of CO 2 contains 6.02 x 10 23 molecules of CO 2. 1 mole of NaCl contains 6.02 x 10 23 molecules of NaCl.
Converting mole particles We use Avogadro s number as a conversion factor to convert between the moles of a substance and the number of particles it contains... The particles can be atoms, molecules, ions, or anything. Replace particles with items if that helps.
Practice Convert 4.00 moles of sulfur to atoms of sulfur. Practice Convert 3.01 x 10 24 molecules of CO 2 to moles of CO 2
moles in chemical formulas We have seen in past chapters that the subscripts in the chemical formula of a compound indicate the number of atoms of each type of element in the compound. For example, aspirin, C 9 H 8 O 4 moles in chemical formulas The subscripts also tell us the number of moles of each element in 1 mole of aspirin.
moles in chemical formulas And because 1 mole aspirin = 6.02 x 10 23 molecules of aspirin moles in chemical formulas We can also use subscripts from the formula, C 9 H 8 O 4, we can write the conversion factors for each of the elements in 1 mole of aspirin.
Practice Propyl acetate, C 5 H 10 O 2, gives the odor and taste of pears. How many moles of C are present in 1.50 moles of propyl acetate? Chapter Seven 7.1 Equations for Chemical Reactions 7.2 Types of Reactions 7.3 Oxidation-Reduction Reactions 7.4 The Mole 7.5 Molar Mass and Calculations 7.6 Mole Relationships in Chemical Equations 7.7 Mass Calculations for Reactions 7.8 Limiting Reactants and Percent Yield 7.9 Energy in Chemical Reactions
7.5 Molar Mass and Calculations Calculate the molar mass for a substance given its chemical formula; use molar mass to convert between grams and moles. A single atom or molecule is much too small to weigh, even using the best scale. In fact, it takes a huge number of atoms or molecules to make enough of a substance for you to see. An amount of water that contains Avogadro s number of water molecules (6.02 x 10 23 ) is only a few sips! So in a lab, we use a scale to weigh out substances in moles.
Molar Mass For any element, the quantity called molar mass is the quantity in grams that equals the atomic mass of that element. Molar Mass We are counting 6.02 x 10 23 atoms of an element when we weigh out the number of grams equal to the molar mass. For example, carbon has an atomic mass of 12.01 on the periodic table. This means 1 mole of carbon atoms has the mass of 12.01g. Then to obtain 1 mole of carbon atoms, we would weigh out 12.01g of carbon. Atomic mass of Carbon: 12.01 amu Molar mass of Carbon: 12.01 g/mol
Using Molar Mass The molar mass of an element is useful to convert moles of an element to grams (or grams to moles). For example, 1 mole of sulfur as a mass of 32.066g 1 mole S = 32.066g of S Molar Mass of a Compound We can calculate the molar mass of a compound by adding up the molar mass of each atom. Example: Calculate the molar mass of lithium carbonate (Li 2 CO 3 )
Practice Calculate the molar mass for salicyclic acid, C 7 H 6 O 3. Examples of 1 mole
Calculations using Molar Mass We can now change from moles to grams (or grams to moles) using the molar mass as a conversion factor. Example: Convert 0.750 moles of silver to grams of silver. Calculations using Molar Mass We can change moles to grams (or vice versa) for a compound using the compound s molar mass. Example: A box of salt contains 737 g of NaCl. How many moles of NaCl are present?
Summary Chapter Seven 7.1 Equations for Chemical Reactions 7.2 Types of Reactions 7.3 Oxidation-Reduction Reactions 7.4 The Mole 7.5 Molar Mass and Calculations 7.6 Mole Relationships in Chemical Equations 7.7 Mass Calculations for Reactions 7.8 Limiting Reactants and Percent Yield 7.9 Energy in Chemical Reactions
7.6 Mole Relationships in Chemical Equations Give a quantity in moles of reactant or product, use a mole-mole factor from the balanced chemical equation to calculate the number of moles of another substance in the reaction. Law of Conservation of Mass In any chemical reaction, the total amount of matter in the reactants is equal to the total amount of matter in the products. - an application of the Law of Conservation of Mass Thus no material is lost or gained as original substances are changed into new substances.
For example, tarnish (Ag 2 S), forms when silver reacts with sulfur to form silver sulfide: 2Ag (s) + S (s) Ag 2 S (s) For every 2 silver atoms, 1 sulfur atom is required Or For every 2 moles of silver atoms, 1 mole of sulfur atoms is required. Because the molar mass can be determined, the moles of Ag, S, and Ag 2 S can also be stated in terms of mass (grams) of each Thus 215.8 g of Ag and 32.1 g of S reacts to form 247.9 g of Ag 2 S. 215.8g + 32.1g = 249.9g The total mass of reactants (249.9 g) equals the total mass of the products (249.9 g).
Info in a balanced equation The various ways in which a chemical equation can be interpreted: Mole-mole factors When iron reacts with sulfur, the product is iron(iii) sulfide: 2Fe (s) + 3S (s) Fe 2 S 3(s) From the balanced equation, we see that 2 moles of iron reacts with 3 moles of sulfur to form mole of iron(ii) sulfide. (Actually, any amount of iron or sulfur may be used, but the ratio of iron reacting with sulfur will always be the same.)
Mole-mole factors From the coefficients, we can write mole-mole factors between any two compounds in an equation. 2Fe (s) + 3S (s) Fe 2 S 3(s) Fe and S: Fe and Fe 2 S: S and Fe 2 S: Mole-mole factors in calculations Whenever you prepare a recipe, you need to know the proper amounts of ingredients and how much the recipe will make. The same is true for chemistry. Now that we have written all the possible mole-mole factors for 2Fe (s) + 3S (s) Fe 2 S 3(s) we will use those mole-mole factors in a chemical calculation.
Practice #1 How many moles of sulfur are needed to react with 1.42 moles of iron? 2Fe (s) + 3S (s) Fe 2 S 3(s) 2molesFe 3molesS 2molesFe 1moleFe 2 S 3 3molesS 1moleFe 2 S 3 Practice #2 How many moles of iron are needed to react with 2.75 moles of sulfur? 2Fe (s) + 3S (s) Fe 2 S 3(s) 2molesFe 3molesS 2molesFe 1moleFe 2 S 3 3molesS 1moleFe 2 S 3
Chapter Seven 7.1 Equations for Chemical Reactions 7.2 Types of Reactions 7.3 Oxidation-Reduction Reactions 7.4 The Mole 7.5 Molar Mass and Calculations 7.6 Mole Relationships in Chemical Equations 7.7 Mass Calculations for Reactions 7.8 Limiting Reactants and Percent Yield 7.9 Energy in Chemical Reactions 7.7 Mass Calculations for Reactions Given the mass in grams of a substance in a reaction, calculate the mass in grams of another substance in the reaction.
When we have the balanced chemical equation for a reaction, we can use the mass of one of the substances (A) to calculate the mass of a second (B). However, the calculations require us to convert the mass of A to the moles of A (using A s molar mass). Then we use the molemole factor to convert moles A to moles B. Then finally use B s molar mass to convert moles B to grams B. molar mass A mole-mole factor molar mass B grams A moles A mole B grams B Practice How many grams of CO 2 are produced when 54.6g of C 2 H 2 is burned? 2C 2 H 2(g) + 5O 2(g) 4CO 2(g) + 2H 2 O (g)
Practice Calculate grams CO 2 produced when 25 g of O 2 reacts. 2C 2 H 2(g) + 5O 2(g) 4CO 2(g) + 2H 2 O (g) Chapter Seven 7.1 Equations for Chemical Reactions 7.2 Types of Reactions 7.3 Oxidation-Reduction Reactions 7.4 The Mole 7.5 Molar Mass and Calculations 7.6 Mole Relationships in Chemical Equations 7.7 Mass Calculations for Reactions 7.8 Limiting Reactants and Percent Yield 7.9 Energy in Chemical Reactions
7.8 Limiting Reactants and Percent Yield Identify the limiting reactant and calculate the amount of product formed from the limiting reactant. Given the actual quantity of product, determine the percent yield for a reaction. Limiting Reactant When you make peanut butter sandwiches for lunch, you need 2 slices of bread and 1 tablespoon of peanut butter for each sandwich. 2 slices bread + 1 Tbl. peanut butter 1 peanut butter sandwich If you have 8 slices of bread and a full jar of peanut butter, you will run out of bread after you make 4 peanut butter sandwiches once the bread is used up, even though there is a lot of peanut butter left in the jar. The number of slices of bread has limited the number of sandwiches you can make.
Limiting Reactant On a different day, you might have 8 slices of bread but only a tablespoon of peanut butter left in the peanut butter jar. You will run out of peanut butter after you make just 1 sandwich and have 6 slices of bread left over. The smaller amount of peanut butter available has the limited the number of sandwiches you can make. Limiting Reactant The reactant that is completely used up is the limiting reactant. The reactant that does not completely react and is left over is called the excess reactant.
Moles from Limiting Reactant In the same way, the reactants in a chemical reactions do not always combine in quantities that allow each to be used up at exactly the same time. In many reactions, there is a limiting reactant that determines the amount of product that can be formed. Identifying Limiting Reactant When we know the quantities of the reactants of a chemical reaction we calculate the amount of product that is possible from each reactant if it were completely consumed. The limiting reactant is the one that runs out first and produces the smaller amount of product.
Practice #1 If 3.00 moles of CO and 5.00 moles of H 2 are the initial reactants, what is the limiting reactant? And how many moles of CH 3 OH can be produced. CO (g) + 2H 2(g) CH 3 OH (g) Practice #2 If 4.00 moles of CO and 4.00 moles of H 2 are the initial reactants, what is the limiting reactant? And how many moles of CH 3 OH can be produced. CO (g) + 2H 2(g) CH 3 OH (g)
Mass from Limiting Reactant The quantities of the reactants can also be given in grams. We combine knowledge from sections 7 and 8 to find the limiting reactant. The calculation to find the limiting reactant is the same as the last two Practice slides. But before we can do it, we must convert grams to moles. - Subscripts in chemical equations are not grams!! -You need to use moles to compare subscripts!! Practice How many grams of CO are formed from 70.0 g of SiO 2 and 50.0 g of C? SiO 2(s) +3C (s) SiC (s) +2CO (g)
Percent Yield In our reactions so far we have assumed 100% efficiency, meaning all of the reactants changed completely to products. While this is ideal, it s hardly reality. Often, reactant is lost when moving from container to container, chemicals aren t 100% pure, and unwanted side reactions use up reactant making unwanted products. Thus 100% of the desired product is not actually obtained. When we do a chemical reaction in the lab, we measure out specific quantities of the reactants. We then calculate the theoretical yield for the reaction. Which is the amount of product (100%) we would expect if all the reactants were converted to the desired products. When the reaction ends, we collect and measure the mass of the product that was actually made. This is called the actual yield. Because the reaction is hardly ever 100% efficient the actual yield is less than the theoretical yield. Percent Yield We can represent the ratio of actual yield to theoretical yield in one number called percent yield. %yield= actual yield theoretical yield x 100
Practice What is the % yield of LiHCO 3 for the reaction if 50.0 g of LiOH gives 72.8 g LiHCO 3. LiOH (s) + CO 2(g) LiHCO 3(s) Chapter Seven 7.1 Equations for Chemical Reactions 7.2 Types of Reactions 7.3 Oxidation-Reduction Reactions 7.4 The Mole 7.5 Molar Mass and Calculations 7.6 Mole Relationships in Chemical Equations 7.7 Mass Calculations for Reactions 7.8 Limiting Reactants and Percent Yield 7.9 Energy in Chemical Reactions
7.9 Energy in Chemical Reactions Given the heat of reaction, calculate the loss or gain of heat for an exothermic or endothermic reaction. Energy in Chemical Reactions Almost every chemical reaction involves a loss or gain of energy. To discuss energy change for a reaction, we look at the energy of the reactants before the reaction and the energy of the products after the reaction. The SI unit for energy is the joule (J) 1 kilojoule (kj) = 1000 joules (J)
Heat of Reaction Energy is often present as heat. The heat of reaction is the amount of heat absorbed or released during a reaction that takes place at constant pressure. A change of energy occurs as reactants interact, bonds break apart, and products form. Another name for heat of reaction is enthalpy change. ΔH = H products -H reactants Exothermic Reactions In an exothermic reaction the energy of the products is lower than that of the reactants. - Heat is released along with the products that form. For example when 1 mole of H 2 and 1 mole of Cl 2 react, 2 mole HCl form. The energy of the HCl is 185kJ less than the H 2 and Cl 2 Because products are lower energy, that energy difference is released as heat. We write ΔH with a negative sign to indicate exothermic.
Exo example: Hot Packs Hot packs are used in medicine, camping gear, emergency kits, etc. Inside the hot pack are two pouches. One pouch has water, the other has CaCl 2 When the boundary is broken between pouches an exothermic reaction takes place and releases heat. CaCl 2(s) CaCl2(aq) Temperatures can increase as much as 66C Endothermic Reactions In an endothermic reaction, the energy of the products is higher than that of the reactants. - Heat is required to begin the reaction and convert reactants to products. For example: 1 mole N 2 and 1 mole O 2 react to form 2 moles NO only when at least 180 kj of heat is added. We write ΔH with a positive sign to indicate endothermic.
Endo example: Cold Packs Cold packs work the same way with two pouches. One pouch as water, the other has NH 4 NO 3. The temperature can drop to 4-5C. NH 4 NO 3(s) + 26kJ NH 4 NO 3(aq) Calculations of Heat in Reactions The value of ΔH refers to the heat change for each substance in the balanced equation. 2H 2 O (l) 2H 2(g) + O 2(g) Δ H = +572 kj means that for this reaction, +572 kj are absorbed by 2 moles of H 2 O to produce 2 moles of H 2 an 1 mole of O 2. We can use ΔH to write conversion factors just like the mole-mole factors from section 7.6.
Calculations of Heat in Reactions Suppose that 9.00 g H 2 O undergoes the reaction. Calculate the heat absorbed. 2H 2 O (l) 2H 2(g) + O 2(g) Δ H = +572 kj Practice How much heat, in kj, is released when nitrogen and hydrogen react to form 50.0g of ammonia (NH 3 )? N 2(g) + 3H 2(g) 2NH 3(g) Δ H = -92.2 kj
Chapter Seven 7.1 Equations for Chemical Reactions 7.2 Types of Reactions 7.3 Oxidation-Reduction Reactions 7.4 The Mole 7.5 Molar Mass and Calculations 7.6 Mole Relationships in Chemical Equations 7.7 Mass Calculations for Reactions 7.8 Limiting Reactants and Percent Yield 7.9 Energy in Chemical Reactions