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Chemical Reactions and Quantities Chapter 7

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Chemical Reactions and Quantities Chapter 7 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 (Ag2S). 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 (O2) in the air to produce a new substance, rust (Fe2O3), which has a reddish- brown 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 (Ag2S). * 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 (O2) in the air to produce a new substance, rust (Fe2O3), 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 (O2) to form carbon dioxide. * Unbalanced equation Chemical equation symbols Cu(s) + S(s) → CuS(s) CaCO3(s) → CaO(s) + CO2(g) Na2SO4(aq) + BaCl2(aq) → BaSO4(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 (H2) reacts with oxygen (O2) to form water: H2(g) + O2(g) Æ H2O(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 H2(g) + O2(g) Æ H2O(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 H2 molecules and two H2O 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: Fe2S3(s) + 6HCl(aq) Æ 2FeCl3(aq) + 3H2S(g) reactants products Fe S H Cl Practice counting atoms Indicate the number of each type of atom in the following balanced chemical equation: 2C2H6(g) + 7O2(g) → 4CO2(g) + 6H2O(g) reactants products C H O Balancing chemical equations When silver tarnishes, the shiny, silver metal (Ag) reacts with sulfur (S) to become the dull, black substance we call tarnish (Ag2S). 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 (O2) in the air to produce a new substance, rust (Fe2O3), which has a reddish-brown color. What is the balanced chemical equation describing this reaction? Balancing chemical equations Balance the following chemical reaction: Na3PO4(aq) + MgCl2(aq) Æ Mg3(PO4)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) + O2(g) Æ 2MgO(s) N2(g) + 3H2(g) Æ 2NH3(g) Cu(s) + S(s) Æ CuS(s) MgO(s) + CO2(g) Æ MgCO3(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) + O2(g) Decomposition Reactions CaCO3(s) → CaO(s) + CO2(g) Fe2S3(s) Æ 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) Æ H2(g) + ZnCl2(aq) Cl2(g) + 2KBr(s) Æ 2KCl(s) + Br2(l) Replacement Reactions In a double replacement reaction, the positive ions in the reacting compounds switch places. Replacement Reactions BaCl2(aq) + Na2SO4(aq) Æ NaOH(aq) + HCl(aq) Æ 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 (O2) to produce carbon dioxide (CO2) and water (H2O) and energy in the form of heat or flame. Fuel + O2 Æ CO2 + H2O *unbalanced Fuel often: methane (CH4), propane (C3H8), or similar. Sometimes also has oxygen atoms in the formula: C3H7O Combustion Reactions Fuel + O2 Æ CO2 + H2O unbalanced Methane gas (CH4) aka natural gas undergoes combustion when used to cook food on a gas stove: CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) + energy Combustion of propane (C3H8): C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(g) + energy Summary of Reaction Types Practice Classify as combination, decomposition, single replacement, double replacement, or combustion reaction. 2Fe2O3(s) + 3C(s) Æ 3CO2(g) + 4Fe(s) Practice Classify as combination, decomposition, single replacement, double replacement, or combustion reaction. 2KClO3(s) → 2KCl(s) + 3O2(g) Practice Classify as combination, decomposition, single replacement, double replacement, or combustion reaction. C2H4(g) + 3O2(g) → 2CO2(g) + 2H2O(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.
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