Chemical Reactions and Equations Chapter 9

Chemical Reaction A process in which one or more substances are converted into new substances with different chemical and physical properties. Not reversible.

Chemical Reactions Examples: Breathing Burning gasoline Baking All chemical reactions involve two types of substances: Reactant – a substance that enters into a chemical reaction Product – a substance that is produced by a chemical reaction

Reasons for Reactions The arrangement of electrons in an atom determines whether it will bond with other atoms and with which atoms it will bond. An atom with a full set of valence electrons will not form bonds. An atom with an incomplete set of valence electrons will bond.

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Chemical Equations 9-2

Chemical Equations Chemical reactions are represented by sentences known as chemical equations. A chemical equation identifies the reactants and products in a chemical reaction.

Chemical Equations A chemical reaction is the process by which one or more substances are changed into one or more different substances It can be represented by an equation. It shows what changes have taken place It also shows the relative amounts of the various elements and compounds that take part in these reactions

Chemical Equations The starting substances in a chemical reaction are the reactants The substances that are formed by the chemical reaction are the products Reactant(g) + Reactants(g)  Products(g) + Products(g)

Chemical Equations The letters in parentheses indicate the physical state of each substance involved. (g) means gas (l) means liquid (s) means solid (aq) means that it is dissolved in water

Word Equations Gives names of the reactants and names of products. Example: calcium + oxygen  calcium oxide + means “reacts with”  means “yields”, indicates direction of reaction.

Formula Equations Writing a chemical equation: First determine the symbols and formulas that describe the reactants and products. Substitute into the word equation. Example: calcium + oxygen  calcium oxide Ca + O2  CaO

Practice Problems Silver (I) nitrate reacts with copper to form copper (II) nitrate and silver. Hydrogen peroxide (H2O2) decomposes to form water and oxygen gas.

Balancing Chemical Equations Law of conservation of mass: Matter is neither created nor destroyed. Total mass of the reactants must be equal to the total mass of the products. For mass to remain constant before and after a chemical reaction, the number of atoms of each element must be the same before and after a chemical reaction.

Balancing Chemical Equations The same number of atoms of each element must appear on both sides of the arrow in a chemical equation. Example: If reactants have 4 hydrogen and 1 carbon, products must also have 4 hydrogen and 1 carbon.

Balancing Chemical Equations Rules for writing balanced equations 1. Determine the reactants and the products 2. Assemble the parts of the chemical equation H2(g) + O2(g)  H2O(l) Reactants Yields Products

Balancing Chemical Equations 3. Write a balanced equation Balancing means showing an equal number of atoms for each element on both sides of the equation Remember nothing is lost or gained! Only Transferred When balancing equations only change the coefficients Never change the subscripts!! Finally make sure that all the coefficients are in the lowest possible ratio.

Balancing Equations Ca + O2  CaO Not balanced. 2 oxygen on the reactants, 1 oxygen in the products. *Equation cannot be balanced by changing subscripts* Changing a subscript changes the identity of the substance.

Balancing Equations To balance an equation correctly we need to use coefficients. Whole numbers written before the formula for reactants and products. Ca + O2  CaO – not balanced 2 Ca + O2  2 CaO – balanced

Balancing Equations Example: Replace words with symbols: Methane + oxygen  carbon dioxide + water Replace words with symbols: CH4 + O2  CO2 + H2O

Balancing Equations CH4 + O2  CO2 + H2O To balance: Start with those elements that occur in only one substance on each side of the equation. **Number of atoms found by multiplying the subscript by the coefficient. Balance by trial and error. Balanced equation should have coefficients in the lowest whole number ratio possible. CH4 + O2  CO2 + H2O

Practice Problems Balance the following equation: H2(g) + O2(g)  H2O(l)

Practice Problems 2H2(g) + O2(g)  2H2O(l) H = 4 H = 4 O = 2 O = 2 The equation is balanced

Practice Problems Balance the following equation: C2H6O + O2  CO2 + H2O

Practice Problems C2H6O + 3O2  2CO2 + 3H2O C = 2 C = 2 H = 6 H = 6 The equation is balanced

Practice Problems Balance the following equations: 1. Zn + HCl ----> ZnCl2 + H2 2. Al + O2 -----> Al2O3 3. Al + CuSO4 ------> Al2(SO4)3 + Cu 4. Li + H2O ------> LiOH + H2

Practice Problems Aluminum reacts with oxygen to produce aluminum oxide. Sodium nitrate reacting with calcium chloride to produce sodium chloride and calcium nitrate Dinitrogen pentoxide reacts with water to produce nitric acid (HNO3)

Writing Complete Chemical Equations A complete reaction must be balanced and include the physical state of each reactant and product. Gas (g), solid (s), liquid (l) Example: CH4 (g) + 2O2 (g)  CO2 (g) + 2H2O (g)

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Classifying Chemical Reactions 9-3

Type of Reactions Four types of chemical reactions. Direct combination (Synthesis) Decomposition Single-replacement Double-replacement

Synthesis Two or more reactants come together to form a single product. Synthesis reactions. A + B  AB Product is always more complex than either of the reactants.

Synthesis Synthesis – two or more substances combine to form one new substance It follows this general form: Element or compound + Element or compound  Compound

Synthesis Examples of synthesis: NH3 + HCl  NH4Cl CaO + SiO2  CaSiO3 2H2 + O2  2H2O 2Na + Cl2  2 NaCl CO2 + H2O  H2CO3

Decomposition Reactions Something breaking down into smaller parts. A reaction in which a single compound breaks into two or more smaller compounds or elements. Identified by only one reactant: AB  A + B AB represents a compound A and B represent elements or simpler compounds.

Decomposition Reactions Decomposition – when a compound is breaking apart or decompose into simpler substances when energy is supplied Energy may be supplied in the form of heat, light, mechanical shock, or electricity It follows this general form: Compound  Two or more elements or compounds

Decomposition Reactions Examples of Decomposition: 2H2O2 → 2H2O + O2 CaCO3 → CaO + CO2 H2CO3 → H2O + CO2 2KClO3 → 2KCl + 3O2 2 H2O  2 H2 + O2 CaCO3  CaO + CO2

Single-Replacement Reaction Single Replacement – one element displaces another in a compound. It follows this general form: A + BC  AC + B Element + Compound  Element + Compound Example of single displacement: 2AgNO3(aq) + Zn(s) → 2Ag(s) + Zn(NO3)2(aq)

Single-Replacement Reaction

Single-Replacement Reaction An uncombined element displaces an element that is part of a compound. Reactants: one element and one compound A + BX  AX + B Generally ionic compounds.

Single-Replacement Reaction Example: Mg + CuSO4  MgSO4 +Cu A more active element will replace a less active element. Figure 9-19 – activity series: Used to predict whether or not a single-replacement reaction will occur. An element can replace any element that is below it. Metals replace metals or hydrogen Nonmetals replace nonmetals Cl2 + 2 KI  ? Cl2 + 2 KI  KCl + I2

Double-Replacement Reaction Double Replacement – is similar to single displacement but it uses two compounds instead of one element It follows this general form: Compound + Compound  Compound + Compound AB + CD → AD + CB

Double-Replacement Reaction Atoms or ions from two different compounds replace each other. Identifying feature: two compounds as reactants and two compounds as products. AX + BY  AY + BX Example: CaCO3 + 2HCl  CaCl2 + H2CO3

Double-Replacement Reaction Examples of Double Displacement: AgNO3 + HCl  AgCl + HNO3 Fe2O3 + HCl  FeCl3 + H2O

Double-Replacement Reaction Conditions: Most reactions will not occur unless the reactants are dissolved in water so that the ions can separate into ions. Reactions are more likely to take place if one of the products is a molecular compound, a precipitate, or a gas.

Combustion Combustion – occurs when a compound burns in air, it is actually reacting with the oxygen in the air. The products of the oxidation of the hydrocarbon under normal conditions are carbon dioxide and water vapor. Combustion reactions are refered to as oxidation reactions. They follow this general form: Hydrocarbon + Oxygen  Carbon Dioxide + Water (MOST USED) Hydrocarbon + Flourine  Carbon Flouride + Hydrogen Flouride

Combustion Examples of Combustion/Oxidation: CH4 + 2O2 → CO2 + 2H2O CH2S + 6F2 → CF4 + 2HF + SF6 CH4 + 2O2 + 7N2 → CO2 + 2H2O + 7N2 + heat

Combustion Not all reactions take one of these five general forms. There are other classes of reactions, but those we will talk about in a later chapter.

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