510 likes | 673 Views
Chemical Reactions Ch. 4. Milbank High School. Sec. 4.1 Balancing Chemical Equations. Objectives Write balanced chemical equations, when given the names or formulas of the reactants and products in a chemical reaction. Balanced Equation. Atoms can’t be created or destroyed
E N D
Chemical ReactionsCh. 4 Milbank High School
Sec. 4.1Balancing Chemical Equations • Objectives • Write balanced chemical equations, when given the names or formulas of the reactants and products in a chemical reaction.
Balanced Equation • Atoms can’t be created or destroyed • All the atoms we start with we must end up with • A balanced equation has the same number of each element on both sides of the equation.
Rules for balancing: • Assemble, write the correct formulas for all the reactants and products • Count the number of atoms of each type appearing on both sides • Balance the elements one at a time by adding coefficients (the numbers in front) - save H and O until LAST! • Check to make sure it is balanced.
Never • Never change a subscript to balance an equation. • If you change the formula you are describing a different reaction. • H2O is a different compound than H2O2 • Never put a coefficient in the middle of a formula • 2 NaCl is okay, Na2Cl is not.
Example H2 + O2 ® H2O Make a table to keep track of where you are at
Example H2 + O2 ® H2O R P 2 H 2 2 O 1 Need twice as much O in the product
Example H2 + O2 ® 2 H2O R P 2 H 2 2 O 1 Changes the O
Example H2 + O2 ® 2 H2O R P 2 H 2 2 O 1 2 Also changes the H
Example H2 + O2 ® 2 H2O R P 2 H 2 4 2 O 1 2 Need twice as much H in the reactant
Example 2 H2 + O2 ® 2 H2O R P 2 H 2 4 2 O 1 2 Recount
Example 2 H2 + O2 ® 2 H2O R P 4 2 H 2 4 2 O 1 2 The equation is balanced, has the same number of each kind of atom on both sides
Example 2 H2 + O2 ® 2 H2O R P 4 2 H 2 4 2 O 1 2 This is the answer Not this
Balancing Examples • _AgNO3 + _Cu ® _Cu(NO3)2 + _Ag • _Mg + _N2® _Mg3N2 • _P + _O2® _P4O10 • _Na + _H2O ® _H2 + _NaOH • _CH4 + _O2® _CO2 + _H2O
Sec. 4.2 Volume Relationships in Chemical Equations • Objectives • Define Gay-Lussac’s Law • Determine how to find volume relationships in a gaseous chemical reaction
Gay-Lussac’s Law • When gases measured at the same temperature and pressure are allowed to react, the volumes of gaseous reactants and products are in small whole-number ratios • 3 hydrogen gas volumes + 1 nitrogen gas volumes = 2 volumes ammonia • (3:1:2) • Practice Exercises Pg. 103
Sec. 4.3Avogadro’s Number • Objectives • Define Avogadro’s Number
Moles (abbreviated: mol) • Defined as the number of carbon atoms in exactly 12 grams of carbon-12. • 1 mole is 6.02 x 1023 particles. • 6.02 x 1023 is called Avogadro’s number. • How large is it?
Sec. 4.4Molecular Masses and Formula Masses • Objectives • Define molecular and formula mass
Molecular Mass • Molecular mass is the average mass of a molecule of a substance relative to that of a carbon-12 atom • sum of masses of the atoms represented in a molecular formula • Formula mass is the same as the molar mass except that it refers to ionic compounds
Sec. 4.5Chemical Arithmetic and the Mole • Objectives • Define and determine molar mass
Molar Mass • Molar mass is the generic term for the mass of one mole of any substance (in grams) • Numerically equal to atomic mass, molecular mass, or formula mass • Expressed as g/mol
Examples • Calculate the molar mass of the following and tell what type it is: • Na2S • N2O4 • C • Ca(NO3)2 • C6H12O6 • (NH4)3PO4
What is a Mole? • You can measure mass, • or volume, • or you can count pieces. • We measure mass in grams. • We measure volume in liters. • We count pieces inMOLES.
Molar Mass • The number of grams of 1 mole of atoms, ions, or molecules. • We can make conversion factors from these. • To change grams of a compound to moles of a compound.
For example • How many moles is 5.69 g of NaOH?
For example • How many moles is 5.69 g of NaOH?
For example • How many moles is 5.69 g of NaOH? • need to change grams to moles
For example • How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH
For example • How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH • 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g
For example • How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH • 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g • 1 mole NaOH = 40.00 g
For example • How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH • 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g • 1 mole NaOH = 40.00 g
For example • How many moles is 5.69 g of NaOH? • need to change grams to moles • for NaOH • 1mole Na = 22.99g 1 mol O = 16.00 g 1 mole of H = 1.01 g • 1 mole NaOH = 40.00 g
Examples • How many moles is 4.56 g of CO2? • How many grams is 9.87 moles of H2O? • How many molecules is 6.8 g of CH4? • 49 molecules of C6H12O6 weighs how much?
4.6 Mole and Mass Relationships in Chemical Equations • Objectives • Determine what a stoichiometric factor is • Solve problems using stoichiometric factors
Stoichiometric Factors • Relates the amounts of any two substances involved in chemical equations
Rules • Write a balanced equation • Determine molar masses • Use molar mass to convert grams to moles of substance • Convert moles of given substance to moles of the substance you wish to find • Use molar mass to convert moles to grams of that substance
Example 4.11 • Example 4.12 • Practice Exercises
Limiting Reagent • Limits or determines the amount of product that can be formed in a reaction • Excess reagent • Reactant that is not completely used up
Example • Sodium chloride can be prepared by the reaction of sodium metal with chlorine gas. 2Na + Cl2 2NaCl Suppose that 6.70 mol Na reacts with 3.20 mol Cl2. What is the limiting reagent? How many moles of NaCl are produced?
Sec. 4.7Structure, Stability, and Spontaneity • Objectives • Determine what exothermic and endothermic reactions are • Define enthalpy
Heat released or absorbed • Heat of reaction • Heat released or absorbed during a chemical reaction • Enthalpy change • Equal to heat of reaction • ∆H • Exothermic • Release heat, -∆H • Endothermic • Absorb heat, ∆H
Sec. 4.8Forward and Reverse Reactions • Objectives • Define activation energy • Recognize the components of a reaction profile
Activation Energy • Minimum energy needed to get the reaction started • AKA energy of activation • Ea • See reaction profiles
Sec. 4.9Reaction Rates: Collisions and Orientation • Reaction rates profoundly affected by temperature, catalysts, and concentration • For a reaction to occur: • Collision of particles • Proper orientation (Fig. 4.13) • Minimum activation energy
Temperature • Generally faster reactions at higher temps • More collisions, more energy • Uses: • Freezing food • Cooking • Organisms • Fevers • Hibernation • Heart surgery
Catalysts • A substance that changes the rate of reaction without being changed itself • Lowers the activation energy (Fig. 4.15) • Increase speed of slow reaction • Enzymes • Biological catalysts • Mediate reactions in living systems
Concentration • More molecules = more collisions • Double concentration = double rate of reaction
Mechanism • Step-by-step process of a reaction • Affected by temp, catalysts, and concentration • Cancer • Break-down in the mechanism of a reaction • Poisons • Knowledge of mechanisms has led to effective treatments
Sec. 4.10Equilibrium in Chemical Reactions • Some reactions proceed in forward and reverse directions at the same time • Double arrow • Equilibrium eventually reached • Rates of forward reaction = rates of reverse reaction