1. Chapter 6 Chemical Reactions Chemical Reactions

2. Chemical Reactions In a chemical reaction, one or more reactants is converted to one or more products In this chapter we discuss three aspects of chemical reactions (a) mass relationships (stoichiometry) (b) types of reactions (c) heat gain and loss accompanying reactions

3. Chemical Equations The following chemical equation tells us that propane gas and oxygen gas react to form carbon dioxide gas and water vapor But while it tells us what the reactants and products are and the physical state of each, it is incomplete because it is not balanced

4. Balancing Equations To balance a chemical equation – begin with atoms that appear only in one compound on the left and one on the right; in this case, begin with carbon (C) which occurs in C 3 H 8 and CO 2 – now balance hydrogens, which occur in C 3 H 8 and H 2 O – if an atom occurs as a free element, as for example Mg or O 2, balance this element last; in this case O 2

5. Balancing Equations Practice problems: balance these equations

6. Balancing Equations Solutions to practice problems – it is common practice to use only whole numbers; therefore, multiply all coefficients by 2, which gives

7. Formula Weight Formula weight: the sum of the atomic weights in atomic mass units (amu) of all atoms in a compound’s formula

8. Formula Weight formula weight formula weight can be used for both ionic and molecular compounds; it tells nothing about whether a compound is ionic or molecular molecular weight molecular weight should be used only for molecular compounds in this text, we use formula weight for ionic compounds and molecular weight for molecular compounds

9. The Mole Mole (mol) – a mole of the amount of substance that contains as many atoms, molecules, or ions as are in exactly 12 g of carbon-12 – a mole, whether it is a mole of iron atoms, a mole of methane molecules, or a mole of sodium ions, always contains the same number of formula units – the number of formula units in a mole is known as Avogadro’s number – Avogadro’s number has been measured experimentally – its value is 6.02214199 x 10 23 formula units per mole

10. Molar Mass Molar mass: Molar mass: the formula weight of a substance expressed in grams Glucose, C 6 H 12 O 6 – molecular weight: 180 amu – molar mass: 180 g/mol – one mole of glucose has a mass of 180 g Urea, (NH 2 ) 2 CO – molecular weight 60.0 amu – molar mass: 60.0 g/mol – one mole of urea has a mass of 60.0 g

11. Molar Mass We can use molar mass to convert from grams to moles, and from moles to grams – calculate the number of moles of water in 36.0 g water

12. Grams to Moles Calculate the number of moles of sodium ions, Na +, in 5.63 g of sodium sulfate, Na 2 SO 4 – first we find the how many moles of sodium sulfate – the formula weight of Na 2 SO 4 is 2(23.0) + 32.1 + 4(16.0) = 142.1 amu – therefore, 1 mol of Na 2 SO 4 = 142.1 g Na 2 SO 4 – the formula Na 2 SO 4 tells us there are two moles of Na + ions per mole of Na 2 SO 4

13. Grams to Molecules A tablet of aspirin, C 9 H 8 O 4, contains 0.360 g of aspirin. How many aspirin molecules is this? – first we find how many mol of aspirin are in 0.360 g – each mole of aspirin contains 6.02 x 10 23 molecules – the number of molecules of aspirin in the tablet is

14. Stoichiometry Stoichiometry: Stoichiometry: the study of mass relationships in chemical reactions – following is an overview of the the types of calculations we study

15. Stoichiometry Problem: how many grams of nitrogen, N 2, are required to produce 7.50 g of ammonia, NH 3 – first find how many moles of NH 3 are in 7.50 g of NH 3 – next find how many moles of N 2 are required to produce this many moles of NH 3

16. Stoichiometry Practice problem (cont’d) – finally convert moles of N 2 to grams of N 2 and now do the math

17. Stoichiometry Practice problems: – what mass of aluminum oxide is required to prepare 27 g of aluminum? 23 – how many grams each of CO 2 and NH 3 are produced from 0.83 mol of urea?

18. Limiting Reagent Limiting reagent Limiting reagent: the reagent that is used up first in a chemical reaction – consider this reaction of N 2 and O 2 – in this experiment, there is only enough O 2 to react with 1.0 mole of N 2 – O 2 is used up first; it the limiting reagent – 4.0 moles of N 2 remain unreacted

19. Limiting Reagent Practice Problem – suppose 12 g of carbon is mixed with 64 g of oxygen and the following reaction takes place – complete the following table. Which is the limiting reagent?

20. Percent Yield Actual yield: Actual yield: the mass of product formed in a chemical reaction Theoretical yield: Theoretical yield: the mass of product that should be formed according to the stoichiometry of the balanced chemical equation Percent yield: Percent yield: actual yield divided by theoretical yield times 100

21. Percent Yield Practice problem: – suppose we react 32.0 g of methanol with excess carbon monoxide and get 58.7 g of acetic acid – complete this table

22. Reactions Between Ions Ionic compounds, also called salts, consist of both positive and negative ions When an ionic compound dissolves in water, it dissociates to aqueous ions What happens when we mix aqueous solutions of two different ionic compounds? – if two of the ions combine to form a water-insoluble compound, a precipitate will form – otherwise no physical change will be observed

23. Reactions Between Ions Example: – suppose we prepare these two aqueous solutions – if we then mix the two solutions, we have four ions present; of these, Ag + and Cl - react to form AgCl(s) which precipitates

24. Reactions Between Ions – we can simplify the equation for the formation of AgCl by omitting all ions that do not participate in the reaction net ionic equation – the simplified equation is called a net ionic equation; it shows only the ions that react spectator ions – ions that do not participate in a reaction are called spectator ions

25. Reactions Between Ions In general, ions in solution react with each other when one of the following can happen – two of them form a compound that is insoluble in water – two of them react to form a gas that escapes from the reaction mixture as bubbles, as for example when we mix aqueous solutions of sodium bicarbonate and hydrochloric acid – an acid neutralizes a base (Chapter 8) – one of the ions can oxidize another (Section 4.7)

26. Reactions Between Ions Following are some generalizations about which ionic solids are soluble in water and which are insoluble – all compounds containing Na +, K +, and NH 4 + are soluble in water – all nitrates (NO 3 - ) and acetates (CH 3 COO - ) are soluble in water – most chlorides (Cl - ) and sulfates (SO 4 2- ) are soluble; exceptions are AgCl, BaSO 4, and PbSO 4 – most carbonates (CO 3 2- ), phosphates (PO 4 3- ), sulfides (S 2- ), and hydroxides (OH - ) are insoluble in water; exceptions are LiOH, NaOH, KOH, and NH 4 OH which are soluble in water

27. Oxidation-Reduction Oxidation: Oxidation: the loss of electrons Reduction: Reduction: the gain of electrons Oxidation-reduction (redox) reaction: Oxidation-reduction (redox) reaction: any reaction in which electrons are transferred from one species to another

28. Oxidation-Reduction Example: if we put a piece of zinc metal in a beaker containing a solution of copper(II) sulfate – some of the zinc metal dissolves – some of the copper ions deposit on the zinc metal – the blue color of Cu 2+ ions gradually disappears In this oxidation-reduction reaction – zinc metal loses electrons to copper ions – copper ions gain electrons from the zinc

29. Oxidation-Reduction – we summarize these oxidation-reduction relationships in this way

30. Oxidation-Reduction Although the definitions of oxidation (loss of electrons) and reduction (gain of electrons) are easy to apply to many redox reactions, they are not easy to apply to others – for example, the combustion of methane An alternative definition of oxidation-reduction is – oxidation: – oxidation: the gain of oxygen or loss of hydrogen – reduction: – reduction: the loss of oxygen or gain of hydrogen

31. Oxidation-Reduction – using these alternative definitions for the combustion of methane

32. Oxidation-Reduction Five important types of redox reactions – combustion: – combustion: burning in air. The products of complete combustion of carbon compounds are CO 2 and H 2 O. – respiration: – respiration: the process by which living organisms use O 2 to oxidize carbon-containing compounds to produce CO 2 and H 2 O. The importance of these reaction is not the CO 2 produced, but the energy released. – rusting: – rusting: the oxidation of iron to a mixture of iron oxides – bleaching: – bleaching: the oxidation of colored compounds to products which are colorless – batteries: – batteries: in most cases, the reaction taking place in a battery is a redox-reaction

33. Heat of Reaction In almost all chemical reactions, heat is either given off or absorbed – example: the combustion (oxidation) of carbon liberates 94.0 kcal per mole of carbon oxidized Heat of reaction: Heat of reaction: the heat given off or absorbed in a chemical reaction – exothermic reaction: – exothermic reaction: one that gives off heat – endothermic reaction: – endothermic reaction: one that absorbs heat – heat of combustion: – heat of combustion: the heat given off in a combustion reaction; all combustion reactions are exothermic

34. End Chapter 6 Chemical Reactions