Chapter 3: Calculations and the Chemical Equation

1. Chapter 3: Calculations and the Chemical Equation The Mole Concept and Atoms

2. Atomic mass unit • 1 amu = 1.661 X 10-24 g • Because the mass of one amu is so small, chemists deal with a much larger number of atoms while working with chemicals

3. Mole • One mole is defined as 6.022 X 1023. • This refers to one mole of anything, eggs, paperclips, atoms. One mole of anything is 6.022 X 1023 items. Much like one dozen of something is 12. • This number, 6.022 X 1023 is called Avogadro’s number, named after the scientist who conducted a series of experiments leading to the “mole concept”.

4. The mole concept • The mole and the amu are related. For atoms, the atomic mass of an element corresponds to the average mass of a single atom in amu And • The mass of a mole of atoms in grams.

5. For example: • The atomic mass of oxygen is 16.00 amu. And • One mole of oxygen atoms (6.022 X 1023 oxygen atoms) has a mass of 16.00 grams

6. Another example • The atomic mass of iron (Fe) is 55.85 amu. And • One mole of iron atoms (6.022 X 1023 oxygen atoms) has a mass of 55.85 grams

7. And yet another example • The atomic mass of radium (Ra) is 226 amu. And • One mole of radium atoms (6.022 X 1023 radium atoms) has a mass of 226 grams

8. Molar mass • The mass of one mol (mole) of atoms in grams

9. Note • One mole of atoms of any element contains 6.022 X 1023 atoms, regardless of the type of element. • The mass of one mole of an element depends on what that element is, and is equal to the atom mass of that element in grams.

10. This means • There are several “conversions” regarding atoms, moles, and mass

11. Converting moles to atoms • How many atoms are in 4 moles of H? 4 moles H X 6.022 X 1023 atoms/mole = 24.088 X 1023 atoms of hydrogen or 2.409 X 1024 atoms • In this case you multiply the number of moles X the number of atoms in each mole.

12. Converting atoms to moles • Calculate the number of moles of copper represented by 3.26 X 1024 atoms. 3.26 X 1024 = 32.6 X 1023 (ok, I did this step to make the math easier.) 32.6 X 1023 / 6.022 X 1023 = 5.413 X 1023 moles of copper. • In this case, to convert atoms to moles, I divide the number of atoms by the number of atoms in one mol (by 6.022 X 1023 )

13. Converting moles of a substance to mass in grams. • What is the mass in grams of 5.6 mol of Neon? • The mass of one mole of Ne is the same as the atomic mass in g (20.18 g) • 5.6 mol X 20.18 g/mol = 100.9 g of Ne

14. Converting grams to numbers of atoms. • How many atoms would be in a gold ring that weighs 25 g? • First, find the number of moles of Gold in 25 g. Gold has an atomic mass of 107.9. • So, 25 g / 107.9 g/mol = 0.2317 mol of gold are in the ring. • Next, 0.2317 mol X (6.022 X 1023) atoms/mol =1.395 X 1023 atoms

15. When dealing with molecules. . . • Like O2 or H2 , double the number of atoms, because there are 2 atoms per molecule. • Remember, one mole of something is 6.022 X 1023 of whatever it is. If it is molecules, it’s6.022 X 1023 of them. If it is atoms, it’s 6.022 x1023 atoms. • If there are 2 atoms per molecule you need to double the number of moles. • 2 X (6.022 X 1023 ) = 12.044 X 1023 or 1.204 X 1024

16. Homework Assignment # 10 • Read p. 119-123. • As you read, complete exercises 1-6. • When you are done reading, answer problems 23-36 on p. 146-147

17. Chapter 4: Calculations and the Chemical Equation Section 4.2: The Chemical Formula, Formula Weight, and Molar Mass

18. Chemical Formula • A combination of symbols of the various elements that make up the compound.

19. Formula Unit • The smallest amount of atoms that provides the following information • The identity of atoms in the compound • The relative numbers of each type of atom • Examples

20. Molecule vs ion pair • Covalent compounds form molecules, and when calculating formula weight all of the atoms in the compound are added together. • Ion pairs (ionic compounds) form crystalline structures. It’s the smallest group of ions that are listed in the formula for these types of chemicals.

21. Formula weight vs Molecular Weight • The sum of all of the atomic weights in the compound in an ionic compound it’s the formula weight. In a covalent compound it’s the molecular weight.

22. Molar Mass • The mass of one mole of the compound or the formula weight in grams. • Examples

23. Conversions using Formula Weight • Finding the number of moles corresponding to a certain number of grams.

24. Conversions using Formula Weight • Finding grams corresponding to a certain number of moles.

25. Homework Assignment #11 • Read p. 123-126. • On p. 147 Exercises 37-58

26. Chapter 4: Calculations and the Chemical Equation Section 3: The Chemical Equation and the Information it Conveys

27. Chemical equation • The shorthand notation for a chemical reaction, where one substance changes chemically into another substance. • An example: burning sugar

28. Reactants • The starting materials that undergo a chemical change

29. Products • The ending materials that are produced by a chemical reaction.

30. Additional information in a chemical reaction • Physical state of the substance (solid, liquid, or gas) • If the reaction occurs • Identifies the solvent, if there is one. (A solvent is the solution the materials are dissolved in, such as water.) • Experimental conditions such as heat, light, or electrical energy added

31. Most importantly • The chemical equation identifies the relative number of moles of reactants and products.

32. According to the Law of Conservation of Mass • Matter cannot be gained or lost in the process of a chemical reaction • The total mass of the products must equal the total mass of the reactants • The chemical equation must be balanced.

33. Features of a chemical reaction. • CaCO3(s) →∆ CaO(s) + CO2(g) The products are on the right of the arrow. Reactants are on the left of the arrow. The arrow indicates the reaction occurs in one direction.

34. Features of a chemical reaction. • CaCO3(s) →∆ CaO(s) + CO2(g) “l” would indicate the substance were a liquid. “s” indicates the chemical is a solid substance “g” indicates the substance is a gas

35. Features of a chemical reaction. • CaCO3(s) →∆ CaO(s) + CO2(g) The ∆ indicates that energy was necessary for the chemical reaction to occur

36. Features of a chemical reaction. • CaCO3(s) →∆ CaO(s) + CO2(g) The main feature of a chemical equation is that it is balanced, with the same number of elements in compounds on both sides of the arrow.

37. The experimental basis of a chemical equation • Evidence for a chemical reaction includes: • The release of a gas resulting in bubbles • The formation of a solid (precipitate) in solution • The production of heat resulting in an increase in temperature • A change in color of a substance

38. The experimental basis of a chemical equation • Sometimes instruments must be used to measure subtle changes that indicate a chemical reaction. • Heat or light absorbed or emitted • Changes in the way a sample behaves in an electrical or magnetic field • Changes in electrical properties

39. Writing Chemical Reactions • Most reactions follow a few simple patterns • Combination reactions • Decomposition reactions • Replacement reactions

40. Combination reactions • Involve the joining or combining of two or more compounds • The general form of the reaction is A + B → AB

41. Combination reactions • Examples include: • Combination of a metal and non-metal to form a salt Ca(s) + Cl2(g)→ CaCl2(s) • Reaction of magnesium oxide and carbon dioxide to produce magnesium carbonate MgO(s) + CO2(g) →MgCO3(s)

42. Decomposition Reactions • Reactions that produce two or more products from a single reactant. • The general form for the reaction is AB → A + B

43. Decomposition reactions • Examples include • The removal of water from a hydrate (a substance that has water molecules linked in it’s structure) CuSO4·5H2O(s) →CuSO4(s) + 5H2O(g) • The heating of calcium carbonate to produce calcium oxide and carbon dioxide gas CaCO3(s) → CaO(s) + CO2(g)

44. Replacement Reactions—Single Replacement • Single replacement reactions is where one atom replaces another in the compound • The general formula is A + BC → AC + B

45. Replacement Reactions • Examples include • The replacement of copper by zinc in copper sulfate forming zinc sulfate Zn(s) + CuSO4(aq) → Zn SO4(aq) + Cu(s)

46. Replacement Reactions—Double Replacement • Two compounds that “switch” atoms with each other • The general formula is AB + CD → AD + CB

47. Replacement Reactions—Double Replacement • Examples include • The formation of salt and water with the reaction of a base and an acid HCl(aq) + NaOH(aq) →H2O(l) + NaCl(aq)

48. Types of Chemical Reactions • There are four main types of chemical reactions • Precipitation reactions • Reactions with Oxygen • Acid-base reactions • Oxidation-reduction reactions

49. Precipitation reactions • A chemical change that produces an insoluble product that will form a solid. Usually the solid can be seen “falling out” of the solution, hence, called precipitation. At other times the solid makes the solution turn from clear to cloudy.

50. Solubility predictions • Na, K, and ammonium compounds are generally soluble. • Nitrates and acetates are generally soluble • Chlorides, bromides, and iodides are generally soluble. However, iodine compounds that contain lead, silver, and mercury are insoluble. • Carbonates and phosphates are generally insoluble. Sodium, potassium, and ammonium carbonates and phosphates are soluble. • Hydroxides and sulfides are generally insoluble. Sodium, potassium, calcium, and ammonium compounds are however soluble.