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Chemisty 2202

Chemisty 2202. How I Will Be Evaluated!. Course Content. Unit 1. Stoichiometry. Say What?. Stoichiometry is defined as the study of chemistry that deals with relative amounts of reactants and products in chemical reactions.

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Chemisty 2202

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  1. Chemisty 2202

  2. How I Will Be Evaluated!

  3. Course Content

  4. Unit 1 Stoichiometry Say What?

  5. Stoichiometry is defined as the study of chemistry that deals with relative amounts of reactants and products in chemical reactions. Gravimetric Stoichiometry is involved with the analysis of masses. Gas Stoichiometry is involved with the volume of gases. Word Alert!

  6. Mass Number = the # of protons in an atom + # of neutrons Chapter 2 (Pg 42) The Mole Recall the structure of the atom Protons, Neutrons, Electrons Word Alert! Atomic Number= the number of protons in an atom

  7. # protons = Mass # - # neutrons # neutrons = Mass # - # protons Note that the equation for mass # may be rearranged to find either the protons or the neutrons in an atom provided that the other two are known.

  8. The Mole Word Alert! Isotopes - isotopes are atoms which have the same number of protons but different numbers of neutrons For example, there are 3 naturally occurring isotopes of carbon – carbon-12, carbon-13 and carbon-14

  9. Isotope Notation The isotopes carbon-12, , carbon-13 and carbon-14 may also be written as: The 6 is called the atomic #and represents the number of protons in the atom. The 12, 13, and 14 are called the atomic mass #and represent the sum of the protons and neutrons in the atom.

  10. Isotope Notation cont… Carbon-14 has 2 more neutrons than carbon-12 and one more neutron than carbon-13. Note that the number of protons remain the same.

  11. The measure of the mass of all other atoms are defined by their relationship to carbon-12. For example: Oxygen-16 has a mass that is 133% of the mass of carbon-12. Thus the mass of an atom oxygen-16 is

  12. % of Isotopes in Nature The % of each isotope in nature is different from each other. Thus when you are trying to find the average atomic mass of an element you must take into account the % of each isotope. In the case of carbon, it is not as simple as adding up the atomic masses of each isotope and dividing by 3.

  13. Mass of an atom is expressed in atomic mass units. (u) or (amu) The atomic mass units are a measure of the mass compared to the mass of the carbon-12 isotope. One atom of carbon-12 is assigned a mass of 12 u. of the mass of one atom of carbon-12 Relative Mass of an Atom(Pg 43)

  14. The average atomic mass of an element is the average of the masses of all an element’s isotopes taking into account their % abundance in nature. Word Alert! Average Atomic Mass Because isotopes occur in nature in different percentages calculating the average is not as simple as adding up the atomic mass of each isotope and dividing by how many isotopes there are.

  15. Average Atomic Mass = (mass of isotope 1 in u)(% abundance) + (mass of isotope 2 in u)(% abundance) + … Calculating Average Atomic Mass Note: - % abundance must be in decimal form. (divide by 100) - the sum of the %’s equal 100% OR in decimal form must equal 1

  16. Your Turn: complete #’s 1 to 4 on pg 45 and #’s 1, 2, 4 on pg 46 Example

  17. The Avogadro Constant and the Mole(Pg 47) 2 6 12 144 500

  18. Example: one mole of copper atoms will contain 6.02 x 1023 atoms Word Alert! Avogadro’s # Link

  19. Formula for changing moles of a substance into molecules, formula units, atoms, ions: Where n = amount in mol NA= Avogadro’s constant (mol-1) (6.02 x 1023 particles/mol) N = number of particles N = n x NA

  20. How many atoms are there in 2 mol of gold atoms? # gold atoms = (# mol)(6.02 x 1023 atoms/mol) = (2 mol)(6.02 x 1023 atoms/mol) = 12.04 x 1023 atoms of gold Note that atoms may be replaced by formula units (for ionic compounds) molecules (for molecular compounds) or ions.

  21. b) Determine the number of atoms in 1.25 mol of NO2. Since there are 3 atoms in 1 molecule of NO2 we need to multiply the number of molecules by 3. Your Turn: Complete #’s 5 to 11 on pages 51/52.

  22. 10. D) How many oxygen atoms are in the sample?

  23. This process involves a rearrangement of the equation used to calculate the particles in a certain number of moles. (N = n x NA) Converting Particles Into Moles

  24. Example Pg 53 How many moles are present in a sample of carbon dioxide, CO2, made up of 5.83 x 1024 molecules? There are 9.68 mol of CO2 in the sample.

  25. Your Turn: Complete #’s 12 to 15 on pages 53.

  26. Periodic Table Molar Mass The molar mass of an element or compound is defined as the mass of one mol of the substance. Symbol is M The Periodic Table gives the average molar mass of all elements. For example, the molar mass of gold is given as 196.97 g/mol. See Periodic Table on back inside cover of text.

  27. Calculating Molar Mass • the molar mass of an element may be determined from the Periodic Table (just look it up) • the molar mass of a compound must take into account the number of each type of atom For example, With water, H2O, there are 2-H’s and 1-O.Thus when you compute the molar mass you would multiply the molar mass of a hydrogen atom by 2 before adding it to the molar mass of oxygen.

  28. Example: Determine the molar mass of Ca3(PO4)2? Complete #’s 16 to 19 Pg 57. There are 3-Ca so 3 x 40.08 g/mol = 120.24 g/mol There are 2-P so 2 x 30.97 = 61.94 g/mol There are 8-O so 8 x 16.00 g/mol = 128.00 g/mol The molar mass of Ca3(PO4)2 is the sum of these 3. Molar Mass of Ca3(PO4)2 = 120.24g/mol + 61.94 g/mol + 128.00 g/mol = 310.18 g/mol The molar mass of Ca3(PO4)2 is 310.18 g/mol.

  29. Problem Type 1: Moles of A to Mass of A With this type of problem you are given moles and have to convert it into the corresponding mass. You must also use the molar mass of the substance.

  30. Example: A flask contains 0.750 mol of carbon dioxide gas (CO2). What mass of CO2 is in this sample? Given: n = 0.750 mol, m = ? We must determine the molar mass of CO2. 1 x 12.01 = 12.01 g/mol 2 x 16.00 = 32.00 g/mol MCO2 = 44.01 g/mol m = (0.750 mol)(44.01 g/mol) = 33.0 g

  31. Your Turn: Complete #’s 20 to 23 on pages 59.

  32. Problem Type 2: From Mass A to Moles A With this type of problem you are given mass and have to convert it into the corresponding moles. You must also use the molar mass of the substance.

  33. Example: How many moles of acetic acid, CH3COOH, are in a 23.6 g sample? We must determine the molar mass of acetic acid. Complete #’s 24 to 27 on pg 60. 4-H means 4 x 1.01 g/mol = 4.04 g/mol 2-C means 2 x 12.01 g/mol = 24.02 g/mol 2-O means 2 x 16.00 g/mol = 32.00 g/mol 60.06 g/mol There are 0.393 mol of acetic acid in 23.6 g.

  34. Converting From Mass to Particles and Vice Versa To convert from one quantity into another you must go through moles. For example to find the # of atoms in 25.0 g of Cu, we must first determine the # of moles in the 25.0 g and then convert it to # of atoms. 3.Particles to Mass Example: What is the mass of 5.67 x 1024 formula units of cobalt(II) chloride, CoCl2. Step 1: mol CoCl2 = (# formula units)( 1 mol ) (6.02 x 1023 f.u.) = (5.67 x 1024 f.u.)( 1 mol ) = 9.42 mol 6.02 x 1023 f.u.

  35. Your Turn: Complete #’s 28 to 33 on page 63. Step 2: Find corresponding mass. We must find the molar mass of CoCl2. (129.83 g/mol) # g = (# mol)(molar mass) # g = (9.42 mol)(129.83 g/mol) = 1.22 x 103 g There are 1.22 x103 gin 5.67 x 1024 f.u. of CoCl2.

  36. 4.Mass to Particles Chlorine gas, Cl2, can react with I2, to form iodine chloride, ICl. How many molecules of iodine chloride are contained in a 2.74 x 10-1 g sample? Step 1: Determine the number of moles of ICl.

  37. Your Turn: Complete #’s 34 to 37 on pages 64 and 2, 3, 4, 5, 6 on pg 65. Step 2: Convert mol ICl into molecules. There are 1.01 x 1021 molecules in 0.247 g of ICl.

  38. Section 2.4: Molar Volume (Pg 56) In this section we will: • state Avogardro’s hypothesis, and explain how it contributes to our understanding of the reactions of gases • define STP and the molar volume of a gas at STP • perform calculations converting between the number of particles, moles, mass and volume of a gas

  39. Your Turn: Complete #’s 34 to 37 on page 64 and #’s 2 to 6 on pg 65. Refer to the concept organizer on pg 64.

  40. communicate your understanding of the following terms: • pressure, • pascal (Pa), • kilopascals (kPa), • standard temperature and pressure (STP), • law of combining volumes, • Avogadro’s hypothesis, • molar volume and • ideal gas

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