Unit 3: Composition of Substances and Solutions
How a chemist bridges the gap between a mass on a balance and a count of particles too small to see, and how that same bridge works for dissolved solutions.
9 topics · 33 questions
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3.1 The Mole and Avogadro's Number
- How it worksHow does one mole of anything end up holding exactly 6.022 times 10 to the 23rd particles?
- Worked exampleFind how many carbon atoms are packed into 12.0 grams of pure graphite.
- Common mistakeWhy isn't the mole itself a unit of mass the way the gram is?
- Going furtherWhat changes when you count formula units of NaCl instead of separate sodium and chloride ions?
Your turn1 of 2
A sample contains 6.33 mol of a pure substance. If the number of particles it contains is written as (coefficient) times 1023, what is the coefficient?3.1 Molar Mass and Formula Mass
- How it worksWalk me through building the molar mass of calcium nitrate from the atomic masses of its elements.
- Worked exampleFind the mass of 0.250 mol of calcium nitrate, Ca(NO3)2.
- Worked exampleFind the number of moles in a 5.00 gram sample of sodium bicarbonate, NaHCO3.
- Common mistakeWhy isn't molar mass simply the total count of protons and neutrons sitting in one atom?
Your turn1 of 2
Find the mass, in grams, of 2.1 mol of calcium nitrate, Ca(NO3)2 (molar mass 164.10 g/mol).3.2 Percent Composition from a Chemical Formula
- How it worksHow does dividing each element's mass contribution by the total molar mass give a percent composition?
- Worked exampleFind the percent composition by mass of hydrogen, carbon, and oxygen in glucose, C6H12O6.
- Common mistakeWhy isn't percent composition the same as the ratio of atoms in the formula?
Your turn1 of 2
Find the percent by mass of carbon in a compound with formula C9H16O11 (use atomic masses C = 12.01, H = 1.008, O = 16.00).3.2 Determining an Empirical Formula from Percent Composition
- How it worksHow does converting percentages to moles and dividing by the smallest value produce a whole-number ratio?
- Worked exampleFind the empirical formula of a compound that is 40.0 percent carbon, 6.7 percent hydrogen, and 53.3 percent oxygen by mass.
- Common mistakeWhy isn't an empirical formula guaranteed to match a compound's actual molecular formula?
Your turn1 of 2
A compound is 55 percent carbon and 10 percent hydrogen by mass, with the remainder oxygen. Using atomic masses C = 12.01 and H = 1.008, find the ratio of moles of hydrogen to moles of carbon.3.2 Molecular Formula from Empirical Formula and Molar Mass
- How it worksHow does dividing the actual molar mass by the empirical formula mass tell you how many times to multiply the formula?
- Worked exampleFind the molecular formula of a compound with empirical formula CH2O and a molar mass of 180.16 g/mol.
- Common mistakeWhy isn't the empirical formula mass automatically equal to the compound's real molar mass?
- Ties togetherHow does the empirical-to-molecular jump connect back to the isomers that share one empirical formula?
Your turn1 of 2
A compound has an empirical formula mass of 48 g/mol and a molar mass of 255 g/mol. What whole-number multiple is the molecular formula of the empirical formula?3.2 Combustion Analysis for Finding a Formula
- How it worksWalk me through how the masses of CO2 and H2O trapped after combustion get converted back into moles of carbon and hydrogen.
- Worked exampleFind the empirical formula of a hydrocarbon that produces 2.20 g of CO2 and 1.20 g of H2O when 0.600 g of it is burned.
- Common mistakeWhy isn't the oxygen in the collected CO2 and H2O automatically assumed to have come from the original sample?
Your turn1 of 2
Burning a hydrocarbon sample produces 7.12 g of CO2 (molar mass 44.01 g/mol). How many moles of carbon were in the original sample?3.3 Molarity and Preparing a Solution of Known Concentration
- How it worksHow does dissolving a measured number of moles in a fixed final volume define molarity precisely?
- Worked exampleFind the molarity of a solution made by dissolving 25.0 g of NaOH in enough water to make 500. mL of solution.
- Common mistakeWhy isn't a solution's molarity the same before and after you add more solvent to it?
- Going furtherWhat happens to a solution's molarity if you spill half of it out without adding or removing any solute?
Your turn1 of 2
Find the molarity of a solution made by dissolving 56 g of NaOH (molar mass 40.00 g/mol) in enough water to make 684 mL of solution.3.3 Dilution and the M1V1 Equals M2V2 Relationship
- How it worksHow does the number of moles of solute staying fixed while volume grows explain why M1V1 equals M2V2?
- Worked exampleFind the volume of 12.0 M HCl needed to prepare 250. mL of a 0.500 M HCl solution.
- Common mistakeWhy isn't diluting a solution changing the total number of moles of solute it contains?
Your turn1 of 2
Find the volume of 14 M stock solution needed to prepare 667 mL of a 4.67 M solution.3.4 Other Units for Solution Concentration: Molality, Mole Fraction, and ppm
- How it worksHow do molality, mole fraction, and parts per million each define concentration using a different denominator?
- Worked exampleFind the molality of a solution made from 5.00 g of glucose dissolved in 100. g of water.
- Worked exampleFind the mole fraction of ethanol in a solution containing 3.0 mol ethanol and 7.0 mol water.
- Common mistakeWhy isn't molality affected by a change in the solution's temperature the way molarity is?
- Going furtherWhat happens to the difference between molarity and molality as a solution becomes extremely dilute?
Your turn1 of 2
Find the molality of a solution made from 35 g of glucose (molar mass 180.16 g/mol) dissolved in 342 g of water.