Unit 3 Review Chemical Calculations

Unit 3 Review: Chemical Calculations

Page 1: Overview

• Introduction to chemical calculations, focusing on precision, accuracy, significant figures, molar calculations, and solutions.

Page 2: Measurement and Significant Digits

• Ruler Measurements:

  • Ruler A: 2.55 (2 certain digits and one estimate).

  • Ruler B: 2.5 (2 significant digits).

• Significant Digits

  • Non-zero digits are always significant.

  • Leading zeros are not significant.

  • Trailing zeros in a decimal are significant.

Page 3: Precision

• Definition of Precision: Number of significant digits measured.

• Example:

  • Graduated cylinder: 10 mL can measure to 0.2 mL.

  • Pipette: 10 mL can measure to 0.1 mL.

• Pipettes are more precise than graduated cylinders.

Page 4: Accuracy vs. Precision

• Accuracy: How close a measurement is to the true value.

• Types of Accuracy: High and low accuracy can also relate to precision.

Page 5: Measurement and Uncertainty

• Choosing the right equipment is crucial for accurate results.

• Last digit in a measurement is uncertain.

Page 6: Certainty in Measurement

• Example of Measuring with Rulers:

  • Ruler one measures to 3 significant digits (2 certain + 1 uncertain).

  • Ruler two only measures to 2 significant digits.

Page 7: Significant Digit Rules

• Rule 1: Digits 1-9 are significant.

• Rule 2: Leading zeros aren't significant.

  • Example: 0.000034 L has 2 significant digits.

Page 8: More Significant Digit Rules

• Rule 3: Zeros between non-zero digits are significant.

  • Example: 205.007 kg has 6 significant digits.

• Rule 4: Trailing zeros in decimals are significant.

  • Example: 5.3400 cm³ has 5 significant digits.

Page 9: Counted Quantities

• Counted quantities have no uncertainty.

  • Example: There are 34 students implies infinite significant digits.

Page 10: Scientific Notation

• Recommended for large/small numbers to maintain significant figures.

Page 11: Rules for Scientific Notation

• 1: One digit in front of the decimal.

• 2: Round to the correct number of significant digits.

• 3: Multiply by correct power of 10.

Page 12: Rounding Rules

• Less than 5: round down.

• More than 5: round up.

• Exactly 5: round to even.

Page 13: Precision in Addition/Subtraction

• Answer's decimal places should match measurement with fewer decimal places.

  • Example: 2.36 + 852.0 + 100.00 = 954.4 mL

Page 14: Certainty in Multiplication/Division

• Answer should match the measurement with the fewest significant digits.

  • Example: A = l x w = 2.36 cm x 25.69 cm = 60.6 cm².

Page 15: Quantitative vs. Qualitative

• Quantitative Analysis: Measures quantities in a sample.

• Qualitative Analysis: Identifies substances without measurements.

Page 16: Conversion Techniques

• Factor Label Method for Metric Conversion.

• Mnemonic: King Henry Died Unusually Drinking Chocolate Milk.

Page 17: Factor Label Method

• Unknown = Given x Conversion Factor.

• Method focuses on cancelling out unwanted units.

Page 18: The Mole (5.1 & 5.2)

• Symbol: n, represents the amount of a substance.

• 1 mole = 6.02 x 10²³ entities (Avogadro’s Number).

Page 19: Using Moles for Counting

• Moles can convert to number of particles using Avogadro's number.

• Conceptual analogy: moles relate to dozens.

Page 20: Finding Moles

• N = number of particles.

  • n = moles.

  • N = n x NA (Avogadro’s Number).

Page 21: Molar Mass Definitions

• Mass of 1 mol of substance (g/mol).

• Example: Molar mass of water (H₂O) = 18.01528 g/mol.

Page 22: Molar Mass Calculations

• To find molar mass, sum atomic masses from periodic table.

• Example for H₂O: 2(H) + O = 18.015 g/mol.

Page 23: Conventions for Molar Calculations

• M = Molar mass, m = mass of sample, n = moles, N = number of particles.

Page 24: Converting from Moles to Mass

• Use formula: m = n x M.

  • Example: m = 0.7500 mol x 44.01 g/mol = 33.01 g of CO2.

Page 25: Converting from Mass to Moles

• Formula: n = mass/molar mass.

  • Example: 23.6 g of acetic acid = 0.393 moles.

Page 26: Solutions and Their Characteristics

• Solutions are homogeneous mixtures of two or more substances.

• Cannot be separated by physical means like filtration.

Page 27: Types of Solutions

• Homogeneous Mixtures: One phase (e.g., air, metal alloys).

• Heterogeneous Mixtures: Multiple phases (e.g., fog, mayonnaise).

Page 28: Solute and Solvent Definitions

• Solute: Substance being dissolved.

• Solvent: Substance that does the dissolving, usually water.

Page 29: Concentration Definitions

• Concentration measures the ratio of solute to solvent: g/L or mol/L.

• Concentrated = high solute; Dilute = low solute.

Page 30: Concentration Calculation Techniques

• Percent concentration can be expressed in volume or weight.

  • E.g., % V/V, % W/V, % W/W.

Page 31: Example Percent Concentration Calculations

• % V/V = (V_solute/V_solution) x 100.

• Example calculation for acetic acid solution:

  • 140 mL of acetic acid in 500 mL total solution = 28% V/V.

Page 32: Extremely Low Concentrations

• Expressed in parts per million (ppm), parts per billion (ppb), etc.

Page 33: Stock Solutions and Dilutions

• Stock solutions are concentrated and diluted for specific use.

• Example: Diluting 12 M HCl to desired concentration.

Page 34: How to Prepare a Standard Solution

• Calculate mass of solid needed based on volume of solution desired.

• Steps to prepare include measuring, dissolving, and transferring to volumetric flasks.

Page 35: Dilutions and Their Calculations

• Use the formula: C1V1 = C2V2 for concentrations and volumes.

• Example provided for calculating final concentrations after dilution.

Page 36: Stoichiometry Overview (6.1 - 6.3)

• Study of mass/mole relationships in chemical reactions.

• Understanding mole ratios through balanced equations is essential.

Page 37: Mole Ratios

• Ratio derived from coefficients in the balanced equation.

• Key for converting between moles and mass in calculations.

Page 38: Limiting and Excess Reagents (6.5)

• Limiting Reagent: Completely consumed and limits amount of product formed.

• Excess Reagent: Remains when reaction is complete.

Page 39: Calculating Percent Yield (6.7)

• Percent yield measures efficiency: % yield = (actual yield/theoretical yield) x 100.

• Factors affecting yield include side reactions, experimental loss, and timing.

Page 40: Example Percent Yield Problem

• Actual yield derived from experiment compared to calculated theoretical yield.

• Discussion about impact on reaction efficiency and common lab issues.