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.