Copy of Unit 6-Calculations in Chemical Reactions (Fancy)

Unit 6 - Calculations in Chemical Reactions (Part 1)

Objectives 1-8

  1. Differentiate between qualitative and quantitative measurements.

  2. Differentiate between accuracy and precision; calculate accuracy.

  3. Determine the number of significant figures in a measurement.

  4. Correctly identify metric prefixes (kilo-, deci-, centi-, milli-, micro-) and their appropriate powers of 10.

  5. Correctly read balances, thermometers, rulers, and graduated cylinders.

  6. Complete unit conversions using dimensional analysis (Factor-Label method).

  7. Calculate percent error in measurements.

  8. Understanding and calculating density.

Objective 1 - Qualitative vs. Quantitative

  • Qualitative Measurements:

    • Descriptive, non-numeric form.

    • Influenced by perception.

  • Quantitative Measurements:

    • Results in numeric form (e.g., temperature change).

    • Not influenced by perception.

Objective 1 - Examples of Qualitative vs. Quantitative

  • Qualitative: After the reaction, the solution is green.

  • Quantitative: The volume of the solution is 232.5 mL.

Objective 2 - Accuracy vs. Precision

  • Precision:

    • Closeness of measurements to each other (Consistent results).

  • Accuracy:

    • Closeness of a measured value to a true value (Correct results).

Practical Scenario 1: Precision Comparison

  • Student A: 72.75g, 73.34g, 73.02g, 73.25g.

  • Student B: 72.01g, 71.99g, 72.00g, 71.98g.

  • More Precise Data: Student B

Practical Scenario 2: Measuring Error

  • Beaker Mass (actual) = 50.62 g, Sugar Mass (recorded) = 19.26 g.

  • Measuring error if laboratory balance was not zeroed out.

Measuring Accuracy (Quantitatively)

  • Percent Error Equation:

    % Error = (Observed value - Expected value) / Expected value * 100

  • Acceptable error is within ±5%.

Example Problem for % Error

  • Roger found density of aluminum = 2.57 g/cm³; Actual density = 2.70 g/cm³.

  • Percent Error: -4.81%.

Objective 4 - Scientific Notation

  • Used for large/small numbers.

  • Written as the product of:

    • A coefficient (1 < coefficient < 10)

    • 10 raised to a power.

Objective 4 - Converting into and out of Scientific Notation

  • Convert between standard and scientific notation:

    • Examples provided for both conversions.

Objective 3 - Significant Figures (Sig. Figs.)

  • Significant Figures: Used to maintain consistency in precision.

    • All non-zero digits are significant.

    • Zeroes between significant figures are significant.

    • Leading zeroes are never significant.

Objective 3 - Determining Sig. Figs

  • Example determination provided for several numeric values.

Objective 4 - Rounding Rules for Sig. Figs

  • If next number < 5, drop it or replace with zero (after decimal).

  • If next number > 5, round up the last sig fig.

Objective 5 - Sig. Figs. in Measurements

  • All digits that can be known precisely plus one estimated digit are significant.

Objective 4 - Sig. Figs. in Calculations

  • Carry extra digits through calculations, then round last.

  • Addition/Subtraction: Round to fewest decimal places of any number.

  • Multiplication/Division: Round to the fewest sig figs present in any number.

Objective 7 - Dimensional Analysis (Factor-Label Method)

  • Converting one unit to another using:

    1. Write given quantity as a fraction.

    2. Insert conversion factors.

    3. Cancel units.

    4. Solve mathematically.

Objective 6 - Metric Prefixes

  • Key metric conversions:

    • 10 dm = 1 m

    • Conversion factors highlighted.

Objective 8 - Finding Density

  • Density Formula: Density = Mass / Volume

  • Units can include g/cm³, g/mL, g/L.

  • Example provided: Density of common substances.

Finding Density Examples

  • Calculation of density based on mass and volume.

  • Density of silver calculated with different mass values.