Objectives and Measurement in Chemistry

Introduction

In chemistry, accurate measurement and understanding of units are foundational skills. This section covers the objectives for introductory chemistry, including reading measurement devices, understanding significant figures, performing unit conversions, and introducing the concept of the mole.

• Reading a graduated cylinder: Essential for measuring liquid volumes accurately.

• Significant figures: Understanding and applying significant figures to reflect measurement precision.

• Unit conversions: Calculating and converting between different units, especially within the metric system.

• The mole: Introduction to counting atoms and molecules using the mole concept.

Metric System and SI Units

Metric System

The metric system is a unified and rational system of measurement created during the French Revolution. It is based on powers of 10 and uses prefixes to indicate the scale of measurement.

• Prefixes: Used to tell how big or small a unit is (e.g., nano-, giga-).

• Examples:

  • Nano (n):

  • Giga (G):

For example, the atomic radius of hydrogen is about 200 picometers, while the distance from the Earth to the Sun is about 150 gigameters.

SI Units (Système International d'Unités)

The SI system is the modern form of the metric system, established in 1960 to standardize measurable quantities in science.

• Base units: Each physical quantity has a standard unit and symbol.

SI Prefix Multipliers

SI prefixes indicate multiples or fractions of units. Here is a summary table:

Significant Figures

Definition and Importance

Significant figures (sig figs) indicate the level of uncertainty in a measurement and reflect the accuracy of the measuring device.

• More significant figures mean higher measurement precision.

• Significant figures are determined by the measuring instrument and estimation.

How to Count Significant Figures

• If a decimal is present, count from the left to the right, starting at the first nonzero digit.

• If a decimal is absent, count from the right to the left, starting at the first nonzero digit.

• For scientific notation, only the significant digits are counted (e.g., has 3 sig figs).

Infinite Significant Figures

• Counting numbers (e.g., 7 students) have infinite sig figs.

• Physical constants and defined values (e.g., speed of light, conversion factors) are considered to have infinite sig figs.

Rules for Calculations

• Multiplication/Division: The final answer can only have as many sig figs as the measurement with the fewest sig figs.

• Addition/Subtraction: The result should be equal to the smallest number of decimal places in the original measurements.

• If there are no decimal places in the numbers, the answer should not have any decimal places after rounding.

Measurement Tools and Estimation

Graduated Cylinders and Burettes

• Always read the bottom of the meniscus for liquid measurements.

• Determine what each line on the scale means.

• Use significant figures to make an accurate estimate.

Unit Conversions and Dimensional Analysis

Converting from One Unit to Another

Unit conversions in chemistry often use the factor-label or dimensional analysis method.

• Set up conversion factors so that unwanted units cancel out, leaving the desired unit.

• Conversion problems typically have infinite sig figs for the conversion factors.

• If the result is more than 1000, use scientific notation. If less than 0.001, also use scientific notation.

Example:

Convert 3.0 ft to inches:

Uncertainty in Measurement

The uncertainty for scaled instruments is half of the smallest division. For example, if the smallest scale on a ruler is 0.1 cm, the uncertainty is ±0.05 cm.

• For electronic balances, the uncertainty is typically ± the last decimal place shown.

Density

Definition and Calculation

Density is the mass per unit volume of a substance.

• Formula:

• Units: g/mL or g/cm3

When calculating the largest possible density, use the largest mass and smallest volume. For the smallest possible density, use the smallest mass and largest volume.

The Mole Concept

Definition and Use

The mole is a counting unit used to quantify atoms or molecules. One mole contains Avogadro's number () of entities.

• 1 mole = atoms or molecules

• 1 mole of an atom (in grams) = atomic mass in grams

Example: 1 mole of Helium = 4.00 g

Using the Periodic Table

• The number under the element's name is the atomic mass (in amu or grams per mole).

• To find the mass of multiple moles, multiply the number of moles by the atomic mass.

Example: What is the mass of 2 moles of He?

Significant Figures in Mole Calculations

• Use the number of significant figures from the periodic table value in calculations.

• For example, 3.03 g H (3 sig figs) × 1 H mole/3.03 g H = 3.00 H mol (3 sig figs).

Molecular Mass vs. Molar Mass

Definitions

• Formula Mass: The mass of a single molecule or formula unit, measured in atomic mass units (amu).

• Molar Mass: The mass of one mole of a substance, measured in grams per mole (g/mol).

Example: For NaCl, the formula mass is the sum of the atomic masses of Na and Cl in amu, while the molar mass is the sum in grams per mole.

Chemical Formulas

• To find the molar mass of a compound, add the atomic masses of all atoms in the formula.

• Example: For CHCl3:

Additional info: The notes also reference the periodic table and the importance of using the correct number of significant figures in all calculations, as well as the use of scientific notation for very large or small numbers.