Chemistry and Measurements: Structured Study Notes for GOB Chemistry

Chemistry and Measurements

Introduction

Chemistry relies on precise measurements to describe and understand the properties of matter. The metric and International System of Units (SI) provide standardized units for scientific measurement, ensuring consistency and accuracy in laboratory and clinical settings.

Units of Measurement

Scientists use the metric system and SI units to measure volume, length, mass, temperature, and time. Understanding these units is essential for interpreting data and performing calculations in chemistry.

• Volume: The space occupied by a substance. SI unit: cubic meter (m3); Metric unit: liter (L). Equalities: 1 L = 1000 mL; 1 L = 1.06 qt; 946 mL = 1 qt.

• Length: The distance between two points. SI unit: meter (m). Equalities: 1 m = 100 cm; 1 m = 39.4 in; 1 m = 1.09 yd; 2.54 cm = 1 in.

• Mass: The quantity of material in an object. SI unit: kilogram (kg); Metric unit: gram (g). Equalities: 1 kg = 1000 g; 1 kg = 2.20 lb; 454 g = 1 lb.

• Temperature: Indicates how hot or cold something is. Metric unit: Celsius (°C); SI unit: Kelvin (K). Water freezes: 0°C; Kelvin scale: starts at 0 K.

• Time: Measured in seconds (s), minutes (min), hours (h), days, years (yr). SI and metric unit: second (s).

Measured Numbers and Significant Figures

Measured numbers are obtained by measuring physical quantities. The last digit is always an estimate, reflecting measurement uncertainty. Significant figures (SFs) indicate the precision of a measurement.

• Significant Figures: All nonzero digits, zeros between digits, and zeros at the end of a decimal number are significant. Zeros before digits are not significant.

• Scientific Notation: Used to clarify which zeros are significant in large numbers.

• Exact Numbers: Obtained by counting or definition (e.g., 1 kg = 1000 g). Exact numbers have unlimited significant figures and are not used to determine SFs in calculations.

Significant Figures in Calculations

When performing calculations, the number of significant figures in the result depends on the operation:

• Multiplication/Division: The answer has the same number of SFs as the measurement with the fewest SFs.

• Addition/Subtraction: The answer has the same number of decimal places as the measurement with the fewest decimal places.

• Rounding Off: If the first digit dropped is 5 or greater, increase the last retained digit by 1; if 4 or less, simply drop the digits.

Prefixes and Equalities

Metric prefixes modify the size of units by powers of ten, allowing for easy conversion between units.

• Prefixes that increase unit size: kilo (k, 1000), mega (M, 1,000,000)

• Prefixes that decrease unit size: milli (m, 0.001), micro (mc, 0.000001), centi (c, 0.01)

• Equalities: Express relationships between units (e.g., 1 km = 1000 m; 1 cm = 0.01 m)

Writing Conversion Factors

Conversion factors are fractions derived from equalities, used to convert between units. They are essential for solving chemistry problems involving unit changes.

• Metric conversion factors: Derived from exact equalities (e.g., $\frac{1\text{ m}}{100\text{ cm}}$ or $\frac{100\text{ cm}}{1\text{ m}}$)

• Metric–U.S. system conversion factors: May involve measured numbers (e.g., $\frac{1\text{ in}}{2.54\text{ cm}}$)

• Percentage, ppm, ppb: Used for very small ratios (e.g., 18% body fat means 18 g per 100 g total mass)

Problem Solving Using Unit Conversion

Unit conversion is a systematic process involving identification of given and needed units, selection of appropriate conversion factors, and calculation.

1. State the given and needed quantities.

2. Write a plan to convert the given unit to the needed unit.

3. State the equalities and conversion factors.

4. Set up the problem to cancel units and calculate the answer.

Example: To convert 2.44 m to centimeters: $2.44\text{ m} \times \frac{100\text{ cm}}{1\text{ m}} = 244\text{ cm}$

Density

Density is a physical property that compares the mass of an object to its volume. It is used to identify substances and solve problems involving mass and volume.

• Formula: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Units: g/mL, g/cm3

• Volume displacement: Used to determine the volume of irregular solids by measuring the change in water level.

• Specific gravity: Ratio of the density of a substance to the density of water (unitless).

Example: If a 48.0-g sample raises water from 25.0 mL to 33.0 mL, the volume displaced is 8.0 mL. Density = $\frac{48.0\text{ g}}{8.0\text{ mL}} = 6.0\text{ g/mL}$

Chemistry Link to Health: Toxicology

Toxicology studies the effects of chemicals on living organisms. The lethal dose (LD50) is the concentration that causes death in 50% of test animals, measured in mg/kg or mcg/kg. Comparing LD50 values helps assess toxicity.

• Example: Parathion (pesticide) has an LD50 of 3 mg/kg, indicating high toxicity.

Summary Table: Common Units and Prefixes

Quantity	SI Unit	Metric Unit	Common Prefixes
Length	meter (m)	meter (m), centimeter (cm)	kilo (km), centi (cm), milli (mm)
Mass	kilogram (kg)	gram (g)	kilo (kg), milli (mg), micro (mcg)
Volume	cubic meter (m3)	liter (L), milliliter (mL)	kilo (kL), milli (mL), micro (mcL)
Temperature	kelvin (K)	Celsius (°C)	—
Time	second (s)	second (s)	—

Summary Table: Significant Figures Rules

Rule	Example
Nonzero digits are always significant	123 (3 SFs)
Zeros between nonzero digits are significant	101 (3 SFs)
Zeros at the end of a decimal number are significant	2.50 (3 SFs)
Zeros before digits are not significant	0.0025 (2 SFs)

Summary Table: Common Conversion Factors

Equality	Conversion Factor 1	Conversion Factor 2
1 m = 100 cm	$\frac{1\text{ m}}{100\text{ cm}}$	$\frac{100\text{ cm}}{1\text{ m}}$
1 L = 1000 mL	$\frac{1\text{ L}}{1000\text{ mL}}$	$\frac{1000\text{ mL}}{1\text{ L}}$
1 kg = 1000 g	$\frac{1\text{ kg}}{1000\text{ g}}$	$\frac{1000\text{ g}}{1\text{ kg}}$
1 in = 2.54 cm	$\frac{1\text{ in}}{2.54\text{ cm}}$	$\frac{2.54\text{ cm}}{1\text{ in}}$