Chapter 1: Chemistry Basics—Matter and Measurement (Study Guide)

Chemistry Basics—Matter and Measurement

1.1 Classifying Matter: Pure Substance or Mixture

This section introduces the fundamental classification of matter, which is essential for understanding chemical processes and properties.

• Matter: Anything that occupies space and has mass.

• Pure Substances: Composed of only one type of substance; can be represented by a single chemical formula or symbol.

• Elements: Simplest form of matter, made up of only one type of atom.

• Compounds: Pure substances formed from two or more elements chemically joined.

• Mixtures: Combination of two or more substances; can be separated into components.

• Homogeneous Mixture: Uniform composition throughout (e.g., salt water).

• Heterogeneous Mixture: Non-uniform composition (e.g., salad).

• Atom: Smallest unit of matter retaining unique properties.

• Example: Air is a mixture; water (H2O) is a compound.

1.2 Elements, Compounds, and the Periodic Table

The periodic table organizes all known elements and provides information about their properties and relationships.

• Periodic Table: A chart listing all elements, each represented by a chemical symbol.

• Groups: Vertical columns; elements in a group share similar chemical behaviors.

• Periods: Horizontal rows; numbered 1 to 7.

• Metals, Nonmetals, Metalloids: Metals are to the left of the staircase line, nonmetals to the right, metalloids border the line (except Al).

• Macronutrients: Elements needed in quantities >100 mg/day (e.g., Na, Mg, K, Ca, Cl).

• Micronutrients: Elements needed in quantities <100 mg/day (e.g., I, F, Fe, Zn).

• Compounds: Pure substances with two or more elements in fixed ratios; represented by chemical formulas (e.g., NaCl).

• Example: Water (H2O) contains two hydrogen atoms and one oxygen atom.

1.3 How Matter Changes

Matter can undergo physical or chemical changes, which are fundamental to chemical reactions and processes.

• Physical Change: Alters the form or state of matter without changing its identity (e.g., melting ice).

• Chemical Change: Changes the chemical identity; involves chemical reactions (e.g., burning wood).

• Chemical Equation: Represents reactants and products; includes physical states (s, l, g, aq).

• Balancing Equations: Ensures the number of atoms is equal on both sides, illustrating the law of conservation of mass.

• Steps to Balance:

  1. Examine the equation for balance.

  2. Add coefficients to balance one element at a time.

  3. Check for smallest possible coefficients.

• Example: $\mathrm{C} + \mathrm{O}_2 \rightarrow \mathrm{CO}_2$

1.4 Math Counts

Mathematical concepts are central to chemistry, including measurement systems, unit conversions, and significant figures.

• SI Units: Standard units for mass (kg), volume (L), length (m).

• Prefixes: Modify units by powers of 10 (e.g., milli-, centi-, kilo-).

• Conversion Factors: Used to convert between units (e.g., 1 dL = 0.1 L).

• Dimensional Analysis: Method for converting units using conversion factors.

• Significant Figures: Digits known with certainty plus one estimated digit; important for reporting measurements.

• Rules for Significant Figures:

  • Addition/Subtraction: Match least decimal places.

  • Multiplication/Division: Match least significant digits.

• Scientific Notation: $C \times 10^n$ where $1 \leq C < 10$.

• Percent: Part per hundred; convert by multiplying by 100.

• Example: 0.25 = 25%

1.5 Matter: The “Stuff” of Chemistry

Understanding mass, volume, density, and energy is crucial for measuring and interpreting chemical phenomena.

• Mass: Amount of material; measured in grams (g).

• Volume: Space occupied; measured in milliliters (mL) or cubic centimeters (cm3).

• Density: Ratio of mass to volume; $d = \frac{m}{V}$

• Specific Gravity: Ratio of sample density to water density; unitless.

• Temperature: Measured in Celsius (°C), Fahrenheit (°F), Kelvin (K).

• Energy: Capacity to do work; measured in joules (J) or calories (cal).

• Specific Heat: Amount of heat to raise 1 g of substance by 1°C.

• States of Matter: Solid (definite shape/volume), liquid (definite volume), gas (no definite shape/volume).

• Example: Water at room temperature is a liquid; ice is a solid.

1.6 Measuring Matter

Measurement accuracy and precision are vital in health and laboratory settings, along with proper unit conversions and dosage calculations.

• Accuracy: Closeness to true value.

• Precision: Consistency of repeated measurements.

• Units: SI/metric and U.S. customary units used in health care.

• Lab Reports: Patient results compared to normal values; units vary.

• Dosage Calculations: Use conversion factors to ensure correct units.

• Drop Units: IV medications measured in drops per mL (gtt).

• Percent in Health: Used for active ingredients, dosing, and nutrition labeling.

• Example: Calculating pediatric dose as a percent of adult dose.

Relevant Image

The following image visually represents the textbook cover for General, Organic, and Biological Chemistry, reinforcing the context of the study material: