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

Chapter 1: Chemistry Basics—Matter and Measurement

1.1 Classifying Matter: Pure Substance or Mixture

Matter is defined as anything that occupies space and has mass. Understanding the classification of matter is foundational in chemistry, as it determines how substances interact and change.

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

• Element: The simplest form of matter, consisting of only one type of atom. Atoms are the smallest units retaining unique properties.

• Compound: A pure substance formed by the chemical combination of two or more elements in fixed ratios.

• Mixture: A physical combination of two or more substances that can be separated by physical means.

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

• Heterogeneous Mixture: Non-uniform composition; components are visibly distinct (e.g., salad).

1.2 Elements, Compounds, and the Periodic Table

The periodic table organizes all known elements based on their properties and atomic structure. It is a critical tool for understanding chemical behavior and relationships.

• Chemical Symbol: One or two-letter abbreviation for an element (e.g., H for hydrogen, Na for sodium from Latin 'natrium').

• Groups: Vertical columns; elements in the same group share similar chemical properties. Numbered 1–18 (IUPAC system).

• Periods: Horizontal rows; numbered 1–7. Elements in the same period have the same number of electron shells.

• Metals, Nonmetals, Metalloids: The staircase line separates metals (left), nonmetals (right), and metalloids (bordering the line except Al).

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

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

• Chemical Formula: Indicates the elements and the number of atoms in a compound (e.g., $\mathrm{H_2O}$ for water, $\mathrm{NaCl}$ for table salt).

1.3 How Matter Changes

Matter can undergo physical or chemical changes. Understanding these changes is essential for describing chemical reactions and processes.

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

• Chemical Change: Alters the chemical identity; new substances are formed (chemical reaction).

• Chemical Equation: Represents a chemical reaction. Reactants are transformed into products. States are indicated as (s), (l), (g), or (aq).

• Balancing Equations: The number of atoms for each element must be equal on both sides, reflecting the law of conservation of mass.

Steps to Balance a Chemical Equation:

1. Examine the equation for balance.

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

3. Check that all elements are balanced with the smallest set of coefficients.

1.4 Math Counts

Mathematical concepts are integral to chemistry, especially for measurement, conversions, and calculations.

• SI Units: Standard units include kilogram (kg) for mass, liter (L) for volume, and meter (m) for length.

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

• Conversion Factors: Ratios used to convert between units (e.g., $1\ \mathrm{dL} = 0.1\ \mathrm{L}$).

• Dimensional Analysis: Systematic approach to unit conversion using conversion factors.

Significant Figures:

• All nonzero digits are significant.

• Zeros are significant only in certain positions (e.g., between nonzero digits, after a decimal point).

• Exact numbers (from counting or definitions) have infinite significant figures.

Rules for Calculations:

• Addition/Subtraction: Result matches the least number of decimal places.

• Multiplication/Division: Result matches the least number of significant digits.

Scientific Notation: Expresses numbers as $C \times 10^n$, where $1 \leq C < 10$ and $n$ is an integer.

Percentages: Percent (%) means per hundred. Convert fractions or decimals to percent by multiplying by 100.

1.5 Matter: The “Stuff” of Chemistry

Understanding the properties and measurements of matter is crucial for laboratory and clinical applications.

• Mass: Amount of material in an object, measured in grams (g).

• Volume: Space occupied by matter, measured in liters (L) or milliliters (mL). $1\ \mathrm{mL} = 1\ \mathrm{cm^3}$.

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

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

• Temperature: Measured in Celsius (°C), Kelvin (K), or Fahrenheit (°F). $K = ^\circ C + 273$; $^\circ F = 1.8 \times ^\circ C + 32$

• Energy: Capacity to do work or supply heat. Measured in joules (J) or calories (cal). $1\ \mathrm{Calorie} = 1000\ \mathrm{calories}$

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

States of Matter:

• Solid: Definite shape and volume; particles tightly packed.

• Liquid: Definite volume, takes shape of container; particles move more freely.

• Gas: No definite shape or volume; particles far apart and move rapidly.

1.6 Measuring Matter

Accurate and precise measurements are essential in chemistry and health sciences. Understanding units and conversions ensures correct dosing and interpretation of laboratory results.

• Accuracy: Closeness of a measurement to the true value.

• Precision: Consistency of repeated measurements.

• SI/Metric vs. U.S. Units: Health professionals must be familiar with both systems.

• Lab Reports: Use metric units; results compared to normal ranges.

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

• Drop Units (gtt): Used in IV medication delivery; drop factor depends on tubing diameter.

• Percents in Health: Used for active ingredient concentration, dosing adjustments, and nutrition labeling.

Key Equations and Relationships

• Density: $d = \frac{m}{V}$

• Temperature Conversions:

  • $K = ^\circ C + 273$

  • $^\circ F = 1.8 \times ^\circ C + 32$

  • $^\circ C = \frac{^\circ F - 32}{1.8}$

• Specific Heat: $q = m \times c \times \Delta T$ where $q$ = heat (J), $m$ = mass (g), $c$ = specific heat (J/g·°C), $\Delta T$ = temperature change (°C)

• Percent: $\text{Percent} = \frac{\text{Part}}{\text{Whole}} \times 100$

Summary Table: Classification of Matter

Type	Description	Examples
Element	One type of atom	Oxygen (O2), Gold (Au)
Compound	Two or more elements chemically combined	Water (H2O), Sodium chloride (NaCl)
Homogeneous Mixture	Uniform composition throughout	Saltwater, Air
Heterogeneous Mixture	Non-uniform composition	Salad, Granite

Summary Table: SI Units and Conversions

Quantity	SI Unit	Common Conversion
Mass	kilogram (kg), gram (g)	1 kg = 1000 g
Volume	liter (L), milliliter (mL)	1 L = 1000 mL
Length	meter (m)	1 m = 100 cm
Energy	joule (J), calorie (cal)	1 cal = 4.184 J