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

Chemistry Basics—Matter and Measurement

1.1 Classifying Matter: Pure Substance or Mixture

Chemistry begins with understanding matter, which is anything that occupies space and has mass. Matter can be classified based on its composition and uniformity.

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

• Element: Simplest form of matter, made of one type of atom (e.g., Calcium, Hydrogen, Carbon).

• Compound: Pure substance made of two or more elements chemically combined in fixed ratios (e.g., Water (H2O), Table Salt (NaCl)).

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

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

• Heterogeneous Mixture: Non-uniform composition (e.g., vegetable soup, chocolate chip cookie).

Example: Gold is an element (pure substance), coffee is a homogeneous mixture, table salt is a compound (pure substance), and vegetable soup is a heterogeneous mixture.

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 foundational tool in chemistry for understanding elements and their relationships.

• Element Symbols: Each element is represented by a one- or two-letter symbol, often derived from English or Latin names (e.g., Na for sodium, Au for gold).

• Groups: Vertical columns with similar chemical properties; labeled 1–18 (IUPAC system).

• Periods: Horizontal rows numbered 1–7.

• Metals, Nonmetals, Metalloids: The periodic table is divided by a staircase line; metals are to the left, nonmetals to the right, and metalloids border the line (except Al).

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

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

Compounds: Pure substances containing two or more elements in fixed ratios, represented by chemical formulas (e.g., H2O, NaCl).

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 ice).

• Chemical Change: Alters the chemical identity; new substances are formed (e.g., burning charcoal).

• Chemical Equation: Represents a chemical reaction, showing reactants, products, and their physical states.

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

Steps to Balance Equations:

1. Examine the equation for balance.

2. Balance one element at a time using coefficients.

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

Example: $\mathrm{CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O}$

1.4 Math Counts

Chemistry relies on precise measurements and mathematical conversions. The SI (metric) system is used for consistency.

• SI Units: Mass (kilogram, kg), Volume (liter, L), Length (meter, m).

• Prefixes: Used to indicate powers of ten (e.g., milli-, centi-, kilo-).

• Conversion Factors: Ratios used to convert between units (e.g., $1\,\mathrm{mg} = 1 \times 10^{-3}\,\mathrm{g}$).

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

• Significant Figures: Digits in a measurement that are known with certainty plus one estimated digit; rules determine which digits are significant.

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

• Percentages: $\text{Percent} = \frac{\text{part}}{\text{whole}} \times 100$

Example: Convert 10,000 cg to g: $10,000\,\mathrm{cg} \times \frac{1 \times 10^{-2}\,\mathrm{g}}{1\,\mathrm{cg}} = 100\,\mathrm{g}$

Example: 0.0055 has 2 significant figures; 1.030 has 4 significant figures.

1.5 Matter: The “Stuff” of Chemistry

Understanding the properties of matter is essential for measurement and application in chemistry and health sciences.

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

• Volume: Space occupied by matter, measured in liters (L) or milliliters (mL); 1 mL = 1 cm3.

• Density: Ratio of mass to volume, $d = \frac{m}{V}$, typically in g/mL.

• 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); conversions are important in science.

• States of Matter: Solid, liquid, and gas, each with distinct particle arrangements and properties.

Example: Calculate density: $d = \frac{1100\,\mathrm{g}}{85\,\mathrm{mL}} = 13\,\mathrm{g/mL}$

Example: Convert 37°C to Kelvin: $K = ^\circ\!C + 273 = 310\,\mathrm{K}$

Property	Solid	Liquid	Gas
Shape	Definite	Adopts container	Adopts container
Volume	Definite	Definite	Fills container
Kinetic Energy	Lowest	Intermediate	Highest
Particle Arrangement	Closely packed	Loosely packed	Far apart
Attraction	Very strong	Strong	Weak

1.6 Measuring Matter

Accurate and precise measurements are crucial in chemistry and health care. Understanding unit conversions and calculations ensures correct dosing and interpretation of data.

• Accuracy: Closeness to the true value.

• Precision: Consistency among repeated measurements.

• SI vs. U.S. Customary Units: Health professionals must convert between systems (e.g., 1 inch = 2.54 cm).

• Dosing Calculations: Use dimensional analysis to determine correct medication dosages.

• Drop Units (gtt): Used in IV medication delivery; conversion factors are essential for calculating drip rates.

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

Property	U.S. Customary Unit	SI/Metric Equivalent	U.S. Customary Equivalent
Mass	Pound (lb)	2.205 lb = 1 kg	1 lb = 16 oz
Volume	Quart (qt)	1.057 qt = 1 L	1 qt = 4 cups
Volume	Fluid ounce (fl oz)	1 fl oz = 29.6 mL	1 cup = 8 fl oz
Volume	Teaspoon (tsp)	1 tsp = 4.93 mL	1 fl oz = 6 tsp
Length	Mile (mi)	1 mi = 1.6 km	1 mi = 5280 ft
Length	Inch (in.)	1 in. = 2.54 cm	1 ft = 12 in.

Example: Calculate acetaminophen dose for a 42 lb child at 7.5 mg/kg: $7.5\,\mathrm{mg/kg} \times \frac{1\,\mathrm{kg}}{2.2\,\mathrm{lb}} \times 42\,\mathrm{lb} = 140\,\mathrm{mg}$

Example: Calculate IV drip time: $1.0\,\mathrm{L} \times \frac{1000\,\mathrm{mL}}{1\,\mathrm{L}} \times \frac{60\,\mathrm{gtt}}{1\,\mathrm{mL}} \times \frac{1\,\mathrm{min}}{100\,\mathrm{gtt}} \times \frac{1\,\mathrm{h}}{60\,\mathrm{min}} = 10\,\mathrm{h}$