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

Classifying Matter

Pure Substances and Mixtures

Matter is anything that occupies space and has mass. It can be classified into two main categories: pure substances and mixtures.

• Pure substances are further classified as elements or compounds.

• Mixtures are classified as homogeneous or heterogeneous.

Definitions

• Element: The simplest type of matter, made up of only one kind of atom.

• Atom: The smallest unit of matter that retains its unique properties.

• Compound: A pure substance composed of two or more elements chemically bonded together.

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

• Homogeneous mixture: Has uniform composition throughout (e.g., salt water).

• Heterogeneous mixture: Composition varies throughout (e.g., salad).

Classification Table

Type	Examples	Characteristics
Element	Copper (Cu)	One type of atom
Compound	Water (H2O)	Two or more elements chemically bonded
Homogeneous Mixture	Brass (copper and zinc)	Uniform composition
Heterogeneous Mixture	Water and copper atoms	Non-uniform composition

Elements, Compounds, and the Periodic Table

The Periodic Table

The periodic table of the elements organizes all known elements based on their properties.

• Each element is represented by a chemical symbol (e.g., H for hydrogen, Na for sodium).

• Symbols may be derived from English or Latin names (e.g., Au for gold from aurum).

Organization

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

• Periods: Horizontal rows, numbered 1–7.

• Metals, Nonmetals, Metalloids: Separated by a staircase line; elements bordering the line (except Al) are metalloids.

Essential Elements for Life

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

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

Chemical Formulas

• Show the elements and the number of atoms in a compound (e.g., $\mathrm{H_2O}$ for water).

How Matter Changes

Physical and Chemical Changes

Matter can undergo changes that are either physical or chemical.

• Physical change: Alters the form or appearance but not the identity (e.g., melting ice).

• Chemical change: Alters the chemical identity; a chemical reaction occurs (e.g., burning wood).

Chemical Equations

• Represent chemical reactions using formulas and symbols.

• Reactants are transformed into products.

• Physical states are indicated: (s)olid, (l)iquid, (g)as, (aq)ueous.

Balancing Chemical Equations

• Law of Conservation of Mass: Matter is neither created nor destroyed.

• Number of atoms of each element must be equal on both sides.

• Add coefficients to balance equations.

Steps to Balance Equations

1. Examine the equation for balance.

2. Balance one element at a time using coefficients.

3. Check for the smallest set of coefficients.

Math Counts

Measurement Systems

Chemistry uses the metric system, specifically the Système International d’Unités (SI).

• Mass: kilogram (kg)

• Volume: liter (L)

• Length: meter (m)

• Prefixes change unit size by powers of 10 (e.g., milli-, centi-, kilo-).

Unit Conversion and Dimensional Analysis

• Use conversion factors to change units (e.g., $1\ \mathrm{dL} = 0.1\ \mathrm{L}$).

• Dimensional analysis involves setting up problems so units cancel, leaving the desired unit.

Significant Figures

• Digits known with certainty plus one estimated digit.

• All nonzero digits are significant; zeros may or may not be, depending on position.

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

Rules for Calculations

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

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

Rounding

• If the digit to be dropped is 4 or less, remove it; if 5 or greater, increase the last retained digit by 1.

• Round only at the end of multi-step calculations.

Scientific Notation

• General form: $C \times 10^n$

• $C$ is the coefficient (1 ≤ $C$ < 10), $n$ is the exponent.

• Positive exponent: number > 1; negative exponent: number between 0 and 1.

Percent Calculations

• Percent (%) = (part/whole) × 100

• Convert fractions or decimals to percent by multiplying by 100.

Matter: The “Stuff” of Chemistry

Mass and Volume

Mass is the amount of material in an object, measured in grams (g). Volume is the space occupied, measured in liters (L) or milliliters (mL).

• 1 mL = 1 cm3

• 5 mL ≈ 1 teaspoon

Density and Specific Gravity

Density is the ratio of mass to volume:

$d = \frac{m}{V}$

• Density of water at 4°C is 1.00 g/mL.

• Specific gravity is the ratio of the density of a sample to the density of water:

$\text{Specific gravity} = \frac{\text{density of sample}}{\text{density of water}}$

• Specific gravity is unitless.

Temperature Scales

• Fahrenheit (°F): Used in the US.

• Celsius (°C): Used worldwide.

• Kelvin (K): SI unit for temperature.

Conversion formulas:

$\text{Kelvin} = \text{Celsius} + 273$

$\text{Celsius} = (\text{Fahrenheit} - 32) \times \frac{5}{9}$

$\text{Fahrenheit} = (\text{Celsius} \times \frac{9}{5}) + 32$

Body Temperature

• Normal: 98.6°F (37.0°C)

• Hyperthermia: >40.0°C (104°F)

• Hypothermia: <35°C (95°F)

Energy and Specific Heat

• Energy: Capacity to do work or supply heat.

• Potential energy: Stored energy.

• Kinetic energy: Energy of motion.

• Law of conservation of energy: Energy is not created or destroyed.

• SI unit: joule (J); 1 calorie = 4.184 J; 1 Calorie (nutritional) = 1000 calories.

Specific Heat

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

$\text{Specific heat (SH)} = \frac{\text{calories}}{\text{grams} \times \Delta T}$

• Water has a high specific heat; metals have low specific heat.

Specific Heat Table (Selected Substances)

Substance	Specific Heat (cal/g·°C)
Water (liquid)	1.00
Iron	0.11
Copper	0.09
Glass	0.20
Aluminum	0.22
Lead	0.03

States of Matter

Matter exists in three main physical states: solid, liquid, and gas.

• Solid: Definite shape and volume; particles are tightly packed and vibrate in place.

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

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

Comparison Table: States of Matter

Property	Solid	Liquid	Gas
Shape	Definite	Adopts container	Adopts container
Volume	Definite	Definite	Fills container
Particle arrangement	Orderly, tightly packed	Less orderly, close	Disordered, far apart
Energy	Lowest	Intermediate	Highest

Measuring Matter

Accuracy and Precision

• Accuracy: How close a measurement is to the true value.

• Precision: How close repeated measurements are to each other.

• Best practice: Take several measurements and average them.

Units and Health Applications

• SI/metric units are standard, but U.S. customary units are also used.

• Common conversions: 1 dL = 100 mL; 1 mmol = 0.001 mol.

• Body weight: measured in pounds (US) or kilograms (pharmaceuticals).

Dosage Calculations

1. Determine units for the final answer.

2. Identify given information.

3. Choose conversion factors to cancel units.

4. Set up the equation so only the desired unit remains.

Drop Units and Percent in Health

• gtt/mL: Drops per milliliter; used for IV drip rates.

• Drop factor depends on IV tubing diameter.

• Percent active ingredient: Used to adjust medication dosages.

• Percent daily value (%DV): Used in nutrition labeling.