Scientific Method

Overview of the Scientific Method

The scientific method is a systematic approach used by scientists to gather information, formulate explanations, and test hypotheses about natural phenomena.

• Observation: Collecting data through senses or instruments (empirical facts).

• Law (Scientific Law): A statement, often mathematical, summarizing repeated experimental observations. It describes what happens but not why.

• Hypothesis: A tentative explanation for observations, used to make predictions and guide further experiments.

• Theory: A well-tested explanation of how nature behaves, subject to modification with new evidence. Theories can never be proven absolutely correct.

Example: The study of gases led to the empirical law (Ideal Gas Law), where R is a constant. The kinetic theory of gases explains this law.

Atomic Theory and Classification of Matter

Atomic Theory

The atomic theory is a foundational model in chemistry, stating that all matter is composed of tiny, indivisible particles called atoms. Atoms are the smallest units retaining the properties of an element.

Matter and Its Properties

• Matter: Anything that has mass and occupies space.

• Mass: The amount of matter in an object; a measure of inertia.

• Weight: The force of gravity acting on an object's mass. Weight varies with location (e.g., Earth vs. Moon), but mass remains constant.

Elements, Compounds, and Chemical Symbols

• Element: A substance that cannot be decomposed into simpler substances by chemical means. Composed of only one type of atom.

• Compound: A substance formed from two or more elements combined in fixed ratios by mass. Can be decomposed into elements by chemical changes.

• Chemical Symbol: One- or two-letter abbreviation for an element (e.g., H for hydrogen, O for oxygen). Used in chemical formulas (e.g., H2O for water).

Pure Substances vs. Mixtures

• Pure Substance: Has a constant composition (elements or compounds).

• Mixture: Contains two or more substances with variable composition. Can be separated by physical means.

Physical and Chemical Changes

Physical Change

A physical change alters the form or state of a substance without changing its chemical composition (e.g., melting ice, dissolving sugar in water).

Chemical Change (Chemical Reaction)

A chemical change results in the formation of one or more new substances with different properties (e.g., burning paper, rusting iron).

Physical vs. Chemical Properties

• Physical Properties: Observed without changing the substance's identity (e.g., color, melting point, density).

• Chemical Properties: Describe how a substance reacts to form new substances (e.g., flammability, reactivity with acid).

Intensive vs. Extensive Properties

• Intensive Properties: Independent of sample size (e.g., density, boiling point).

• Extensive Properties: Dependent on sample size (e.g., mass, volume).

Measurement and SI Units

Observations and Measurements

• Quantitative Observations: Involve numbers and units (e.g., 25.0 g).

• Qualitative Observations: Descriptive, non-numeric (e.g., color, appearance).

All measurements must include a unit and are subject to uncertainty.

International System of Units (SI)

The SI system is the standard for scientific measurements. The seven base units are:

Derived units are formed from base units (e.g., area in m2, volume in m3).

Common Non-SI Units and Conversions

• Length: 1 cm = 0.01 m; 1 mm = 0.001 m

• Volume: 1 L = 1000 mL = 1000 cm3

• Mass: 1 kg = 1000 g

SI Prefixes

Laboratory Measurements

• Length: SI unit is meter (m); commonly use cm or mm in lab.

• Volume: SI unit is m3; commonly use liter (L) or milliliter (mL).

• Mass: SI unit is kilogram (kg); commonly use gram (g).

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

Temperature Scales and Conversions

• Celsius (°C): Water freezes at 0°C, boils at 100°C.

• Kelvin (K): SI unit; absolute zero is 0 K. Water freezes at 273.15 K, boils at 373.15 K.

• Fahrenheit (°F): Water freezes at 32°F, boils at 212°F.

Conversion formulas:

Accuracy, Precision, and Significant Figures

Accuracy and Precision

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

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

Significant Figures

Significant figures reflect the precision of a measurement. Rules for counting significant figures:

• All nonzero digits are significant.

• Zeros between nonzero digits are significant.

• Leading zeros are not significant.

• Trailing zeros are significant if there is a decimal point.

Examples:

• 19.0000 has 6 significant figures.

• 0.0005650850 has 7 significant figures.

Rounding Rules

• If the digit to be dropped is greater than 5, round up.

• If less than 5, leave the last digit unchanged.

• If exactly 5, round to the nearest even digit.

Scientific Notation

Used to express very large or small numbers and clarify significant figures. Format: where .

Significant Figures in Calculations

• Multiplication/Division: Result has as many significant figures as the measurement with the fewest significant figures.

• Addition/Subtraction: Result has as many decimal places as the measurement with the fewest decimal places.

Exact numbers (from definitions or counting) have infinite significant figures and do not limit the precision of calculations.

Dimensional Analysis (Factor Label Method)

Using Conversion Factors

Dimensional analysis uses conversion factors to change units. A conversion factor is a fraction equal to one, derived from an equality between units.

• General formula:

• Units cancel algebraically, just like numbers.

Example: Convert 68.0 in to cm using (exact).

Density and Specific Gravity

Density

Density () is the ratio of mass to volume. It is an intensive property, independent of sample size.

• Formula:

• Common units: g/mL (liquids/solids), g/L (gases)

Density can be used to convert between mass and volume.

Specific Gravity

Specific gravity is the ratio of the density of a substance to the density of water (unitless):

• Formula:

Temperature Effects on Density

• Most substances expand slightly when heated, decreasing density.

• For solids and liquids, this effect is usually small.

Summary Table: Classification of Matter

Key Equations

• Ideal Gas Law:

• Density:

• Temperature Conversions:

Practice and Application

• Classify matter as element, compound, homogeneous mixture, or heterogeneous mixture.

• Identify physical vs. chemical changes and properties.

• Use dimensional analysis for unit conversions.

• Apply significant figure rules in calculations.

• Calculate density, mass, or volume given two of the three variables.

Additional info: This summary integrates and expands upon the main points from the provided lecture slides, adding definitions, examples, and formulas for clarity and completeness.