The Chemical World

What is Chemistry?

Chemistry is the science that seeks to understand how matter behaves by studying what atoms and molecules do. The properties of materials are determined by their constituent atoms and molecules, and changing the arrangement or type of these particles changes the properties of the material.

• Matter: Any substance that has mass and volume; everything we touch or feel.

• Atoms: The fundamental building blocks of matter, extremely small (about 2 nanometers in size).

• Molecules: Groups of atoms bonded together; their characteristics depend on the types and arrangements of atoms.

• Chemical bonds: Attachments that hold atoms together in molecules.

• Example: Both graphite and diamond are made of carbon, but their different atomic arrangements give them distinct properties.

Chemicals in Everyday Life

Everything around us is composed of chemicals, not just dangerous substances. Air, water, and DNA are all chemicals.

• Common misconceptions associate chemicals only with pollutants or poisons (e.g., lead in water, carbon monoxide).

• In reality, chemicals make up all matter, including safe and essential substances.

Measurement and Problem Solving

The Scientific Method

The scientific method is a systematic approach to understanding why things happen, relying on observation and repeated experimentation.

• Observation: Measurement or description of some aspect of the physical world.

• Hypothesis: A tentative, falsifiable explanation for an observation.

• Experiment: Controlled test to validate or invalidate a hypothesis.

• Scientific Law: Summarizes the results of many observations and predicts future ones (e.g., Law of Conservation of Mass).

• Scientific Theory: Well-supported explanation for phenomena, based on multiple hypotheses and observations (e.g., Atomic Theory).

Example: Lavoisier's Law of Conservation of Mass states that matter is not created or destroyed in a chemical reaction.

Analyzing and Interpreting Data

• Quantitative measurements: Include numbers (e.g., mass, volume), often presented in tables, graphs, or charts.

• Qualitative measurements: Descriptions without numbers (e.g., color, texture).

• Data analysis involves looking for patterns or trends and drawing conclusions.

Scientific Notation

Scientific notation is used to express very large or small numbers conveniently.

• Format: Decimal part (between 1 and 10) × 10n (exponential part).

• Large numbers: Positive exponents (move decimal left).

• Small numbers: Negative exponents (move decimal right).

• Example:

Significant Figures and Uncertainty

Significant figures (sig figs) reflect the precision of a measurement. The last reported digit is uncertain.

• All nonzero digits are significant.

• Interior zeros (between nonzero digits) are significant.

• Trailing zeros after a decimal point are significant.

• Leading zeros are not significant; they only locate the decimal point.

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

Rules for Rounding:

• Round down if the last digit dropped is 4 or less; round up if 5 or more.

• In multi-step calculations, round only the final answer.

Significant Figures in Calculations:

• Multiplication/Division: Result has the same number of sig figs as the factor with the fewest sig figs.

• Addition/Subtraction: Result has the same number of decimal places as the value with the fewest decimal places.

Units of Measurement

The International System of Units (SI) is based on the metric system. Prefix multipliers (e.g., kilo-, milli-, micro-) are used to express different scales.

• Example: 1 kilogram (kg) = 1,000 grams (g)

Unit Conversions and Dimensional Analysis

Unit conversions use conversion factors (fractions with equal quantities in different units) to change from one unit to another. Dimensional analysis treats units algebraically, ensuring consistency throughout calculations.

• Always write numbers with their units.

• Use a solution map to outline the steps needed for complex conversions.

• For multistep conversions, each step should cancel the previous unit and introduce the next.

Density as a Conversion Factor

Density is the ratio of mass to volume and can be used to convert between mass and volume.

• Formula:

• Example: A sample with mass 260.5 g and volume 250.0 cm3 has density .

Matter and Energy

What is Matter?

Matter is anything that has mass and occupies space. It is composed of atoms and molecules. Not all things we encounter are matter (e.g., light, heat, sound are forms of energy).

States of Matter

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

• Solid: Fixed volume and rigid shape; particles vibrate but do not move past each other. Can be crystalline (ordered) or amorphous (disordered).

• Liquid: Fixed volume but takes the shape of its container; particles are close but can move past each other.

• Gas: No fixed volume or shape; particles are far apart and move freely. Gases are compressible.

Classifying Matter by Composition

• Pure Substances: Composed of only one type of atom or molecule.

  • Element: Cannot be broken down by chemical means (e.g., oxygen, gold).

  • Compound: Composed of two or more elements in fixed proportions (e.g., water, table salt).

• Mixtures: Composed of two or more substances combined in variable proportions.

  • Heterogeneous mixture: Composition varies from one region to another (e.g., oil and water).

  • Homogeneous mixture: Uniform composition throughout (e.g., saltwater, sweetened tea).

Physical and Chemical Properties

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

• Chemical property: Observed only by changing the substance's composition (e.g., flammability, acidity, reactivity).

Physical and Chemical Changes

• Physical change: Alters appearance but not composition (e.g., melting, boiling, cutting).

• Chemical change: Alters composition, resulting in new substances (e.g., rusting, combustion).

• Reactants: Substances present before a chemical change.

• Products: Substances formed as a result of a chemical change.

Law of Conservation of Mass

Mass is neither created nor destroyed in ordinary chemical reactions. The total mass of reactants equals the total mass of products.

• Example: Burning 58 g of butane with oxygen produces 176 g of carbon dioxide and 90 g of water. Total mass before and after is 266 g.

Energy

Energy is the capacity to do work. It is conserved and can be transformed from one form to another.

• Kinetic energy: Energy of motion.

• Potential energy: Energy due to position or composition.

• Thermal energy: Associated with the random motion of particles.

• Chemical energy: Stored in the composition and positions of particles.

• Electrical energy: Associated with the flow of electric charge.

• SI unit: Joule (J); 1 calorie (cal) = 4.184 J; 1 kilowatt-hour (kWh) = J.

Energy in Physical and Chemical Changes

• Exothermic reaction: Releases energy (e.g., combustion of TNT).

• Endothermic reaction: Absorbs energy.

• Systems with high potential energy tend to change to lower potential energy, often releasing energy in the process.

Temperature and Heat

• Temperature: Measure of the thermal energy of a substance.

• Heat: Transfer of thermal energy due to temperature difference.

• Temperature Scales:

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

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

  • Kelvin (K): Water freezes at 273 K, boils at 373 K; absolute zero is 0 K.

• Conversions between scales are possible using formulas.

Atoms and Elements

Atomic Theory

John Dalton's atomic theory forms the basis of modern chemistry:

• All matter is composed of tiny particles called atoms.

• Atoms are indivisible and indestructible (later modified).

• Atoms of a given element have the same properties, including mass.

• Atoms of different elements have different properties.

• Atoms combine in simple whole-number ratios to form compounds.

The Nuclear Atom

• J.J. Thomson: Discovered the electron, a negatively charged particle much smaller than the atom.

• Plum Pudding Model: Electrons embedded in a sphere of positive charge.

• Ernest Rutherford: Gold foil experiment showed that atoms have a small, dense, positively charged nucleus, with electrons dispersed around it.

• Nuclear Model: Most of the atom's mass and all positive charge are in the nucleus; electrons occupy the surrounding space.

Subatomic Particles

• Protons: Positively charged, found in the nucleus.

• Neutrons: Neutral, found in the nucleus.

• Electrons: Negatively charged, found outside the nucleus.

• Atomic mass unit (amu): Defined as 1/12 the mass of a carbon-12 atom.

• Atoms are electrically neutral because the number of protons equals the number of electrons.

Elements and the Periodic Table

• Atomic number (Z): Number of protons in the nucleus; defines the element.

• Periodic table arranges elements by increasing atomic number.

• Elements are identified by name, symbol, and atomic number.

• First 30 elements' names and symbols should be memorized.

Organization of the Periodic Table

• Dmitri Mendeleev: Arranged elements by increasing mass, leading to the periodic law (properties recur periodically).

• Elements are classified as metals, nonmetals, or metalloids.

• Groups (columns): Elements with similar properties (e.g., Group 1A: alkali metals).

• Periods (rows): Horizontal rows in the table.

Ions and Predicting Ion Charges

• Cations: Positively charged ions (loss of electrons).

• Anions: Negatively charged ions (gain of electrons).

• Ions form to achieve the same number of valence electrons as the nearest noble gas.

• Valence electrons are the outermost electrons, important for chemical bonding.

• Group number indicates the number of valence electrons for main group elements.

Isotopes

• Atoms of the same element with different numbers of neutrons are called isotopes.

• Mass number (A): Sum of protons and neutrons ().

• Isotopes have different mass numbers but the same atomic number.

Average Atomic Mass

The atomic mass of an element is a weighted average of the masses of its naturally occurring isotopes.

• Formula:

• Fractions are the percent natural abundances converted to decimals.