Introduction to Chemistry

What is Chemistry?

Chemistry is the scientific study of matter, its properties, composition, structure, and the changes it undergoes during chemical reactions. It bridges physics and biology and is fundamental to understanding the material world.

• Matter: Anything that has mass and occupies space.

• Chemical Change: A process where substances are transformed into different substances.

• Physical Change: A change in state or appearance without altering composition.

The Scientific Method

Steps and Terminology

The scientific method is a systematic approach to research and discovery in science.

• Observation: Gathering data about phenomena.

• Hypothesis: A tentative explanation for observations, testable by experiments.

• Theory: A well-substantiated explanation of some aspect of the natural world, based on a body of evidence.

• Law: A concise statement that summarizes observed behavior, often expressed mathematically.

• Experiment: A controlled procedure to test a hypothesis.

Matter and Its Classification

Types of Matter

• Pure Substance: Matter with a fixed composition and distinct properties (e.g., elements, compounds).

• Mixture: A combination of two or more substances where each retains its identity. Can be separated by physical means.

• Element: A substance that cannot be broken down into simpler substances by chemical means.

• Compound: A substance composed of two or more elements chemically combined in fixed proportions.

• Homogeneous Mixture (Solution): Uniform composition throughout (e.g., salt water).

• Heterogeneous Mixture: Non-uniform composition (e.g., salad, sand in water).

Example: Air is a homogeneous mixture; salad is a heterogeneous mixture.

Phases of Matter and Phase Changes

• Solid: Definite shape and volume.

• Liquid: Definite volume, takes shape of container.

• Gas: No definite shape or volume; expands to fill container.

• Phase Transitions: Melting, freezing, vaporization, condensation, sublimation, deposition.

Physical vs. Chemical Changes

• Physical Change: Alters physical properties without changing composition (e.g., melting ice).

• Chemical Change: Alters the chemical composition (e.g., rusting iron).

Separation of Mixtures

• Filtration: Separates solids from liquids.

• Distillation: Separates based on boiling points.

• Chromatography: Separates based on movement through a medium.

• Decanting: Pouring off a liquid to separate it from a solid.

Measurement and Uncertainty

Accuracy, Precision, and Significant Figures

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

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

• Uncertainty: The degree of doubt in a measurement, typically one digit beyond the smallest scale division.

• Significant Figures (Sig Figs): Digits in a measurement that are known with certainty plus one estimated digit.

Rules for Significant Figures:

• All nonzero digits are significant.

• Zeros between nonzero digits are significant.

• Leading zeros are not significant.

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

Reporting Sig Figs in Calculations:

• Multiplication/Division: Result has as many sig figs as the measurement with the fewest sig figs.

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

Exact Numbers

• Numbers from definitions or counting (e.g., 1 dozen = 12) have infinite significant figures and do not limit the number of sig figs in calculations.

Scientific Notation

• Used to express very large or small numbers.

• Format: where and is an integer.

• Converting between decimal and scientific notation is essential for clarity and correct sig figs.

SI Units and Metric Prefixes

• SI Base Units: Meter (m) for length, kilogram (kg) for mass, second (s) for time, kelvin (K) for temperature, mole (mol) for amount of substance, ampere (A) for electric current, candela (cd) for luminous intensity.

• Common Prefixes: kilo- (k, ), centi- (c, ), milli- (m, ), micro- (, ), nano- (n, ).

Conversions and Rounding

• Use conversion factors to switch between units (e.g., , exactly).

• Round up if the next digit is 5 or more; round down if less than 5.

Density

• Definition: Density is mass per unit volume.

• Formula:

• Units: g/cm3 or kg/L.

• Used to identify substances and for conversion between mass and volume.

Temperature Conversions

• Formulas for converting between Celsius, Fahrenheit, and Kelvin will be provided on the exam.

Atoms, Molecules, and Ions

Definition and Structure

• Atom: The smallest unit of an element that retains its chemical properties.

• Molecule: Two or more atoms chemically bonded together.

Fundamental Laws

• Law of Conservation of Mass: Mass is neither created nor destroyed in a chemical reaction.

• Law of Definite Proportions: A compound always contains the same elements in the same proportion by mass.

• Law of Multiple Proportions: When two elements form more than one compound, the ratios of the masses of the second element that combine with a fixed mass of the first element are ratios of small whole numbers.

Atomic Theory and Models

• Dalton's Atomic Theory:

  1. Elements are composed of tiny, indivisible particles called atoms.

  2. Atoms of the same element are identical; atoms of different elements are different.

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

  4. Chemical reactions involve rearrangement of atoms; atoms are not created or destroyed.

• Thomson's Model: Plum pudding model; electrons embedded in a positively charged sphere.

• Rutherford's Gold Foil Experiment: Discovered the nucleus; most of the atom is empty space, with a dense, positively charged nucleus.

Subatomic Particles

• Proton: Charge +1, relative mass ≈ 1 amu.

• Neutron: Charge 0, relative mass ≈ 1 amu.

• Electron: Charge -1, relative mass ≈ 1/1836 amu.

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

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

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

Atomic Notation

• General format: where X is the element symbol, A is the mass number, Z is the atomic number.

• Given atomic notation, you can determine the number of protons, neutrons, and electrons (for neutral atoms and ions).

Ions and Isotopes

• Ion: An atom or molecule with a net electric charge due to loss or gain of electrons.

• Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

• Isotope: Atoms of the same element (same Z) with different numbers of neutrons (different A).

• Average Atomic Mass: Weighted average of all naturally occurring isotopes of an element.

• All atomic masses are derived relative to carbon-12.

Periodic Table and Element Classification

• Metals: Left and center of the table; good conductors, malleable, ductile, shiny.

• Nonmetals: Right side; poor conductors, brittle, dull.

• Metalloids (Semimetals): Border between metals and nonmetals; intermediate properties.

• Groups: Vertical columns (e.g., alkali metals, alkaline earth metals, halogens, noble gases).

• Periods: Horizontal rows.

• Main Group (Representative) Elements: Groups 1, 2, and 13-18.

• Transition Elements: Groups 3-12.

Chemical Bonds and Nomenclature

Ionic vs. Covalent Compounds

• Ionic Compounds: Formed between metals and nonmetals; transfer of electrons; composed of cations and anions.

• Covalent (Molecular) Compounds: Formed between nonmetals; sharing of electrons.

Monoatomic and Polyatomic Ions

• Monoatomic Ions: Single atom ions (e.g., Na+, Cl-).

• Polyatomic Ions: Ions composed of multiple atoms (e.g., NO3-, SO42-).

• Memorize the list of common polyatomic ions provided in class.

Types of Cations and Naming

• Type I Cations: Metals that form only one type of positive ion (e.g., Na+, Ca2+).

• Type II Cations: Metals that can form more than one type of positive ion (e.g., Fe2+, Fe3+).

• Naming: Type I: name of element + 'ion'; Type II: name of element + (Roman numeral) + 'ion' (e.g., iron(III) ion).

Naming Compounds

• Binary Ionic Compounds (Type I): Metal (Type I) + nonmetal; name metal, then nonmetal with '-ide' ending.

• Binary Ionic Compounds (Type II): Metal (Type II) + nonmetal; name metal with Roman numeral, then nonmetal with '-ide'.

• Binary Molecular (Type III) Compounds: Two nonmetals; use prefixes (mono-, di-, tri-, etc.) to indicate number of atoms.

• Acids: Naming depends on the anion; binary acids (hydro- + base name + -ic acid), oxyacids (based on polyatomic ion name).

Determining Formulas from Names

• Use the names and charges of ions to write correct chemical formulas.

• Balance charges to ensure neutrality.

Lab Techniques and Measurement

Density, Accuracy, Precision, and Significant Figures in the Lab

• Be able to perform and interpret density calculations from lab data.

• Understand how to measure liquids and lengths accurately, including reading the meniscus and estimating the last digit.

• Apply significant figure rules to all measurements and calculations.

Tables

Common Conversion Factors

Classification of Matter

Subatomic Particles

Additional info: Some details, such as the full list of polyatomic ions and SI prefixes, are referenced but not included in the original file. Students should refer to their class handouts for these lists.