chemistry final study guide

I. Units and Measurements

• SI Units: The International System of Units (SI) is used for measurements. Base units include:

  • Second (s) for time

  • Kilogram (kg) for mass

  • Kelvin (K) for temperature

  • Meter (m) for length

  • Derived units are combinations of base units, e.g., cubic meter (m³) for volume.

• Prefixes: Common prefixes include:

  • Centi- (c, 10⁻²)

  • Milli- (m, 10⁻³)

  • Kilo- (k, 10³)

  • Mega- (M, 10⁶)

  • Nano- (n, 10⁻⁹)

  • Pico- (p, 10⁻¹²)

• Temperature Scales: The Kelvin scale is the SI unit for temperature.

• Unit Conversions: Dimensional analysis is used to convert between units. Examples:

  • 1 kilogram = 1,000 grams

  • 1 megaliter = 1,000,000 liters

  • 1 meter = 100 centimeters

• Base vs. Derived Units:

  • Base units: Fundamental measurements (e.g., meter).

  • Derived units: Combinations of base units (e.g., m³).

• Density: Density is mass per unit volume. Formula:

  • Volume = Mass / Density

• Problem Solving: A three-step process is often used for solving measurement-related problems.

• Celsius to Kelvin Conversion: To convert degrees Celsius to kelvins, add 273.15 to the Celsius temperature.

• Scientific Notation: Expresses numbers as a coefficient multiplied by a power of 10.

• Significant Figures & Rounding: Significant figures represent the precision of a measurement. Rules for rounding and determining significant figures depend on the calculation.

• Accuracy and Precision:

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

  • Precision: The reproducibility of measurements.

• Data Representation: Common graph types include bar graphs and line graphs.

II. Matter—Properties and Changes

• Matter: Anything that has mass and volume.

• Substances: Matter with uniform and unchanging composition.

• Physical Properties: Observed without changing chemical composition (e.g., color, density, melting point, boiling point).

• Chemical Properties: Observed only by changing chemical composition (e.g., reactivity, flammability, rusting).

• States of Matter:

  • Solids: Definite shape and volume.

  • Liquids: Definite volume but take the shape of their container.

  • Gases: Take both the shape and volume of their container.

• Physical Changes: Alter form or appearance without changing chemical composition (e.g., boiling, melting, freezing, crushing).

• Chemical Changes: Result in new substances with different properties (e.g., burning, rusting, decomposition).

• Conservation of Mass: Mass is neither created nor destroyed in chemical reactions.

• Chemical Equations: Represent reactants and products in a chemical reaction.

• Mixtures:

  • Heterogeneous: Visible differences in composition (e.g., sand-water mixture).

  • Homogeneous (Solutions): Uniform composition (e.g., salt-water mixture).

• Separation Techniques:

  • Filtration: Separates solids from liquids.

  • Distillation: Separates substances based on boiling points.

  • Crystallization: Forms pure solids from solutions.

  • Chromatography: Separates substances based on movement through a medium.

• Elements and Compounds:

  • Elements: Substances that cannot be broken down into simpler substances.

  • Compounds: Chemical combinations of two or more different elements.

• Laws of Composition:

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

  • Law of Multiple Proportions: If two elements form multiple compounds, the masses of one element combine with a fixed mass of the other in a simple whole-number ratio.

III. The Structure of the Atom

• Early Atomic Models:

  • Ancient philosophers and John Dalton contributed to early atomic theory.

  • Dalton's theory proposed that atoms are indivisible, which was later disproved.

  • The plum pudding model and Rutherford's nuclear model advanced atomic understanding.

• Subatomic Particles:

  • Atoms contain protons (positive), electrons (negative), and neutrons (neutral).

  • Protons and neutrons are in the nucleus; electrons orbit outside the nucleus.

• Atomic Number: Number of protons in an atom; identifies the element. Equal to the number of electrons in a neutral atom.

• Isotopes: Atoms of the same element with different numbers of neutrons.

• Mass Number: Total number of protons and neutrons in an atom.

• Atomic Mass: Weighted average of the masses of an element's naturally occurring isotopes.

• Radioactivity: Spontaneous emission of radiation from an unstable nucleus. Types:

  • Alpha: Helium nuclei.

  • Beta: High-energy electrons.

  • Gamma: High-energy electromagnetic waves.

IV. Electrons in Atoms

• Light and Electromagnetic Radiation:

  • Behaves as a wave (wavelength, amplitude, frequency).

  • Energy is quantized (photoelectric effect).

• Atomic Emission Spectra: Unique sets of colored lines produced when atoms emit light.

• Models of the Atom:

  • Bohr Model: Electrons orbit the nucleus at specific energy levels.

  • Quantum Mechanical Model: Treats electrons as particles and waves.

  • Heisenberg Uncertainty Principle: Exact location and momentum of an electron cannot be simultaneously known.

• Electron Configurations:

  • Describe electron arrangement in an atom.

  • Governed by the Aufbau principle, Pauli exclusion principle, and Hund’s rule.

• Valence Electrons: Outermost electrons, represented in electron-dot structures.

V. The Periodic Table and Periodic Law

• Development of the Periodic Table:

  • Mendeleev: Arranged elements by increasing atomic mass.

  • Moseley: Arranged elements by increasing atomic number.

  • Periodic Law: Element properties recur periodically when arranged by atomic number.

• Organization:

  • Periods: Rows.

  • Groups: Columns.

  • Representative Elements: Groups 1, 2, and 13–18.

  • Transition Elements: Groups 3–12.

• Element Classification:

  • Metals: Conduct electricity, malleable, ductile.

  • Nonmetals: Poor conductors, brittle.

  • Metalloids: Properties of both metals and nonmetals.

• Important Groups:

  • Group 1: Alkali metals.

  • Group 2: Alkaline earth metals.

  • Group 17: Halogens.

  • Group 18: Noble gases.

• Periodic Trends:

  • Atomic Radius: Increases down a group, decreases across a period.

  • Ionic Radius: Follows a similar trend to atomic radius.

  • Ionization Energy: Energy required to remove an electron; increases across a period, decreases down a group.

  • Electronegativity: Ability to attract electrons in a bond; increases across a period, decreases down a group.

  • Octet Rule: Atoms tend to have eight electrons in their outermost energy level.

• Blocks in the Periodic Table:

  • s-block: Groups 1 and 2.

  • p-block: Groups 13–18.

  • d-block: Groups 3–12.

  • f-block: Lanthanides and actinides.