Chemistry Basics Quiz

Chapter 1: The Basics of Chemistry

• Chemistry is the study of matter, its properties, and the mechanisms involved in its changes.
• Matter is any substance that has mass and occupies space.
• The fundamental building blocks of matter are atoms, which cannot be chemically broken down into smaller components.
• Atoms combine to form larger molecules and compounds.
• Matter comes in three states: gas, solid, and liquid.

States of Matter

• Solids: Atoms are closely packed together, relatively stationary, slightly compressible, have a definite shape and volume.
• Liquids: Atoms are more spread out, have intermediate mobility, slightly compressible, no definite shape, but definite volume.
• Gases: Atoms are very far apart, high velocities, highly compressible, no definite shape or volume.

Classification of Matter

• Matter:
  • Pure substance: Cannot be physically separated.
    • Element: Cannot be decomposed chemically.
    • Compound: Can be decomposed chemically.
  • Mixture: Can be physically separated.
    • Homogeneous mixture: Uniform composition (solution).
    • Heterogeneous mixture: Non-uniform composition.

The Scientific Method

• Develop theories that are consistent with available data.
• Make observations about nature using experiences, thoughts, and readings.
• Ask interesting questions.
• Formulate hypotheses based on observations and questions.
• Develop testable predictions from hypotheses.
• Gather data to test predictions and verify results.

Units of Measurement

• Property | English units | SI units
• --|---|---
• Mass* | pounds (lbs) | kilogram (kg)
• Length* | yard (yd) | Meter (m)
• Volume* | gallon (gal) | Cubic Meter (m^3)
• Temperature* | Fahrenheit (°F) | Kelvins (K)

Temperature Scales

• Conversion formulas for temperature scales: °F = (°C) + 32 °C = (°F - 32)/9 9 * K =°C + 273.15

Significant Figures

• Report all measurements to one digit past the last digit of certainty.
• The number of digits in a measurement indicates the precision of the measurement.
• Zeros between non-zero digits are significant
• Leading zeros are not significant
• Trailing zeros after a decimal point are significant

Exact Quantities

• Quantities whose precision is considered infinite are known as exact numbers
• Equality statements or conversion factors

Dimensional Analysis

• Use conversion factors to express a given unit in a different unit
• The magnitude of the units in the conversion factor has to be the same as the given unit; therefore, the given unit has to be on the bottom and the new unit on the top
• The magnitude of the units on the top and bottom has to be the same

Density

• Density is the ratio of mass to volume.
• Density = mass/volume
• Density, is a physical property, unique to each substance, and is temperature dependent.
• Water Density at 20°C = 0.99823 g/mL

Chapter 2: Matter at the Atomic Level

• Atoms are extremely small, fundamental particles of matter.
• Atoms of the same element are identical to each other
• Atoms in chemical reactions are not changed into other elements; atoms are neither created or destroyed.
• Compounds are formed when atoms of more than one element combine in a fixed, whole-number ratio.
• Atomic theory
• J.J. Thomson and the electron
• Rutherford's gold-foil experiment and the nuclear model of the atom
• Atomic numbers
• lons
• Atomic mass and isotopes
• The periodic table
• The mole and Avogadro's number
• Radioactivity
• Subatomic particles
• Symbols of elements, and their relations to atomic numbers and mass numbers
• Isotopes and their percent natural abundance
• Average mass calculation
• Properties of metals, nonmetals and metalloids
• Using the Periodic Table to predict properties of elements and formations of ions

Important Concepts

• Atomic number (Z): The number of protons in an atom.
• Mass number (A): The sum of protons and neutrons in an atom.
• Isotopes: Atoms of the same element with different mass numbers (different numbers of neutrons).
• Isotopic abundance: The percentage of each isotope in a naturally occurring sample of an element.
• Atomic mass: The weighted average of the masses of all the naturally occurring isotopes of an element.
• Atom: The smallest unit of an element that retains the properties of that element.
• Mole: A unit that represents 6.022 x 10^23 particles (atoms or molecules).
• Molar mass: The mass of one mole of a substance in grams/mole

Chapter 3: Molecules, Compounds, and Their Composition

• Introduction: Importance of compound formulas and naming.
• Chemical formulas representing compounds.
• Differentiating ionic and covalent compounds
• Naming ionic and molecular compounds
• Molar mass calculation
• Mass percent calculation
• Calculating empirical formula
• Calculating molecular formulas

Chapter 4: Chemical Reactions and Stoichiometry

• Introduction: Defining chemical reactions and their relevance to understanding chemical systems.
• Chemical Equations provide a representation of reactants and products.
• Coefficents are used when balancing chemical equations for conservation of mass, showing the number of moles of each reactant and product in a chemical reaction.
• Stoichiometry: relates the amounts of compounds or elements involved in a reaction.
• Using chemical equations/coefficients to calculate the number of moles of reactants/products needed or formed in a reaction.
• Limiting reactants: The reactant that determines the amount of product that can be formed.
• Excess reactants: The reactant that is not completely consumed in the reaction.
• Calculating the percent yield of a chemical reaction and the relationship between theoretical and actual yields.
• Types of chemical reactions to be aware of: - double replacement (exchange/metathesis - precipitation - acid-base neutralization - gas -forming
• Reactions in aqueous solutions:
  • solutions/solvents/solutes
  • molarity calculations including using molarity in stoichiometric calculations

Chapter 5: Thermochemistry

• Introduction to Thermochemistry
• Energy Types (Kinetic and Potential)
• Law of Conservation of Energy
• System and Surrounding
• Heat and Work relationships
• Enthalpy of reaction
• Specific Heat measurement
• Calorimetry (Coffee cup and Bomb)
• Hess's Law

Chapter 6: Early Quantum Theory - The Nature of Light and Matter

• Introduction (Classical Mechanics vs Quantum theory)
• Electromagnetic Radiation
• Wavelength and Frequency
• The Electromagnetic Spectrum
• Atomic Spectroscopy
• Planck's constant
• Bohr Model
• De Broglie Wavelength
• The Uncertainty Principle
• Quantum Mechanics
• Quantum Numbers
• Using Quantum Mechanics to describe electron behavior in atoms

Chapter 8: Representing Valence Electrons: Lewis Structures of Atoms, Ions, and Molecules

• Introduction to Electron Configurations, Lewis Structures
• Lewis Structures of selected atoms
• Writing Lewis Structures
• Formal Charges
• Resonance
• Exceptions to the Octet Rule