Atomic Theory and Structure

Dalton's Atomic Theory

Dalton's atomic theory laid the foundation for modern chemistry by describing the nature of atoms and their role in chemical reactions.

• Atoms of a given element have the same mass and properties that distinguish them from atoms of other elements.

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

• Atoms are indivisible in chemical processes; they cannot be created or destroyed.

• Each element is composed of tiny indestructible particles called atoms.

• Exception: The concept of a nucleus and subatomic particles was not part of Dalton's original theory. Modern understanding includes protons, neutrons, and electrons.

Example: Water (H2O) is formed by combining two hydrogen atoms and one oxygen atom in a fixed ratio.

Atomic Mass Unit (amu)

The atomic mass unit is a standard unit for measuring atomic and molecular masses.

• Definition: 1 amu is defined as 1/12 the mass of a carbon-12 atom.

• Hydrogen atom: Has only one proton.

• Carbon atom: Contains six protons and six neutrons.

• Mass of electrons: Negligible compared to protons and neutrons.

Formula:

Isotopes

Isotopes are atoms of the same element that have different numbers of neutrons.

• Same element: Same number of protons, different number of neutrons.

• Isotopic notation: , where A = mass number, Z = atomic number, X = element symbol.

Example: Carbon-12 and Carbon-14 are isotopes of carbon.

Periodic Table and Electron Configuration

Periodic Table Organization

The periodic table arranges elements by increasing atomic number and groups elements with similar properties.

• Metals: Located on the left side; typically form cations.

• Nonmetals: Located on the right side; typically form anions.

• Groups: Vertical columns; elements in the same group have similar chemical properties.

• Periods: Horizontal rows.

Example: Group 2A elements are called alkaline earth metals.

Electron Configuration

Electron configuration describes the arrangement of electrons in an atom's orbitals.

• Notation: Uses numbers and letters to indicate energy levels and sublevels (e.g., 1s2 2s2 2p6).

• Hund's Rule: Electrons fill orbitals singly before pairing.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of quantum numbers.

Example: The electron configuration for potassium (K) is .

Chemical Bonds and Ions

Ions and Isotopes

Ions are atoms or molecules that have gained or lost electrons, resulting in a net charge.

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

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

• Isotopes: Atoms with the same number of protons but different numbers of neutrons.

Example: has 21 electrons (24 protons minus 3 positive charge).

Measurement and Significant Figures

Scientific Notation and Significant Figures

Scientific notation is used to express very large or small numbers. Significant figures indicate the precision of a measurement.

• Scientific notation:

• Significant figures: All nonzero digits and zeros between them or after a decimal point are significant.

Example: The number 10.010 has five significant figures.

Metric Prefixes and Unit Conversion

Metric prefixes are used to express units in powers of ten.

• Milli-:

• Kilo-:

• Micro-:

Example: 1 cubic meter = cubic centimeters.

Density and Calculations

Density is the mass per unit volume of a substance.

• Formula:

• Units: g/cm3 or kg/m3

Example: An object with mass 1.840 kg and volume 0.0015 m3 has a density of .

States of Matter and Changes

Physical and Chemical Changes

Physical changes do not alter the chemical composition of a substance, while chemical changes result in new substances.

• Physical change: Melting, boiling, dissolving.

• Chemical change: Burning, rusting, metabolism.

Example: The burning of natural gas is a chemical change.

Mixtures and Pure Substances

Mixtures can be separated by physical means, while pure substances cannot.

• Homogeneous mixture: Uniform composition (e.g., sugar dissolved in water).

• Heterogeneous mixture: Non-uniform composition.

• Pure substances: Elements and compounds.

Example: Gasoline is a homogeneous mixture.

Energy and Thermochemistry

Energy Concepts

Energy is the capacity to do work or produce heat. It cannot be created or destroyed (law of conservation of energy).

• Kinetic energy: Energy of motion.

• Potential energy: Stored energy due to position.

Formula: (heat transfer equation)

Specific Heat Calculations

Specific heat is the amount of heat required to raise the temperature of one gram of a substance by one degree Celsius.

• Formula:

• Variables: q = heat (J), m = mass (g), c = specific heat (J/g°C), = change in temperature (°C)

Example: A 15.0 g lead ball heated with 40.5 J of heat and specific heat of 0.128 J/g°C:

Light and Electromagnetic Spectrum

Properties of Light

Light travels as electromagnetic waves and has properties such as wavelength and frequency.

• Speed of light: m/s

• Photon: A packet of light energy.

• Wavelength: Determines color and energy of light.

Electromagnetic Spectrum

The electromagnetic spectrum includes all types of electromagnetic radiation, arranged by wavelength.

Example: Radio waves have the longest wavelength; gamma rays have the shortest.

Bohr Model of the Atom

The Bohr model describes electrons in specific, quantized orbits around the nucleus.

• Electrons absorb energy: Move to higher energy levels.

• Electrons emit energy: Fall to lower energy levels, releasing photons.

• Limitation: The model does not explain all atomic behavior, especially for multi-electron atoms.

Formula: (energy of a photon)

Additional info:

• Some questions reference specific quiz numbers and multiple-choice formats, indicating these are exam or homework questions.

• All content is relevant to introductory college chemistry, covering atomic theory, measurement, periodic table, energy, and matter.