chem lecture 2

Chapter Overview

  • Course: CHEM1307-003

  • Instructor: Sarah E. Ring, TTU Fall 2024

  • Schedule: Tuesday/Thursday, 2:00 PM - 3:20 PM

Pure Substances vs Mixtures 2.1

Pure Substance

  • Definition: Consists of only an element or a compound.

  • Characteristics:

    • Constant composition and properties throughout.

    • Every sample taken has exactly the same chemical makeup.

  • Types:

    • Atoms (elements): Smallest identifiable particles.

    • Molecules: Two or more of the same atom chemically bound together.

    • Compounds: Pure substances made of two or more elements in a fixed ratio.

Mixture

  • Definition: Consists of two or more pure substances that are not chemically bonded.

  • Characteristics:

    • Each substance retains its own properties.

    • Can be homogeneous (uniform composition) or heterogeneous (distinct different components).

Fundamental Chemical Laws 2.2

Law of Conservation of Mass

  • Principle: Mass is neither created nor destroyed in a chemical reaction.

  • Key Concept: The total mass of reactants equals the total mass of products.

  • Formula: Mass of products = Mass of reactants.

Example Problem

  • Combustion Reaction of Propane:

    • 55.5 g of propane reacts with 96.0 g of oxygen, producing water and carbon dioxide.

    • If 63.2 g of water is produced, calculate the mass of carbon dioxide produced.

Law of Definite Composition

  • Definition: A specific compound is composed of the same elements in the same mass fractions, regardless of the source.

  • Example: For glucose (C6H12O6), the mass composition is:

    • Carbon: 40%

    • Hydrogen: 7%

    • Oxygen: 53%

  • Important Note: Altering the ratios results in a different substance.

Example Problem

  • H3PO4 Analysis:

    • Given 3.03 g H, 30.97 g P, 64.00 g O, determine:

      • Mass fraction

      • Percent by mass of each element.

Dalton’s Atomic Theory 2.3

Postulates

  1. All matter consists of atoms; atoms cannot be created or destroyed.

  2. Atoms of one element cannot be converted into those of another.

  3. Atoms of an element are identical in mass and properties, different from other elements.

  4. Compounds are formed from specific ratios of atoms (law of definite composition).

Erroneous Aspects of Dalton’s Theory

  • Recent discoveries have disproved some postulates:

    1. Atoms can be divided.

    2. Atoms can be created or destroyed.

    3. Isotopes exist (different atoms can vary within the same element).

    4. Elements exhibit periodic properties.

Models of the Atom 2.4

J.J. Thompson

  • Discovery: Electron

    • Characteristics: Negatively charged and has a very small mass.

    • Method: Used cathode ray tubes; demonstrated electrons could be removed to create electric charge.

Robert Millikan

  • Discovery: Charge and Mass of the Electron

    • Experiment: Used oil droplets in an electric field to measure charge.

    • Findings: Droplets were attracted to the positive plate, counteracting gravity.

Ernest Rutherford

  • Discovery: Nucleus

    • Experiment: Used alpha particles directed at gold foil.

    • Results: Some particles deflected back, some passed through, indicating a dense nucleus.

Atomic Structure & Isotopes 2.5

Subatomic Particles

  • Types:

    • Protons: Positive charge, 1 amu, located in the nucleus.

    • Neutrons: Neutral charge, 1 amu, located in the nucleus.

    • Electrons: Negative charge, 0.0005 amu, orbiting the nucleus.

Atomic Notation

  • Nuclear notation: XA

    • A = mass number (protons + neutrons)

    • Z = atomic number (number of protons)

  • Isotope Example:

    • Notation: 4He (mass number 4, atomic number 2).

Hyphen Notation

  • Element symbol notation: Element name - mass number.

  • For example: Uranium-235, Carbon-12, Carbon-13, Carbon-14.

Example Problems

  • Use APE MAN method to determine protons, neutrons, and electrons for isotopes:

    • U-235, Ca-41, Mo-96, and provide calculations for Carbon-14.

Atomic Mass & Percent Abundance 2.6

Atomic Mass

  • Definition: Average mass of isotopes, weighted by their natural abundance.

  • Formula: Sum of (Mass of isotope x % abundance of isotope).

Example Problems

  • Silver (Ag) and Iron (Fe): Calculate the atomic mass using given isotopic mass and abundance data.