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Chapter 3: Periodic Table and Elements

Introduction to Chapter 3: The Periodic Table and Elements

  • Preparation for Lecture:
    • Students are advised to pause the video and retrieve the formula sheet from lecture materials.
    • Printing the formula sheet, especially the periodic table, is recommended for note-taking during the lecture.

Elements: Building Blocks of Matter

  • Definition:
    • Elements are a type of pure substance, defined as the fundamental building blocks of everything.
    • All matter is composed of combinations of approximately 100 basic substances, known as elements.
  • Number of Elements:
    • Currently, 118 elements have been discovered and isolated, all reflected on the periodic table.
    • Roughly \frac{3}{4} to \frac{2}{3} of these elements are naturally found.
    • The remaining elements are synthetic (man-made in a lab), often radioactive, unstable, and highly reactive.
  • Historical Discovery:
    • The discovery of elements dates back to ancient times.
    • A graphic shows periods of increased discovery, especially when technology advanced, indicating progress in scientific analysis.
  • Simplest Pure Substance:
    • Elements are the simplest type of pure substance and cannot be broken down further by chemical means.
    • They are the backbone of all substances and matter.
  • Composition:
    • An element is composed solely of atoms of the same type (e.g., hydrogen is made only of hydrogen atoms, carbon only of carbon atoms).
    • The image of graphite-like carbon shows a sample of the element large enough to be weighed, indicating its macroscopic presence.

Chemical Symbols

  • Uniqueness: Each element on the periodic table has a unique symbol.
  • Correlation:
    • Most symbols correlate directly to the element's English name (e.g., H for hydrogen, O for oxygen, N for nitrogen, C for carbon).
    • Some symbols derive from ancient Greek or Latin names:
      • Gold (Aurum in Latin/Greek) has the symbol Au.
      • Lead (Plumbum in Latin) has the symbol Pb (seen in terms like "unleaded fuel").
  • Formatting:
    • Chemical symbols are typically one or two-letter abbreviations.
    • The first letter is always capitalized.
    • If a two-letter symbol, the second letter is lowercase.
    • Example: Tungsten has the symbol W, despite no "w" in its English name, often joked to be due to running out of letters during discovery.
  • Required Memorization:
    • A list of 42 to 45 commonly used elements and their symbols are required for memorization.
    • These are frequently encountered in introductory chemistry.
    • On exams/quizzes, unmemorized elements will be provided with their symbols if needed.
    • The periodic table on the formula sheet only displays symbols, necessitating memorization to correlate symbols with names.
    • Recommendation: Use flashcards (symbol on one side, name on the other) for memorization.
    • Periodic tables will be provided on quizzes for questions requiring their use.

Atoms: The Smallest Particle

  • Definition: The smallest particle of an element that retains the properties of that element is an atom.
  • Identity:
    • An element of carbon is composed only of carbon atoms.
    • All atoms of a certain type are similar to one another but different from all other types.
    • With 118 known elements, there are 118 different types of atoms (e.g., hydrogen atoms are different from carbon atoms).
  • Elemental Forms:
    • Most elements exist as individual atoms (e.g., carbon exists as single C atoms).
    • Diatomic Molecules: Some elements naturally exist as diatomic molecules (two of the same atom bonded together).
      • This is their most stable and naturally occurring state.
      • Examples include elemental hydrogen (H2), oxygen (O2), nitrogen (N2), chlorine (Cl2), fluorine (F2), bromine (Br2), and iodine (I_2).
      • Students should memorize these diatomic elements, potentially by noting them on flashcards.
    • Other Forms: Some elements exist in more complex molecular forms (e.g., sulfur exists as S_8).

Dalton's Atomic Theory (Five Postulates)

  • Postulate 1: All matter is made up of small particles called atoms (now known: 118 types).
  • Postulate 2: All atoms of a given type are similar to one another and significantly different from all other types.
  • Postulate 3: The number and arrangement of different types of atoms in a pure substance determines its identity (relevant to fixed ratios in compounds).
  • Postulate 4: A chemical change is a combination, separation, or rearrangement of atoms to form new substances (foundation of chemical reactions).
  • Postulate 5: Only whole atoms take part in or result from any chemical reaction (no partial atoms).

Structure of the Atom: Discovery and Models

  • JJ Thompson (1897): Discovery of the Electron
    • Experiment: Used a cathode ray tube to investigate a gas discharge tube.
    • Observation: A beam (cathode ray) within the tube bent towards a positive magnet pole.
    • Conclusion: The beam was made of negatively charged particles.
    • Discovery: Named these particles electrons, the first subatomic particle (smaller than the atom, but a building block of it).
    • Plum Pudding Model: Proposed to describe the atom's structure post-electron discovery.
      • Described as a uniform, relatively positive sphere.
      • Negatively charged electrons were embedded or scattered throughout this positive sphere, like plums in pudding, held by attraction.
  • Ernest Rutherford (1911): Discovery of the Nucleus and Proton
    • Experiment: Gold Foil Experiment, designed to test Thompson's plum pudding model.
      • Shot high-energy, positively charged alpha particles (\alpha particles) at a thin piece of gold foil.
    • Expected Result (based on Plum Pudding): All alpha particles should pass straight through a uniformly distributed positive mass without significant deflection.
    • Actual Result:
      • Most particles did pass straight through (indicating mostly empty space).
      • Some particles were deflected at large angles, and a very few bounced straight back (indicating a dense, positively charged obstruction).
    • Conclusion: The plum pudding model was incorrect.
    • Rutherford's Model:
      • Proposed a small, dense core at the center of the atom, called the nucleus.
      • The nucleus contains most of the atom's mass and all its positive charge.
      • Called the positively charged particles within the nucleus protons (second subatomic particle discovered).
      • The rest of the atom is mostly empty space with electrons scattering around the nucleus.
  • Niels Bohr and Quantum Models:
    • Bohr proposed a model depicting electrons in orbits, similar to planets around a sun (solar system analogy).
    • Later, Schrodinger and Heisenberg refined this, showing electrons exist more as a "cloud" or "stars" around the nucleus, not fixed orbits.
    • Electrons hover around the nucleus due to opposite charges and participate in chemical reactions based on their distance from the nucleus.
  • James Chadwick: Discovery of the Neutron
    • Discovery (after Rutherford): Chadwick discovered the neutron (third subatomic particle).
    • Reasoning: To account for the additional mass in the nucleus (beyond just protons) and maintain the overall neutral charge of the atom, there must be a neutral particle with mass within the nucleus.
    • Properties: Neutral charge, mass similar to a proton.
    • Experiment: Oil Drop Experiment (though Chadwick's neutron discovery used different experiments involving alpha particles and beryllium).

Overview of Subatomic Particles

  • Electron (e^-):
    • Charge: Approximately -\text{1} (negative).
    • Mass: Very lightweight, essentially negligible (0.00055 \text{ AMU}, roughly \frac{1}{1836} of a proton or neutron).
    • Location: Outside the nucleus, in an electron cloud.
    • Discoverer: JJ Thompson.
  • Proton (p^+):
    • Charge: Approximately +\text{1} (positive).
    • Mass: Approximately 1 \text{ AMU} (atomic mass unit).
    • Location: Inside the nucleus.
    • Discoverer: Ernest Rutherford.
  • Neutron (n_0):
    • Charge: Neutral (0).
    • Mass: Approximately 1 \text{ AMU}.
    • Location: Inside the nucleus.
    • Discoverer: James Chadwick.

Atomic Mass Units (AMU)

  • Purpose: Since subatomic particles have extremely small masses, AMU is used for the atomic mass scale.
  • Definition:
    • Reference point: One carbon-12 atom is designated as 12 \text{ AMU}.
    • Therefore, 1 \text{ AMU} is defined as \frac{1}{12} of the mass of one carbon-12 atom.
    • 1 \text{ AMU} \approx 1.66 \times 10^{-24} \text{ g}.
  • Relevance: Primarily represents the number of protons and neutrons in the dense nucleus, where most atom's mass resides.

Atomic Number (Z) and Mass Number (A)

  • Atomic Number (Z):
    • Definition: The number of protons in an atom's nucleus.
    • Identification: Uniquely identifies an element. All atoms of the same element have the same atomic number.
    • Location on Periodic Table (Formula Sheet): The whole number located above the element symbol.
  • Mass Number (A):
    • Definition: The sum of the number of protons and neutrons in an atom's nucleus.
    • Unit: Expressed in atomic mass units (\text{AMU}).
    • Characteristic: Always a whole number (no fractional protons or neutrons).
    • Calculation of Neutrons: Number of neutrons = Mass Number (A) - Number of Protons (Z).
  • Neutral Elements:
    • If an element is neutral (no positive or negative charge on its symbol), the number of protons equals the number of electrons.
    • All elements shown on the periodic table are assumed neutral.

Isotopes

  • Definition: Atoms of the same element (same atomic number/protons) that have different mass numbers due to a differing number of neutrons.
  • Characteristics:
    • Same atomic number (e.g., all magnesium isotopes have Z = 12\text{ protons}).
    • Different mass numbers (e.g., magnesium-24 (12\text{ neutrons}), magnesium-25 (13\text{ neutrons}), magnesium-26 (14\text{ neutrons})).
    • Different number of neutrons (calculated as Mass Number - Atomic Number).
  • Impact: Neutrons do not affect the identity or overall chemical nature of the atom, but they contribute to its mass.
  • Notation: Isotopes are designated by element name-mass number (e.g., Chlorine-35, Chlorine-37).

Atomic Weight (Weighted Isotopic Average)

  • Concept: Elements often exist naturally with multiple isotopes, each having a different mass and relative abundance.
  • Definition: Atomic weight is the average of all isotope masses for a particular element, taking into account their natural percent abundances.
  • Location on Periodic Table: The decimal number listed below the element symbol on the periodic table.
  • Calculation of Weighted Average:
    • Atomic Weight = \Sigma (\text{Isotope Mass} \times \frac{\text{Percent Abundance}}{100}) for each isotope.
    • Example (Chlorine): Chlorine exists as Chlorine-35 (approx. 75\text{%} abundance) and Chlorine-37 (approx. 25\text{%} abundance).
      • The atomic weight of chlorine (35.45 \text{ AMU}) is closer to 35 because of the higher abundance of Chlorine-35. This also explains why hydrogen's atomic weight is 1.008 \text{ AMU} instead of exactly 1 \text{ AMU}, due to small amounts of hydrogen-2 and hydrogen-3 isotopes.
  • Distinction from Mass Number:
    • Mass Number: Designates #protons + #neutrons, always a whole number, no decimals.
    • Atomic Weight: Weighted average of isotopic masses, found on the periodic table, usually has decimals.

Homework Questions

  • Homework 1: Fill in a table with element name, symbol, atomic number, protons, electrons, neutrons, and mass number. Assume all elements are neutral, so #electrons = #protons. Use the periodic table for assistance.
  • Homework 2: Calculate the atomic weight of Gallium (Ga) given two isotopes, their isotopic masses, and their relative percent abundances.
    • Gallium-69: Mass = 68.93 \text{ AMU}, Abundance = 60.11\text{%}.
    • Gallium-71: Mass = 70.92 \text{ AMU}, Abundance = 39.89\text{%}.
    • Full work must be shown for credit. The answer can be checked against the periodic table.