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Modelling the Atom & Bohr-Rutherford Diagrams

Learning Objectives:

• Understand the historical changes and advancements in atomic models from the early theories of matter to modern quantum mechanics.

• Identify and describe the major components of atoms, including protons, neutrons, and electrons.

• Calculate average atomic mass based on isotopic abundance and mass.

Skills to Develop:

• Describe the significant experiments and discoveries that influenced atomic theory up to 1913, focusing on contributions by key scientists.

• Utilize standard atomic notation effectively for problem-solving, including understanding the implications of charge and isotopes.

Warm-Up Activity

Task:

Create a labeled sketch of your perception of an atom, including all known parts (nucleus, electron shells, etc.) and their charges. For additional review, engage with an interactive activity online that illustrates atomic structure.

Dalton’s Model of the Atom

Description:

• Known as the Billiard Ball Model, Dalton represented atoms as solid, indivisible spheres that formed the basis for further scientific inquiry.

• Atoms of different elements were proposed to be distinct and varied in mass.

Dalton’s Atomic Theory

Core Principles:

1. Matter is composed of indivisible atoms. (This was later revised as atoms are now known to be comprised of subatomic particles: protons, neutrons, and electrons.)

2. All atoms of a given element are identical in mass and properties. (This has been revised to factor in isotopes, which have differing numbers of neutrons and therefore mass.)

3. Compounds are formed from whole number combinations of different types of atoms, demonstrating the law of definite proportions.

4. Chemical reactions involve rearrangement, combination, or separation of atoms without altering their intrinsic properties.

Thomson’s Experiment

Objective:

• To infer the fundamental aspects of atomic structure based on the behavior of electrons and their interactions.

Key Notes:

• Conducted experiments with and without an electromagnetic field to observe particle behavior.

Thomson’s “Plum Pudding” Model

Concept:

• Atoms are electrically neutral; hence, the positive charge must balance the negative charges of electrons.

• Thomson introduced the idea of a positively charged spherical cloud with negatively charged electrons embedded within it, marking the first proposal of internal atomic structure.

• His work was pivotal as he was the first to demonstrate that atoms could be divided into smaller components (the electron).

Rutherford’s Experiment

Notable Findings:

• In a groundbreaking experiment aiming alpha particles at a thin gold foil, most particles passed through; however, a notable few were deflected at large angles.

• This deviation led to the inference of a concentrated, positively charged mass within the atom, contradicting Thomson’s model.

Rutherford’s Model of the Atom

Also Known As: Planetary Model

Characteristics:

• Electrons orbit a positively charged nucleus similarly to planets orbiting the sun.

• This model identified the nucleus as the central point of mass and positive charge, with the majority of the atom’s volume being empty space.

Chadwick’s Contribution

Key Discoveries:

• Identified the neutron as an uncharged yet massive particle residing in the nucleus, which explained discrepancies in atomic mass calculations and the stability of the nucleus.

Summary of Classical Chemistry Figures

Key Contributors:

• J.J. Thomson - Discovered the electron and proposed the plum pudding model, marking the beginning of modern atomic theory.

• Henri Becquerel - Investigated radioactive emissions, contributing to the understanding of atomic decay.

• Marie & Pierre Curie - Notably discovered radioactive elements such as polonium and radium, advancing the study of nuclear chemistry.

• Robert Millikan - Conducted the oil-drop experiment, successfully determining the charge and mass of electrons.

• Ernest Rutherford - Developed the nuclear model of the atom and discovered protons through his gold foil experiment.

• James Chadwick - Discerned the existence of neutrons, enriching the understanding of atomic structure.

Spectroscopy & Emission Spectra

Concepts:

• Continuous Spectrum: Produced by white light containing all frequencies, showing a seamless gradient of colors.

• Line Spectra: Produced by isolated elements, these spectra display specific frequencies corresponding to the energy levels of electrons transitioning in the atom's structure and are valuable in identifying elements.

Bohr’s Modifications to Rutherford’s Model

Key Observations:

• Electrons inhabit distinct energy levels or shells, with each level correlating to specific colors of emitted light during transitions.

• Energy can be absorbed or released, resulting in electrons moving between these quantized energy levels, contributing to the understanding of atomic emission and absorption spectra.

The Bohr-Rutherford Model

Structure:

• Each electron shell represents an energy level, with a maximum capacity of 2n² electrons, where n is the shell level.

• Valence electrons reside in the outermost shells, while inner shells fill preferentially based on energy levels.

Schrodinger’s Quantum Mechanical Model

Features:

• Electron shells consist of subshells where electrons are most likely to be found, described using shapes and orientations (s, p, d, f).

• The model represented electron locations as probabilistic clouds rather than fixed orbits, emphasizing the uncertainty principle in quantum mechanics.

Representing Elements

Symbols:

• Each element has a designated one or two-letter symbol, derived primarily from its English or Latin name.

Atomic Number:

• Represents the number of protons in an atom's nucleus; in neutral atoms, it also indicates the number of electrons.

• The charge is determined by the difference between protons and electrons: equal counts equal a neutral charge; differing counts yield either anions (negatively charged) or cations (positively charged).

Isotopes:

• Variants of the same element with different numbers of neutrons, resulting in varied atomic masses; isotopes can have significant implications in fields like medicine and archaeology.

Standard Atomic Notation

Format:

• Comprises the chemical symbol, the atomic mass, and the atomic number: specifics of isotopes are delineated via mass numbers (e.g., 12C for Carbon-12).

Average Atomic Mass Calculation

Formula:

• AMU = ( %1 × mass1 + %2 × mass2 + ... + %n × massn) / 100This calculation allows for the determination of an element's average atomic mass based on the relative abundances of its isotopes.

Examples of Average Atomic Mass Calculations:

• Rubidium: 85Rb (72.2%) and 87Rb (27.8%): Average mass = 85.56 amu.

• Uranium: Isotopes 234U (0.01%), 235U (0.71%), and 238U (99.28%): Average mass = 237.98 amu.

• Titanium: Common isotopes: 46Ti (8.0%), 47Ti (7.8%), 48Ti (73.4%), 49Ti (5.5%), 50Ti (5.3%): Average mass = 47.92 amu.

Challenge: Finding Isotopic Abundance

Task:

Identify isotopic abundances for a new element with distinct isotopes given specific isotopic masses and the average atomic mass; apply average atomic mass calculations appropriately to solve the problem.