Atomic Structure, Rutherford's Experiment, and Subatomic Particles

Rutherford's Gold Foil Experiment and the Discovery of the Nucleus

  • Alpha emitters are devices that shoot alpha particles, which are defined as helium nuclei.

  • Alpha particles are radioactive particles that will be explored in more detail in later discussions.

  • During the experiments, an alpha emitter shot particles at a thin piece of gold foil.

  • The gold foil was surrounded by a strip of photosensitive film, which produced a visible spot whenever an alpha particle struck it.

  • Experimental observations revealed that most alpha particles passed directly through the gold foil, creating a spot on the backside.

  • However, a small portion of the particles scattered and hit different spots along the foil at various angles.

  • Rutherford reasoned from these results that the previous atomic models were inaccurate.

  • He concluded that an atom must contain a tight, dense nucleus that holds most of the mass, while the rest of the atomic volume is primarily free space.

  • This experiment directly led to the discovery of the nucleus and the understanding that atoms are composed of electrons, protons, and neutrons.

Subatomic Particle Composition and Mass

  • The neutron was theorized to exist as a stabilizing particle. Because protons are positively charged, they naturally repel one another; the neutron helps overcome these repulsions to keep the nucleus intact.

  • The specific mass of the three primary subatomic particles are as follows:

    • Electron: 9.11×10−31 kg9.11 \times 10^{-31}\,kg

    • Proton: 1.673×10−27 kg1.673 \times 10^{-27}\,kg

    • Neutron: 1.675×10−27 kg1.675 \times 10^{-27}\,kg

  • Key observations regarding particle mass:

    • The neutron is the heaviest of the three particles.

    • Masses of particles are determined by the magnitude of their negative exponents; a smaller negative exponent identifies a larger mass.

    • Protons and neutrons are several orders of magnitude heavier than electrons.

    • The mass of a proton is approximately 1,0001,000 times greater than that of an electron.

    • A conceptual comparison of this mass difference is similar to comparing a gram to a kilogram, or a cantaloupe to a car.

Atomic Scale and Volume

  • The nucleus contains nearly all of the atom's mass, but standard drawings of atoms are almost never to scale.

  • If a true-to-scale model of an atom were drawn with a nucleus the size of a small object, the electron cloud would be massive. For example, if the nucleus were sized to a specific point, the electron cloud would stretch the distance from I-64 to Main Street.

  • Virtually all the volume of an atom is comprised of the electron cloud.

  • Electrons are so light and move so fast that they exhibit specific properties where their position is not easily measurable; they are often described as being "smeared out" or existing in multiple places simultaneously (quantum behavior).

  • Summary of atomic roles:

    • All volume is due to electrons.

    • All mass is concentrated in the nucleus (protons and neutrons).

Atomic Characterization and Notation

  • Atoms are characterized by three primary identifiers:

    • Atomic Number (ZZ): The number of protons in the nucleus.

    • Atomic Symbol: The shorthand abbreviation for the element.

    • Atomic Mass: The weighted average of mass numbers of an element's isotopes by their natural abundances.

  • Mass Number (AA): The total count of protons and neutrons in a specific atom.

    • Formula: A=protons+neutronsA = \text{protons} + \text{neutrons}

    • Notation: Uppercase AA denotes the mass number.

  • To determine the number of neutrons in an atom, subtract the atomic number (ZZ) from the mass number (AA):

    • Neutrons=A−Z\text{Neutrons} = A - Z

Mathematics of Weighted Averages and Isotopes

  • The atomic mass listed on the periodic table is a weighted average, not a simple average.

  • A simple average is calculated as:

    • Average=1n∑i=1nxi\text{Average} = \frac{1}{n} \sum_{i=1}^{n} x_i

  • In a weighted average, a weighting factor is applied to values. The critical rule for weighted averages (normalization) is that the sum of all weighting factors must equal one.

  • Natural abundance determines the weighting factors based on how common specific isotopes are in nature.

  • An isotope is defined as an atom with a specific mass number. Different isotopes of the same element have the same number of protons but a different number of neutrons.

Examples of Isotopes and Atomic Calculations

  • Hydrogen Isotopes:

    • Protium (Hydrogen-1): Has 1 proton and 0 neutrons. Mass number A=1A = 1.

    • Deuterium (Hydrogen-2): Has 1 proton and 1 neutron. Mass number A=2A = 2.

    • Tritium: A radioactive isotope that glows and was historically used in watch faces.

  • Chlorine Isotopes:

    • Chlorine-35 and Chlorine-37 exist naturally in a roughly 3:13:1 ratio (75%75\% to 25%25\%).

    • The weighted average results in the atomic mass of 35.4535.45.

    • If you were to pick a single atom from a sample of chlorine, you would never find one with a mass of exactly 35.4535.45; it would be either 3535 or 3737. The value 35.4535.45 represents the average mass of a large group of atoms.

  • Fluorine (FF):

    • Atomic Number (ZZ): 99

    • Atomic Mass: 18.99818.998

    • Mass Number (AA): 1919 (rounded from atomic mass)

    • Protons: 99

    • Neutrons: 1010 (calculated as 19−919 - 9)

    • Electrons: 99 (in a neutral atom, electrons equal protons because protons are +1+1 and electrons are −1-1).

  • Beryllium (BeBe):

    • Atomic Number (ZZ): 44

    • Atomic Mass: 9.0129.012

    • Mass Number (AA): 99 (rounded)

    • Protons: 44

    • Neutrons: 55 (calculated as 9−49 - 4)

    • Electrons: 44 (assuming neutrality).

Practical Application and Study Requirements

  • Rule for Mass Number: When specific isotope information is not provided, round the atomic mass from the periodic table to the nearest integer to find the most common mass number (AA).

  • Changing Identity vs. Charge:

    • Changing the number of protons changes the identity of the element.

    • Changing the number of electrons changes the charge (creating an ion).

    • Changing the number of neutrons results in a different isotope.

  • Polyatomic Ions: Preparation for future lessons requires memorizing common polyatomic ions from Table 2.5, including their names, formulas (symbols and numbers), and charges.

  • Upcoming Quiz Details:

    • Covers unit conversions: metric to English, area, and volume.

    • Specific conversion factors (e.g., pennyweights to ounces) will be provided, but the student must know how to perform the conversion.

    • Calculators must be non-graphing (e.g., TI-30X IIS).

Questions & Discussion

  • Question Regarding Smoke Detectors: How do smoke particles affect the sensor?

    • Answer: Smoke particles move between the eye and the emitter in a detector, blocking the eye from seeing the alpha particles, which triggers the alarm. The alarm also sounds if the battery is low.

  • Question Regarding Particle Mass: Which subatomic particle is the heaviest?

    • Answer: The neutron is the heaviest because it has the smallest negative exponent (10−2710^{-27}) compared to −31-31 and the highest numerical value (1.6751.675 vs 1.6731.673) among the heavy particles.

  • Question Regarding Chlorine Mass Probability: If I reach into a sack of chlorine atoms, what is the probability that I pull out one with exactly 35.4535.45 mass?

    • Answer: Zero probability. Every individual atom is either mass 3535 or mass 3737. The value 35.4535.45 is a statistical average for a group of atoms.

  • Discussion on Calculators: The speaker noted a preference for the TI-30X IIS and joked about the presence of a Casio user in every class.

  • Symbols of the Periodic Table Groups

    • Group 1: Alkali Metals
        Hydrogen (H)
        Lithium (Li)
        Sodium (Na)
        Potassium (K)
        Rubidium (Rb)
        Cesium (Cs)
        Francium (Fr)

    • Group 2: Alkaline Earth Metals
        Beryllium (Be)
        Magnesium (Mg)
        Calcium (Ca)
        Strontium (Sr)
        Barium (Ba)
        Radium (Ra)

    • Group 3-12: Transition Metals
        Scandium (Sc)
        Titanium (Ti)
        Vanadium (V)
        Chromium (Cr)
        Manganese (Mn)
        Iron (Fe)
        Cobalt (Co)
        Nickel (Ni)
        Copper (Cu)
        Zinc (Zn)
        Yttrium (Y)
        Zirconium (Zr)
        Niobium (Nb)
        Molybdenum (Mo)
        Technetium (Tc)
        Ruthenium (Ru)
        Rhodium (Rh)
        Palladium (Pd)
        Silver (Ag)
        Cadmium (Cd)
        Hafnium (Hf)
        Tantalum (Ta)
        Tungsten (W)
        Rhenium (Re)
        Osmium (Os)
        Iridium (Ir)
        Platinum (Pt)
        Gold (Au)
        Mercury (Hg)
        Rutherfordium (Rf)
        Dubnium (Db)
        Seaborgium (Sg)
        Bohrium (Bh)
        Hassium (Hs)
        Meitnerium (Mt)
        Darmstadtium (Ds)
        Roentgenium (Rg)
        Copernicium (Cn)

    • Group 13: Other Metals
        Thallium (Tl)
        Lead (Pb)
        Bismuth (Bi)
        Nihonium (Nh)

    • Group 14: Carbon Group
        Carbon (C)
        Silicon (Si)
        Germanium (Ge)
        Tin (Sn)
        Lead (Pb)
        Flerovium (Fl)

    • Group 15: Nitrogen Group
        Nitrogen (N)
        Phosphorus (P)
        Arsenic (As)
        Antimony (Sb)
        Bismuth (Bi)
        Moscovium (Mc)

    • Group 16: Chalcogens
        Oxygen (O)
        Sulfur (S)
        Selenium (Se)
        Tellurium (Te)
        Polonium (Po)
        Livermorium (Lv)

    • Group 17: Halogens
        Fluorine (F)
        Chlorine (Cl)
        Bromine (Br)
        Iodine (I)
        Astatine (At)
        Tennessine (Ts)

    • Group 18: Noble Gases
        Helium (He)
        Neon (Ne)
        Argon (Ar)
        Krypton (Kr)
        Xenon (Xe)
        Radon (Rn)
        Oganesson (Og)