Chapter 3

Atoms and the Periodic Table

Chapter 3 suggested problems: 42, 46, 48, 50, 54, 60, 62, 64, 66, 68, 72, 74, 78, 88, 92, 104

Interesting Links:

  1. Dalton's atomic theory

    1. Elements are made of atoms

    2. All of the atoms in an element are identical

    3. The atoms of an element are different from the atoms in any other element

    4. The atoms of two or more elements can combine to form compounds

    5. Law of constant composition: compounds always have the same ratios of atoms regardless of where the compound is found

    6. Atoms are not created or destroyed in chemical reactions

  2. The structure of the atom

    1. Atoms are made of three types of subatomic particles
       
      charge
      relative mass
      absolute mass (kg)
      proton
      +1
      1
      1.6726 x 10-27
      neutron
      0
      1
      1.6749 x 10-27
      electron
      -1
      1/1800
      9.109 x 10-31

    2. Protons and neutrons are held together in the nucleus at the center of the atom

      1. Despite mutual repulsion of like-charged particles they are held together by Strong Force

    3. Electrons orbit the nucleus relatively far away; the atom is mostly empty space

      1. Electrostatic attraction (i.e. electromagnetic attraction) occurs between the nucleus and orbiting electrons

      2. The number of electrons determines an atom's chemical reactivity

    4. The number of protons in the nucleus is unique for each type of atom (i.e., the atoms in each element); the number of protons in an element is invariable - if the number of protons changes, the element itself changes

      1. Atomic number - the top number for each element in the Periodic Table

      2. Abbreviated "Z"

    5. The number of neutrons in a nucleus can vary for each element

      1. Isotopes: substances with the same number of protons but different numbers of neutrons

        1. Atomic mass - the sum of the protons and neutrons in an isotope

        2. 126C 136C 146C

        3. 3517Cl 3717Cl

      2. Mass of individual atoms is expressed in amu - atomic mass units - 1/12 the mass of single a 12C isotope (1.661 x 10-27 kg)

      3. The bottom number in the Periodic Table for each element is the atomic weight i.e. the weighted average of the atomic masses of all of the isotopes

        1. 12C - 98.9%, 13C - 1.1%, 14C - ~0%
          (0.989)(12) + (0.11)(13) + (0)(14) = 12.011 amu

        2. 35Cl - 75.78%, 37Cl - 24.22%
          (0.7578)(35) + (0.242)(37) = 35.48 amu

        3. 50Cr - 4.345%, 52Cr - 83.789%, 53Cr - 9.501%, 54Cr - 2.315%
          (0.0.4345)(50) + (0.83789)(52) + (0.09501)(53) + (0.02315)(54) = 52.03 amu

    6. In their elemental state elements are electrically neutral, i.e., in each atom the #electrons = #protons

      1. Ions - atoms with an electrical charge

      2. Cations - have lost electrons, since (#electrons < #protons) have net positive charge

      3. Anions - have gained electrons, since (#electrons > #protons) have net negative charge

      4. Substance cannot gain protons, they can only gain or lose electrons

      5. Examples

        1. Mg and Mg2+

        2. As and As3-

        3. NH4+ and SO42- - molecular ions

      6. The Octet Rule and the concept of "isoelectronic"

        1. Isoelectronic: having the same number of electrons and hence, the same electron configuration, as another element

        2. The Octet rule: elements "want" to be isoelectronic with the nearest noble gas

        3. Elements become isoelectronic with the nearest noble gas by gaining or losing electrons (i.e. by forming cations or anions)

        4. The number of electrons an atom gains or loses depends on how far away it is from the nearest noble gas

    7. Atomic symbols - amanXcharge

  3. Radiation

    1. Ionizing - comes from radioactive materials (Uranium, plutonium, etc.), consists of atomic nuclei, electrons, or energy

    2. Electromagnetic - due to disturbance of electromagnetic fields - of very large items (stars) and very small items (atoms, nuclei)

      1. Electromagnetic radiation consists of small, massless packets of energy called photons that travel in waves

      2. Characteristics of waves

        1. Wavelength - crest to crest distance, can range from kilometers (radiowaves) to picometers (x-rays)

        2. Frequency - how many waves per second (Hz)

        3. Amplitude - how intense (high) the wave is

        4. c = (wavelength) x (frequency)

      3. Visible light is just a small fraction of the electromagnetic spectrum

        1. Radiofrequency (30 cm and greater)

        2. Microwave (3 mm- 30 cm)

        3. Infrared (1000 nm - 3 mm)

        4. Visible (400 nm - 800 nm)

        5. Ultraviolet (300 nm - 3 nm)

        6. X-rays and gamma rays (shorter than 3 nm)

  4. Fundamentals of electrons in atoms

    1. In neutral atoms (i.e. atoms in their elemental state) the number of protons is equal to the number of electrons

    2. Electrons do not orbit the nucleus in planetary fashion; the position of an electron is described in terms of statistical probability (as calculated using quantum mechanics)

    3. Electrons cannot exist just anywhere with respect to the nucleus; they can only be found at certain specific distances from the nucleus (as calculated using quantum mechanics)

    4. These distances correspond to energies; the further an electron is from the nucleus the greater its energy

    5. Areas of high probability are called orbitals

    6. The discrete distances at which electrons can be found from the nucleus can be broken down into shells, subshells, and orbitals

      1. Shell

        1. Describes distance from nucleus

        2. As n increases, distance increases

        3. The spacing between shells is not linear

      2. Subshell

        1. Each shell contains as many subshells as its number

        2. Describes shapes and orientations of orbitals

      3. The different orbitals are at roughly the same distance from the nucleus but not quite

      4. Any orbital can only hold a maximum of two electrons
        shell (n)
        subshell (l)
        # orbitals
        #electrons /orbital
        total electrons
        1
        s
        1
        2
        2
        2
        s, p
        1 + 3 = 4
        2
        8
        3
        s, p, d
        1 + 3 + 5 = 9
        2
        18
        4
        s, p, d, f
        1 + 3 + 5 + 7 = 16
        2
        32
        5
        s, p, d, f, g
        1 + 3 + 5 + 7 + 9 = 25
        2
        50
        6
        s, p, d, f, g, h
        1 + 3 + 5 + 7 + 9 + 11 = 36
        2
        72
        7
        s, p, d, f, g, h, i
        1 + 3 + 5 + 7 + 9 + 11 + 13 = 49
        2
        98

      5. Electrons can have one of two spin states - spin up or spin down

    7. Ground state and excited electrons

      1. Electrons in their lowest energy state are said to be in the ground state

      2. When an electron absorbs energy (either thermal or electromagnetic) it may become excited and can move to higher unoccupied orbitals

      3. g, h, i subshells are virtual and only contain excited electrons

      4. The difference in energy between the various energy levels of shells and subshells correspond to specific amounts of energy which in turn correspond to specific wavelengths of light (E = h x frequency; h = Plank's constant)

      5. Atomic spectra are based on these transitions and are unique for each element (e.g. see Spectra of Gas Discharges)
        Most common elements in solar spectrum listed in order of decreasing abundance (from Spectra of Gas Discharges )
        Atomic No. Element Emission Lines 4000-7000 Å
        1 5
        2 23
        3 24
        8 73
        6 27
        7 84
        10 75
        12 54
        14 109
        16 39
        26 235
        13 38
        20 78
        18 159
        11 90
        36 75
        38 54
        54 139
        56 92

  5. Electron configurations and the Periodic Table

    1. Ground state electron configurations - always start from the bottom and build up (three types: full configuration, Noble gas core, electron arrow diagrams)

      1. H

      2. He

      3. Li

      4. Mg

      5. B

      6. C

      7. N

      8. O

      9. F

      10. Ne

      11. 3rd period elements

    2. Note the arrangement of the Periodic Table with respect to electron configuration

      1. s, p, d, f blocks

      2. Periods and groups give the "address" of the element's highest electron

      3. No electrons have the same set of quantum numbers

    3. The electron configurations of ions: examples

    4. Inner shell and outer shell (valence) electrons

  6. Atomic properties and Periodic trends

    1. Atomic size

    2. Metallic character

    3. Ionization energy and electron affinity

    4. Electronegativity
[Chemistry 1110 Index Page]