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Modern Atomic Theory Review

  • light is a form of electromagnetic radiation

    • properties of both waves and particles:

      • wavelength (λ) - the distance between adjacent wave crests, meters

        • red light (750 NM) has longest wave length

        • violet light (400 NM) has shortest wave length

        • 1 NM = 1 * 10^-9 meters

      • frequency (v) - number of cycles or crests that pass through a stationary point in one second

      • amplitude - the height of the wave from zero to crest

    • wavelength and frequency are inversely/indirectly related - the shorter the wavelength, the higher the frequency

    • speed of light: 2.998 * 10^8 meters/second = λv

  • electromagnetic radiation

    • light can be viewed as a stream of particles

      • particle of light is a photon

      • photon - a single packet of light energy

        • has specific wavelength, determines what light we see

        • wavelengths of spectral lines are characteristics of the element

          • make up atomic emission spectra

        • no two elements have the same emission spectra

      • amount of energy carried in the packet depends on the wavelength of the light - the shorter the wavelength, the greater the energy

        • light waves that carry more energy in their crests are closer

        • violet light carries more energy per photon than red light

    • the photoelectric effect - the emission of electrons from a metal when light shines on the metal

      • quantum of energy - the minimum quantum of energy that can be lost or gained by an atom

      • quantized: an electron has to absorb/emit a specific amount of energy to move from one energy level to another

        • ground state: the normal energy level any given electron occupies

        • excited state: the energy level an electron occupies when it has absorbed the specific quantum of energy to move up to that level

      • Planck’s Law - E=Hv

        • E - energy, joules

        • H - Planck’s constant, 6.626 * 10^-34 J*S

        • v - frequency

  • Bohr’s Model

    • Niels Bohr changed Rutherford’s model to include newer discoveries about how the energy of an atom changes when the atom absorbs/emits energy

      • proposed electron is found only in specific circular paths/orbits around the nucleus ❌

        • incorrect - if the orbits were truly circular, the electron would spiral into the nucleus

      • each possible electron orbit has a fixed energy - energy level

      • each orbit is a specific distance from the nucleus and at each specific energy

        • impossible for an electron to exist between orbits

      • amount of energy is directly related to the frequency → wavelength

  • de Broglie: proposed “electrons be considered as waves confined to the space around an atomic nucleus”

  • Heisenberg Uncertainty Principle

    • Werner Heisenberg

    • states that it is impossible to determine simultaneously both the position and velocity of an electron

    • “we cannot know both the position and speed of a particle, such as a photon or electron, with perfect accuracy”

      • Schrödinger Wave Equation

        • Erwin Schrödinger developed an equation that treated electrons as waves

        • Quantum Theory - describes mathematically the wave properties of electrons

          • electrons exist in certain regions called orbitals

          • orbitals - 3D regions around the nucleus that indicate the probable location of an electron

            • represent probability maps showing a statistical attribution of where the electron is likely to be found

        • 4 Wave Properties

          • Energy Level: Principal Quantum Numbers - number specifying the principle shell of orbital

            • n - indicates the energy level

            • energy increases with principal quantum number

            • maximum of 7 energy levels

            • n^2 - how many orbitals in any energy level

            • 2n^2 - maxim. number of electrons possible in any energy level

          • Sub Level: Shapes of Quantum Mechanical Orbitals

            • letter indicates subshell of orbital, specifies shape

            • possible letters - s, p, d, f

            • electrons are more likely to be found closer to the nucleus than farther away

          • Orbital: Orientation

            • s - 1 orbital

            • p - 3 orbitals

            • d - 5 orbitals

            • f - 7 orbitals

          • Spin: clockwise or counterclockwise

ENERGY LEVEL

SUB-LEVEL

# ORBITALS (n^2)

ELECTRONS (2n^2)

n=1

1s

1

2

n=2

2s 2p

4

8

n=3

3s 3p 3d

9

18

n=4

4s 4p 4d 4f

16

32

  • Electron Configuration

    • arrangement of electrons in an atom and the way in which the electrons are arranged in various orbitals around the nucleus

    • Aufbau Principle

      • the electrons will fill the orbitals in a very specific order

      • lowest → highest energy

      • The Diagonal Rule

    • Pauli Exclusion Principle

      • an individual orbital may describe at most TWO electrons

      • in order to occupy the orbital, the two electrons must have opposite spins: ⬆⬇

    • EXAMPLES

      • Carbon 6e- : 1s^2, 2s^2, 2p^2

      • Aluminum 13e- : 1s^2, 2s^2, 2p^6, 3s^2, 3p^1

    • Noble Gas Configuration

      • Aluminum 13e- : [Ne] 3s^2, 3p^1

  • Hund’s Rule

    • orbitals in the same sub-level must all fill with one electron before a second electron is added to any of the orbitals: (n)p^4 - ⬆⬇ ⬆ ⬆

    • the “single” electrons will all have the same spin direction

    • Orbital Diagram

  • Valence and Core Electrons

    • valence electrons are the electrons on the outermost energy level

    • the noble gases always have full valence shells

    • Selenium 34e- : 1s^2, 2s^2, 2p^6, 3s^2, 3p^6, 3d^10, 4s^2, 3d^10, 4p^4

    • Silicon 14e- : 1s^2, 2s^2, 2p^6, 3s^2, 3p^2

    • the chemical properties of elements are largely determined by the number of valence electrons they contain

    • properties vary in a periodic fashion because the number of valence electrons is periodic

  • Atomic Physical Properties

    • Atomic Size ⬇⬅

      • left-right decreases ⬅

        • across a period, the amount of protons in the nucleus increases which has a stronger pull on the electrons, causing them to move closer to the nucleus

      • top-bottom increases ⬇

        • size of the orbital increases with increasing principal quantum shell number

        • electrons occupying the outermost orbitals are farthest from nucleus

    • Ionization Energy ⬆➡

      • amount of energy needed to remove a single electron

      • left-right increases ➡

        • electrical pull on electrons from the # protons in nucleus causes increases amount of energy needed

      • top-bottom decreases ⬆

        • electrons in outermost orbitals are less affected by the electrical pull from nucleus

    • Electronegativity ⬆➡

      • ability of an atom of an element to attract electrons when the atom is in a compound

      • left-right increases ➡

      • top-bottom decreases ⬆

Modern Atomic Theory Review

  • light is a form of electromagnetic radiation

    • properties of both waves and particles:

      • wavelength (λ) - the distance between adjacent wave crests, meters

        • red light (750 NM) has longest wave length

        • violet light (400 NM) has shortest wave length

        • 1 NM = 1 * 10^-9 meters

      • frequency (v) - number of cycles or crests that pass through a stationary point in one second

      • amplitude - the height of the wave from zero to crest

    • wavelength and frequency are inversely/indirectly related - the shorter the wavelength, the higher the frequency

    • speed of light: 2.998 * 10^8 meters/second = λv

  • electromagnetic radiation

    • light can be viewed as a stream of particles

      • particle of light is a photon

      • photon - a single packet of light energy

        • has specific wavelength, determines what light we see

        • wavelengths of spectral lines are characteristics of the element

          • make up atomic emission spectra

        • no two elements have the same emission spectra

      • amount of energy carried in the packet depends on the wavelength of the light - the shorter the wavelength, the greater the energy

        • light waves that carry more energy in their crests are closer

        • violet light carries more energy per photon than red light

    • the photoelectric effect - the emission of electrons from a metal when light shines on the metal

      • quantum of energy - the minimum quantum of energy that can be lost or gained by an atom

      • quantized: an electron has to absorb/emit a specific amount of energy to move from one energy level to another

        • ground state: the normal energy level any given electron occupies

        • excited state: the energy level an electron occupies when it has absorbed the specific quantum of energy to move up to that level

      • Planck’s Law - E=Hv

        • E - energy, joules

        • H - Planck’s constant, 6.626 * 10^-34 J*S

        • v - frequency

  • Bohr’s Model

    • Niels Bohr changed Rutherford’s model to include newer discoveries about how the energy of an atom changes when the atom absorbs/emits energy

      • proposed electron is found only in specific circular paths/orbits around the nucleus ❌

        • incorrect - if the orbits were truly circular, the electron would spiral into the nucleus

      • each possible electron orbit has a fixed energy - energy level

      • each orbit is a specific distance from the nucleus and at each specific energy

        • impossible for an electron to exist between orbits

      • amount of energy is directly related to the frequency → wavelength

  • de Broglie: proposed “electrons be considered as waves confined to the space around an atomic nucleus”

  • Heisenberg Uncertainty Principle

    • Werner Heisenberg

    • states that it is impossible to determine simultaneously both the position and velocity of an electron

    • “we cannot know both the position and speed of a particle, such as a photon or electron, with perfect accuracy”

      • Schrödinger Wave Equation

        • Erwin Schrödinger developed an equation that treated electrons as waves

        • Quantum Theory - describes mathematically the wave properties of electrons

          • electrons exist in certain regions called orbitals

          • orbitals - 3D regions around the nucleus that indicate the probable location of an electron

            • represent probability maps showing a statistical attribution of where the electron is likely to be found

        • 4 Wave Properties

          • Energy Level: Principal Quantum Numbers - number specifying the principle shell of orbital

            • n - indicates the energy level

            • energy increases with principal quantum number

            • maximum of 7 energy levels

            • n^2 - how many orbitals in any energy level

            • 2n^2 - maxim. number of electrons possible in any energy level

          • Sub Level: Shapes of Quantum Mechanical Orbitals

            • letter indicates subshell of orbital, specifies shape

            • possible letters - s, p, d, f

            • electrons are more likely to be found closer to the nucleus than farther away

          • Orbital: Orientation

            • s - 1 orbital

            • p - 3 orbitals

            • d - 5 orbitals

            • f - 7 orbitals

          • Spin: clockwise or counterclockwise

ENERGY LEVEL

SUB-LEVEL

# ORBITALS (n^2)

ELECTRONS (2n^2)

n=1

1s

1

2

n=2

2s 2p

4

8

n=3

3s 3p 3d

9

18

n=4

4s 4p 4d 4f

16

32

  • Electron Configuration

    • arrangement of electrons in an atom and the way in which the electrons are arranged in various orbitals around the nucleus

    • Aufbau Principle

      • the electrons will fill the orbitals in a very specific order

      • lowest → highest energy

      • The Diagonal Rule

    • Pauli Exclusion Principle

      • an individual orbital may describe at most TWO electrons

      • in order to occupy the orbital, the two electrons must have opposite spins: ⬆⬇

    • EXAMPLES

      • Carbon 6e- : 1s^2, 2s^2, 2p^2

      • Aluminum 13e- : 1s^2, 2s^2, 2p^6, 3s^2, 3p^1

    • Noble Gas Configuration

      • Aluminum 13e- : [Ne] 3s^2, 3p^1

  • Hund’s Rule

    • orbitals in the same sub-level must all fill with one electron before a second electron is added to any of the orbitals: (n)p^4 - ⬆⬇ ⬆ ⬆

    • the “single” electrons will all have the same spin direction

    • Orbital Diagram

  • Valence and Core Electrons

    • valence electrons are the electrons on the outermost energy level

    • the noble gases always have full valence shells

    • Selenium 34e- : 1s^2, 2s^2, 2p^6, 3s^2, 3p^6, 3d^10, 4s^2, 3d^10, 4p^4

    • Silicon 14e- : 1s^2, 2s^2, 2p^6, 3s^2, 3p^2

    • the chemical properties of elements are largely determined by the number of valence electrons they contain

    • properties vary in a periodic fashion because the number of valence electrons is periodic

  • Atomic Physical Properties

    • Atomic Size ⬇⬅

      • left-right decreases ⬅

        • across a period, the amount of protons in the nucleus increases which has a stronger pull on the electrons, causing them to move closer to the nucleus

      • top-bottom increases ⬇

        • size of the orbital increases with increasing principal quantum shell number

        • electrons occupying the outermost orbitals are farthest from nucleus

    • Ionization Energy ⬆➡

      • amount of energy needed to remove a single electron

      • left-right increases ➡

        • electrical pull on electrons from the # protons in nucleus causes increases amount of energy needed

      • top-bottom decreases ⬆

        • electrons in outermost orbitals are less affected by the electrical pull from nucleus

    • Electronegativity ⬆➡

      • ability of an atom of an element to attract electrons when the atom is in a compound

      • left-right increases ➡

      • top-bottom decreases ⬆

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