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Chapter 3 - Atoms, Electrons, and Periodic Table

Chapter 3.1: The Nuclear Atomic Model

  • John Dalton’s atomic theory: atoms are tiny, solid, and indestructible spheres

    • No longer applies

      • Atoms can be divisible

  • Dmitri Mendeleev’s periodic table: listed known elements increasing atomic mass

    • Similar chemical properties were listed in the same column

    • The theory has undergone many changes

  • Joseph John Thomson: discovered the existence of a negatively charged particle (known today as an electron)

    • His model of the atom contained electrons, but the entire sphere carried a uniform positive charge

  • Ernest Rutherford: discovered that the atom is mainly made up of empty space but has a concentrated positive charge in the center

    • This is also known as the discovery of the atomic nucleus

    • Nuclear model: diagram of an atom that pictures electrons in motion around the atomic nucleus

      • Also known as the planetary model

  • The Discovery of neutron help explain why the atomic nucleus did not fly apart or what the total mass of an atom was

  • Niels Bohr: developed a similar atomic model to Rutherford but helped explained that electrons had certain restrictions:

    • Atom has only specific energy levels, called stationary states

    • While in a stationary state atoms do not emit energy

    • Atom can change stationary states when it emits or absorbs specific quantities of energy that is equal t the difference in energy between two stationary states

  • Planck suggested that matter, at the atomic level, can absorb or emit only discrete quantities of energyQuantum of energy: these specific quantities

  • Photons: particle-like energy, (ex. light)

  • Emission spectra: when atoms in an excited state emit photons when they fall to lower energy levels

  • Absorption spectrum: results when atoms absorb photons of certain wavelengths and are excited from lower to higher energy levels

Chapter 3.2: The Quantum Mechanical Model of the Atom

  • Quantum mechanical model of the atom: a model of the atom that described atoms as having wave-like properties

  • Uncertainty principle: impossible to know the position and momentum of an object beyond a certain measure of precision

  • Orbitals: wave functions that give information about an electron’s energy and location within an atom

  • Ground state: electron-density probabilities for the lowest energy level

  • Quantum numbers: Principle quantum number (n): a positive whole number that describes energy level of atomic orbital and relative size Larger n value means larger energy level and electron is farther away from the nucleus Second quantum number/Orbital-shape quantum number (l): describes the orbital shape, a positive integer from 0 to (n-1); refers to to the energy sublevels within each principal energy level;l=0: sl=1: p l=2: dl=3: fTo find energy sublevel/type of orbital, you combine principle quantum number with the letter of orbital shape (s,p,d, or f) Third Quantum number/Magnetic quantum number (m1): values ranging from -1 to +1 and indicates the orientation of orbital in the same around the nucleus

Chapter 3.3: Electron Configurations and Periodic Trends

  • Fourth Quantum number/Spin quantum number (ms): specifies the direction of where the electron is spinning

    • Can either be +½ or -½

  • Pauli exclusion principle: only two electrons of opposite spin could occupy an orbital

    • No two electrons in an atom have the same four quantum numbers

  • Electrons do not actually occupy orbitals, instead, calculate probability densities for electrons

  • Electron configuration: shorthand notation of number and arrangement of electrons

    • Atom’s chemical property associated with ground-state electron configuration

  • Aufbau principle: build up the ground state electronic structure for each atom

  • Orbital diagram: uses a box for each orbital in each given principal energy level

    • Empty box: orbital with no electrons

    • Arrow up/down: orbital with an electron that spins in that direction

    • Two arrows mean filled orbital

  • Hund’s rule: no two electrons can have the same set of quantum numbers

    • The orbital diagram must have at least one before adding a second electron to the orbital

  • Some elements have different electron configurations from predicted

    • Cr, Cu have electron configurations that result in more stable electrons, so 4s has 1 electron filled instead of 2

  • Atomic radius: measurable property between nuclei of bonded

    • Molecules atomic radius is half the distance between nuclei of identical atoms that bonded together

    • As n increases, higher probability of finding electrons far away from the nucleus since atomic volume is larger

    • Effective nuclear charge: the net force of attraction between electrons and the nucleus

      • Down group: inner electrons increase which increases shielding effect

      • Across period: as you go across decrease in atomic size results in effective nuclear charge increasing

  • Ionization energy: the energy needed to remove one electron from ground state gaseous atom

    • Decreases as you go down a group due to the atomic radius (increase in distance between nucleus and valence electrons)

    • Increases across a period due to the atomic radius (effective nuclear charge)

Electron affinity: change in energy that occurs when an electron is added to a gaseous atom

Chapter 3.1: The Nuclear Atomic Model

  • John Dalton’s atomic theory: atoms are tiny, solid, and indestructible spheres

    • No longer applies

      • Atoms can be divisible

  • Dmitri Mendeleev’s periodic table: listed known elements increasing atomic mass

    • Similar chemical properties were listed in the same column

    • The theory has undergone many changes

  • Joseph John Thomson: discovered the existence of a negatively charged particle (known today as an electron)

    • His model of the atom contained electrons, but the entire sphere carried a uniform positive charge

  • Ernest Rutherford: discovered that the atom is mainly made up of empty space but has a concentrated positive charge in the center

    • This is also known as the discovery of the atomic nucleus

    • Nuclear model: diagram of an atom that pictures electrons in motion around the atomic nucleus

      • Also known as the planetary model

  • The Discovery of neutron help explain why the atomic nucleus did not fly apart or what the total mass of an atom was

  • Niels Bohr: developed a similar atomic model to Rutherford but helped explained that electrons had certain restrictions:

    • Atom has only specific energy levels, called stationary states

    • While in a stationary state atoms do not emit energy

    • Atom can change stationary states when it emits or absorbs specific quantities of energy that is equal t the difference in energy between two stationary states

  • Planck suggested that matter, at the atomic level, can absorb or emit only discrete quantities of energyQuantum of energy: these specific quantities

  • Photons: particle-like energy, (ex. light)

  • Emission spectra: when atoms in an excited state emit photons when they fall to lower energy levels

  • Absorption spectrum: results when atoms absorb photons of certain wavelengths and are excited from lower to higher energy levels

Chapter 3.2: The Quantum Mechanical Model of the Atom

  • Quantum mechanical model of the atom: a model of the atom that described atoms as having wave-like properties

  • Uncertainty principle: impossible to know the position and momentum of an object beyond a certain measure of precision

  • Orbitals: wave functions that give information about an electron’s energy and location within an atom

  • Ground state: electron-density probabilities for the lowest energy level

  • Quantum numbers: Principle quantum number (n): a positive whole number that describes energy level of atomic orbital and relative size Larger n value means larger energy level and electron is farther away from the nucleus Second quantum number/Orbital-shape quantum number (l): describes the orbital shape, a positive integer from 0 to (n-1); refers to to the energy sublevels within each principal energy level;l=0: sl=1: p l=2: dl=3: fTo find energy sublevel/type of orbital, you combine principle quantum number with the letter of orbital shape (s,p,d, or f) Third Quantum number/Magnetic quantum number (m1): values ranging from -1 to +1 and indicates the orientation of orbital in the same around the nucleus

Chapter 3.3: Electron Configurations and Periodic Trends

  • Fourth Quantum number/Spin quantum number (ms): specifies the direction of where the electron is spinning

    • Can either be +½ or -½

  • Pauli exclusion principle: only two electrons of opposite spin could occupy an orbital

    • No two electrons in an atom have the same four quantum numbers

  • Electrons do not actually occupy orbitals, instead, calculate probability densities for electrons

  • Electron configuration: shorthand notation of number and arrangement of electrons

    • Atom’s chemical property associated with ground-state electron configuration

  • Aufbau principle: build up the ground state electronic structure for each atom

  • Orbital diagram: uses a box for each orbital in each given principal energy level

    • Empty box: orbital with no electrons

    • Arrow up/down: orbital with an electron that spins in that direction

    • Two arrows mean filled orbital

  • Hund’s rule: no two electrons can have the same set of quantum numbers

    • The orbital diagram must have at least one before adding a second electron to the orbital

  • Some elements have different electron configurations from predicted

    • Cr, Cu have electron configurations that result in more stable electrons, so 4s has 1 electron filled instead of 2

  • Atomic radius: measurable property between nuclei of bonded

    • Molecules atomic radius is half the distance between nuclei of identical atoms that bonded together

    • As n increases, higher probability of finding electrons far away from the nucleus since atomic volume is larger

    • Effective nuclear charge: the net force of attraction between electrons and the nucleus

      • Down group: inner electrons increase which increases shielding effect

      • Across period: as you go across decrease in atomic size results in effective nuclear charge increasing

  • Ionization energy: the energy needed to remove one electron from ground state gaseous atom

    • Decreases as you go down a group due to the atomic radius (increase in distance between nucleus and valence electrons)

    • Increases across a period due to the atomic radius (effective nuclear charge)

Electron affinity: change in energy that occurs when an electron is added to a gaseous atom