
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 energy; Quantum 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 energy; Quantum 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