Atomic Structure and the Periodic Table Flashcards
Define “n”
N is the quantum number describing energy. For example, the 1 is 1s, or the 2 in 2p. It goes in descending order.
Define “l”
L is the shape of the orbital, and can be 0 to n-1. L=0 is the s orbital, and so on.
Define “m of l”
M of l is the rotation of the shape. It is -l, or positive l. For example, if l is 2, then m of l can be -2, -1, 0, 1, or 2.
Define M of s
This is the spin. 1/2, or -1/2. Nothing more than that. Electrons have to have opposite spins.
The Aufbau Principle
When an atom is acquiring electrons they are placed in orbitals (sub-shells) of descending energy. This means that the lower energy levels are filled before the higher energy levels, as they are closer to the nucleus.
The Pauli Exclusion Principle
No TWO electrons can have the same quantum numbers.
Hund’s Rule
When electrons are added to a sub shell, it’ll occupy an empty orbital. In other words, each orbital will at least have one if at least one doubles up.
Diamagnetism
All the electron’s spins are paired, all subshells are filled. Since there are no electrons that would be attracted to something, the element would not be affected by magnetic fields.
Paramagnetism
The element does not have all of its orbitals paired, and a subshell is incomplete. Since a stray electron is noted, there is going to be considerable effects from magnetic fields.
Electrons and Energy
As the electrons grow farther in distance from the nucleus, their potential energy grows.
Planck’s Constant Conversion
kg x m^2/sec^2
Isoelectronic Elements (Ions)
Have the same configuration as the nearest noble gas, metals lose electrons, nonmetals gain electrons.
Thompson’s Experiment
Cathode Ray Experiment: Provided evidence of electrons, and provided the basis of a model for the atom with protons and electrons.
Rutherford’s Gold Foil Experiment
Rutherford fired alpha particles at a thin sheet of gold foil, new experiment showed that the nucleus was in the middle.
Describe Atomic Radius Trend in Relation to Periodic Table
The Atomic Radius decreases as one moves to the right along a period (Effective Nuclear Charge), and increases as you move down a family/group (Nuclear shielding).
Effective Nuclear Charge
Filling the outermost “s” and “p” orbitals, meaning that the nucleus is going to pull those electrons in closer and decrease the radius.
Nuclear Shielding
More energy levels, with electrons occupying the orbitals, making the atomic nucleus increase.
Ionic Radius
The “more positive the charge”, the smaller the radius. The more negative a charge, the more larger the radius. Decreases going across period, increases top to bottom.
Ionization Energy
The energy needed to remove one mole of electrons from one mole of atoms.
Describe Ionization Energy in Relation to Periods
Increases left to right (metals towards nonmetals). Metals want to lose electrons, and nonmetals want to gain electrons to become anions (with s and p orbital filled up)
Describe Ionization Energy in Relation to Groups
Decreases as one moves down a group. The addition of energy levels lessens the hold on the outer “s” and “p” electrons, making them more easy to remove. In other words, the nucleus doesn’t have as much of an attraction to the electrons because there’s a distance between them.
Electron Affinity
An atoms attraction to electrons, metals generally want to lose electrons, having a low affinity for “additional” electrons. Nonmetals want to gain electrons, and have a higher affinity.
Trend of Affinity (Highest and Lowest)
Fluorine has the highest affinity (electronegative value), and Francium has the lowest.
Electronegativity
Is the pull of the nucleus of one atom on the electrons of other atoms. Moving left to right increases electronegativity, and moving up to down decreases the electronegativity.
Covalent Bonds
Bonds between atoms that are shared. Only nonmetals.
Polar Covalent Bonds
When there is a difference in electronegativity between two different nonmetals. Unequal sharing, normally between two different nonmetals.
Non-Polar Covalent Bonds
No difference in electronegativity, usually between same atoms. Equal sharing, for example N2.
Ionic Bonds
Always between a nonmetal and a metal, metals give their valence electrons to the nonmetal, and the two are held together with electrostatic attraction. Large difference in in electronegativity.
Valence Electron (List how many are in Na)
An electron of an atom, located in the outermost shell (valence shell) of the atom, that can be transferred to or shared with another atom. Na has 1 valence electron.
