Unit 1 - Structure and Properties Flashcards
Principal Quantum Number
-indicates relative size and energy of atomic orbitals
-represents energy level, or shells, the electrons can occupy in an atom
-as n increases, orbital size increases
-max number of electrons per energy level: 2n^2
Orbitals (definition, different quantum numbers)
represent probability clouds that show a statistical distribution of where the electron is likely found
Described by 4 quantum numbers:
-principal quantum #
-angular momentum (secondary) quantum #
-magnetic quantum #
-spin quantum #
Angular Momentum (secondary) Quantum Number (definition, how many electrons held by each shape)
-energy sublevels are contained within the principal energy levels
-gives shape of the subshell (s,p,d,f)
-L = 0 up to (n-1)
s orbitals: L=0, 1 s orbital per subshell, can hold total of 2e-
p orbitals: L=1, 3 p orbitals per subshell, can hold total of 6e-
d orbitals: L=2, 5 d orbitals per subshell, can hold total of 10e-
f orbitals: L=3, 7 f orbitals per subshell, can hold total of 14e-
energy increases from s->f orbitals as they get further from nucleus
Magnetic Quantum Number
-represents the orientation of the subshells
-has values of -L to +L
Spin Quantum Number
-represents the spin of the electron
-has values of +1/2 or -1/2
Heisenberg Uncertainty Principle
it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time
Pauli Exclusion Principle
-No two electrons in an atom can have the same FOUR quantum numbers
-means only e- of opposite spin can occupy an orbital
Aufbau Principle (both parts)
- Place electrons into the orbitals in order of increasing energy level
- Each subshell of orbitals in the same energy level must be completely filled before proceeding to the next orbital or series of orbitals
Hund’s Rule (both parts)
- Whenever electrons are added to orbitals of the same energy sublevel, each orbital recieves one electron before pairing occurs
- When electrons are added singly to seperate orbitals of same energy level, the electrons must all have the same spin
Ground State vs Excited State of electron
Excited state is when valence electron absorbs energy, commonly in the form of jumping from its ground state (original energy level or orbital) to an empty orbital of a higher energy shell
Isoelectronic Ions
-Atoms and ions that have the same number of electrons
-elements coming before and after noble gases form isoelectronic ions by gaining or losing electrons until they have the same # of electrons as the noble gas
Valence Shell Electron Pair Repulsion Theory/VSEPR (definition, key ideas)
Allows us to predict the geometry of molecules based on the repulsion of electron pairs. There are five basic geometries, and 10 subtypes for if lone pairs are present, total 15 geometries.
Key Ideas:
1. e- pairs (both bonding and lone pairs around central atom) repel each other electrostatically
2. The molecular shape is determined by the positions of the electron pairs when they are a maxium distance apart
Types of Electron Pairs
-Bonding pairs or lone pairs (non-bonding pairs)
-If all e- pairs are identical, then the molecular shape is the same as the e- pair arrangement (CCl4)
-If all e- pairs are NOT identical, then e- pair arrangements is only an approximation of the shape of the molecule (CCl3F)
—Three reasons for this: lone pairs occupy more space, space occupied by bonding pairs related to EN of the atom, double and triple bonds occupy more space
Ionic Bond
-Metal and non-metal atom
-TRANSFER of electrons from metal to non-metal
-cation and anion
-ΔEN is 1.7 and above
Polar Covalent Bond
-electrons are shared UNequally because 1 atom attracts them more strongly than the other atom
-dipole created, region of slightly positive charge, region of slightly negative charge
-ΔEN is 0.5 to anything LESS THAN 1.7
Non-Polar Covalent Bond
-electrons are shared EQUALLY between the atoms
-ΔEN is anything LESS THAN 0.5
Electronegativity
defines how strongly an atom attracts a pair of electrons it shares with another atom in a covalent bond
Criteria for entire molecule to be polar
More than one polar bonds and they DO NOT cancel each other out
INTRAmolecular Forces (definition, characteristics)
-forces WITHIN a molecule (ie bonds)
-can be formed or broken by CHEMICAL changes
-influence CHEMICAL properties
-relatively STRONG
INTERmolecular Forces (definition, characteristics)
-forces BETWEEN molecules
-can be formed or broken by PHYSICAL changes
-influence PHYSICAL properties (ie melting point, state, etc)
-relatively WEAK
Dipole-Dipole Forces (definition, strength, traits)
Caused by: dipoles of POLAR molecules that position their positive and negative ends near each other
Experienced by: Polar molecules
Strength depends on: distance between dipoles
Increase as bond polarity increases (compare ΔEN)
Hydrogen Bonds (definition, strength)
Caused by: Hydrogen atom bonded to a highly electronegative atom (O, N, F) attracted to a partially negative atom
Experienced by: water, ammonia, hydrogen flouride
Strength depends on: polarity of bonds and closeness of molecules
London/Dispersion Forces (definition, strength, traits)
Caused by: non-polar molecules electrons are in motion and produce temporary “dipole-like” arrangement of charge. Instantaneous dipoles created by movement of e-
Experienced by: All substances, only mentioned if other forces are absent
Strength depends on: Molar Mass
Increase as number of electrons increases, equal for isoelectronic molecules, NON-POLAR molecules ONLY experience LDF
Rank of Intermolecular Forces Strength, Other Qualities (ranking, m.p, and b.p)
Strength Ranking: Ionic>H bonding>Dipole-Dipole>LDF
The stronger the IMF, the greater the energy required to overcome the forces, and therefore increased melting and boiling point.
“like dissolves like” also still applies
