Chem Exam 4
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88 Terms
Relates to the size and energy of the orbital
Represented as a positive whole number, cannot be 0
As this value increases, the size and energy of the orbital increase
Relates to the shape of the orbital
Must be a positive whole number, can be zero but cannot be greater than n-1
L number and corresponding letters: 0 → s, 1 → p, 2 → f, 4 → g
Provides the orientation of the orbital
Ranges from -l to +l, it can be positive, negative, or zero
Provides the orientation of the electron spin
Can be -½ (spin down) or +½ (spin up)
Have 2 lobes, like dumbbells
Have 3 different orientations (vertical, horizontal, diagonal)
They have 4 lobes
5 different orientations, 1 looks like a ring around a stick
Zero probability of finding electron density
Total number of nodes in different orbitals can be defined using n-1
Cr, Mo, Cu, Ag, Au
One of the s electrons moves up to the d subunit
Not excited state configurations, they are ground state with increased stability
An electron transitions from a lower energy to higher energy orbital or a neighbor degenerate orbital
Ex: N ground state: 1s2, 2p2, 2p3 N excited state: 1s2, 2p2, 2p2, 3d1
Take away from highest n value
If multiple orbitals have same n value, start by taking away electrons from highest n and l value
Increases down a group because the n value of the electrons in the outermost orbital increases
Decreases across a period because of an increase in effective nuclear charge
Anion > neutral atom > cation
Increasing electrons increase electrons repulsion leading to less attraction to nucleus and expansion of the size of the ion
Atoms and ions with identical number of electrons
Look at number of protons, higher protons = smaller ionic radii because Zeff creates more positive charge and constricts size
The ease of distorting the electron density of an atom/ion
The larger the atom/ion size, the greater the polarizability
The change in energy when gaseous ions are packed together to form an ionic solid
M+ (g) + X- (g) → MX (s)
The energy required to separate one mole of a crystalline ionic solid into its gaseous ions
MX (s) → M+(g) + X-(g)
As ionic radii increases, the release of energy decreases (less negative, less favored) because ions cannot get as close to each other
Decrease in magnitude/strength of lattice energy
Increase in ion charge results in an increase of release of energy (more favored, more negative)
Increase in magnitude/strength of lattice energy. Ion charge has a more significant role than ion size
The energy required to remove an electron from an atom or ion in the gaseous state
The energy to remove the first electron is the first ionization energy, the second electron is the second ionization energy
Reported as a positive value since energy is required to remove as electron
General equation: X (g) → X+(g) + e-
First ionization energy: Li (g) → Li+(g) + e-
Second ionization energy: Li+(g) → Li2+(g) + e-
Decreases down a group because n increases
Increases across a period because Zeff increases
Switch groups 2 and 13 → Group 2 = higher because of full s subshell, group 13= lower since electron in p is easier to remove
Switch groups 15 and 16 → group 15= higher since half filled subshell is stable, group 16= lower since p electron is easier to remove because it reduces repulsions and results in half filled shell
The energy associated with adding an electron to a gaseous atom
Typically a negative number since energy is released as a result of attraction between incoming electron and nucleus (Zeff)
X(g) + e- → X-(g)
Becomes more negative across a period because Zeff increases so more energy is being releases
Becomes more positive down a group because orbital size increased and the electron has less interaction with nucleus and less energy is released
Switch groups 1 and 2 → group 1 = more negative because of full s shell, group 2= more positive because adding electron goes to higher subshell
Switch groups 14 and 15 → group 14 = more negative because adding electrons gives half-full p shell, group 15= more positive because adding electrons causes repulsions and less interactions with the nucleus
The energy required to remove an electron from an anion with a -1 charge which is opposite of electron attachment enthalpy
Some resources call electron attachment enthalpy electron affinity
Increases across a period and decreases down a group
N, O, F, and Cl have specially high values
Noble gases excluded
Fluorine = most electronegative and francium = least electronegative
Hydrogen exists between B and C
Start filling in electrons first then pair them
Unpaired electrons are likely to form bonds
Draw lewis symbol
Take away or add electrons
Put chemical symbol in brackets and write charge on the outside
2 electrons shared between 2 atoms
1 double bond = 2 bonding pairs
Write general skeletal structure, central atom is the least electronegative, hydrogens are always terminal
Distribute electrons giving octets to as many as possible
Polyatomic ions follow same rule but must have brackets surrounding it
An atom surrounded by 8 electrons
Hydrogen may only have 2 electrons in their valence shells
Same rules as covalent
C is usually central atom and forms 4 bonds
Take note of how molecules are written
Have an unpaired electron pair in their lewis structures making them very reactive
Count valence electrons to see if its an odd number, if compound has a dot next to it means there is an unpaired electron
Molecules from with atoms that do not have a full octet of electrons
Typically seen in beryllium, boron, and aluminum
Only seen in elements of the 3rd period
Can hold more than 8 electrons
Charge assigned to atoms in a lewis structure assuming bonding electrons are shared equally between atoms
Assumes an electron own all of its bonding electrons and one half of its bonding electrons
Sum of all formal charges in neutral molecules should be 0
Sum of all formal charges in ion should be equal the charge of ion
More than 1 valid lewis structure
Atom arrangement is always the same, only electron density changes
Formal charges closer to 0 on individual atoms are between than smaller/larger ones
Negative charges should reside on most electronegative charges and positive charges on the less electronegative atoms
Bond order of single bond = 1
Bond order of double bond = 2
Bond order of triple bond = 3
As bond order increases, strength increases and length decreases
Single bonds are the longest and weakest while triple bonds are the shortest and longest
More electrons shared results in more overlap between electron clouds and shorter bonds
More electrons shared required more energy to break so it is stronger
3 energies- nuclei-electron attractive forces, electron-electron repulsive forces and nuclei-nuclei repulsive forces
The bottom of the curve represents bond length and strength