TTU Chem 1307 Exam 3 Notecards

Sublimation Energy

energy require to change state from solid to gas directly

Angular Momentum Number (l)

An integer range from 0-(n-1) and indicates shape of orbital, multiple possible values of l

ex: if n=3, l= 0,1,2.
if n=1, l=0.

Aufbau Principle

e- occupy the lowest energy orbitals first (1s2, then 2s2, then 2p6, then 3s2...etc.)

Pauli Exclusion Principle

1. no two electrons in the same atom can have the same set of four quantum numbers
2. No more than 2 e- can occupy an orbital
3. e- that occupy same orbital must have different values for ms (ms= +1/2 and ms= -1/2)

Periodic Trends: Atomic Radius

Highest in lower left corner, lowest in top right corner

Lattice Energy

(only have to know definition) The enthalpy change that takes place when a gaseous ion forms a solid ion

Execptions to Electron Configuration

Copper - 4s1 3d10
Chromium - 4s1 3d5

Change of Enthalpy (ΔH)

Used for heat transfer/capacity questions. Quantifies heat flow in or out of system.
ΔH = H(products) - H(reactants)

Exothermic

Heat flows out of system (-ΔH)

Endothermic

Heat flows into system (+ΔH)

Heat calculation formula

q=mc∆T
q=heat absorbed or released by substance
m=mass
c=specific heat capacity
ΔT= change in temp. (final-initial)

Specific Heat of H2O

c=4.184

Calorimetry Solving Process

q=mcΔT, slove for water first, then use the q for the alloy equation. Flip sign (-/+) at the end of solving process.

Bomb Calorimetry

energy of reaction = ΔT - Heat capacity of Calorimeter

Hess's Law

using two reactions to achieve desired equation
rules: 1.if you reverse equation, reverse sign of ΔH associated with that equation, 2. if you multiple/divide equation, divde/multiple ΔH by the same

Standard Enthalpy of Formation

ΔH(rxn) = Sum of ΔH products (ΔHfproducts) - Sum of ΔH reactants (ΔHfreactants)
element in standard form (Xe, H2, etc.) has an sum of ΔHf of zero
Most compounds have -ΔHf

Wavelength (λ)

The distance between any point on a wave and the corresponding point on the next crest(or trough) of the next wave.

Frequency (v)

The number of cycles of the wave in a give amount of time (ex:5.65x1o^14 Hz)

Amplitude

The height of crest, or depth of trough, of a wave from the middle line.

c = v x λ

Speed of light = frequency x wavelength

Nanometers --> Meters conversion

10^-9

Speed of light (c)

3.0 x 1o^8 m

E = hv = hc/λ

energy of light (J)
h= planks constant
c= speed of light

Plank's Constant (h)

6.626 x 10^-34

Concept of Photoelectric Light

electrons are emitted as electromagnetic radiation(eg.light) hits the surface of a metal

De Broglie's Equation

λ=h/mv
*mass is in Kg here, only here

Bohr's Model of the Atom

"electron cloud model" Electrons occupy a definite amount of orbitals that require specific energies to occupy.

Emission

Electrons moving from a higher shell to a lower shell emit energy(light)

Absorption (Evolving) (Excited)

Electrons moving from a lower energy shell to a higher shell absorb energy to do so.

Energy Level Formula

En = RH(1/n^2)
n = energy level outmost e is occupying, quantum number
RH = Rydberg's Constant

Rydberg's Constant

-2.178 x 10^-18

Heisenberg Uncertainty Principle

It is impossible to simultaneously know an electron's location AND speed of orbit

Principle Quantum Number (n)

a positive integer(1,2,3..) indicating size and relative distance from nucleus, specifies the energy level

Magnetic Quantum Number (ml)

Integer range from -l to 0 to +l, indicates orientation of the orbital in the space around the nuclues
ex: if l=2, ml= -2, -1, 0, 1, 2.

Level (Shell)

denoted by (n), defined further by subshells

Sublevel (Subshell)

designates the orbital shape, denoted by l quantum number
l= 0 -- s sublevel
l= 1 -- p sublevel
l= 2 -- d sublevel
1= 3 -- f sublevel

Orbital shapes

s = spherical shaped
p = dumbbell shaped
d = clover leaf shaped

Hund's Rule

when filling orbitals, e- go into separate orbitals with parallel spins until all of the orbitals are occupied by one e-, then they pair up.

Valence Electrons

Electrons in the outermost energy level of an atom
ex: [Ne] 3s23p4 (everything following [Ne] is valence b/c it is not full)

Periodic Trends: Ionization Energy

Lowest IE in left bottom corner, highest in right upper corner.

Periodic Trends: Electron Affinity

Amount of energy required to add an e- to a gaseous ion/atom
Lowest in bottom left corner, highest in top right corner

Ionic Bonds

Metal and nonmetal, >1.7 electronegativity difference

Polar Covalent Bond

Nonmetals, 0.7-1.7 electronegativity difference

(non-polar) Covalent Bond

Nonmetals >0.7 electronegativity difference

Electronegativity

The ability of an atom in a molecule to attract shared electrons to itself
ex: H=2.1
B=2.0
C=2.5
N=3.0
O=3.5
S=2.5
F=4.0

Bond Polarity

Difference in EN, results from an unequal distribution of electrons

Molecular Polarity (dipole moement)

When the entire atom has a partially positive and partially negative charge.

Indicated by lewis structure, symmetric=nonpolar

Isoelectric Ions

Ions with the same number of electrons in the configurations after incorporating ionic charges. ex. Al^3+ and O^2- are isoelectric

Ion size decreases as atomic number(z) increases when comparing isoelectric ions

Ion size rules

1. Cations (positive ion) is smaller than parent atom
2. Anions (negative ions) are larger than parent atom (more e-)
3. Size generally decreases down a group

Dissociation Energy

energy required to break apart a compound (breaks bond)

Heat of formation

heat energy released or evolved when an electron is added to a neutral atom(Hf)

Exceptions to Octet rule

more than octet- Anything after P can have more than 8 e- if it is the central atom

Less than octet- boron

Resonance

occurs when two or more Lewis structures are valid and stable for a particular molecule

Transition metals lose e- from --- orbital before --- orbital

s orbital before d orbital

Formal charge

the difference between the number of valence electrons on a free atom and the number of electrons on the atom in the molecule (use lewis structures)

Bond order

number of chemical bonds between a pair of atoms in a molecule
single
double
triple

Bond energy

the energy required to break a bond
breaking bond - positive enthalpy
creating bond - negative enthalpy