Chem Semester 1 Final

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Cation

metals lose e-
always smaller than original molecule

Anion

nonmetals gain e-
always larger than original molecule

Atomic/Ionic Radius

increases as you go down because more e- shells

decreases as you go across because stronger nuclear charge pulls e- in closer

Ionization Energy

decreases as you go down because it's easier to remove an e- farther from the nucleus

increases as you go across because the e- are closer to the nucleus so it's harder to remove them

Electronegativity

decreases as you go down, atoms become more metallic and want to get rid of e-

increases as you go across, atoms become more nonmetallic and want to gain e- to form an octet

*noble gases excluded because they have an octet

Exceptions to Ionization Energy

all of 3A and 6A because of e- pairing
(consult orbital diagram notes for further explanation)

Coloumbic Forces

1. charge of ion
2. size/atomic radius of ion

high charge + small radius = larger lattice energy and stronger attraction and vice versa

Ionic Compounds

- formed when metal transfers VE to nonmetals
- ions arrange themselves in a crystal lattice
- held together via strong electrostatic attractions
- high melting point
- if dissolved in a polar solvent, will form an electrolytic solution

Covalent Compounds

- formed when nonmetals share VE to achieve octet
- covalent bonds form molecules
- held together via IMF, not chemical bonds
- low melting points
- never forms an electrolytic solution

*network covalent compounds don't dissolve

Substitutional Alloy

formed between 2 atoms of similar atomic radii

Interstitual Alloy

formed from 2 metals of vastly different radii
- smaller atom fills up empty spaces between larger atoms and results in a more durable structure

Bond Strength

triple > double > single

stronger bond = shorter length

2 domains: 2 bonding pairs

linear
180°
sp

3 domains: 3 bonding pairs

trigonal planar
120°
sp2

3 domains: 2 bonding pairs, 1 nonbonding pair

bent

4 domains: 4 bonding pairs

tetrahedral
109.5°
sp3

4 domains: 3 bonding pairs, 1 nonbonding pair

trigonal pyramid

4 domains: 2 bonding pairs, 2 nonbonding pairs

bent

ΔHrxn=ΣΔHbroken + ΣΔHformed

bond enthalpy: energy absorbed/released in a rxn

need the lewis structure for this one

ΔHrxn > 0 --> endothermic rxn
ΔHrxn < 0 --> exothermic rxn

Formal Charges

FC = # of VE in free atom - # of bonds - # of unbonded e-

Resonance Structures

occurs when e- shift to form alternative lewis structures

Resonance Stability

- all atoms must have complete octet
- formal charges kept @ minimum
- if formal charges are present, the negative charge is on the more electronegative element

Net Ionic Equations

1. write the balanced chemical equation
2. write the complete ionic equation
3. cancel out the common ones basically, then you get the net ionic equation

Limiting/Excess Reagents

1. write the equation
2. figure out what's limiting: convert grams to moles then use balanced chemical equation to figure out what's limiting
3. calculate how much was used and how much is left

Molarity

M = n/v

n = moles of solute

v = volume of solution in L

Solution Stoichiometry

idk man just figure out what you’re given and do the stoichiometry

Preparing Solutions: From Scratch

use M = n/v

Preparing Solutions: From Stock Solution

use M1V1 = M2V2

Factors That Influence Pressure

1. volume: lower volume = higher pressure
* does not increase particle speed
2. moles of gas: more moles of gas = higher pressure
3. temperature: higher temp = higher pressure

Gas Laws

P1V1 = P2V2

V1/T1 = V2/T2

P1/T1 = P2/T2

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PV = nRT

* remember the difference between both R constants; this one is 0.08206 L\*atm / mol\*K

ideal gas conditions: high temp & low pressure

Mole Fractions

you can either use PV = nRT with only the moles of the gas you want, or calculate total pressure then multiply by the fraction of moles of gas

Kinetic Molecular Theory

* volume of individual gases is negligible relative to the total volume
* kinetic energy is only affected by temperature

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graph is molecular speed vs fraction of molecules

lower temp: shift left w/ steeper curve

higher temp: shift right w/ less steep curve

Effusion/Diffusion

any atom/molecule with a smaller mass will effuse/diffuse faster

Polar Molecules

polar bonds: nonmetal bonded to F, O, N, Cl

* fluorine is the most electronegative element

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a symmetrical molecule or a molecule with no polar bonds is nonpolar

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an asymmetrical molecule with at least 1 polar bond is polar

IMFs

* **not** a chemical bond

strong IMF = low vapor pressure = high melting/boiling point

Ionic Bonds

* formed when a positive ion bonds w/ a negative ion via electrostatic attraction

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Ion Dipole: attraction between a polar molecule & ion

Hydrogen Bonds

* strongest IMF
* only occurs in polar molecules with H is bonded to F, O, or N and is attracted to lone pairs on another molecule
* any molecule capable of H-bonding is also capable of D-D

Dipole-Dipole

* all polar molecules are capable of exerting these
* occurs when lone pairs on more EN atom are attracted to the partially positive atom on the other molecule

London Dispersion Forces (including shape of molecule)

* present in all molecules
* having more electrons makes it more polarizable (easier to separate charges) which means a stronger LDF
* the larger the molecule, the stronger the LDF

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if 2 molecules have the same formula, the one with a more linear structure will have a stronger LDF due to more surface area of interaction

Empirical Formulas

1: write mass %s as grams

2: convert to moles

3: divide all values by the smallest mole value

4: round to the nearest whole number unless it ends in 0.20, 0.25, 0.33, or 0.5

5: divide the molar mass of the compound by the empirical mass to get the ratio

e.g. molar mass is 162.1, empirical mass is 81 so since 162.1/81 = 2, the molecular formula will be C10H14N2 instead of C5H7N

Hydrated Salts

always written as salt \* XH2O where X = whole number

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1: determine mass of dry salt and of H2O

2: convert both to moles

3: divide mole values by the smallest mole value, which should always be the salt

Orbital Diagrams

just make sure one arrow is up and one is down

* more stable when they are either 1/2 full or completely full
* make sure each orbital has at least 1 electron before pairing up any

Electromagnetic Radiation

C = λν

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C = speed of light: 3.0 x 10^8 m/s

λ = wavelength in m (convert nanometers to m)

ν = frequency in 1/s or Hz

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6\.02 x 10^23 photon/mol

Planck's Equation

E = hν

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E = energy per photon in J/photon

h = Planck’s Constant: 6.63 x 10^-34 J*****s

ν = frequency in 1/s or s^-1

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6\.02 x 10^23 photon/mol

Thermochemistry

system: the object which heat measurement is in reference to

surrounding: everything that isn’t the system

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endothermic: system temp increases, surrounding temp decreases

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exothermic: system temp decreases, surrounding temp increases

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* exothermic: __**ex**__udes heat

Molar Enthalpy in relation to energy

ΔHsol = q/n

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q = energy in KJ

n = moles of solute

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measured in KJ/mol

Phase Change Using ΔHfus and ΔHvap

ΔHfus: melt/freeze 6.01 KJ/mol

ΔHvap: vaporize/condense 40.7 KJ/mol

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melting & vaporizing are endothermic so positive ΔH

freezing & condensing are exothermic so negative ΔH

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1: bring to phase change temp (either 0 or 100C) using

q = MCΔT

2: phase change calculations using q = nΔHvap/fus

3: bring to final temp using q = MCΔT

4: add up all the energy changes

Enthalpy of Rxn: Stoichiometry
A+B -> C+D

endo: energy is a reactant

exo: energy is a product

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ΔH is measured in KJ/ __**mol rxn**__

Enthalpy of Formation

ΔHrxn = nΣΔHproducts - nΣΔHreactants

* pure elements will always have ΔHformation of 0 KJ/mol

Hess's Law

just cancel things out, reverse the signs if necessary, and add them all up

Entropy

ΔS = nΣΔSproducts - nΣΔSreactants

* usually measured in J/mol\*K
* positive ΔS = increasing disorder (favorable)
* negative ΔS = decreasing disorder

Factors That Affect Entropy

1. temperature: higher temp = higher ΔS due to more molecular movement
2. # of particles: more moles = higher ΔS
3. size of molecule: larger molecule = more atoms = higher ΔS due to more vibrational motion
4. physical state: gas has highest ΔS

Gibbs Free Energy (all equations)

ΔG = ΔH - TΔS

* Gibbs is in KJ/mol so convert ΔS to KJ instead of J

ΔG = -RTln(K)

* this R constant is 8.314 J/mol\*K or 0.008314 KJ/mol\*K

ΔG = nΣΔGproducts - nΣΔGreactants

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\-ΔG means spontaneous rxn & means K>1

\+ΔG means nonspontaneous rxn & means K

Spontaneity Table

\-ΔH and + ΔS = spontaneous @ all temps

\+ΔH and -ΔS = nonspontaneous @ all temps

\-ΔH and -ΔS = spontaneous only @ low temps

\+ΔH and +ΔS = spontaneous only @ high temps

Reaction Quotient

Q = \[C\]^c\[D\]^d/\[A\]^a\[B\]^b

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Q is unitless and measures the ratio of products to reactants at all conditions but equilibrium

Equilibrium Expression

Keq = \[C\]^c\[D\]^d/\[A\]^a\[B\]^b

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Keq is unitless and measures the ratio of products to reactants at equilibrium

Rates of Rxn

basically slope; measured in M/s

Criteria for Rxn to Occur

* particles must collide in order to react
* particles must collide in correct orientation (Tetris)
* particles must collide w/ enough energy to meet activation energy requirement

Rates of Rxn Graphs

activation energy: reactants to peak

ΔH: reactants to products

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endo: reactants lower than products

exo: reactants higher than products

Factors Affecting Rxn Rate

1. concentration: higher concentration = faster rxn
2. temperature: higher temp = faster rxn
3. surface area: exposing more surface area = faster rxn (basketball vs ping pong balls example)
4. catalyst: adding a catalyst = faster rxn (duh)
* catalyst lowers activation energy

Q < K

too many reactants!
shift right to produce more products

Q > K

too many products!
shift left to produce more reactants

negative ΔG correlates with:

spontaneous reaction
more products than reactants present
K > 1

positive ΔG correlates with:

nonspontaneous reaction
more reactants than products present
K < 1

Rate Laws

Rate (M/s) = k[A]^x[B]^y
A and B are the concentration/molarity of the reactants
x and y are the rate orders
k is the rate constant with units dependent upon order

If the overall order (sum of x and y) is 0, the units of k are M/s

What affects kinetic energy?

temperature

What are the units of k?

0th order: M/s

1st order: 1/time

2nd order: 1/M\*time

Integrated Rate Law for 0th Order:

\[x\] = -kt + \[x\]initial

Integrated Rate Law for 1st Order:

ln\[x\] = -kt + ln\[x\]initial

Integrated Rate Law for 2nd Order:

1/\[x\] = kt + 1/\[x\]initial

What do you use integrated rate laws for?

if you are only given time and concentration data and are asked to determine rate order