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missing lewis acid/base

Heterogenous Equilibria

when equilibrium involves more than one phase (concentration between solids and gasses or other)

Pure Solids and Pure Liquids

In Heterogenous Equilibria, these two do not affect the equilibrium constant

Gas Equilibrium (Kp)

This equilibrium is measured through its partial pressure, its ICE table uses pressure

Relationship Between Kc and Kp (eq)

Kp = Kc * (RT)^Δn or Kp = ([products]/[reactants]) * (RT)^Δn (Δn is products - reactants)

Reaction Quotient (Qc)

used for predicting direction of reaction, left or right. Uses initial concentrations.

Qc and Kc

Qc > Kc - left
Qc < Kc - right

Qc = Kc - at equilibrium

Reaction Quotient Eq (Qc)

Qc = [C]^cini[D]^dini / ([A]^aini[B]^bini)

Le Chatelier’s Principle

Any change in equilibrium causes a shift in the system to reduce the effect of the change until equilibrium is reached again; reaction will always shift to the lesser side to balance itself

Le Chatelier’s and Pressure (Increase)

Increase Pressure = Decrease Volume; the system will shift to the side with LESS moles of gas

Le Chatelier’s and Pressure (Decrease)

Decrease Pressure = Increase Volume; the system will shift to the side with MORE moles of gas

Le Chatelier’s and Pressure (same moles)

If there is the same number of moles on both sides, equilibrium is NOT AFFECTED BY PRESSURE

Le Chatelier’s and Temp (increase)

adding heat = adding reactants; shift to products

Le Chatelier’s and Temp (Decrease)

taking away heat = taking away reactants; shift to reactants

Le Chatelier’s and Catalysts

A catalyst will make reaching equilibrium faster but not affect it

Le Chatelier’s and Concentration

Equilibrium will always go to the side with the lower amount; more product means go to reactants and vice versa.

Arrhenius Acid

A substance that INCREASES the concentration of H3O+ in water; a substance when combines with water will make the solution acidic

Arrhenius Base

A substance that INCREASES the concentration of OH- in water; a substance when combined with water will make the solution basic

Bronsted Acid

A substance that can donate a proton (H+)

Monoprotic

A Bronsted Acid that can only donate one proton (H+)

Polyprotic

A Bronsted Acid that can donate 2+ protons (H+)

Bronsted Base

A substance that can accept a proton (H+)

pH scale

Less than 7 = Acidic; Greater than 7 = Basic; Equals 7 = Neutral

pH to [H3O+]

pH = -log[H30+] , 10^(-pH) = [H3O+] (note: same is with pOH, just switch pH with pOH and [H+] with [OH-])

pH to pOH

14 - pH = pOH or 10^(-14)/[H+]=[OH-]

Strong Acids and Strong Bases

Both Ionize 100%; HBR, HI, HCL, H2SO4, HNO3, HCLO4

Autoionization of Water

Water will ionize by itself sometimes; H20(l) + H2O(l) <-> H3O+ + OH- Kc = 1*10^-14

Buffer Solution

A solution that is made from a weak acid and conjugate base or a weak base and conjugate acid that resists pH change when a SA or SB is added (note: the SB will react to completion)

Henderson-Hasselbalch Equation

pH = pKa + log([A-]/[HA])

Titration curve Regions

Initial Point, Buffer Region (plateau), Equivalence Point (exponential), After Equivalence Point

Direction of the acid-base reaction

The Direction of the acid-base reaction always favors the weaker acid and base; “The stronger acid plus the stronger base produces the weaker acid and weaker base”.

Dilution

Increased dilution = more water added = more reactants added = shift right to products

Initial Point (Titr, Crv)

Use ICE Table with initials to find pH

Buffer Region (Titr Crv)

find added moles, use BCE table to to get new initials, use ICE table with new initials to determine pH (1st)

Equivalence Point

the point where HA moles = SB moles

Buffer Region Eq

½ * Volume @ Equivalence Point

Equivalence Point (Titr Crv)

Find added moles, use BCE table to find new initials, use ICE table to find pH (2nd)

After Equivalence Point

Subtract Equivalence Point Volume from Current Volume, convert new volume to moles, divide by current volume which give [OH-] or [H3O+], convert [OH-] or [H3O+] to pH

Thermochemistry

the study of heat released or absorbed during chemical reactions

Δ Internal Energy =

Δ Energy Final - Δ Energy Initial

Two Components of E

Heat(q) and Work(w)

Heat (q)

the exchange of thermal energy between the system and surroundings.

Work (w)

the energy exchange when a force, F, moves an object a certain distance

q < 0 means:

it means loss of heat from system, Exothermic

q > 0 means:

it means system gains heat, Endothermic

Specific Heat

the amount of heat required to change 1 gram of a substance by one degree Celsius or Kelvin; J/g℃

Heat Capacity

the amount of heat required to change a substance by 1 degree Celsius (no grams); J/℃

Heat Capacity =

C = q/ΔT

Specific Heat =

q = mCsΔT

State Functions

properties that are pathway independent

Types of State Functions:

Enthalpy ΔH, Law of Hess, Internal Energy ΔE, Pressure (p), Temperature (t), Volume (v)

Work (w) =

q - ΔE

calorie (cal)

4.184 J

1 C (nutritional cal)

1000 calories or 4.184×10³ J

First Law of Thermodynamics

Energy cannot be created nor destroyed

Energy Transfer equation

q System = - q Surrounding

Why the energy transfer equation is like that:

Heat in the universe is constant because of the First Law of Thermodynamics, energy, aka heat, cannot be created nor destroyed. This implies the heat absorbed from a reaction will equal the opposite of the heat given from the environment, or vice versa.

Thermal Equilibrium

when there is no heat flow between two systems in contact and the q is equal in both; this is like the problems where a hot object is placed on a cooler one

Enthalpy

heat that is given off or absorbed by a system at constant pressure for one mole of a reaction

ΔHrxn =

Σ n*ΔH(products) - Σ n*ΔH(reactants)

Enthalpy of Reaction (ΔHrxn)

the amount of heat released/ absorbed from a reaction

Law of Hess

If we have a reaction (Equation A) with an unknown ΔHrxn, this allows us to manipulate, combine, and cancel out multiple known equations so that their overall reaction matches Equation A. By summing their ΔHrxn values accordingly, we can determine the ΔHrxn of Equation A.

State Functions and Law of Hess

The ability for this to equal ΔHrxn regardless of a different path makes it a state function

Rules of Law of Hess

Multiplication, Division, Reversing

Standard Enthalpy of Formations

ΔH°_f; the enthalpy of a reaction at standard conditions (1 atm, 25℃); note: pure elements (O₂, C(graphite)) ΔH°_f = 0

Standard Enthalpy of Formations =

ΔH°_f = Σ np*ΔH°_f products - Σ nr*ΔH°_f reactants

Shielding

the effect of full electron shells blocking outer valence e- from full nuclear charge

Shielding Trends

More shielding = Larger Atomic Radius & More Metallic Character ; Less shielding = Larger IE and EA

Atomic Radius Trends

down = more orbitals, left = weaker nuclear charge on valence e- (more shielding) so wider orbitals; wants to be Francium

Atomic Radius Cations

want to gain e-: + charge; smaller size; larger proton to neutron ratio

Atomic Radius Anions

wants to gain e-: - charge; larger size; smaller proton to neutron ratio

Ionization Energy Trends

up = less shielding: easier to lose electron → increases ; right = wants electron to complete octet → increases; more valence e- = greater __; wants to be Fluorine

Ionization Energy Def

the energy required to remove an electron from a neutral atom (energy to make a cation); harder to remove valence electron = higher this

Electron Affinity Def

the energy to released or absorbed from adding an electron/ making an anion

Electron Affinity Trends

up = less shielding: stronger nuclear charge: less energy to add e-; right = less shielding: stronger nuclear charge: less energy to add an e-; wants to be Fluorine

Metallic Character Trends

down = increased; left = increased; wants to be Francium

Isoelectronic Species/ Particles

different elements with the same number of electrons
ex) Ne, O^2-, N^3-, Na⁺, Mg²⁺ (all have same # of e-)

Born-Haber Cycle

the utilization of the Law of Hess to determine Lattice Energy

Bond Length (Lewis Structures)

the average distance between the two nuclei of covalently bonded atoms, the shorter the distance the stronger the bond

Electronegativity def

the more ___ an atom is, the stronger its attraction to electrons are; a measure of the attractive force an atom has for a shared electron

Electronegativity Trends

Up = Increases; Right = Increases; wants to be Fluorine

Valence Electrons

the e- in the outermost shell; Total amount of ___ divided by 2 is the total amount of bonds for Lewis Structure

Exceptions to Octet Rule

Incomplete: Beryllium, Boron; Expanded: SF₆, Odd: NO₂

Molecule Polarity Determination

all nonpolar bonds = nonpolar; nonsym & 1 polar = polar; sym and all polar = nonpolar

Bond Polarity for Covalent

Nonpolar bond = <0.4; Polar Bond = > 0.5

London Dispersion Force (LDF)

very weak IMF, this is present in all bonds

Dipole-Dipole

medium strength IMF, present in only polar bonds

H-Bond

second strongest IMF, present in only molecules with Hydrogen bonded with Nitrogen, Hydrogen, Fluorine, and Oxygen

Surface Tension

the IMFs at the surface of a liquid between molecules form a tension; directly related to strength of IMF

Viscosity

the resistance of a liquid to flow; directly related to IMF strength

Evaporation

the change from a liquid state to a gaseous state; indirectly related to IMF strength

Vapor Pressure

the pressure exerted by the vapor above a liquid; inversely related to IMF strength

Boiling Point

the temperature at which the vapor pressure equals the atmospheric pressure

Normal Boiling Point

the temperature at which vapor pressure equals atmospheric pressure of 1atm

Deposition and Sublimation

The phase change from solid to gas and vice versa

Evaporation and Condensation

The phase change from liquid to gas and vice versa

Fusion and Freezing

The phase change from solid to liquid and vice versa

Heat of Vaporization (ΔHvap)

the quantity of heat needed to convert a liquid at its boiling point to the gaseous state; given as J/mol or J/g

Heat of Fusion (ΔHfus)

the quantity of heat needed to convert a solid at its melting point to the liquid state; given as J/g

Temperature and Evaporation

the rate of evaporation increases as temperature increases → direct relationship: more temp = more evaporation of liquid

How to find number of atoms in a unit cell

count: the faces (1/2), edges (1/4), corners (1/8), and whole (1)