C200_ General Chemistry equation Sheet

Dilutions

M1V1=M2V2 or C1V1=C2V2

colligative properties

properties that depend on the concentration of solute particles but not on their identity

* Examples of these include boiling point depression, freezing point depression, osmotic pressure and vapor pressure

Amphoteric

Compounds that can react as both an acid and a base, depending on the reaction conditions

Percent error

(A-T)/T x 100

absorbance and beer lambert law

ecl (e=molar extinction coefficient; c=samples concentration; l= path length

*if you know the absorbance of a solution, you can determine its concentration

1 mol = how many atoms

6.023x10^23

combustion reaction

(hydrocarbon or alcohol) + 02 --> co2 to H20

Emprical Formula

most reduced form of a molecular formula

Energy of a photon equation

Planck's constant x speed of light / frequency (HC/F)

*H=6.63x10^-34

*c=3.0x10^8

f= photons frequency

quantum numbers

n (principle number), l(azimuthal), ml(magnetic), ms (spin)

*n= shell (distance from the nucleus)

*l= subshell (type of orbital)

-l=0 (s); 1(p); 2(d); 3(f)

*Ml=(-1....1) (determined by the l value )

ms= 1/2 & -1/2

paramagnetic vs diamagnetic

unpaired electrons attracted to magnets

dia: all paired and slightly repelled by magnets

relationship between energy, frequency and wave length

high energy = high frequency = low wave length

ionic bond properties

(metal + nonmetal)

generally have high melting & boiling points, tend to be hard, brittle solids; tend to conduct electricity only when melted or dissolved in water----not in solid state

covalent bond properties

(between two nonmetals)

sharing of electrons between with elements with similar electronegativity.

* high melting and boiling poiints, hard and do no conduct electricity

molecular bonds

(2 or more non metals) low melting points and do not conduct electricity

Ion size

more electrons = bigger size (anion) -

less electrons = smaller size (cation +)

isoelectronic series

comparing elements with the same number of electrons.

The compound with the smallest amount of protons will be the largest in size

ionization energy

The amount of energy required to remove an electron from an atom

trend: like electronegativity but includes noble gasses

electronegativity

how thirst an atom is for electrons

up and to the right excluding noble gasses

electron affinity

the energy given off when an atoms gains an elecctron (exothermic reaction)

*Cl has highest electro negativity

temperature c to k

k= c+ 273

Volume

1cm^3=1ml =1cc

pressure

force/area (1atm = 760 torr or mmhg)

Ideal gases behave

most ideally under low pressures and high temperatures

at STP 1 mole of any gas =

22.4 L

boyles law (gas law)

V is inversely related to pressure. ( V= 1/p)

Charles law (gas law)

V and temperature are related

Avagadro's Law

V=N

Combined Gas Law

P1V1/n1T1 = P2V2/n2T2

Ideal gas law

PV=nRT

R=.0821

At STP

P=1atm and T =273k

Density

P(MM)/RT = m/v

r=.0821 L*ATM/ mol k

Dalton's Law of Partial Pressures

Total pressure of a gas is equal to the sum of the partial pressure of the component gases

Ptotal= Pa + Pb...

In terms of gas A : Pa=Xa P total

Xa= mol fraction of gas A

Graham's law of effusion

Effusion refers to the movement of gas particles through a small hole. Graham's Law states that the effusion rate of a gas is inversely proportional to the square root of the mass of its particles.

r1/r2 = √(M2/M1) M=mola

r mass

r=rate of effusion

intramolecular forces vs Intermolecular forces

Intramolecular forces are the forces that hold atoms together within a molecule.

Intermolecular forces are forces that exist between molecules.

intermolecular forces

are forces that exist between molecules.

1). hydrogen bonding (strongest)

2) dipole dipole

3) dispersion forces (london disp/ van der wall) (weakest)

hydrogen bonding

Hydrogen atom bonded to either an O, N, or F

Dipole-Diople Forces

NON METAL molecules with polar bond causing the molecule to have a partial negative and positive charge

They line up in a complementary to each other and have big electronegativity differences between the two

dispersion force

no significant difference in electronegativity

Dispersion forces are in all molecules and are responsible for molecules boiling points

intermolecular forces and boiling points

responsible for boiling points. stronger the intermolecular forces leads to :

1) higher boiling point

2)high heat of vaporization

3) higher viscosity

4) higher surface tension

5) LOWER vapor pressure

Intramolecular forces

forces are the forces that hold atoms together within a molecule.

Metallic bonds, Ionic bonds, Polar covalent bonds, covalent bonds,

Phase changes

Freezing: crystalization

melting : fusion

Sublimation; fusion and vaporization : consumer heat therefore are endothermic reactions with a positive +ΔH and +ΔS (entropy, increasing disorder)

Deposition; condensation; crystallization are exothermic -ΔH because they release heat and have a -ΔS (because more order is being established)

Boiling point and vapor pressure

boiling point is when the vapor pressure = external pressure.

high boiling point = low vapor pressure

Phase Diagram - triple point

unique for each substance, point where all three phases are in equilibrium

Phase diagram - critical point

Point where liquid and gas are no longer distinguishable

Molarity (M)

moles of solute/liters of solution

Molality (m)

moles of solute/kg of solvent

Solubility of ionic compounds in water rules

most group one metal cations NO3-; clo4-;c2h302- and nh4+ salts are soluble

most ag2+; pb2+; s2-: oh-; hg2 (2+); co3 2-; po4 2- are insoluble

Solubility of solids and gasses (henrys law)

Pa=Kh(A)

pa= partial pressure of gas a

Kh= henerys law constant; changes with each problem

A= concentration of gas a

Freezing point depression

ΔTf= -iKfm

i= vant hoff factor

Kf= freexing point depression constant

m= molality

Boiling point elevation

ΔTb=- iKbm

i= vant hoff factor

Kb= boiling point depression constant

m= molality

Vapor pressure depression

Psoln=X solv P solv

Psoln= new VP of solution

x solv = mol fraction of solvent (percent of solvent in new solution)

P solv = VP of pure solvent

Osmotic pressure

pi=iMRT

i= vant hoff

M= molarity

r= .0821

T= temp

Pi= (n/v)RT

(n/v) = M

chemical kinetics

the study of how fast a reaction occurs or the reaction rate

Thermodynamics

tells us whether a reaction will occur but not how fast or slow it will occur

General Rate rate

rate= k (A)^m (B)^n

k= rate constant

m&n: determined experimentally

Rate constant units

Zeroth order: M/s (m^1 x s^-1)

First order: 1/s ( x s^-1)

Second order: 1/ M *S (m^-1 x s^-1)

Third order: 1/ M^2 *s (m^-2 x s^-1)

Integrated rate law units (Graphs)

Collison theory and arrhenius equation

K = Ae ^-Ea/RT

ea= activation energy

A= frequency factor

r= 8.314 J/mol k

k=rate constant

Equilibrium constant expressions

Kc & Kp = Products/ reactants

Keq= kForward/ kreverse

Keq Rules

K>1 = products are favored at equilibrium

k<1= reactant are favored at equilibrium

solubility product constant (Ksp)

Ksp= products / reactants

pH & pOH

pH = -log (H+)

pOH= -LOG (oh-)

ph +poh = 14

(h+) & (oh-)

(H+)= 10^-PH

(OH-)= 10^-POH

(H+)(OH-)= 1*10^-14

Weak acids

HA + H2O <----> h30+ + A-

ka= (h30+)(A-)/ (HA)

*ka= acid dissociation constant

(H+) = SR(ka)(HA)

Weak base

A- + H2O <----> HA + OH-

kB= (HA)(OH-)/ (A-)

*kB= BASE dissociation constant

(OH-) = SR(kB)(A-)

pKa and pKb

pKa=-log(ka)

pKb=-log(kb)

pka+pkb = 14

Kw=ka x kb =1

Ka and Kb strength

larger Ka = smaller pka = stronger acid

larger kb= smaller pka = stronger base

Arrhenius acid

h+ donor in water

arrhenius base

OH- donor in water

Bronsted-Lowry acid

h+ donor

Bronsted-Lowry base

H+ acceptor

lewis acid

electron pair acceptor

lewis base

electron pair donor

conjugate acid-base pair

acid --> base (remove hydrogen and the charge decreases by one_

base---> acid (adds 1 hydrogen and increases the charge by 1)

Trends in acid strength

the more oxygen atoms means the more acidic it is due to resonance.

if the number of oxygens is the same the the more electronegative heteroatom = more acidic

Neutralization reactions and normality

*reactions that occur between acids and base; which always makes h20 and a salt

NaMaVa=NbMbVb

Na= number of hydrogen groups the acid can donate

Nb= number of oh groups the base can donate

Buffers and henderson hasselbalch

buffer= is a solution that resists ph change. it is made from a weak conjugate acid/ conjugate base pair

ph= pka +log (A-)/(HA)

Enthalpy

the amount of heat energy a substance contains

HA + H2O <----> h30+ + A-

ΔH > 0 -endothermic

ΔH < 0 exothermic

Enthalpy of formation (ΔHf)

Δhf = nΔH(products) - nΔH(reactants)

n= coefficient from the balanced reaction

first law of thermodynamics (law of conservation of energy)

ΔE= q + w

q= heat

w=work

ΔE= change in internal energy

pressure-volume work

w= -p Δ V

ΔV= change in volume

p= external pressure

calorimetry thermal energy

q= -Ccalorimetry ΔT

Cal= specific heat of calorimeter

heat curves and thermal energy

q=mc ΔT or q=mc ΔH(fusion or vaporization)

m=mass

c= specific heat

ΔT= change in heat (final - initial )

entropy formula

a measure of how disordered something is

ΔS= Sum(n)Sproducts- Sum(n) S reactants

Sgas> Sliquid> Ssolid

saq> Ssolid

Srxn> 0: when there is an increase in the number of moles of gas

Bond Disociation Energy

ΔH= Sum ΔH (products) - S ΔH (reactants)

=Sum ΔHbroken - Sum ΔH formed

making bonds = exothermic (ΔH-)

breaking bonds= endothermic (ΔH+)

Gibbs free energy equation

a measure of spontaneity

ΔG^0 = ΔH - TΔS

ΔG^0= standard conditions

ΔH= enthalpy

ΔS= entropy

T= temperature in kelvins

Gibs free energy equation (Keq standard and nonstandard conditons)

ΔG= ΔG^0 + R(T)(Ln Q)

ΔG^0= -RTLnKeq

ΔG^0= standard conditions

Keq= equilibrium constant

Q= reaction quotient

R= 8.314 J/mol k

standard cell potential (E°cell) (reduction potentials)

E^o cell = E^o reduction + E^o oxidation

E^o = E^o cathode + E^o anode

Gibs free energy (ΔG) rules

ΔG= - (spontaeous)

ΔG= + (nonspontaenous)

ΔG= 0 at equilibrium

Gibs free energy relation with enthalapy (ΔH) and entropy (ΔS)

Relating Gibbs Free Energy and K

Redox reactions

Reactions that involve the transfer of electrons from one element to another.

Reduction-oxidation reaction

(OIL RIG)

-3------> 3 (OXIDATION)

3+-------> 3 (REDUCTION)

Reducing agent (reductant) in redox reactions

is the element that GIVE UP e- (gets oxidized) is the REDUCING AGENT

Oxidizing agent (oxidant) in redox reactions

is the reactant that gets reduced

Balancing redox reactions

Break the reactions into (1/2) reactions

1 should be the oxidation reaction and the other should be the reduction equation

anode

is the site of oxidation

cathode

is the site of reduction

Balancing Redox Reactions (under acidic conditions)

1). identify what is being oxidized and what is being reduced

2). seperate the oxidation reaction from the reduction reaction (1/2) reactions

3). balance all the atoms the ARE NOT HYDROGENS OR OXYGENS

4). balance the oxygens by adding H2o to which side needs it

5). balance the hydrogens by adding h+ to the side the needs it

6). add e- wherever necessary to balance the charge on each side of the equation

7). as needed, add integer to your half reactions to make the number of moles of e- equal in each of your half reactions.

8). the final balanced number of e-s is the number of overall electrons per mole that get transferred in this particular redox reaction

Balancing Redox reactions (under basic conditions)

1). identify what is being oxidized and what is being reduced

2). seperate the oxidation reaction from the reduction reaction (1/2) reactions

3). balance all the atoms the ARE NOT HYDROGENS OR OXYGENS

4). balance the oxygens by adding H2o to which side needs it

5). balance the hydrogens by adding h+ to the side the needs it

6). add the same number of OH- to both side of the equation that you added to your half reactions.

- combine the h+ and oh- on the same side of the equation to form H2o. and eliminate excess h20

7). add e- wherever necessary to balance the charge on each side of the equation

8). as needed, add integers to your half reactions to make the number of moles of e- equal in each of your half reaction.

9) the final balanced number of e-s is the number of overall electrons per mole that get transferred in this particular redox reaction