GenChem Final Rev3

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Solution made of solvent (liquid) and solute (dissolved solid)
Acid dissolves in water and generates hydronium
Name the 7 strong acids HCl; HBr; HI; HNO3; HClO3; HClO4; H2SO4
Strong acids and strong bases _______ _______ dissociate completely
Base dissolves and generates hydroxide or accepts H+
Name the 8 strong bases NaOH; KOH; LiOH; RbOH; CsOH; Ca(OH)2; Sr(OH)2, Ba(OH)2
Electrolytes conduct electricity; salts; dissolves in water to yield solution that conducts electricity; acids are electrolytes; if strong acid= strong electrolyte
salts negative/positive charge; ions move electrons to allow conduction of electricity
Precipitation two things soluble in solution and get together to create solid that is no longer soluble; usually involve ionic compound
Hydration water molecules surround solute particles in aqueous solution
Rule #1: All Group _______ and _______ are _______ Group 1; NH4+; soluble
Rule #2: All _________, __________, _________ are soluble acetates, nitrates, chlorates/perchlorates
Rule #3: All Group 7 salts are ________; Cl-, Br-, and ____-; EXCEPT; ______, _____ , _____ soluble; I-; Ag+ Hg2 2+; Pb2+
Rule #4: All _______ salts are _______ EXCEPT _______, ______ , ______, _____, _____, ____ Ag+ Hg2 2+; Pb2+; Ca2+; Sr2+; Ba2+
Rule #5; Insoluble salts are ____, ____ , _____, _____, unless with Group 1 CO3 2- ; PO4 3-; CrO4 2- ; S 2-;
Rule #6: All _______ salts are ________ EXCEPT Group ____ and _______ hydroxide; insoluble; 1; Ba 2+
Spectator ion hangs in solution and doesn't do anything; appears both on reactant and product side in ionic equation
Neutralization reaction is also known as acid-base reaction
Arrhenius acid (H+); base – dissolved to produce OH-
Bronstead acid = anything that produces H+ (proton donor) ; Base = anything that accepts H+ (proton acceptor)
Proton hydrogen atom that has lost its electron
Oxidation loss of electron; gives up electron to become positively charged; electron gained on right side
Reduction gain of electron; gains electron to go down in oxidation state; REDUCTION OF CHARGE; electron gains to left side
Reducing agent gets oxidized ; donates electrons
oxidizing agent gets reduced ; accepts electrons
Monoprotic acids acid with one ionizable hydrogen atom
Diprotic acids acid with two ionizable protons (H2SO4)
Triprotic acids acid molecule with three ionizable protons (H3PO4)
Redox reaction electrons are transferred form one reactant to another
oxidation state charge any atom would have if electrons were transferred completely; increase in oxidation state = being oxidized ; reduce in oxidation state = being reduced
Disproportionation reaction when element undergoes both oxidation and reduction in same reaction
Molarity mol solute/volume of solution (L)
If something is dissolved in water, it breaks up into its _______ ions
Dilution process of preparing less concentrated solution from a more concentrated one; take something with higher concentration, move to another vessel and add water
Dilution equation M1V1 = M2V2
pH = -log[H3O+ or molarity] ; neutral = 7; acidic pH less than 7, basic pH more than 7
H3O+ = 10^-pH
Titration Solution of accurately known substance (standard solution) is added gradually to another solution of unknown substance until chemical reaction between the two are complete
Equivalence point point in titration where reaction is complete
End point point at which indicator color changes
indicator substance that has a distinctly different color in acidic and basic media
Exothermic gives off heat; heat leaves system; burning of wood; causes surroundings to gain heat
Endothermic Heat enters system; instant ice pack; feelings cold ; causes surrounding to lose heat
Thermodynamics study of interconversion of heat and other kinds of energy
Open system mass and energy can move from surroundings to system (vice versa)
Closed system mass does not exchange with surroundings
Isolated system no exchange of mass or energy
State functions volume , temp, pressure, energy but NOT heat; only depend on beginning and end but not path taken
First Law of Thermodynamics U sur +U sys = 0; cannot create more energy; can only change form
change in U equation Work + q (heat) ; negative = exothermic ; positive = endothermic ; units = Joules
Work is done on sytem +
Work is done by system on surrounding
Endothermic +
Exothermic
Enthalpy Measurement of heat; heat exchanged between system and surroundings at constant pressure; H
Equation for work w = -P(change in Volume); units = atmL
Enthalpy of reaction equation Enthalpy of prod – Enthalpy of reactions
Elements in standard state , Enthalpy = 0
Heat transfer of energy (thermal) between two bodies that are at different temperatures
Thermochemistry study of heat associated with chemical reactions
State of system values of all relevant macroscopic properties (comp, energy, temp, pressure)
State functions properties that determined by state of system, independent of how state was achieved
Calorimetry measurement of heat changes
Specific Heat (s) Energy required to raise 1 gram of substance by 1C; J/gC
Heat capacity amount of energy required to raise substance by 1C; J/C
Equation for heat q (Joules) = sm(change in T); m = mass in grams ; T = temp in C
Other equation for heat q (joules) = C(change in Temp) ; C = heat capacity
q sys = = sm(change in T sys) = – C (change in T sys)
Hess's Law Change in enthalpy that occurs when reactants are converted to products in a reaction is the same whether reaction takes place in one step or in a series of steps
When flipping the equation…. remember to multiply by -1
Standard enthalpy heat changes that result when one mol of compound is formed from its constituent elements in standard states
standard enthalpy of formation of any elements in most stable form zero
Bond enthalpy enthalpy change associated with breaking a particular bond in one mole of gaseous molecules
Bond enthalpy equation sum of bonds broken – sum of bonds made
Lattice energy amount of energy required to convert one mole of ionic solid to its constituent ions in the gas phase; useful measure of an ionic compound's stability;
Four properties of gases 1. fluid = take on shape of container 2. compressible3. densities are lower than solids and liquids4. mixtures of gases are homogenous
Boltzman and Maxwell Theory of Gases 1. Gases are made of particles that are separated by large distances; 2. Gases are in constant random motion3. Gases do not interact4. Energy of gas is proportional to absolute Temp
Diffusion random motion and frequent collisions that leads to mixing of gases
Effusion Particles move to vacuum
Graham's Law rates of diffusion and effusion are inversely proportional to square root of molar mass of gas
Pressure Force / area; SI unit = newtons (N); area = m^2; Pressure = pascal
Pressure = (height m)(gravity m/s^2)(Density kg/m^3)
Boyles Law P1V1 = P2V2 at constant temp
Charles/ Guy-Lassac Law V1T2 = V2T1 at constant pressure
Avogadro's Law V1N2 = V2N1 at constant pressure and temp
Combined Gas Law P1V1/T1 = P2V2/T1
Ideal Gas Law PV = nRT ; P = atm; V = L; T = K; R = 0.08206; n = mol
Standard temp and pressure 0C / 273.15 K and 1 atm
Density equation mp / RT
Molecular weight equation mw = dRT/P
Gases behave at ______ temp, and ______ pressure high; low
Mole Fraction number of moles of component divided by total number of moles in mixture; mole fraction of mixture is always less than 1; sum = 1; dimensionless; Xt = Xa + Xb + Xc
Root mean square speed molecular speed that is inversely proportional to molecular mass
RMS equation = squ root (3RT/M) ; R = 8.314 J/Kmol; M = molecular mass in kg/mol; T = Kelvin

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