chemistry quantit

percent yield

(actual yield/theoretical yield) x 100

mass percent

(mass of element/mass of compound) x 100

stoichiometric calculations

****Calculating the mass of a product from the mass of a reactant requires: •The mole ratio from the balanced chemical equation. •Molar mass of the reactant. Molar mass of the product.

aqueous insoluble compounds

**All hydroxides (OH−) insoluble except for: •Group IA, Ca(OH)2, Sr(OH)2, and Ba(OH)2 **All sulfides (S2−) insoluble except for: •Group IA and NH4+, CaS, SrS, and BaS **All carbonates (CO32−) insoluble except for: IA and NH4+ **All phosphates (PO43−) insoluble except for: IA, NH4+

solution concentration: molality (m)

**Changes in boiling point/freezing point of solutions depends on molality: **for typical solutions: molality > molarity

solutions of volatile compounds

**For mixtures containing more than one volatile component: •Partial pressure of each volatile component contributes to total vapor pressure of solution. •Ptotal = X1P1° + X2P2° + X3P3° ... •Where Xi = mole fraction of component i, and Pi° = equilibrium vapor pressure of pure volatile component at a given temperature.

potential energy

**Potential: due to position or composition: •Can be converted to work: PE = m × g × h »m = mass, g = force of gravity, and h = vertical distance -chemical energy is a form of potential energy Potential energy is a state function: determined only by position or condition -Depends only on the difference between initial and final state of the system. Independent of path between states

factors affecting vapor pressure

**Temperature: •Higher temperature = higher kinetic energy; more molecules with sufficient Ek to overcome attractive forces in liquid phase. **Surface Area: •Higher surface area = greater number of molecules in position to enter gas phase. **Intermolecular Forces: •Stronger forces = higher Ek needed to enter gas phase.

reverse osmosis

**Use of high pressure to move solvent across membrane from region of high solute concentration to region of lower solute concentration. **Application: desalination/water purification.

combustion reaction

-Reactions between oxygen (O2) and another element in a compound. Hydrocarbons •Molecular compounds composed of only hydrogen and carbon. •"Organic" compounds. •Combustion products are CO2 and H2O.

aqueous solubility rules

-Soluble Cations: •Group I ions (alkali metals) and NH4+ -Soluble Anions: •NO3− and CH3COO− (acetate) •Halides (Group 17) »Exceptions: Ag+, Cu+, Pb2+, Hg22+ •Sulfates (SO42−) »Exceptions: Pb2+, Hg22+, Ca2+, Ba2+, Sr2+ -Combining anions/cations not listed above will result in formation of an insoluble compound.

non-electrolyte

-Substances in which no ionization occurs. There is no conduction of electrical current. -Examples: Aqueous solutions of sugar, ethanol, ethylene glycol.

mole

1 mole = 6.022 × 1023 particles

mass % to empirical formula

1. Convert mass% to moles. 2. Find smallest whole number mole ratio, which is the empirical formula, dividing each by the smallest # of moles.

methods of determining ΔHrxn

1.From calorimetry experiments: •ΔHrxn = −Ccal ΔT 2.From enthalpies of formation: •ΔHrxn° = ΣnpΔHf°(products) − ΣnrΔHf°(reactants) •ΔHf° values for substances in the Appendix. 3.Using Hess's Law.

Calculations using Hess's Law

1.If a reaction is reversed, ΔH sign changes. N2(g) + O2(g) → 2NO(g) ΔH = 180 kJ 2NO(g) → N2(g) + O2(g) ΔH = −180 kJ 2. If the coefficients of a reaction are multiplied by an integer, ΔH is multiplied by that same integer. 6NO(g) → 3N2(g) + 3O2(g) ΔH = 3(−180 kJ) ΔH = −540 kJ

Measuring Heats of Reaction (ΔH rxn)

A bomb calorimeter is a constant-volume (no PV work) device used to measure the heat of a combustion reaction. Heat produced by reaction = heat gained by calorimeter -ΔE ~ -ΔH = qcal = CcalΔT

standard solution

A solution of known concentration (also called the titrant)

Titration

A volumetric method to determine the concentration of an unknown substance by reacting it with a standard solution

fuel values

Amount of energy (in kJ/g) produced from a combustion reaction. example •CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) •ΔHcomb = 802.3 kJ/mol •Fuel value = (802.3 kJ/mol)∙(1 mol/16.04 g) = 50.02 kJ/g

food values

Amount of energy produced when food is burned completely: •Determined by bomb calorimetry. Nutritional Calorie = 1 kcal = 4.184 kJ

Bond Length

Bond length depends on: •Identity of the atoms •Number of bonds between them

units of energy

Calorie (cal): •The amount of heat necessary to raise the temperature of 1 g of water 1°C. Joule (J): •The SI unit of energy. •4.184 J = 1 cal Energy = heat and/or work (same units!).

Calorimetry

Calorimetry = measurement of heat. • Typically, measuring change in heat that accompanies a physical change or chemical process. A calorimeter = device used to measure the absorption or release of heat by a physical or chemical process. •A closed system! −qsystem = qcalorimeter

Precipitation

Combination of non-soluble cation with non-soluble anion = precipitate! ex PBI2(s)

Ideal Gas Law

Combined Gas Law : (P x V)/(n x T) = constant=R rearranges to: PV=nRT R = universal gas constant = 0.08206 L∙atm mol−1K−1 P = pressure (in atm) V = volume (in liters) n = moles T = temperature (in Kelvin)

stock solution

Concentrated solution (i.e., high solute-to-solvent ratio).

saturated solution

Contains the maximum amount of solute that can dissolve in a given volume at a given temperature

Molar concentration units

Define the amount of solute in a solution: amount of solute/amount of solvent OR amount of solute/amount of solution

Energy

Energy: Heat vs Work

enthalpy and change in enthalpy

Enthalpy (H) = E + PV Change in Enthalpy (ΔH) = ΔE + PΔV (constant P) ΔH = change in enthalpy; energy flows as heat at constant pressure: ΔH = qP = ΔE + PΔV ΔH > 0, Endothermic; ΔH < 0, Exothermic Add subscripts to indicate ΔH for specific process or part of the universe; e.g., ΔHvap, ΔHrxn, ΔHsys.

surroundings

Everything in the universe that is not part of the system.

fuel density

For liquid fuels, energy released in kJ/L

endothermic process

Heat flows into system from surroundings (q > 0). vaporization = endothermic Endothermic processes absorb heat (energy) from the surroundings (energy enters the system, q > 0) i.e. The system increases in energy (+ q)

exothermic process

Heat flows out of system to surroundings (q < 0). condensation = exothermic Exothermic processes release heat (energy) to the surroundings (energy leaves the system, q < 0) i.e. The system decreases in energy (- q)

Hess's Law: A Thermodynamic Cycle

Hess's Law of constant heat of summation: •The ΔH of a reaction that is the sum of two or more reactions is equal to the sum of the ΔH values of the constituent reactions.

internal energy

Internal energy of a system = sum of all KE and PE of all components of the system. •Different types of molecular motion contribute to overall internal energy: (a) translational, (b) rotational, and (c) vibrational.

kinetic energy

Kinetic: due to motion of the object •KE = 1/2 mu^2 (m = mass, u = velocity)

First law or thermodynamics

Law of Conservation of Energy: •Energy can be neither created or destroyed. •Can be converted from one form to another. »Potential energy → kinetic energy »Chemical energy → heat (q) q is heat, typically units of J, kJ, or kJ/mol: Note sign convention: qsystem = -qsurroundings Energy of the universe is constant! Universe= system + surroundings •energy gained or lost by a system must equal the energy lost or gained by surroundings.

heat capacities

Molar heat capacity (cp) is the heat required to raise the temperature of 1 mole of a substance by 1°C at constant pressure. •q = ncpΔT (cp = J / (mol∙°C)) Specific heat (cs) is the heat required to raise the temperature of 1 gram of a substance by 1°C at constant pressure. •q = ncsΔT (cs = J / (g∙°C)) Heat capacity (Cp) is the quantity of heat needed to raise the temperature of some specific object by 1°C at constant pressure

mole fraction

Ratio of the # of moles of a given component in a mixture to the total # of moles in a mixture

molar mass from colligative properties

Rearrange Colligative Property Relationships •ΔTf, ΔTb: ΔTf= Kf*m=Kf*(g/M)/ kg solvent M=Kf*g / ΔTf*(kg solvent) Osmotic Pressure pi=MRT=(g/M)RT M=(g *RT)/ pi •Typically used only for non-electrolyte solutes. •Osmotic pressure most common application.

Heat

The energy transferred between objects that are at different temperatures, naturally flows from hot to cold

vapor pressure lowering

The lowering of vapor pressure of a solvent by the addition of a nonvolatile solute to the solvent. one of the colligative properties of solutions

system

The part of the universe that is the focus of a thermodynamic study. •Isolated / Open / Closed

Enthalpy of formation (ΔHf)

The standard enthalpy of formation, ΔHf°: •The enthalpy change of the formation reaction. •A formation reaction is the process of forming 1 mole of a substance in its standard state from its component elements in their standard states. •For example, formation reaction for water: »H2(g) + ½ O2(g) → H2O(l) »ΔHrxn = ΔHf°(H2O) (The standard state of a substance is its most stable form under 1 bar pressure (~ 1 atm) and at 25°C.)

Universe

Universe = System + Surroundings Change in Universe = 0 so... 0 = System + Surroundings System = - Surroundings

Calculating molality

a) Mass of solute and solvent. b) Mass of solute/ volume of solvent. c)Volume of solution

Water is amphiprotic

acts as acid or base

supersaturated solution

contains more dissolved solute than a saturated solution at the same temperature

Clausius-Clapeyron Equation

for a given solvent at two temperatures

solubility

g solute/ 100 mL solvent

equivalence point

in a titration, the point at which enough titrant has been added to fully neutralize the analyte: [H+] = [OH-]

end point

in a titration, when the indicator changes color

Clausius-Clapeyron Equation

ln(Pvap)=-(delta Hvap)/R * (1/T)+C Plot of ln(P) vs 1/T yields straight line: •Slope = −ΔHvap/R •Intercept = constant

Fractional distillation

process to separate volatile components of a mixture Vapor phase enriched in more volatile component: •As vapor condenses, evaporates again, achieve greater enrichment. •Ideally, separate and collect pure components. •Greater vapor pressure solvent rises to the top of the column faster.

water as base

proton acceptor (base)

water as acid

proton donor (acid)

Ideal Solutions

solutions that obey Raoult's Law

Law of Conservation of Mass

sum of the masses of the reactants of a chemical equation is equal to the sum of the masses of the products.

normal boiling point

temperature at which the vapor pressure of the liquid is equal to 1 atm

combination reaction

two or more substances combine to form one product. example: sulfur trioxide + water= sulfuric acid

Work

w=F x d •Work (w) is done when a force (F) moves an object through a distance (d).

change in internal energy

ΔE = q + w: •ΔE = change in system's internal energy •q = heat • w = work Work: •w = −PΔV » where P = pressure, ΔV = change in volume = Vf - Vi. •Work done by the system is energy lost by the system (added to the surroundings), hence the negative sign.

molecular formula

•Actual molar ratio of elements in a compound. •Equal to some multiple of empirical formula: •Need empirical formula and molecular mass.

Heat of reaction

•Also known as enthalpy of reaction (ΔHrxn). •The heat absorbed or released by a chemical reaction

Strong Acids/Bases:

•Completely ionized in aqueous solution (i.e., strong electrolytes). •All other acids assumed to be weak. ex strong acids: HCl, HBr, HI, HNO3, H2SO4, HClO4

chemical equations

•Describe proportions of reactants (the substances that are consumed) and products (the substances that are formed) during a chemical reaction. •Describe the changes on the atomic level. •Physical state of products/reactants: »(s) = solid; (l) = liquid; (g) = gas; (aq) = aqueous soln.

Thermochemistry

•Energy in the form of heat consumed or produced by chemical reactions. •2H2(g) + O2(g) → 2H2O(l) + energy (heat)

Solutions

•Homogeneous mixtures of two or more substances: » solvent = substance present in the greatest quantity »solutes = substances dissolved in the solvent (i.e., the other ingredients in the mixture, includes gases) aqueous solutions--> water is solvent

ionic equations

•Ionic species represented as dissolved ions: H+(aq) + Cl−(aq) + Na+(aq) + OH−(aq) → Na+(aq) + Cl−(aq) + H2O(l)

Molarity

•Moles of solute (moles/L = n/V = Molarity) •Molarity is expressed as capital "M"

osmosis

•Movement of solvent through semi-permeable membrane from region of low solute concentration to region of higher solute concentration.

Strong electrolytes

•Nearly 100% dissociated into ions. •Conduct current efficiently. ***Examples: Solutions of NaCl, HNO3, HCl NaCl → Na+(aq) + Cl−(aq)

weak electrolyte

•Only partially dissociate into ions. •Slightly conductive. ****Examples: Vinegar (aq. solution of acetic acid); tap water. •CH3CO2H ⇌ CH3CO2−(aq) + H+(aq)

Dilution

•Preparation of dilute solution (i.e., low solute-to-solvent ratio) by adding solvent to a given volume of stock solution. Number of moles of solute remains constant •M1V1 = M2V2

vapor pressure

•Pressure exerted by a gas in equilibrium with its liquid. •Rates of evaporation and condensation are equal. •Pressure exerted is called the vapor pressure. •In general, vapor pressure of solution is lower than vapor pressure of pure solvent. rate of evaporation < rate of condensation

osmotic pressure (pi)

•Pressure required to halt flow of solvent through membrane due to osmosis. •π = iMRT (M = molarity of solution)

molecular equations

•Reactants/products written as undissociated molecules: HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

net ionic equation

•Remove spectator ions (ions present in same form on both reactants and products side of chemical equation.) dissolved ions.

empirical number

•Simplest whole-number molar ratio of elements in a compound.

precipitate

•Solid product formed from reactants in solution. •AgNO3(aq) + NaCl(aq) → NaNO3(aq) + AgCl(s) -Can predict formation of precipitates based on solubility "rules." •Precipitation reactions can be written using net ionic equations.

net ionic equations

•Soluble ionic compounds: •Strong electrolytes »ionize completely in aqueous solution. •Can write total and net ionic equations.

colligative properties

•Solution properties that depend on concentration of solute particles, not the identity of particles. previous example: vapor pressure lowering consequences: change in boiling pt, and freezing pt of solution

limiting reactants (reagents)

•Substance that is completely consumed in the chemical reaction. •Determines the amount of product that can be formed during the reaction. •Identified by: » number of moles of reactants »stoichiometry of balanced chemical equation

Theoretical yield

•The calculated amount of product formed based on the amount of limiting reactant.

Actual yield

•The measured amount of product formed.

Bond Order

•The number of bonds between two atoms: » 1 for a single bond » 2 for a double bond » 3 for a triple bond

Stoichiometry

•The number of moles of reactant/products needed for conservation of mass. Indicated in chemical equation by coefficients.

balanced chemical equation

•Total mass of each element on the reactant side must equal the total mass of each element on the product side. •Total charge of reactant side must equal the total charge of product side.

Thermodynamics

•Transformation of energy from one form to another.

Weak acids/bases

•dissociate only partially (weak electrolytes). •Use double arrow ( ⇌ ) to indicate incomplete reaction to form ions. •Acid: CH3CO2H ⇌ CH3CO2−(aq) + H+(aq) •Base: NH3 + H2O ⇌ NH4+(aq) + OH−(aq) •Net ionic equation for neutralization reaction are the same for strong or weak acids and bases: H+ + OH− → H2O

van't Hoff factor (i)

•number of ions in formula unit. • e.g., NaCl, i = 2 ΔTb = i∙Kb∙m & ΔTf = i∙Kf∙m Deviations from theoretical value (always use whole number) due to ion pair formation.

Raoult's Law

•vapor pressure (P) of solution is proportional to mole fraction of solvent (non-volatile solute). •Psolution = XsolventPsolvent°

calculating energy through a change in state Molar heat of fusion (melting):

•ΔHfus = heat needed to convert 1 mole of a solid at its melting point to 1 mole of liquid. •q = nΔHfus

calculating energy through a change in state Molar heat of vaporization (evaporation):

•ΔHvap = heat needed to convert 1 mole of a liquid at its boiling point to 1 mole of vapor. •q = nΔHvap

Boiling Point Elevation (ΔTb):

•ΔTb = Kb∙m Kb = boiling point elevation constant of solvent; m = molality

Freezing Point Depression (ΔTf):

•ΔTf = Kf∙m •Kf = freezing-point depression constant; m = molality.