CHEM 120 Final Review (Version 2), Introduction to General, Organic & Biological Chemistry, Chamberlain College of Nursing

CHEM 120 Final Exam Review

CHAPTER 4: GASES
 Pressure is force exerted per unit area. Units: Pascal (Pa) = N/m 2
 Atmospheric pressure: force exerted by atmosphere = m × g, also known as Barometric pressure
 A manometer measures the pressure of a gas with respect to that of the atmosphere, P = d × h × g
 If Pgas > Patm, Pgas = Patm + ΔP. If Pgas < Patm, Pgas = Patm – ΔP
 Boyle’s Law – at constant temperature, the volume decreases as pressure increases V α 1/P
 Charles’s Law – the volume of a fixed amount of an ideal gas at constant pressure is proportional to its
absolute temperature (units in K) V α T, all gases reach zero volume at –273.15°C
 Avogadro’s Principle – the volume occupied by an ideal gas (at given T and P) depends only on the # of
molecules present (not the type), equal # of molecules occupy equal volumes
 The ideal gas equation: PV = nRT
 STP: 273.15K (0°C) and 1 bar, the volume occupied by 1 mol of an ideal gas is 22.7 L
 When a gas is compressed: the molecules are confined in a smaller volume so density increases
 When a gas is heated: increased pressure increases volume occupied by gas so density decreases

Application of Gas Laws
 P1V1 = P2V2
 P1 / T1 = P2 / T2
 V1 / T1 = V2 / T2
 P1V1 / n1T1 = P2V2 / n2T2
 Molar mass: PV = mRT / M M = mRT / PV m = MPV / RT
 Density: d = m / V = nM / V = MP / RT

Gases in Chemical Reactions
 Gases react in volumes proportional to small whole numbers
 Assuming the reactants and products are at the same T & P: volumes are proportional to stoichiometric
coefficients in a chemical equation
 Mole fraction: fraction of molecules that are gas A in a mixture χ A = nA / ntotal, χA + χB = 1
 Partial pressure: nA = PAV/RT nA / ntot = PA / Ptot = χA
 The total pressure of a mixture of gases is the sum of the partial pressures

Kinetic Molecular Theory
 Large particles all following Newtonian motion, separated by a large distance
 The collision between each particle is elastic; total kinetic energy is constant
 Root mean square velocity (urms): close to the average speed of the particles
 Most probable speed (um): highest # of molecules have this speed
 Average speed (uav): average speed of the molecules, little higher than u m

Non-Ideal Gases
 van der Waals equation takes into account intermolecular attraction and excluded volume
 P = (nRT / V-nb) – (n2/v2 a)
 Conditions for ideal gas: high temperature, low pressure
 Conditions for non-deal gas: low temperature, high pressure




CHAPTER 5: THERMODYNAMICS
 Energy is the capacity to do work, work is done when a force acts over a distance

,  Heat is the quantity of energy (q) transferred between a system and its surrounds
 First Law: the total change in the system’s internal energy (Δu) is the sum of the energy transferred as
heat or work, Δu = q + w
 State function: a property that only depends on the current state of the system (e.g. density)
 Both work and heat aren’t state functions, Δu is a state function
 w = F × d or w = ma × d (units in Joules)
 Δu = qp + w = qv (qp is heat of reaction when P constant, qv is heat of reaction when V constant)
 Calculating work done: pΔV = RTΔn

Measuring Heats of Reactions
 Heat capacity (C): amount of heat required to change the temp of a system by 1 K
 Specific heat capacity (c): amount of heat required to change temp of 1 g of substance by 1 K
 Molar heat capacity (Cm): amount of heat required to change temp of 1 mol of substance by 1 K
 Heat capacity: q = CΔT Specific heat: q = mcΔTmolar heat: q = nCmΔT
 Calorie (cal): amount of heat required to change the temp of 1 g of water by 1°C
 Under constant V: if the heat capacity is known then the increase in temperature gives us the heat of the
reaction qcal = CΔT = –qrxn
 Under constant P: qH2O = – qrxn in most cases qcal ~ 0, –msolidCsolidΔTsolid = mH2OCH2OΔTH2O

Using Hess’s Law to Determine ΔH (Enthalpy)
 Extensive properties: depends on the size or amount of material in system, ex. mass, volume
 Intensive properties: don’t depend on the size or amount of material in system, ex. temp, density
 Enthalpy change is extensive
 ΔH = Δu + pΔV
 The overall change in enthalpy is the sum of the enthalpy changes in the individual steps
 ΔHtotal = ΔH1 + ΔH2 + ΔH3 + … + ΔHn

Standard Enthalpies
 Standard enthalpies of reactions (ΔrH°): enthalpy change of reaction measured at STP
 Standard enthalpies of formation (ΔfH°): enthalpy change that occurs in formation of 1 mol of a substance
in the standard state, ΔfH° = 0 for pure elements in their standard state
 ΔrH° = Σ nΔfH°products – Σ nΔfH°reactants




CHAPTER 14: KINETICS
 Chemical kinetics measures the rate at which the concentration of a substance changes over time
 Reaction rate: expressed as rate of disappearance of reactants, rate of appearance of products