Chemistry - Chapter 6

energy

law of conservation of energy

capacity to do work or produce heat, a state function

energy can be converted from one form to another, but can be neither created nor destroyed

types of energy

radiant energy: comes from the Sun and is Earth’s primary energy source

thermal energy: energy associated with the random motion of atoms and molecules

chemical energy: energy stored within the bonds of chemical substances

nuclear energy: energy stored within the collection of neutrons and protons in the atom (nucleus)

potential energy: energy available by virtue of an object’s position, energy due to position or composition

kinetic energy: energy due to the motion of the object and depends on the mass of the object and its velocity

heat vs. temperature

involves the transfer of energy between two objects due to a temperature differences, not a state function

a measure of thermal energy, not equal to thermal energy

work

state function

force acting over a distance, not a state function

a property that does not depend in any way on system’s past or future, only the present, energy is a state function, work and heat are not

universe

surroundings

system

contains surroundings and system

include everything else in the universe

part of universe where attention is at

endothermic reaction

heat flow is into system

thermal energy is transferred from surroundings to system

energy is absorbed from surroundings

system is positive

if temperature increases, rate of reaction increases

(energy on the reactants side of a chemical reaction)

exothermic reaction

heat flow is out of system

thermal energy is transferred from system to surroundings

energy is released from system

system is negative

if temperature increases, rate of reaction increases

(energy on the products side of a chemical reaction)

thermodynamics vs. thermochemistry

the study of energy and its interconversions, the study of the effects of work, heat, and energy on a system

the study of heat change in chemical reactions

heat energy equation

Q = m c △T

• Q: energy transferred (J)

• m: mass (g)

• c: specific heat capacity (4.184 for water, 0.451 for iron, 0.900 for aluminum)

• △T: temperature change (C)

• J / C = calories

• convert if unit for temperature and calories are not the same

4 quantities of thermodynamics

temperature (T)

internal energy (U)

entropy (S)

heat (Q)

classical thermodynamics vs. statistical thermodynamics

concerns the relationships between bulk properties of matter

seeks to explain those bulk properties of matter in terms of constituent atoms

laws of thermodynamics according to British scientist C.P. Snow

1. you can’t win

2. you can’t break even

3. you can’t get out of the game

1st law of thermodynamics

energy can be converted from one form to another, but can be neither created or destroyed

an extension of the law of conservation of energy

the change in internal energy of a system is equal to the heat added to the system minus the work done by the system

△U = Q ± W

2nd law of thermodynamics

entropy of an isolated system always increases, heat flows spontaneously from a hot object to a cold object (spontaneously meaning without the assistance of external work)

every effort put forth, no matter how efficient, will have a tiny bit of waste

3rd law of thermodynamics

no system can reach absolute zero

this is why the Kelvin temperature scale is used, not only is the internal energy proportional to temperature, but you never have to worry about dividing by zero in an equation

0th law of thermodynamics

two systems in thermal equilibrium with a third system and in thermal equilibrium with each other

internal energy

the sum of the potential and kinetic energies of all the “particles” in the system

thermodynamic quantities

consists of two parts: number gives the magnitude of change, sign indicates the direction of flow

sign reflects the system’s point of view

endothermic process: q is positive

exothermic process: q is negative

system does work on the surroundings: W is negative

surroundings do work on the system: W is positive

1st law terminology

adiabatic: no heat transferred (Q = 0) (specifically for closed, isolated, insulated systems)

isothermal: constant temperature

isobaric: constant pressure

isochoric: constant volume

adiabatic process

a process that transfers no heat

when a system expands adiabatically, W is positive (system does work) so △U is negative

when a system compresses adiabatically, W is negative (work is done on system) so △U is positive

Q = 0

△U = ±W

isothermal process

a constant temperature process

any heat flow into or out of the system must be slow enough to maintain thermal equilibrium

for ideal gases only

△T = 0, △U = 0

±Q = -/+W

isobaric process

a constant pressure process

△U, W, and Q are generally non-zero, but calculating the work done by an ideal gas is straightforward

ex. water boiling in a saucepan

W = ± P x △V

U = ±Q ±W

isochoric process

a constant volume process

when the volume of a system doesn’t change, it will do no work on its surroundings

ex. heating gas in a closed container

△V = 0, W = 0

△U = ± Q

W = ±P x △V

w: atm x L

expanding is -, compressing is +

1 atm x L = 101.3 J

enthalpy

H, the heat content of a chemical system, used to quantify the heat flow into or out of a system

process that occurs at constant pressure, state function

△H = H_products - H_reactants

enthalpy change

the amount of heat released or absorbed when a chemical reaction occurs at constant pressure, kJ/mol

endothermic reaction

energy is absorbed

energy is a reactant of the reaction

reaction vessel becomes cooler and temperature decreases

energy of products > energy of reactant

sign of △H is positive

ex. solid to liquid, liquid to gas

exothermic reaction

energy is released

energy is a product of the reaction

reaction vessel becomes warmer and temperature increases

energy of reactants > energy of products

sign of △H is negative

ex. liquid to solid, gas to liquid

calorimetry

the science of measuring heat, measures the thermal energy (heat) exchanged between the reaction (the system) and the surroundings

by measuring the change in temperature of water

△E

specific heat capacity

the energy required to raise the temperature of one gram of a substance by one degree Celsius

molar heat capacity

the energy required to raise the temperature of one mole of a substance by one degree Celsius

constant volume calorimetry

device called a constant volume (bomb) calorimeter

system: sample is burned in oxygen gas

surroundings: calorimeter

q_cal = C_cal x  △T

q_rxn = -q_cal

use mcat when heat capacity of calorimeter is not mentioned, but when it is use the equation (m = amt of water in mL)

whatever the calorimeter received, the reaction gave away

use heat capacity to find sig figs

△E_rxn can also be expressed per mole of reactant

divide q_rxn by moles of reactant

data from certain experiments would not give the correct molar heats of combustion because calorimeters would not usually have that much energy

constant pressure calorimetry

constant pressure (usually 2 styrofoam coffee cups)

reaction takes place in solution inside inner cup

volume is not constant due to evaporation

q_soln = m_soln x c_{s, soln} x △T

add mass when “dissolved” or “stirred” only if temps are the same, if temps diff then do equations separate and then add

q_rxn = -q_soln

△H_rxn = -q_soln / mol_soln

standard enthalpy of formation (△H_f^0)

the heat change that results when one mole of a compound is formed from its elements at a pressure of 1 atm and a temperature of 298 K

standard enthalpy of formation of any element in its most stable form is zero

they are compiled in huge tables of thermodynamic quantities

calculating △H for a reaction

aA + bB > cC + dD

△H_rxn^0 = c (△H_f^0 (C) ) + d  (△H_f^0 (D) ) - (a  (△H_f^0 (A) ) + b (△H_f^0 (B) ) )

• balance equation

• add the products first then add the reactants, THEN subtract

• ex. I₂, H = 0

• write if exothermic or endothermic

Hess’s Law

when reactants are converted to products, the change in enthalpy is the same whether the reaction takes place in one step or in a series of steps

calculating △H for multiple reactions

work backward from the required reaction, using the reactants and products to decide how to manipulate the other given reactions at your disposal

reverse any reactions as needed to give the required reactants and products

multiply reactions to give the correct numbers of reactants and products

change the sign of and multiply △H as needed