Final Exam Review - Summary General Chemistry II PDF

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CHEM*1050 FINAL EXAM NOTES Thermodynamics: the science of the interconversion of different forms of energy Thermochemistry: the study of heat given off or absorbed in a chemical reaction. Chemical reactions involve a transfer of energy from the system and the surroundings. System: part of the universe under observation. Surroundings: everything else The state of the system: is characterized by a set of variables. These can be classified as Intensive or extensive. Intensive: independent of size. (size does not matter) Extensive: dependent on size (size does matter) Internal Energy (U): The sum of the internal kinetic and potential energy of the system. There are two ways to change U: 1. Heat (q) 2. Work (w)

But we can only measure changes in U, not the U at a specific time. The first Law of Thermodynamics: We cannot create or destroy energy, but it can be transferred in or out of a system. Usually q will move from hot to cold. Temperature: an indication of heat content. Objects with the same temperature have a q=0. If temperatures are different, heat flows from high temperature to low temperature. Exothermic Reaction: is q is released in a reaction. Implies that the bond energies of all the reactants is less than the bonding energies of the products. Most combustion reactions are exothermic! (SIGN IS NEGATIVE!) Endothermic Reaction: if q is absorbed in a reaction (SIGN IS POSITIVE!) Enthalpy Change: the heat given off or absorbed at constant pressure! (∆H) and qp=∆H “Open System”- both mass and energy may leave and enter “Closed System” – energy can be exchanged but no mass can enter or leave. “Isolated System” neither mass nor energy may enter or leave

Specific Heat Capacity (c): The energy required to heat a specific material by 1 K. This is an extensive property. Units J/K.

𝑐=

𝑄 𝑚∆𝑇

𝑄 = 𝑚𝑐∆𝑇

Heat capacity (C): the heat capacity of the whole system, regardless of what material is in it. This is an intensive property. Units J/Kmol

𝑄 = 𝐶∆𝑇 In each case we should indicate how the heat capacity is measured. Under constant pressure use Cp and under constant volume use Cv. They also depend on the phase of the substance (ie. C for ice is not the same as C for water) Constant Pressure Calorimeter – Coffee Cup Made up of 2 Styrofoam cups, a lid, a stirrer and a thermometer.

Use these equations!

Constant Volume Calorimeter- Bomb Calorimeter • • • •

Carries out combustion reactions at a constant volume Heat produced raises the temperature of the bomb and of the surrounding liquid. The number of gaseous products formed is not the same as the quantity of the gaseous reactants. Treat the combustion reaction as the system and the entire bomb as the surroundings . Use this to calculate the heat released when one more of the substance is burned! Qbomb= The heat capacity of the bomb x change in temperature

Qbath=The heat capacity of the liquid x the mass of the liquid x the change in temperature

First Law of Thermodynamics:

What about the work being done when the number of moles in the products does not equal the number of moles in the reactants? Compression causes POSITIVE WORK Expansion causes NEGATIVE WORK ∆Vsys=Vp-Vr

The Internal Energy (U) of an ideal gas is dependant only on temperature. Therefore, if ∆T=0 then ∆U=0 as well. That also means that q=-w. Reactions at Constant Volume (Bomb Calorimeter):

𝑪=

𝒒𝒔𝒖𝒓𝒓 −𝒒𝒔𝒚𝒔 −∆𝑼𝒔𝒚𝒔 = = ∆𝑻 ∆𝑻 ∆𝑻

Reactions at a constant Pressure (Coffee Cup): • • • •

The change in volume (∆V) does not equal 0 anymore! Work does not equal 0! W=-Pext∆V Q=∆U + Pext∆V = ∆Hsys

In summary:

When constant volume

When constant pressure

Thermochemical equations must include: • • •

Phases ∆H values with signs (stoichiometric value) Convention: what temperature is this at? If not stated, assume 298K (room temp)

Rule one of Thermochemistry: Thermochemical equations describe the energy changes associated with one mole something, therefore when a thermochemical equation is multiplied by a factor, so is ∆H

Rule Two of Thermochemistry: If we reverse the direction of a reaction, we change the sign of the ∆H value from negative to positive or positive to negative. This includes phase transitions Rule Three of Thermochemistry: Thermochemistry and Thermodynamics is only concerned with the energy differences between the initial and final states, not what is in between. Therefore, we can add, subtract thermochemical equations to produce other equations using Hess’s Law.

Changes of Physical State A. Enthalpy of Fusion (melting) • Enthalpy change when we fuse (melt) 1 mole of something isothermally • Always positive • For the reverse (freezing, solidification) we need to make sure that the sign is negative, as we are “cooling” the system B. Enthalpy of Vapourization • Enthalpy change when we vapourize 1 mole of liquid • Always positive • Reverse process is condensation and is negative! C. Enthalpy of Sublimation • The enthalpy when we sublime 1 mole of a solid directly to gas • Always positive • Is usually equal to the fusion values plus the vapouration values.

Enthalpy changes for chemical reactions: A. Enthalpy of Combustion • Enthalpy change when we burn 1 mole of something completely in oxygen. • Almost always negative, as energy is almost always released Standard Enthalpy of Formation: The amount of energy needed to form a product in its standard state.

Note: the enthalpy of formation for H+(aq) is equal to 0 Bond Enthalpies: The energy required to break a mole of bonds in the gas phase only. AB→ →A+B ALWAYS ENDOTHERMIC. Products are neutral species not ions. Always make sure there is only one BOND in your final answer.

In general the formula to use is:

But be careful! We are used to products-reactants and this is the opposite! As Bond Dissociation Energy gets larger, bond order increases and the bond length gets shorter. Lattice Enthalpy: change in energy that occurs when an ionic crystalline solid is separated into isolated ions in the gas phase. Always positive. The magnitude of the lattice enthalpy depends on the magnitudes and signs on the charges. Basically, the larger the radius of the atom the weaker the lattice energy, and the greater the ABSOLUTE VALUE of the charge, the greater the lattice energy.

Spontaneous Reaction: a reaction that will occur without external intervention. The reverse reaction is not spontaneous at all. Some spontaneous processes: A. B. C. D.

Chemical reactions Heat flow from hot to cold objects Mixing of two ideal gases (∆H=0) The expansion of an ideal gas in a vacuum (∆H=0)

Consider heat flow from hot to cold: Heat is the product of two terms: • •

Driving factor (T) Quantity or capacity factor (Entropy (S))

𝑄 = 𝑇∆𝑆𝑠𝑦𝑠 ∆𝑆𝑠𝑦𝑠 =

𝑞𝑟𝑒𝑣 𝑇

Where qrev is the heat supplied reversibly Entropy: A measure of randomness/disorder. It measures how dispersed the energy of a system is among the different ways that a system can contain energy. There are two situations: Positional: distribution of species in a space Thermal: distribution of energy among a species or the distribution of species over energy levels.

The Second Law of Thermodynamics: in any spontaneous reaction the entropy of the universe increases. There are 3 scenarios here: Spontaneous process: the entropy of the universe is greater than 0 Equilibrium: The entropy of the universe stays the same Non-spontaneous: The entropy of the universe is less than 0

Third Law of Thermodynamics: A perfectly crystalline substance at 0 Kelvin has S=0

Some reactions are considered to be “Enthalpy Driven” (∆Ssys is small and ∆Hsys is less than 0) and others are “Entropy Driven” (∆Ssys is less than 0 and ∆Hsys is small) If ∆Hsys is less than 0 and ∆Ssys is greater than 0 the reaction will ALWAYS BE SPONTANEOUS If ∆Hsys is greater than 0 and ∆Ssys is less than 0 the reaction will ALWAYS BE NON-SPONTANEOUS

Gibbs Free Energy (∆Gsys):

Remember that ∆Suniv is equal to the ∆Ssys+∆Ssurr

Q is defined by this

And K is the equilibrium constant When the K value is positive, G is negative When the K value is negative, G is positive

When K is greater than Q spontaneous from left to right. When K is less than Q, spontaneous from right to left, (not spontaneous as written) When Q =1 the reactants and products are in their standard states

Oxidation Reaction: when the element loses electrons (one or many) Reduction Reaction: when an element gains electrons (one or more) Oxidizing Agent: causes another species to be oxidized by accepting/taking the electrons from it -usually it is being reduced. Reducing Agent: causes another species to be reduced by donating electrons to it – usually being oxidized. How to balance an equation: 1. 2. 3. 4.

Balance the atoms EXCLUDING HYDROGEN AND OXYGEN Balance the Oxygen with water Balance the Hydrogen with H+ Balance the charge by adding electrons a. In basic only – add the same number of OH as you have H+ to both sides! On one side you will create water with the H and OH and the other side will just have residual OH Make sure when you “unite” the reactions that the number of electrons are the same!!!!

Voltaic Cells: an electrochemical cell in which a spontaneous reaction will occur and generate an electrical current. In this set up, the anode is where the oxidation takes place, which is on the left of the conventional notation. The cathode is where the reduction reaction takes place, which is on the right in the conventional notation. To calculate the Ecell use the following equation but remember E0 is NEGATIVE!

𝑬𝒄 = 𝑬𝟎 + 𝑬𝒓 If Ec is positive, the reaction is spontaneous, if not than its non-spontaneous. In the event that you are given the Eo and are asked for the Ecell with a bunch of concentrations/pressures, use this.

Electrolytic Cells: where electrochemistry drives as otherwise non-spontaneous reaction How to solve how much mass/number of moles is deposited: 1. Remember that 96 485C is faraday’s number (F), so divide this by the charge in Amps (A). 2. Multiply by the ratio of moles : electrons 3. Multiply by molar mass factor 4. Multiply by time

Typical Rate Law: Rate: k[A]order Remember if there are multiple terms, the overall order of reaction is the sum of the exponents. Rate of Disappearance: Always the reactants. How fast are they disappearing? Basically means the rate of the reaction with respect to the reactants. BE CAREFUL WHEN THERE ARE COEFFIENCIENTS! Rate of Appearance: Always the products. How fast are they appearing? Basically means the rate of the reaction with respect to the products.

Zeroth Order

First Order

Second Order...