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R's Chem 2 > Exam 3 - Equilibria > Flashcards

Exam 3 - Equilibria Flashcards

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1
Q

Explain the Common Ion Effect

A

The presence of a common ion suppresses the ionization of a weak acid or weak base.
We have the initial concentration of the reactant but we now also have the initial concentration of a product. Polyprotic acids are an example of this.

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2
Q

Henderson Hasselbalch equation

A

pH= pKa + log( CB/WA)
Can only use if you can assume x is small
used to calculate the ratio of base to acid in order to reach a specified pH. Or to find pH when given amounts.

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3
Q

Buffer

A

A solution that contains a weak base and conjugate acid or a weak acid and conjugate base . Buffer solutions resist changes to pH.

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4
Q

What makes a “good” buffer?

A

WA/CB or WB/CA ratio is between 10-1-0.1
must be a lot of WA/CB or WB/CA
the pKa or pKb is near the target pH

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5
Q

Titration

A

The slow addition of one solution of a known concentration (called a titrant) to a known volume of another solution of unknown concentration until the reaction reaches neutralization, which is often indicated by a color change

Uses a balanced equation and mole to mole ratio

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6
Q

Three types of titraitons

A

Acid Base- unknown acid with known base
Redox- oxidizing agent with a reducing agent
Precipitation-reactants that form an insoluble product

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7
Q

Types of acid based titrations

A

Strong Acid, Strong Base
Weak Acid, Strong Base
Weak Base, Strong Acid

In this class we will NEVER titrate a weak acid and weak base. We want to force the equilibria to one side, both weak will never go to completion.

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8
Q

Indicator

A

A weak acid or base that changes color when ionized or de ionized.
Used to identify when we have reached the endpoint.

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9
Q

How to choose an indicator

A

Similar to choosing a buffer: we want an indicator with a pKa or pKb near the pH where the A/B ratio changes

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10
Q

How do you calculate the titration of a Strong Acid with a Strong Base?

A

Stoichiometric relationship
Get both molarities into moles (M x L). Subtract the larger mole from the smaller (which is the limiting agent), and convert back into Molarity (moles/total L). -log of that Molarity. Make sure you know if you are calculating pH or pOH due to larger concentration of H or OH

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11
Q

How do you calculate the titration of a weak acid with a strong base?

A

Stoichiometric and equilibrium relationship, has a buffer capacity
Get both molarities into moles (M x L). Subtract larger mole from the smaller (limiting agent). Convert back into Molarity (moles/ total L). Now you need an ICE table since a weak acid will not fully dissociate. Need a Ka/Kb to solve! If weak acid, switch to Kb after the half equivalence point.

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12
Q

Equivalence Point

A

where equal molar amounts of acid and base have been combined

“Amount of moles of titrant equals the amount of moles of unknown”

the point at which the acid has been neutralized completely by the base

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13
Q

Endpoint

A

The point at which the color of the indicator changes. The point at which we know all of the unknown has been reacted

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14
Q

Molar Solubility (M)

A

M= moles/Liter
The number of moles of solute in 1 liter of saturated solution (mol/L)

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15
Q

Solubility Product Constant (Ksp)

A

The equilibrium constant that indicates to what extent a slightly soluble ionic compound dissolves in water
the equilibrium of insoluble ionic compounds. (Ksp= [product][product], no reactant since its solid)

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16
Q

What determines the solubility of a product?

A

The Ksp.
Ksp < 1.0 x 10^ -5 is insoluble
Ksp from 1.0 x 10^ -5 to 1.0 x 10^ -3 is slightly soluble

K > 1.0 x 10^ -3 is soluble

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17
Q

Selective Precipitation

A

A technique of separating ions in an aqueous solution by using a reagent that precipitates one or more of the ions, while leaving other ions in solution

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18
Q

Complex Ion

A

Compounds that can form a series of covalent bons while still maintaining a charged state. Usually form when metal ions dissolve in water.
Very Large Keq: Kf

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19
Q

Ligand

A

Ions or neutral molecules that bond to a central metal atom or ion. Ligands act as Lewis bases (electron pair donors)

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20
Q

Polydentate Ligand

A

Ligand that contributes or accepts multiple lone pairs. Metal central atom, usually a transition metal. Ex: EDTA

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21
Q

Formation Constant (Kf)

A

The equilibrium constant that indicates to what extent complex-ion formation reactions occur

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22
Q

Calculate the pKa given a Kb

A

1.0 x 10^-14 / Kb = Ka.
-log(Ka)=pKa

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23
Q

What is the importance of the 1/2 equivalence point

A

It represents the point at which exactly half of the acid in the buffer solution has reacted with the titrant. The half equivalence point is relatively easy to determine because at the half equivalence point, the pKa of the acid is equal to the pH of the solution.
“where moles of WA = moles of CB or vice versa”—- because log(1)=0, pH=pKa

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24
Q

Label on a graph the endpoint and equivalence point

A

Endpoint is color change
Equivalence point is end of chemical reaction

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25
Q

Identify a weak acid/ strong base titration curve from a strong acid/weak base titration curve

A

In a weak acid-strong base titration, the pH is greater than 7 at the equivalence point. In a strong acid-weak base titration, the pH is less than 7 at the equivalence point

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26
Q

What is the ratio for mixing buffers?
WA/CB
WA/SB
SB/WA

A

WA/CB 1:1
WA/SB 2:1
WB/SA 2:1

27
Q

Explain the graph of a strong acid strong base titration

A
28
Q

Explain the graph of a weak acid strong base titration

A
29
Q

Explain the graph of a strong acid weak base titration

A
30
Q

What is the pH at the equivalence point of:
strong acid strong base
weak acid strong base
weak base strong acid

A

strong acid strong base: 7.00
weak acid strong base: above 7.00
weak base strong acid: below 7.00

31
Q

What is Ksp? How does our ice table change

A

Ksp= solubility product.
Ksp= [prod] [ prod]
solid –> [aq] + [aq ] so ICE table is shifted, ignoring the reactant.

The smaller the Ksp the less soluble it is

32
Q

If Q is less than or equal to Ksp, will precipitate form?

What if Q is greater than Ksp?

A

Q </= Ksp No, precipitate does not form

Q> Ksp = Yes, precipitate forms

33
Q

What is Oxidation Reduction (Redox)

A

A transfer of electrons. Must take into account oxidation states and elections

34
Q

An Oxidized substance has _______ electrons while a reduced substance has _____electrons

A

An Oxidized substance has ___Lost____ electrons while a reduced substance has __Gained___electrons

35
Q

When an oxidation state INCREASES from reactants to products we say that the substance has been _______

When an oxidation state DECREASES from reactants to products we say that the substance has been _______

A

When an oxidation state increases from reactants to products we say that the substance has been __oxidized_____

When an oxidation state DECREASES from reactants to products we say that the substance has been __reduced_____

36
Q

If a substance RELEASES electrons we would call that substance a _________ agent.
If a substance GAINS electrons we would call that substance a ________ agent.

A

If a substance released electrons we would call that substance a reducing agent.
If a substance gains electrons we would call that substance an oxidizing agent.

37
Q

Label the reducing and oxidizing agents and find oxidations states
Zn(s) + 2HCl = ZnCl2 + H2

A
38
Q

When balancing redox reactions, what two things must each step have before moving on?

A

Must be balanced by mass and by charge.
“free” Electrons should cancel

39
Q

Steps to balancing a redox reaction

A
  1. Identify oxidation states
  2. Identify what is reducing step and oxidizing step.
  3. Balance in acidic solution : Balance by mass and by electrons, adding H+ and H2O as needed
  4. Balance in basic solution: change the H+ to H2O by adding OH- to that side, then add the same amount of OH- to other side, ensure balanced by mass and electrons
  5. make electrons equal in both half reactions. Add and cancel equations.
40
Q

Voltaic cells (Galvanic cells)

A

thermodynamically spontaneous: have a negative delta G.
Ex: battery usage.

41
Q

Electrolytic Cells

A

involve reactions that are thermodynamically nonspontaneous (positive delta G) but have been provided an extra “push” from an external source in order to proceed.
Ex: recharging batteries

42
Q

Anode

A

Site of oxidation of the reducing agent.
Loses electrons

43
Q

Cathode

A

Site of reduction of the oxidizing agent.
Gaines Electrons

44
Q

Salt Bridge

A

Used to allow ions to migrate towards anodes and cathodes as chemical potential changes due to electron transfer.
Usually K+ and NO3-

45
Q

Draw a voltaic cell transfer with all parts

A

Always indicate anode first, then salt bridge, then cathode.
ex: Fe/Fe+2 l KNO3 l MnO4-/Mn2+e
anode bridge cathode

46
Q

Why do we need a salt bridge

A

because as the anode electrons are moving to the cathode they are decreasing the charge. When it becomes too negative, electrons do not want to continue in that direction (opposites attrack, lechatliers principle..) so the salt bridge keeps the balance to allow more electrons to move right.

47
Q

Electrochemical Cell Potential

What is its unit of measure?

A

potential energy residing in an electrochemical reaction
Ex: water building up pressure behind a damn. Higher the damn, the higher the potential energy
Unit of measure is Voltage

48
Q

Electrochemical Cell

A

The difference in electric potential between the cathode and the anode
Ecell= Ered + Eoxidation

Red Cat pleases An Ox
(Reduction/Cathode plus Anode/Oxidize)

49
Q

Electromotive force (EMF)

A

a difference in potential that tends to give rise to an electric current.

50
Q

Voltage vs Current

A

Voltage is the force pushing charges.
Current is the flow rate of charges.

Water example: Voltage is the size of pipe water is flowing through. Larger the current the more electrons are flower. The larger the voltage the more energy that current can have.

51
Q

How to predict if a reaction is spontaneous

A

ΔG = −nFEc
n= moles of electrons transfered
F= Faradays constant 96,500 C/mol
Ec=electrochemical standard cell potential in volts

Since n and faradays are always +, the sign of Ec determined spont.
- Ec = + G, so non spont.
+ Ec = - g, so spont

52
Q

Finding Ecell

A

Ecell= Ered + Eoxi
Use table 19.1 to find E of ox and red. recognize forward vs backward reactions and change sign as needed.

53
Q

What is the Nernst Equation and why do we need it?

A

An equation relating the emf of a galvanic cell with the standard emf and the concentrations of reactants and products

54
Q

How does changing Q effect Ecell? Use the Nernst equation to predict direction of shift

A

increase reactants pushes equation forward/spontaneous/ Increase Ecell to make negative G
Decrease reactants pushes equation in reverse, non spontaneous, lower Ecell positive G.
Increase products pushes to reverse, non spont, lower ECell positive G
Decrease products pushes forward, spontaneous, higher ecell, negative G
E> E standard when reactants are more than products
E< E standard when reactants are less than products

55
Q

If Ecell is positive the reaction is ________
If Ecell is negative the reaction is ______

A

+ =spontaneous because it results in a - G
- = non spontaneous

56
Q

What does it mean when Ecell = E^0cell?

A

when Ecell equals the Estandard cell we are at equilibrium

57
Q

Florine, at the top right of the PT, is a strong what in terms of activity series and potential?

A

Florine is a strong oxidizing agent (steal electrons) because it has the largest EN. It will get reduced and has highest reduction potential (top of the list)

58
Q

What equations are needed when asked to find K given a reaction? (Electrochem)

A
  1. Ecell = Ean+ Ecath
  2. G= -nF Ecell
  3. G= -RT lnKeq but manipulated

OR use manipulated Nernst :
10 ^ n*Ecell / 0.0592

59
Q

What 3 things does the Nernst equation allow us to predict?

A

The electrochemical cell potential under any conditions
the direction the electrochemical cell will proceed when Prod > Reac
The direction the electrochemical cell will proceed when Prod < Reac

60
Q

What are the two types of rechargeable batteries and the two types of non rechargeable batteries?

A

Rechargeable: Ni-Cd and Ni-Metal hybrid
Nonchargeable: Dry Cell and Lead Acid

61
Q

What is electrolysis

A

A form of electrochemistry that requires an outside source of electrons to drive the redox reaction that otherwise would be nonspontaneous.
ex: H2O Ecell is -2.06. In order for this to be spontaneous (since - Ecell is +G and non spon) an oustide source must provide at least 2.06 V.

62
Q

Voltaic vs electrolytic cells

A

Voltaic are spontaneous
Electrolytic are non spontaneous

63
Q

label the following electrochemical cell notation:
Zn/Zn+2 // H+/H2

A

Zn –> Zn+2 + 2e- Oxidizing step/ Anode
// salt bridge
2H+ + 2e—> 2H Reduction step/Cathode

64
Q

List of Strong Acids and Bases

A

Acids: HCl, HBr, HI, HNO3, HClO4, H2SO4
Bases: LiOH, NaOH, KOH, RbOH, Sr(OH)2, Ba(OH)2

R's Chem 2 (5 decks)

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