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                                                                                                                     *

 

Using kto find the pH of a Base

kw=[H+][OH]

NaOH & KOH - Strong bases - donate 1 mole of OH- ions per mole of base

  • the value of kw @ 298K us 1.0 X 10-14 mol2dm-6
  • find the pH of 0.1 moldm-3 NaOH @298K
    • [OH-] = 0.1moldm-3 => [H+] = kw/[OH-] = 1.0 X 10-14/.01
    • pH= -log10 1.0 X 10-13 = 13

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     *     

 

Diprotic Acids

They release 2 protons when it dissociates

ex. H2SO4(l)+water -> 2H+(aq) + SO4-2(aq)

H+ = 0.2 moldm-3    so....pH = -log10[0.2] = 0.70

 

pH of sulfuric acid 0.25 moldm-3

[H+] = 2 X 0.25 = 0.5 => pH = -log10[.5] = 0.30

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*    

 

pH Definition

  • The measure (from 0-14) of Hydrogen Ion Concentration
    • 0 = Very Acidic
    • 7 = Neutral
    • 14 = Very Alkaline (base/basic)

*Expressed in -log10    --->    pH = -log10[H+] ex. pH = -log10[0.01] = 2

Or...

[H+] = 10-pH       ex. [H+] = 10-1.52 = 0.03moldm-3 = 3 X 10-2moldm-3

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*      

 

 

Acids & Bases

  • Protons are transferred (Acid -> Bases) when A&B react
  • Acids can olny get rid of protons when there is a base to accept them
    • ex. HA(aq)+B(aq) ⇔BH+(aq)+ A-(aq)
  • If acid is addded to water the water acts as the base and accepts the proton
    • HA(aq) + H2O ⇔ H3O+(aq) + A-(aq)

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*      

 

Bases: Strong & Weak

  • Strong Bases - ionise almost completely in water too
    • ex. sodium hydroxide - NaOH(s) + Water ⇒ Na+(aq) + OH-(aq)

  • Weak Bases - only slightly dissociate in water
    • ex. Ammonia - NH3(aq) + H20 ⇔ NH4+(aq) + OH-(aq)

​*Equilibrium lies well over to the left

 

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*          

Acids: Strong & Weak

- Acid releases a proton - A base accepts a proton

Strong Acids - dissociate (or ionise) almost completely in  water - nearly all the H+ ions will be released

(Hydrochloric acid)      ex.  HCl(g) + Water ⇒ H+(aq) + Cl-(aq)

Weak Acids - dissociate only very slightly in water - so only small numbers of H+ ions are formed

(Ethanoic or citric)    ex.  CH3COOH(aq) ⇔CH3COO-(aq) + H+(aq)

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*    

Kinetics: Rate Equation

Rate = k [A]n + [B]m  

[A] - Concentration of A 

n - rate order 

k - Rate Constant 

*Overall order of reaction is n + m

* Increase in temperature will increase k

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*   

Orders of Reaction

* Order can be determined only by experiment using the Method of Initial Rates (the rate for a short time @ the beginning of the reaction is measured @ several different concentrations of reactants)

  1. First Order - X1 = Rate doubles when reactant doubles
    • X2 = X2; X3 = X3; ...
  2. Second Order - X2 = Rate is (x4) when the reactant doubles
    • X2 = X4 (22) ; X4 = X16; .....
  3. Zero Order - X0 = rate stays the same regardless of reactant 
    • X2 = 1; X4 = 1; .....

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*     

Kinetics: Concentration of Catalysts

  • Increase concentration of catalysts = Increase rate of reaction
  • Increase concentration of reactants in a solution, on avg. the particles will be closer together
    • closer ⇒ collide more often ⇒ more collisions ⇒ more chances to react
  • if gases are involved, and increase in pressure of the gas works the same way
  • Catalysts increase rate of reactions too by providing an alternative reaction pathway w/ a lower activation energy​

*Catalyst is chemically unchanged @ the end

 

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*   

Kinetics: Physical State of Reactants

  • Particles must collide to react
    • in the right direction, facing the right way
    • must collide w/ the min. amt. of kinetic energy

(Collision Theory)

*Liquids & Gases best as particles move

*Increase in temp. = particles have more kinetic energy = faster reactions @ activation energy more particles have enough energy @ 35ºC > 25ºC 

 

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*   

Kinetics: Physical State of Reactants

  • Solids - particles very close together.
    • High density & incompressible.
    • Particles vibrate about a fixed point & can't move freely
  • Liquids - Similar density to a solid & is virtually incompressible
    • Particles move freely & randomly w/in the liquid
  • Gas - particles have lots more energy & are much further apart
    • Density is pretty low & it's very compressible
    • Particles move freely, diffuse quickly, no alot of attraction between them

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*   

Gibbs free energy (1)

  • Free enthalpy, Gibbs energy, or Gibbs function
  • ΔG = ΔH - TΔS
    •  H - Heat energy in the system (kj)
    • S - Measure of Entropy (J/kmol)
    • T - temp (K)​
      • Equation to determine how likely a reaction is to take place spontaneously

* If a reaction will take place it reduces Gibbs free energy (ΔG < 0)

*Gibbs energy is reduced if H is reduced

*Gibbs energy is reduced if S is increased

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*   

Gibbs free energy (2)

- Reactions most likely to happen if:  ΔH < 0 & ΔS > 0

Endo & Exothermic

  • Endo - energy taken in so heat in system is increased            ΔH > 0 = unfavourable
  • Exo - energy given off so heat in system is reduced                      ΔH < 0 = favourable

* Endothermic reaction can still take place if it results in a large enough increase in entropy (ΔS>0)

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*   

Gibbs free energy (3)

* at a phase change (gas → water )

  • ΔG = 0
  • ΔG = positive # = not spontaneous
  • ΔG = negative # = spontaneous

ex. ΔG = ΔH -TΔS

  • +      +   : ΔG= negative if temp high
  • -       -   : ΔG= negative if temp low
  • +      -   : ΔG= positive (non spontaneous) always
  • -       +  : ΔG= negative always

 

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*   

Feasibility of Reactions

  • More Negative ( or less positive) EØ value moves left
  • More Positive (or less negative) EØ value moves right

 

Ex. Fe(OH)3(s)+ e- ⇔ Fe(OH)2(s) + OH-(aq)     EØ = -0.56V

O2(g) + 2H2O + 4e- ⇔ 4OH-(aq)     EØ = +0.40V

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*   

Electrode Potential

(conditions affecting it)

  •  Half cell reactions are reversible
    • equilibrium position is affected by changes in:
      1. Temperature
      2. Pressure
      3. Concentration
  • Standard Conditions are:
    1. Temp - 25°C (298K)
    2. Pressure - 100kPa
    3. Concentration - 1.00 moldm-3

 

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*   

Electrode Potential

(standard elctrochemical cell drawings)

  • the potential difference between the electrode & its solution

Eøcell = (Eøright side - Eøleft side)

Zn/Cu cell short hand

Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)                       (Zn2+(aq) + 2e- ⇔ Zn(s) )

—Charges go this way——►                  (Cu2+(aq) + 2e- ⇔ Cu(s))

Reduced| Oxidised || Oxi. | Red. 

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Substitution & Elimination (1)

(Rules & what influences then)

Most important factor: type of Halokane

  • Primary - mostly substitution
  • Secondary - both substitution & elimination
  • Tertiary - mostly elimination
    • Can be influenced by changing conditions​

The Solvent = proportion of ethanol to water

  • more water = more substitution
  • more ethanol = more elimination

Concentration - of sodium or potassium hydroxide solution

  • higher concentration = higher elimination

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Substitution & Elimination

Reactions

Substitution - the halogen is replaced by an  -OH group to give alcohol

  • CH3CHCH⇒ NaOH ⇒ CH3CHCH3 + NaBr
    • |                                          |
    • Br                                       OH

Elimination - also in the presence of Sodium &/or potassium

- hydrogen  bond is removed from one of the end carbon atoms toghether w/ brownie from centre one

  • ​CH3CHCH3 + NaOH ⇒ CH2=ChCH3 + NaBr + H2O
    • ​|
    • Br

 

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SN1  Reactions

SN1 - Nucleophillic substitution 

  • S = Substitution ; N = Nucleophilic ; 1 = the initial stage involves 1 species
  • Faster mechanism
  • best with tertiary halokanes
    • ​ex. R3C - X ⇒ R3C + ⇒ R3C - Nu
      • ​                 ↑
      •                 Nu

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SN2 Reactions

SN2 - best with Primary Halokanes (initial stage 2 species)

  •                            X 
  • ex. R - X → R<          → R - Nu
  •                            Nu

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Fischer Projections

- 2-D representation of a 3-D organic molecule

- Used for carbohydrates but not non-carbohydrates       ex. Fischer

  • Rules                                                                             H        O
    • Carbon Chain - vertical, C1 at top                         \      //
    • Horizontal bonds project toward the viewer          C
    • vertical bonds project away                                      |
    •                                                                           H —  C — OH
    •                                                                                     |
    •                                                                         OH — C — H
    •                                                                                     |
    •                                                                            H — C — OH
    •                                                                                     |
    •                                                                                    H

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Optically Active

  • Rotate plane - polarised light
    • normal light vibrates in all directions but plane - polarised light only vibrates in one direction
  • One enantiomer rotates it in a clockwise direction & the other rotates it in an anti clockwise direction

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Racemates (Isomers)

  • Racemic mixture
  • Contains equal quantities of each enatiomer of an optically active compound
  • don't show any optical activity
    • ​they cancel each other's light-rotating effect
  • ​made typically by reactive 2 achrial things together to get a racemic mixture as the chances of each enantiomer is equal
    • ​ex.    H                                                     CL                                    H
    •           |                                                        |                                       |
    •          C              +    Cl   →  HCl +          C                  or                C   
    •     /   |   \                                      /    |    \                        /   |   \
    • CH3     |      H                               CH3            H               CH3      |     Cl 
    •         C2H3                                            C2H3                               C2H3

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Chirals

- Carbon w/ 4 different groups attached

  • ex.  H   H        O
  •         |     |        /
  •  H—C—C*—C                              * = Chiral Centre
  •         |     |        \
  •        H   OH      OH

Enantiomers: 

  •              H                         |                     H
  •               |                          |                      |
  •              C                         |                     C
  •          /    |    \                    |                  /    |    \
  • HOOC   |      CH3                |          H3C      |      COOH
  •             OH                       |                    OH

 

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Optical Isomers

- type of stereoisomerism - same structural formula but atoms are arranged differently

-Chiral - (asymmetric) carbon atom is an optical isomer that has 4 different groups attached to it. The groups attached to it. The groups can be arranged in 2 different ways so that 2 different molecules are made

* called enantiomers or optical isomers 

  • ex.     H              |               H
  •            |               |                |
  •           C              |               C
  •         /  |  \            |            /   |   \
  •  HOOC   CH3           |    H3C     |   COOH
  •          OH                 |             OH 

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Stereoisomers

- Same structural formula but a different arrangement in space

- w/ a double bond molecules can't twist around so you get either E or Z isomers

  • E (entgegen(opposite)) -  H            CH3
  •                                                \        /
  •                                                  C=C                   E-but-2-ene
  •                                                 /       \
  •                                            H3C       H
  • Z (zusammen (together)) - H3C      CH3
  •                                                                  \        /
  •                                                    C=C                 Z-but-2-ene
  •                                                    /      \
  •                                                  H       H

 

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Isomer Formulas

  • General Formula -                         CnH2n+1OH
  • Empirical (simplest ratio) -            C4H10O
  • Molecular (actual #) -                     C4H10O
  • Structural-                                        CH3CH2CH2CH2OH
  •                         H  H  H   H
  •                          |    |    |     |    
  • Displayed  H–C–C–C–C–O–H
  •                          |    |     |    |    
  •                         H  H   H  H 
  • Skeletal - shows the bonds of carbon only, w/ any functional group
  •                   /\/\OH

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Isomers

  • Molecules with the same molecular formula but different molecular structures

3 Types

  • Chain Isomers - straight & branched carbon skeletons (
    • ex. butane    |    |    |    |
    •                H– C–C–C–C–H
    •                       |    |    |    |
  • Positional Iso - same skeleton but w/ a functional group attached
    • ex. l-Chlorobutane     |    |     |    | 
    •                               H–C–C–C–C–Cl 
  • Functional group Isomers - same atoms arranged into functional groups
    • ex. propanone   H   O  H
    •                              |    ||   |  
    •                        H–C–C–C–H
    •                              |          |

 

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Periodic Table:

Atomic Radius & Ionisation

- Atomic radius decreases across a period

- ↑ # of Protons = ↑ positive charge of nucleus = ↑ pull of centre = ↓ of radius

- Ionisation ↑ across a period

  • ↑ attraction between outer shell & nucleus