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  • Created by: alida
  • Created on: 11-05-14 12:05

Storing Carbon Dioxide

Carbon dioxide = dissolves in oceans = soluble in water

Large amounts of carbon dioxide can be stored in oceans

Increasing carbon dioxide levels from human activity = contributes to global warming

Carbon dioxide levels increasing

= more is dissolved in the oceans

= helps slow down global warming

Sea is naturally alkaline, carbon dioxide (acidic) dissolves

= makes sea more acidic

= changes the overall pH of the oceans

= harms species, ecosystems, food chains

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Reducing carbon dioxide emission

1. Use less fuel

Increasing efficiency of engines and industrial machines

= Less energy is needed

= less Carbon dioxide is released

2.    Use different fuels

- Nuclear energy

produce a lot of energy

leaks of radiation = health problems

carbon dioxide is produced during mining, transporting, processing

- hydrogen

hard to store, only produces clean products, easily available

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Removing Carbon dioxide

1. photosynthesis

=   reforestation

= more trees

= more photosynthesis

= more carbon dioxide removed from the atmosphere

2. Carbon capture and storage

= collect carbon dioxide from power stations

= pumped into porous rocks under the sea

= stored


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Buffer = solution that resist changes in pH when small volumes of acid/alkali are added

Formed of:

- Acid buffers = weak acid + one of its salts

- Alkali buffers = weak base + one of its salts

2 assumptions:

1. All the  A⁻ ions (anions) come from the salt 

      - contribution of anions from the acid = negligible compared to the fully ionised salt

2. The concentration of HA in the buffer remains unchanged

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Two factors which buffers depend on:

   1.  The Value of Ka

   2. The ratio of [salt] : [acid]

      - Ratio provides a 'fine tuning' of a buffers pH

      - for an effective buffer, the ratio must not be too large or too small


All buffers contain a large amount of proton donors and proton acceptors

= addition of acid/alkali react with the proton donors/acceptor 

= maintains the pH within limits  

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     HA (aq)      <=========>      H⁺(aq)     +       A⁻

Conjugate acid                                           conjugate base

small amounts of acid are added:

= added H+ ions are removed with salt ions present in large amounts

= equilibrium shifts to the left

= maintaining the pH

small amounts of alkali added:

= weak acid dissociates

= produced more H+ ions

= react with OH- ions

= equilibrium shifts to the left

= maintains the pH

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Calculations with buffers

Ka = the acidity constant/ acid dissociation constant

greater the value of Ka = stronger the acid


Ka =   H⁺    x     [salt]  /  [acid]   


      =   H⁺    x     [A⁻]   /  [HA]


[H+] =  Ka     x     [acid]   /   [salt]


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Buffers in action

carbon dioxide dissolving in water:

CO₂ (aq) + H₂O (l) <== ==>   H₂CO₃ (aq)

Carbon dioxide dissolved in oceans/limestone rock

= acts as a buffering system

carbonic acid = H₂CO₃

= acts as a weak acid

limestone = CO₃²⁻

= acts as an anion

= as atomospheric carbon dioxide level increase

= dissolve to form carbonic acid

= pH can be maintained

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entropy change (ΔS) for a system: difference between the entropies of the reactants and products in an equation for a reaction

ΔSsys = ΣΔS(products) - ΔΣS(reactants)

Entropy change for the surroundings:

ΔS = -ΔH x 1000/T            

T = Kelvin = +273 degrees ΔH = Joules

ΔSTotal = ΔSsys + ΔSsurr

ΔS total > 0 = positive = reaction will occur spontaneously

ΔS total = 0 = reaction is at equilibrium

entropy increases = endothermic reactions occur spontaneously

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Lattice Enthalpy

Ionic compounds dissolve in water = ionic lattice is broken up = ions are seperated = become hydrated

Energy is used to break up the lattice Energy is given out when the ions are hydrated

Lattice enthalpy: Enthalpy change when 1 mole of a solid is formed from its seprate gaseous ions

The size of ΔHLE = shows the strength of the bonds in the lattice

Bigger charge/smaller ions=stronger ionic bonds=more-ve= ΔHLE

Ionic radius of an element depends on:

- Nuclear charge/ atomic number = bigger the nuclear charge = the smaller the ions are

- Number of full energy levels = more level = bigger the ion

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Enthalpy change of hydration

ΔHhyd = enthalpy change when an aqueous solution is formed from 1 mole of gaseous ions

= always negative = involves an ion-dipole forming


Greater the ΔHhyd value:

 = the stronger the ion dipole attractions

= greater number of water molecules surrounding the ion


The bigger the charge/smaller the size of the ions

= stronger the ion-dipole bonds

= larger the enthalpy of hydration

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Enthlapy change of solution

ΔH solution = enthalpy change when 1 mole of a solution dissolves to form a dilute solution


2 step process : breaking down the lattice, and then hydrating the gaseous ions produced ( water is the solvent)


ΔH solution= ΔHhyd( cation) + ΔHhyd (anion) - ΔHLE


ΔH solvation = enthalpy change when a solution is formed from 1 mole of gaseous ions using a solvent other than water

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enthalpy change of solution

ΔH solution = negative or slighlty positive

= solid will dissolve due to entropy changes being favoured


If ΔHsolution = large and positive

= solid will not dissolve

= too much energy is needed to break the bonds

 = occurs in ionic solutes in non-polar solvents

= -Hsolv = tiny

= little attraction between ions and solvent

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Acids and Bases

weak acids/bases= partially dissociate in aqueous solutions


Acids  = not all acid molecules donate their protons = equilibrium lies to the left (side of the acid)


Bases = not all bases accept protons = equilibrium lies to the left


Stronger the Acids = the more the equilibrium lies to the right-hand side

Stronger the Base = the more the equilibrium lies to the right - hand side

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Acids and Bases

Acids = proton donors= donate H+ ions

Bases = proton acceptors = accept H+ ions

Strong acids = fully dissociate in aqueous solution = all acid molecules donate their protons

= we can assume that the conc of acid put into the solution = conc of H+ ions in the solution IF the acid = monoprotic acid

monoprotic acid = release one proton per molecule = only have one hydrogen

Strong bases = fully dissociate in aqueous solution = all base molecule accept H+ ions

For strong bases/ acid = equilibrium lies to the right

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Calculations with Acids/Bases

Ka = acid dissociation constant

Greater the value of Ka = stronger the acid

weak acids = smaller Ka value

pKa = -logKa

Ka= 10^-pKa

pH = -log[H+(aq)]


Kw = ionic product of water = [H+][OH-]

Kw = same value at a given temperature


neutral solution = [H+] = [OH-]

Kc = [H+][OH-] / [H2O]

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Assumption for weak acids =

1. Few molecules dissociate = conc of acid = conc given

2. Assume for every 1 mole [H+] = 1 mole of base

HA and A- = conjugate pairs = HA = conjugate acid of A- (vice versa)

Equilibrium with conjugate pairs = set up when a base dissolves in water

In water:

Water and H3O+ = conugate pair

Water and OH- = conjugate pair

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Ionisation Enthalpy

Ionisation enthalpy = energy needed to remove one mole of electrons from one mole of gaseous ions to form one mole of gaseous cation

Increases across a period and up a group


1. Atomic radius

Further away the outer shell from the +ve nucleus = less attraction=lower IE

2. Nuclear charge

More protons = higher attraction between nucleus and outer e' = higher the IE

3. Shielding

More shells = less attraction = lower the IE

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Hydrogen Bonding

Strongest type of intermolecular bonding

- Between H and an electronegative atom (F, O, N)

- electronegative atom attracts electrons away from the H

substances won't dissolve in water unless they can form H- bond with water

The bonds in the solvent and solute break --> new bonds between the solvent and solute form

Won't dissolve if the bonds to be broken are stronger than those that will be formed

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Hydrogen Bonding

Ionic substances: won't dissolve in non-polar solvents

= non polar molecules

= do not interact strongly enough with ions to pulls them away from the ionic lattice

= electrostatic forces between ions

= much stronger than the bonds that could form between ion and solvent

Most covalent substance: only dissolve in non-polar solvents

= bonds between covalent molecules

= weak

= can be broken by non-polar solvents

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Hydrogen Bonding

Ionic solids = dissolve in polar solvents (water)

- The delta +ve H atom = attracted to -ve ions

- the delta -ve O atom = attracted to +ve ions

= ions seperate from the ionic lattice

= become surrounded by water molecules

= ion- dipole forms = each ion-dipole is weak

= total attraction from all the water molecules

= strong to overcome the strong ionic attraction in the lattice

= hydration ( solvent isn't water = solvation)

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Properties of water due to H bonds

1. High boiling point

- H bonds = strong = lot of energy needed to overcome the H bonds in order for them to break (much higher BP than other Gp6 hydrides)

2. High specific heat capacity

- takes a lot of energy to raise its temperature

- large masses of water circulating the ocean = can carry a large amount if energy = distribute it across the worlds oceans

3. High enthalpy of Vaporisation

= amount of energy needed to change a substance from its standard state --> vapour

= water = more energy in order to evaporate

4. Ice is less dense than water

- as water freezes all H bonds that can form are formed = wasted space = less dense

- ice melts = H bonds break = lattice breaks down = water can 'fill the gaps' = increase in density

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