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04 Kémiai reakciók

2014.04.09

 

Chemical Reaction
 
Chemical reaction: a process that leads to the transformation of one set of chemical substances to another: a new substance is made
Changing of energy
·         Exothermic: heat is given out (temperature of the surroundings goes up) ΔHreaction is negative
·         Endothermic: heat is taken in (temperature of the surroundings goes down) ΔHreaction is positive
ΔHreaction (standard enthalpy of reaction): amount of energy change if the reaction occurs by the equation
ΔHformation (standard enthalpy of formation): amount of energy change if the compound is formed by the equation
Hess’ Law: enthalpy change for any chemical or physical process is independent of the pathway or number of steps (if a chemical change takes place by several different routes, the overall enthalpy change is the same)
 
Criterions of chemical reactions
Collision of the particles is needed
·         Homogenous reaction: reactants are in the same state (gas mixtures, solutions)
·         Heterogenous reactions: reactants are in different states (reaction occurs on the surface)
Increasing the concentration – number of the collisions increases
Successful collision: particles are in a special position (M fig. 91/2)
In a successful collision: reactants are joined, old bonds get weaker, new bonds are creating (Active state)
 
Activation energy (Ea): is needed old bondings to loosen
 
Reaction rate
Slow reaction: rusting, rot
Fast reaction: burning, explosion
Rate ~ concentration of reactants
r = Δc/Δt                     (r=rate, c=concentration, t=time)
H2 + Cl2 = 2HCl
r ~ [H2][Cl2]                                (rate is directly proportional to the concentration of the reactants)
r = k [H2][Cl2]             (k=constant, depends on the reaction)
Reaction rate decreases during the reaction (concentration of the reactants decreases)
Rate depends on:
·         Concentration of reactants
·         Temperature: temperature increases – kinetic energy of the particles increases – rate increases
·         Pressure (in gases): particles are closer – number of collisions increases – rate increases
·         Surface area
·         Stirring up
·         Catalysts: connecting to the reactant particles - activation energy is decreased (old bondings are loosened, less energy is enough to break them)
·          
Equilibrium
CO2 + H2O = H2CO3
Reaction is reversible: occurs in two directions (forwards: reaction 1, backwards: reaction 2)
Dynamic equilibrium: rates of the reactions are the same
r1 = r2
k1 [CO2][H2O] = k2 [H2CO3]
k1 / k2 = [H2CO3] / [CO2][H2O]
k1 / k= K (equilibrium constant)
Meaning of K:
Value of K is small: amount of k2 is larger – reactants are dominant
Value of K is large: amount of k1 is larger – products are dominant
How to effect the equilibrium?
3H2 + N2 = 2NH3
·         Increase the concentration of reactants
·         Decrease the concentration of protucts
·         Increase the pressure
·         Medium temperature
·         Catalyst
 
Acid-base reactions
Acids:
HCl + H2O = H3O+ + Cl-
In water acids increase the concentration of H3O+ (hydronium) ions
Acids are proton (H+) donors
Important acids:
·         Hydrochloric acid (strong) - HCl
·         Sulphuric acid (strong) - H2SO4
·         Nitric acid (strong) - HNO3
·         Phosphoric acid (medium) - H3PO4
·         Acetic acid (medium) – CH3COOH
·         Formic acid (medium) – HCOOH
·         Carbonic acid (weak) – H2CO3
Dissociacion constans (Ka): see ChFY p. 385, M 103
·         Value of Ka is small: weak acid
·         Value of Ka is large: strong acid
Bases:
NH3 + H2O = NH4+ + OH-
In water bases increase the concentration of OH- (hydroxide) ions
Bases are proton (H+) acceptors
Important bases:
·         Ammonia (weak) - NH3
·         Sodium hydroxide (strong) - NaOH
·         Potassium hydroxide (strong) - KOH
 
Conjugated acid-base pairs:
HNO3 + H2O = NO3- + H3O+
Acid 1       Base 2     Base 1       Acid 2
Water:
It can be acid or base (it depends on the partners) – amphoteric
Autoprotolysis:
H2O + H2O = H3O+ + OH-
Equilibrium constant: Kw, see M 104
Kw = [H3O+][OH-] = 10-14 (mole/dm3)2
 pH
We can express the acidic-alcalic conditions of a solution with numbers
pH = -lg[H3O+]
A pH az oxóniumionok koncentrációjának 10 hatványaként felírt alakjának hatványkitevője ellenkező előjellel
[H3O+] = 10-1, pH = 1
[H3O+] = 10-7, pH = 7
[H3O+] = 10-14, pH = 14
Indicators
Change of color shows the acidic-alcalic conditions
Methyl orange
Methyl red
Phenolphtalein
Natural indicators: litmus, flowers, tea
 
Neutralisation
 acid + base = salt + water
HCl + NaOH = NaCl + H2
Tasks:
H2SO4 + NaOH =
HNO3 + KOH =
H3PO4 + NaOH =
HNO3 + NH4OH =
Hydrolysis
salt + water = acid + base
Na2CO3 + H2O = H2CO3 + NaOH
Weak acid and strong base: alcalic solution
NH4Cl + H2O = HCl + NH4OH
Strong acid and weak base: acidic solution
NaCl + H2O = NaOH + HCl
Strong acid and strong base: neutral solution
 
Reduction, oxidation
Oxidation: loss of electrons
Reduction: gain of electrons
Reducing agent: gives electrons
Oxidizing agent: accepts electrons
Redox reaction: oxidation and reduction occurs at the same time
Oxidation state, oxidation number
Oxidation state (oxidation number): number of lost or gained electrons is shown
Oxidation number rules:
The oxidation number of
·         any uncombined element is 0.
·         any simple ion equals in charge
The sum of the oxidation numbers
·         in a compound equals 0.
·         in any polyatomic ion equals the charge of the ion
Important oxidation numbers:
·         Group 1 metals: +1
·         Group 2 metals: +2
·         Al: +3
·         H: +1 (except in hydrides: -1)
·         O: -2
·         Halogenes: usually -1

How to equalize equations: in a reaction the overall reaction number change is zero. (Ratio of the