IB Chemistry Topic 9
Terms in this set (34)
a chemical reaction in which changes in the oxidation state occur
Colour change Potassium Dichromate (K2Cr2O7)
K2Cr2O7(aq) is orange
So with Cr3+(aq), which is dark green, the solution changes from orange to green.
Loss of electrons
Gain of hydrogen
Oxidation occurs when there is an increase in oxidation state of an element
Gain of electrons
Loss of hydrogen
Reduction occurs when there is a decrease in oxidation state of an element.
shows what happens to each of the reactants, allowing for the examination of the movement of electrons
Key Oxidation Numbers
- elements in their free state have an oxidation number of zero
- oxidation number usually matches to group number.
oxidation numbers to remember:
H = +1 (in metal hydrides (NaH) it = -1)
O = -2 (except in peroxides, where it = -1, or in OF2 where it = +2)
Li, Na, K = +1
Mg, Ca = +2
F, Cl = -1
monoatomic ions have same oxidation as charge.
eg. Na+ = +1 and Ca^2+ = +2
Oxidation number definitions
the number of electrons lost by an atom of that element in a compound.
positive is lost, negative is gained.
Sum of oxidation numbers in a compound must = 0.
In a polyatomic ion, the oxidation numbers must add to equal the charge (eg. NH4+ = +1)
In covalent compounds, the negative oxidation is given to the most electronegative atom.
The sign MUST BE WRITTEN FIRST
ion with a positive charge
ion with negative charge
polyatomic and monoatomic ions
ions (charged) with one or two atoms/elements
mono = Na+
poly = NH4+
F > O > N > Cl > Br > I > S > C > H >>> metals
When a more reactive metal takes the place as ions, making the previously ionic metal return to its element state.
the substance that is oxidised
It's oxidation state becomes more positive
More reactive metals are stronger reducing agents
the substance that is reduced
It's oxidation state becomes more negative
More reactive non-metals are stronger oxidising agents than less-reactive non-metals.
A more reactive non-metal is able to oxidise the ions of a less reactive non-metal (the more reactive one is able to gain the electrons so is able to be reduced)
determine the unknown concentration of a substance in a solution
It is a titration between an oxidising agent and a reducing agent, allowing for the transfer of electrons from the reducing agent to the oxidising agent.
finds the equivalence point where they have reacted by transferring electrons.
Redox titration of Iron with Magnanate(VII)
KMnO4 is the is the oxidising agent, because Fe is the more reactive metal
5Fe^(2+) + MnO4 --> 5Fe^(3+) + Mn^(2+) + 4H2O
Turns from purple to colourless
The point where the colour becomes colourless is the equivalence point.
Biological oxygen demand
the quantity of oxygen needed to oxidize organic matter in a sample of water over a 5 day period at a specified temperature.
measure of degree of pollution
High levels mean low levels of oxygen --> unhealthy
Dissolved oxygen content
a measure of the health and quality of a body of water
Can be used to measure dissolved oxygen content of water
Uses redox titration
1. Dissolved oxygen in the water is fixed by the addition of Manganese(II) salt (MnSO4). Causes oxidation of Mn(II) to a higher oxidation such as Mn(IV)
2Mn^(2+) + O2 + 4HO- --> 2MnO2 + 2H2O
2. Acidified iodide ions (I-) are added to the solution, and oxidised by Mn(IV) to I2
MnO2 + 2I- + 4H+ --> Mn^(2+) + I2 + 2H2O
3. The iodine produced is then titrated with the sodium thiosulfate.
2S2O3^(2-) + I2 --> 2I- + S4O6^(2-)
Every 1 mole of O2 in the water uses 4 moles of S2O3^(2-)
Voltaic Cell vs Electrolytic Cell
oxidation at anode (+)
reduction at cathode (-)
oxidation at anode (-)
reduction at cathode (+)
Voltaic (generate electricity from chemical reactions)
Electrolytic (drive chemical reactions using electrical energy)
Voltaic Half Cells Explained
Generate electricity from spontaneous redox reactions.
- The reaction involves oxidation of one metal into its ions, and another is reduced into its atom state. This releases energy (exothermic) because bonds are being broken.
- The two reactions are separated into two half-cells, with a bridge that allows the transport of electrons between them.
- when a metal strip is added, the metal atoms will lose electrons to form ions, creating an electrode potential when there is a charge separation with the surface of the metal and the ions in the solution.
- At the same time, the ions are gaining electrons and forming its atom state again, resulting in an equilibrium.
- The equilibrium is happening between both metals, and the reactivity will determine the position of the equilibrium.
Converts the energy released from a spontaneous exothermic reaction into electrical energy.
Is two joined half-cells connected by an external wire.
Oxidation occurs at the anode (more reactive metal)
Reduction occurs at the cathode (less reactive metal)
Anode is -
Cathode is +
External wire - anions ALWAYS move from anode to cathode.
Salt bridge - anions from cathode to anode. Balance potential charge difference by movement of anions and cations, maintaining electrode potential.
Voltaic Cell diagram convention
single line between the atom and the ion state.
two lines indicate salt bridge
aqueous solutions next to the salt bridge
anode of left and cathode on right, so electrons flow left to right.
Zn(s) | Zn^(2+)(aq) || Cu^(2+)(aq) | Cu (s)
the more reactive metal will have a more negative electrode potential in its half cell.
Electrolysis of NaCl in molten state
1:1 mole ratio
2Na+ + 2 e- --> 2Na (negative electrode = cathode, reduced)
2Cl- -2e- --> Cl2 (positive electrode = anode, oxidised)
--> 2Na+ + 2Cl- --> 2Na + Cl2
Electrode cell potential
= Ered - Eox
* Ered is the more positive standard electrode potential
* Eox is the more negative standard electrode potential.
Zn/Zn2+ and Fe/Fe2+
--> total electrode cell potential
= -0.45 + 0.76
= 0.31 V
Electrolytic Cell Explained
The reverse of a voltaic cell
An external source of electricity drives non-spontaneous redox reactions.
- As the electrical current is passed through the electrolyte, redox reactions occur at the electrode, removing the charges on the ions an forming products that are electrical neutral (ions are disCHARGED)
- Electrodes do not take part in the reaction, and are just made from conducting substances.
- The power source pushes electrons towards the (-) cathode and into the electrolyte
- Electrons are released at the (+) anode and returned to the power source.
- The current moves through the mobile ions (not electrons) and are removed from the solution by the reaction occurring at each electrode.
- To extract positive ions into metals, they have to be reduced (making them not ions).
- Ions move to their oposite charge, so cations move towards the (-) cathode, and the anions move to the (+) anode.
- REDUCTION AT CATHODE (-)
- OXIDATION AT ANODE (+)
The power comes from a DC power source (drawn with big line (+) and little line (-)
The anode is +
The cathode is -
Electrodes are connected to the power source and put in the electrolyte, not touching
Uses of electrolysis
Extracts metals from their ores, when there are no other metals strong enough to do it, so it must be done directly by electrons (electrical current)
A liquid - can be molten ionic compound, or a solution of an ionic compound.
As the electrical current is passed through the electrolyte, redox reactions occur at the electrode, removing the charges on the ions an forming products that are electrical neutral
is in the data booklet
Higher on reactivity series --> loses electrons more easily --> oxidised more easily--> is the stronger reducing agent
Electrolysis in Aqueous
water can be oxidised to oxygen at the anode
water can be reduced to hydrogen at the cathode
Aqeous solution electrolysis
The gasses produces are the ions more likely to be discharged.
use discharge series.
ions: Na+, Cl-, OH-, H+
discharged: Cl- and H+
so gasses produced are Cl2 and H2
Ox at anode: 2Cl- - 2e- --> Cl2
Red at cathode: 2H+ + 2e- --> H2