I'm having a hard time with Redox and how to figure out which is redox
Select all of the reactions that are redox reactions.
I) Ca + 2H2O Ca(OH)2 + H2
II) CaO + H2O Ca(OH)2
III) Ca(OH)2 + H3PO4 Ca3(PO4)2 + H2O
IV) Cl2 + 2 KBr Br2 + 2 KCl
I and II
II and III
I and IV
III and IV
What is the oxidizing agent in the following reaction?
Zn(s) + NO31-(aq) Zn(OH)42-(aq) + NH3(aq)
Zn(s)
NO31-(aq)
Zn(OH)42-(aq)
NH3(aq)
And no this isnt homework just "practice probs" and no answer...i don't know how to do them or understand them thanx
look at 1) and iv) Both have valences that are different on each side.
Oxidizing agent=reduced species=thing that gained electrons
Zn lost electrons
N went from +5 (In NO3^-1) to +3 in NH3, so N gained electrons.
So then because the valences are diff on each side that means its a redox reaction??
And so N gained eletrons does that mean it's the NO3 ^-1 that is the oxidizing agent or the N in NH3??
this stuff is tripping me up.....can ou explain it any more for me? sorry...im a little slow with chemistry..
To determine which reactions are redox reactions, you need to identify whether there is a change in the oxidation state of any elements in the reaction. Here's how you can approach each reaction:
I) Ca + 2H2O → Ca(OH)2 + H2
First, write the oxidation states of each element:
Ca: 0 → +2
H: +1 → 0
O: -2 → -2
H: +1 → 0
Since the oxidation state of Ca changes from 0 to +2, it is being oxidized, and the hydrogen in water is being reduced from +1 to 0. Therefore, reaction I is a redox reaction.
II) CaO + H2O → Ca(OH)2
Ca: +2 → +2
O: -2 → -2
H: +1 → +1
O: -2 → -2
In this reaction, there is no change in the oxidation state of any element. Therefore, reaction II is not a redox reaction.
III) Ca(OH)2 + H3PO4 → Ca3(PO4)2 + H2O
Ca: +2 → +3
P: +5 → +3
O: -2 → -2
H: +1 → +1
O: -2 → -2
In this reaction, the oxidation state of calcium changes from +2 to +3, and the oxidation state of phosphorus changes from +5 to +3. Therefore, reaction III is a redox reaction.
IV) Cl2 + 2 KBr → Br2 + 2 KCl
Cl: 0 → 0
K: 0 → 0
Br: -1 → 0
Br: -1 → 0
In this reaction, the oxidation state of bromine changes from -1 to 0, and chlorine remains at 0 oxidation state. Therefore, reaction IV is a redox reaction.
Based on the analysis above, the redox reactions are:
I) Ca + 2H2O → Ca(OH)2 + H2
III) Ca(OH)2 + H3PO4 → Ca3(PO4)2 + H2O
IV) Cl2 + 2 KBr → Br2 + 2 KCl
Moving on to the second question, to determine the oxidizing agent in the given reaction, you need to look for the element or compound that gets reduced, as the oxidizing agent causes the reduction to occur.
Zn(s) + NO3^-(aq) → Zn(OH)4^2-(aq) + NH3(aq)
In this reaction, the nitrogen in NO3^- changes its oxidation state from +5 to 0 in NH3, indicating it is being reduced. Therefore, the oxidizing agent is NO3^-(aq).
I hope this explanation helps you understand the concepts of redox reactions and identifying oxidizing agents. Feel free to ask for further clarification if needed!