Oxidation Number for H2O . Oxidation state of h2o . Oxidation state of water. H2O oxidation numbers
Oxidation Number for H2O . Oxidation state of h2o . Oxidation state of water. H2O oxidation numbers

Download presentation

Presentation is loading. Please wait.

Published byKatherine Miles Modified over 5 years ago

1
CHAPTER 20 Oxidation-Reduction Reactions LEO SAYS GER

2
20.2

3
Section 20.2 – Oxidation Numbers
Objectives Determine the oxidation number of an atom of any element in a pure substance. Define oxidation and reduction in terms of a change in oxidation number, and identify atoms being oxidized or reduced in redox reactions.

4
Remember Oxidation Numbers

5
Oxidation Numbers What is an oxidation number:
An oxidation number is a positive or negative number assigned to an atom to keep track of electron transfers & electron sharing

6
Rules for Assigning Oxidation Numbers
The oxidation number of any uncombined element is zero. 2.The oxidation number of a monatomic ion equals its charge. +1 -1 2Na + Cl2 2NaCl

7
Rules for Assigning Oxidation Numbers
3. The oxidation number of O in compounds is -2, Exceptions: in hydrogen peroxide: H2O2 where it is -1, in compounds with the more electronegative fluorine F2O +1 -2 +1 -1 -1 +2 F2O H2O H2O2 2(+1) + 1(-2) = 0 2(-1) + 1(+2) = 0 2(+2) + 2(-2) = 0

8
Rules for Assigning Oxidation Numbers
4. The oxidation number of H in compounds is +1, Exception: in metal hydrides it is -1. +1 -2 +1 -1 H2O NaH Warning! Don’t get too bogged down in these exceptions. In most of the cases you will come across, they don’t apply!

9
Rules for Assigning Oxidation Numbers
5. For any neutral compound, the sum of the oxidation numbers of the atoms in the compound must equal 0. +1 -2 +2 -2 +1 Ca(OH)2 H2O (+2) + 2(-2) + 2(+1) = 0 Ca O H 2(+1) + (-2) = 0 H O

10
Rules for Assigning Oxidation Numbers
6. For a polyatomic ion, the sum of the oxidation numbers in the formula is equal to the ionic charge of the ion ? -2 ? -2 SO42- NO3- X + 3(-2) = -1 N O X + 4(-2) = -2 S O THUS X = +6 THUS X = +5

11
20.1

12
Section 20.1 – The Meaning of Oxidation and Reduction
Objectives: Define oxidation and reduction in terms of loss or gain of oxygen and the loss or gain of electrons State the characteristics of a redox reaction Describe what happens to iron when it corrodes

13
Oxidation and Reduction
We already learned why salt is spread on roads during winter months Lowers the freezing point of water preventing the buildup of slippery ice But…what about the affect of the salt that clings to the metallic parts of cars? RUST The corrosion of metal is one example of an oxidation-reduction reaction

14
Other Examples of Oxidation-Reduction Reactions
When methane burns in air, it oxidizes and forms oxides of carbon and hydrogen One oxide of carbon is carbon dioxide Elemental iron slowly oxidizes into compounds such as iron (III) oxide, or rust Hydrogen peroxide also releases oxygen when it decomposes – when you pour it on an open cut Bleaching stains in fabrics is an oxidation-reduction reaction

15
Oxidation and Reduction
What is an oxidation-reduction reaction? Classic Definition: A reaction where oxygen is transferred from one substance to another. In an oxidation-reduction reaction: The substance gaining oxygen is oxidized The substance losing oxygen is reduced. Oxidation & reduction always occur simultaneously An abbreviation for oxidation-reduction reactions is redox reactions

16
Redox Reactions Modern concepts of redox reactions include reactions that do not even involve oxygen Redox reactions are currently understood to involve the shift of electrons between reactants Oxidation Complete or partial loss of electrons, or gain of oxygen Reduction Complete or partial gain of electrons, or loss of oxygen

17
Why call it Oxidation? With the exception of fluorine, oxygen is the most electronegative element When oxygen bonds with atoms of different elements, the electrons from the other atom shift toward oxygen When a substance gains electrons, it acting like oxygen

18
Oxidation and Reduction
Until now our view of a reaction: Mg + S MgS Reaction viewed as Redox: Mg S → Mg S2- The magnesium atom changes to a more stable magnesium ion by losing 2 electrons, and is thus oxidized The sulfur atom is changed to a more stable sulfide ion by gaining 2 electrons, and is thus reduced.

19
Oxidation and Reduction
Mg + S MgS The over all process is represented as the two component processes below: Mg Mg e- S + 2e- S2- Oxidation – loss of electrons Reduction – gain of electrons

20
Example 2: Oxidation and Reduction
Each sodium atom loses one electron – oxidation Each chlorine atom gains one electron – reduction

21
LEO says GER Gain Electrons = Reduction Lose Electrons = Oxidation
Sodium is oxidized Gain Electrons = Reduction Chlorine is reduced

22
Remember Oxidation – reduction reactions involve a transfer of electrons. LEO – GER Lose Electrons Oxidation Gain Electrons Reduction

23
Practice Problem Determine what is oxidized and what is reduced in each reaction. 2Na + S Na2S 4Al + 3O2 2Al2O3 Na is oxidized (Na1+ S2-) S is reduced Al is oxidized (Al3+ O2-) O is reduced

24
Practice Problem Identify these processes as either oxidation or reduction. S2- S + 2e- Zn2+ + 2e- Zn Oxidation – is loss Reduction – is gain

25
Redox in Covalent Compounds
It is easy to see the loss and gain of electrons in ionic compounds, but what about covalent compounds? (complete electron transfer does not occur) Remember when we did bonding, we looked at Electronegativity. Electronegativity = strong pull on the electron. If the electrons are shared between two atoms, the atom with the higher Electronegativity has the electron most of the time. So the oxidation number will show the electron belonging to the atom with the higher Electronegativity (In reality the e- is not completely transferred, but it helps us keep track of where the e- is)

26
Redox in Covalent Compounds
It is easy to see the loss and gain of electrons in ionic compounds, but what about covalent compounds? (complete electron transfer does not occur) Consider: 2H2 + O H2O Oxygen is highly electronegative In H there is a shift of bonding electrons away from H Hydrogen is oxidized In O there is a shift of electrons toward O Oxygen is reduced loss gain

27
Oxidation and Reduction
In some reactions involving covalent reactants and products, the partial electron shifts are less obvious A few general guidelines: Gain of oxygen – oxidation Loss of oxygen – reduction Loss of hydrogen by a covalent compound – oxidation Gain of hydrogen by a covalent compound – reduction Increase in oxidation number – oxidation Decrease in oxidation number – reduction

28
Terminology for Redox Oxidation Reduction Oxidizing Agent
Loss of electrons Increase in oxidation number Increase in oxygen Oxidizing Agent Electron acceptor Species that is reduced Reduction – Gain of electrons – Decrease in oxidation number – Decrease in oxygen Reducing Agent – Electron donor – Species that is oxidized

29
Trends in Oxidation and Reduction
Active metals: Lose electrons easily Are easily oxidized Are strong reducing agents Active nonmetals: Gain electrons easily Are easily reduced Are strong oxidizing agents

30
Oxidation – Number Changes in Chemical Reactions
An increase in oxidation number = oxidation A decrease in oxidation number = reduction Determining all of the oxidations in a chemical equation allows us to determine the oxidizing agent and reducing agent (which species is oxidized and which species is reduced) Determine the oxidation numbers in the following chemical equation. 2AgNO3 + Cu Cu(NO3)2 + 2Ag +1 +5 -2 +2 +5 -2 Copper is oxidized Silver is reduced

31
20.3

32
Section 20.3 – Balancing Redox Equations
Objectives: Describe how oxidation numbers are used to identify redox reactions. Balance a redox equation using the oxidation-number-change method. Balance a redox equation by breaking the equation into oxidation and reduction half-reactions, and then balance using the half-reaction method.

33
Identifying Redox Equations
In general, all chemical reactions can be assigned to one of two classes: oxidation-reduction, in which electrons are transferred: Single-replacement, combination, decomposition, and combustion All other reactions… No electron transfer Double-replacement and acid-base reactions

34
Identifying Redox Equations
In an electrical storm, oxygen and nitrogen react to form nitrogen monoxide: N2(g) + O2(g) → 2NO(g) Is this a redox reaction? If the oxidation number of an element in a reacting species changes, then that element has undergone either oxidation or reduction; therefore, the reaction as a whole must be a redox reaction. YES!

35
Balancing Redox Equations
It is essential to write a correctly balanced equation that represents what happens in a chemical reaction Many redox reactions are too complex to be balanced by trial and error Fortunately, there are two systematic methods based on the fact that the total electrons gained in reduction equals the total lost in oxidation The methods: Using oxidation-number changes Use half-reactions

36
Using Oxidation-Number Changes
Type of “chemical bookkeeping” to keep track of electron transfers Balance a redox equation by comparing the increases and decreases in oxidation number Start with the skeleton equation Fe2O3 + CO Fe + CO2 (unbalanced) Assign oxidation numbers to all the atoms in the equation Fe2O3 + CO Fe + CO2 +3 -2 +2 -2 +4 -2

37
Using Oxidation-Number Changes
2. Identify which atoms are oxidized and which are reduced Iron decreases in oxidation number from +3 to 0, a change of -3. Iron is reduced Carbon increases in oxidation number from +2 to +4, a change of +2. Carbon is oxidized 3. Use brackets to connect the atoms that undergo oxidation and another to connect those that undergo reduction. Write the oxidation number change on the line. Fe2O3 + CO Fe + CO2 +2 oxidation -3 reduction As the equation is written, the number of electrons transferred in oxidation, does not equal the number of electrons transferred in reduction

38
Using Oxidation-Number Changes
4. Make the total increase in oxidation number equal to the total decrease in oxidation number by using appropriate coefficients Fe2O3 + CO Fe + CO2 Using the numbers you multiplied the oxidation number changes by, add appropriate coefficients to the equation Fe2O3 + 3CO 2Fe + 3CO2 5. Finally, make sure that the equation is balanced for both atoms and charges. If necessary, finish balancing the equation by inspection 3 X (+2) = + 6 2 X (-3) = – 6

39
Using Oxidation-Number Changes
Balance this redox reaction by using the oxidation-number-change method K2Cr2O7 + H2O + S KOH + Cr2O3 + SO2 3 X + 4 oxidation = + 12 +6 +3 +4 4 X – 3 reduction = – 12 2 K2Cr2O7 + H2O + S KOH + Cr2O3 + SO2 2 3 4 2 3

40
Rules shown on page 651 – bottom
Using Half-Reactions A half-reaction is an equation showing just the oxidation or just the reduction that takes place They are then balanced separately, and finally combined 1. Write unbalanced equation in ionic form 2. Write separate half-reaction equations for oxidation and reduction 3. Balance the atoms in the half-reaction 4. Add enough electrons to one side of each half-reaction to balance the charges 5. Multiply each half-reaction by a number to make the electrons equal in both 6. Add the balanced half-reactions to show an overall equation 7. Add the spectator ions and balance the equation Rules shown on page 651 – bottom

41
You can’t have one… without the other!
Reduction (gaining electrons) can’t happen without an oxidation to provide the electrons. You can’t have 2 oxidations or 2 reductions in the same equation. Reduction has to occur at the cost of oxidation GER!

42
Using Half-Reactions Consider the reduction of Ag+ ions with copper metal. Cu + Ag give–> Cu Ag

43
Using Half-Reactions Step 1: Divide the reaction into half- reactions, one for oxidation and the other for reduction. Ox Cu —> Cu2+ Red Ag+ —> Ag Step 2: Balance each element for mass. Already done in this case. Step 3: Balance each half-reaction for charge by adding electrons. Ox Cu —> Cu e- Red Ag+ + e- —> Ag

44
The equation is now balanced for both mass & charge
Using Half-Reactions Step 4: Multiply each half-reaction by a factor so that the reducing agent supplies as many electrons as the oxidizing agent requires. Reducing agent Cu —> Cu e- Oxidizing agent 2 Ag e- —> 2 Ag Step 5: Add half-reactions to give the overall equation. Cu Ag > Cu Ag The equation is now balanced for both mass & charge

45
Zn0 + H+1N+5O3-2 Zn+2(N+5O-23)3 + N+4O-22 + H2+1O-2
Practice Problem Balance the following equation using the half reaction method Zn + HNO3 Zn(NO3)3 + NO2 + H2O Step 1: Identify the oxidized and reduced species by assigning oxidation numbers Zn0 + H+1N+5O3-2 Zn+2(N+5O-23)3 + N+4O-22 + H2+1O-2 Step 2: Write the half reactions Oxidation: Zn0 Zn+2 + 2e- Reduction: N+5 + 1e- N+4 2( N+5 + 1e- N+4) 2N+5 + 2e- 2N+4 Step 3: Balance species oxidized and reduced in equation Zn + 2HNO3 Zn(NO3)3 + 2NO2 + H2O Step 4: Balance remaining atoms Nitrogen is not balanced because some are not reduced, balance H and O Zn + 4HNO3 Zn(NO3)3 + 2NO2 + 2H2O

46
Choosing a Balancing Method
If the oxidized and reduced species appear only once one ach side of the equation and no acids and bases are involved in the reaction – oxidation number change usually works Sometimes the same element is both oxidized and reduced – reactions when this occurs are best balanced by the half-reaction method Equations for reactions that take place in acidic or alkaline solution are best balanced by the half-reaction method

Similar presentations

© 2023 SlidePlayer.com Inc.
All rights reserved.

You are watching: CHAPTER 20 Oxidation-Reduction Reactions LEO SAYS GER.. Info created by PeakUp selection and synthesis along with other related topics.