The University of Adelide Experiment 8F Reduction Oxidation Chemistry HW

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Question Description

Question 1

a) Which metals from Table 8.1 are considered VERY ACTIVE (react with water and aqueous hydrochloric acid), MODERATELY ACTIVE (react with aqueous hydrochloric acid but not water) or have LOW ACTIVITY (do not react with either)?

VERY ACTIVE:

MODERATELY ACTIVE:

LOW ACTIVITY:

b) For the metals that reacted with water and/or aqueous hydrochloric acid, write out an appropriate redox reaction equation and assign oxidation numbers to ALL elements (HINT: Use Equations 8.10 and 8.11 in your laboratory manual in order to guide you).

c) From your equations above, identify the species that are the oxidising agents and reducing agents in each redox reaction. Provide an explanation for your choices. (HINT: Think about the oxidation numbers of each element and how they change).


Question 2 a) Why do we not test each metal against a solution containing that metal? What would you observe if you did?

b) Using your results complete the following redox equations. If no reaction occurred, then write ‘no reaction’ like in Equation 8.8 of your laboratory manual. Mg(s) + Pb(NO3)2(aq) →

Pb(s) + MgSO4(aq) →
Equation 8.13

Equation 8.14 c) What is the oxidation number of magnesium on the left and right-hand sides of Equation 8.13? Has magnesium been oxidized or reduced? Explain your answer.

d) Explain the outcome of the reaction in Equation 8.14.

e) Complete the Metal Activity Series below by considering the numbers you wrote in the final column of Table 8.2 (ignore Metal ‘Pg’ for the moment).

> > Metal Pg > >
Most Easily Oxidised
Least Easily Oxidised


Question 3

a) Complete the Halide Activity Series below by considering the numbers you wrote in the final column of Table 8.3. (HINT: Refer to page 7 of your Manual for more information).

> >
Most Easily Reduced
Least Easily Reduced

b) From the Activity Series above, complete the following sentence:
″The ability to reduce elements in Group 17 _____________ as you go down the group.″

c) Based upon yours answers in Question 3a) and the periodic trend described in Question 3b), write out the full Halogen Activity Series below. Be sure to include Fluorine (F) and Astatine (At) in your answer.

> > > >
Most Easily Reduced
Least Easily Reduced




Unformatted Attachment Preview

Laboratory Manual ACTIVITY SERIES EXPERIMENT 8F Activity Series (This experiment is done in pairs.) Useful background reading (this is not compulsory but may be helpful): “Chemistry Core Concepts”. 1st Edition. John Wiley & Sons Australia Ltd. Section 12.1, 12.3 (pages 522-523), 12.4 (page 531); Worked examples 12.1, 12.2, 12.3, 12.4, 12.5; Practice exercises 12.1, 12.2, 12.3, 12.4, 12.9 “Chemistry Core Concepts”. 2nd Edition. John Wiley & Sons Australia Ltd. Section 12.1, 12.3 (pages 705-706), 12.4 (page 716); Worked examples 12.1, 12.2, 12.3, 12.4, 12.5; Practice exercises 12.1, 12.2, 12.3, 12.4, 12.9 Where does this practical fit in? This experiment will introduce you to Redox (Reduction-Oxidation) chemistry. In molecules, electrons are the “glue” that holds atoms together. Therefore, Chemists are very interested in reactions which involve the transfer of electrons between reacting species, which is the fundamental basis of Redox chemistry. In this experiment, you will perform a series of ‘Single Replacement’ reactions; a type of redox reaction for metals and halogens. From your results, you will to develop an Activity Series for metals and halogens, which will involve ordering both respective species in decreasing order of reactivity. Upon construction of this activity series, you will then use the information it provides to predict the result of other redox reactions. Learning objectives (remember these are different to the scientific objectives i.e. they should NOT appear in the AIM or CONCLUSION sections of your practical book): On completion of this practical, you should have:  An understanding of what a Redox reaction involves.  An appreciation for how assigning oxidation numbers to elements in a compound allow oxidation and reduction processes to be identified.  Understand how we are able to use an Activity series to predict Redox reactivity. EXPERIMENT 8F-1 Laboratory Manual ACTIVITY SERIES Introduction Redox reactions are an important class of chemical reaction whereby electrons are transferred from one substance to another. An example of a redox reaction is the reaction between sodium (Na) and chlorine (Cl2) to give sodium chloride (NaCl). This reaction involves a transfer of electrons from sodium to chlorine (Equation 8.1). 2Na + Cl2 → 2NaCl Equation 8.1 A redox reaction involves two simultaneous processes - an oxidation reaction and a reduction reaction. We define oxidation as the loss of electrons and reduction as a gain of electrons. In Equation 8.1, sodium is the species that is undergoing oxidation, hence we say that it is being ‘oxidised’. In the same equation, chlorine is the species that is undergoing reduction, hence we say that it is being ‘reduced’. To better highlight these two different processes that are occurring simultaneously, it is convenient to write out the half-equations that describe each individual process. These are just equations that show the individual redox steps that are occurring simultaneously. Equation 8.2 is the half-equation that shows that sodium is undergoing oxidation. Equation 8.3 is the half-equation that shows that chlorine is reduced. Na → Na+ + e- (oxidation – loss of an electron) ½ Cl2 + e- → Cl- (reduction – gain of an electron) Equation 8.2 Equation 8.3 An alternative (yet equally important) way to interpret what is happening in a redox reaction is to consider which species in the reaction are oxidising agents and reducing agents. The oxidising agent is the species that has the ability to remove electrons from another substance, thus being reduced as it does so. The reducing agent is the species that has the ability to give electrons to another substance, thus being oxidised as it does so. In Equation 8.1, sodium is acting as a reducing agent because it is able to donate electrons to chlorine, thus undergoing oxidation in the process. Additionally (and from an alternative perspective!), chlorine is acting as the oxidising agent because it is removing electrons from the sodium, thus undergoing reduction in the process. The language of chemistry… 1. A simple way to remember which process is oxidation and which is reduction is to use the simple acronym ‘OIL RIG’. Oxidation Reduction Is Is Loss Gain 2. An oxidising agent (or oxidant) oxidises a substance and is in the process reduced. A reducing agent (or reductant) reduces a substance and is in the process oxidised. A reducing agent (or reductant) reduces a substance and is in the process oxidized. EXPERIMENT 8F-2 Laboratory Manual ACTIVITY SERIES Assigning Oxidation States One method that is used to identify which substance has been oxidised and which has been reduced in a redox process is to assign oxidation numbers. Using the example shown in Equation 8.1, the following oxidation numbers or oxidation states can be given to each element: Oxidation numbers → 0 0 +1 -1 2Na + Cl2 → 2NaCl Equation 8.4 The oxidation state of sodium has gone from 0 to +1, which indicates a loss of one electron. This confirms that it has been oxidised. The oxidation state of chlorine has gone from 0 to 1, indicating that is has gained one electron and has therefore been reduced. There are some simple rules that can be used to work out the oxidation number for each element in a compound. These rules are given below. Rules for Assigning Oxidation Numbers: 1. The oxidation number of elements in their naturally occurring form is 0. 0 0 Na Cl2 2. The oxidation number of a monoatomic ion is equal to its charge. +2 -1 2+ Ca F3. The sum of oxidation states of all atoms in: a) a neutral molecule is 0. b) an a charged molecule is equal to the charge on the molecule. H2O Ox. No. of H = +1 Ox. No. of O = -2 SO42Ox. No. of S = +6 Ox. No. of O = -2 (2 x +1) + -2 = 0 +6 + (4 x -2) = -2 4. When part of compounds: a) Group I metals (eg. Li, Na, K) have an oxidation state of +1. b) Group II metals (eg. Be, Mg, Ca) have an oxidation state of +2. EXPERIMENT 8F-3 Laboratory Manual ACTIVITY SERIES 5. In their compounds, non-metals are assigned oxidation numbers according to the table below: Nonmetal Fluorine, F Hydrogen, H Oxygen, O Group 17 Group 16 Group 15 Oxidation No. -1 +1 -2 -1 -2 -3 Please note, that this table is hierarchical. That is, if a particular compound contains two or more of these elements, then the element HIGHER in position on the table will dominate in oxidation number assignment. For example, let’s look at H2O2: H2O2 Ox. No. of H = +1 Overall Charge of Compound = 0 Therefore, Ox. No. of O has to be -1 (Due to Rules 3a and 5) Predicting Redox Reactivity It is possible to predict whether or not a redox reaction will take place when two substances react together. To do this, we need to know the redox reactivity of each element. That is, we need to know the rank of the elements involved in order of most easily oxidised to least easily oxidised. For example, it is known that zinc (Zn) has a greater tendency to lose electrons (and therefore be oxidised) than copper (Cu). This can be shown experimentally if a strip of metallic zinc is dipped into a solution of copper sulfate (CuSO4). A reddish-brown layer of copper will be deposited on the zinc strip, which indicates that the copper ions (Cu2+) from the copper sulfate are undergoing reduction (and therefore, are being reduced). Cu2+(aq) + 2e- → Cu(s) (reduction) Equation 8.5 When the solution is analysed more thoroughly, it is also found that there are now zinc ions (Zn2+) in solution. This arises from the zinc metal undergoing oxidation (and therefore, is being oxidised). Zn(s) → Zn2+(aq) + 2e- (oxidation) Equation 8.6 EXPERIMENT 8F-4 Laboratory Manual ACTIVITY SERIES By combining the oxidation and reduction half-equations (Equations 8.5 and 8.6, respectively), an equation can be written to represent the overall redox reaction that is taking place (Equation 8.7 shown on the next page): Cu2+(aq) + 2e- → Cu(s) (reduction) Zn (s) → Zn2+(aq) + 2e- (oxidation) Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) (overall redox reaction equation) Equation 8.7 Alternatively, we could put a strip of copper in a solution of zinc sulfate (ZnSO4). However, because zinc has a greater tendency to lose electrons than copper, it will not be reduced (i.e. gain electrons). In this experiment there will be no change to the copper metal indicating there has been no reaction (Equation 8.8). Cu(s) + Zn2+(aq) → NO REACTION Equation 8.8 These two experiments are examples of Single Replacement Reactions. If there is replacement of one metal in solution by another (e. g. zinc replaces copper in Equation 8.7), it shows that the replacing metal is more active. By performing a series of Single Replacement Reactions we can establish an Activity Series. This will allow us to rank different metals according to their tendency to lose or gain electrons relative to other metals. As we have seen, zinc has a greater tendency to lose electrons compared to copper and it would therefore be higher on the activity series. The general equation for a Single Replacement Reaction is shown below in Equation 8.9. A(s) + BC(aq) → B(s) + AC(aq) Equation 8.9 If A is a more active species than B then displacement will occur and A will be oxidized just as zinc was in Equation 8.7. If B is the more active species, then no reaction will occur (as in Equation 8.8). Establishing an Activity Series of Metals Very active metals can be identified by their ability to reduce hydrogen in water. In a replacement reaction, a very active metal such as rubidium (Rb) would replace hydrogen in water as shown in Equation 8.10. Notice that the oxidation number of the element hydrogen (H) has gone from the +1 to 0 which shows a gain of an electron. The oxidation number for rubidium has gone from 0 to +1 which shows a loss of an electron. Therefore, we can say that rubidium has a greater tendency to lose electrons than hydrogen. 0 ↓ +1 -2 ↓ ↓ +1 -2 +1 ↓ ↓↓ 0 ↓ EXPERIMENT 8F-5 Laboratory Manual ACTIVITY SERIES 2Rb(s) + 2H2O(l) → 2RbOH(aq) + H2(g) Equation 8.10 The Group 1 Alkali metals, like rubidium, are very active and are often stored under inert oils to prevent them from reacting with oxygen and moisture in the atmosphere. For safety reasons, we will not be directly testing these metals in the experiment. Moderately active metals cannot replace hydrogen from water, but are able to replace hydrogen from an aqueous acid such as hydrochloric acid (HCl). The replacement reaction between manganese (Mn) and hydrochloric acid is shown below in Equation 8.11. Once again, hydrogen is reduced and the manganese is oxidised as indicated by the oxidation numbers that have been assigned to each element. 0 +1 -1 +3 -1 0 ↓ ↓↓ ↓ ↓ ↓ 2Mn(s) + 6HCl(aq) → 2MnCl3(aq) + 3H2(g) Equation 8.11 In the reaction described in Equation 8.11, hydrogen gas (H2(g)) is produced. In reality, if we were to perform this reaction in the lab, we would see bubbles of hydrogen gas in the solution. This is a visual indication that a redox reaction has taken place and is an observation you should take note of when performing your own experiments. To compare the reactivity of low activity metals relative to each other, each metal must be placed into a metal salt solution of the particular metal that the reactivity is being compared to. For example, we have already shown that zinc is more active than copper in Equation 8.7 and Equation 8.8 above. Since zinc replaces copper (Equation 8.7) but copper does not replace zinc (Equation 8.8), zinc is deemed to be the more active metal. In this experiment, you will establish an activity series for the following metals: Copper (Cu) Zinc (Zn) Magnesium (Mg) Lead (Pb) In Part One, you will react each metal with water and aqueous hydrochloric acid in order to observe how reactive each metal is relative to one another. In Part Two, you will react each metal with metal salt solutions of the other metals in order to complete the activity series. EXPERIMENT 8F-6 Laboratory Manual ACTIVITY SERIES Halogen Reactivity The elements in Group 17, known as the halogens, react with most other elements on the Periodic Table. Their excellent reactivity can be rationalized by considering that each halogen only needs one electron to fill its valence shell. In other words, the halogens have a tendency to gain electrons and become reduced. We have already seen an example of chlorine being reduced by sodium in Equation 1.1. Group 17 9 F 17 Cl 35 Br 53 I 85 At Just as we can investigate the reactivity of metals, we can also compare the redox reactivity of the halogens. The halogens tend to exist as diatomic molecules in their naturally occurring form (e. g. F2, Cl2, Br2, I2). We can determine the activity series for the halogens by investigating their ability to undergo single replacement reactions when the naturally occurring form of a particular halogen (X2) is reacted with a halide-containing compound (denoted Y- to indicate that ‘X’ and ‘Y’ are different halogen elements). As in the case of the single replacement reactions performed with various metals in Part Two, a more active halogen will displace a less active halogen in solution as shown in the Equation 8.12 below: 0 -1 -1 0 X2(aq) + 2Y (aq) → 2X (aq) + Y2(aq) Equation 8.12 Notice in Equation 8.12 that the more active halogen (X) changes oxidation number from 0 to -1, indicating that it has been reduced and has gained one electron. The less active halogen (Y) has been oxidized as shown by the change in oxidation number of -1 to 0. In Part Three of the experiment, you will determine the relative reactivity of chlorine (Cl2), bromine (Br2) and iodine (I2) by reacting each with solution of potassium chloride (KCl), potassium bromide (KBr) and potassium iodide (KI) (obviously without reacting two solutions containing the same halide). You will then establish a Halogen Activity Series and predict the reactivity of fluorine (F2) and astatine (At2) based upon your observations. EXPERIMENT 8F-7 Laboratory Manual ACTIVITY SERIES Experimental Reminder - students should work in pairs. It is important that you work carefully in order to gain worthwhile experience from this class. PART ONE SINGLE REPLACEMENT REACTIONS WITH WATER AND AQUEOUS HYDROCHLORIC ACID Experiment 8F equipment Test tubes (12) and rack Additional equipment and location On the bench: All metals and corresponding metal salts, hydrochloric acid 2M, potassium iodide, potassium chloride, potassium bromide. In fume cupboard: Halogenated waters with residue disposal On sink: residue disposal for metals and solutions containing metals Please ensure this is correct and your area is clean and tidy before leaving the laboratory! Hazardous substances Hydrochloric Acid HCl corrosive and irritant Lead Pb toxic and teratogen Magnesium Mg highly flammable EXPERIMENT 8F-8 Laboratory Manual ACTIVITY SERIES PROCEDURE 1 Label 4 test tubes ‘Copper’, ‘Zinc’, ‘Lead’ and ‘Magnesium’, respectively. Put approximately 3 mL of deionised water in each one. 2 Drop a small piece of copper in the test tube labelled ‘Copper’. Make note of any observations in Table 8.1 (i.e. any bubbles of gas evolved, increase in temperature, colour changes etc.). If you observe no obvious reaction in the first 30 seconds, wait 10 minutes and then check the test tube again. 3. Repeat step 2 for each of the three remaining metals. Be sure to place the correct metal into its labelled test tube. As with Copper, if you observe no obvious reaction in the first 30 seconds, wait 10 minutes and then check the test tubes again. 4 Repeat Steps 1-3 but instead of deionised water, use 3 mL of 2M hydrochloric acid. Make note of any observations in Table 8.1. 5 For any metals that did not react with deionised water, decant the water from the test tube, tip out the metal pieces onto paper towel and dry them and keep aside for use in Part Two. For the metals that did not react with 2M hydrochloric acid, top up the test tube with water, decant water from the test tube, re-fill the test tube with water again, decant the water, tip out the metal pieces onto paper towel, dry them and return to the servery (if there is any left!). Table 8.1 Metal Reaction with H2O Reaction with HCl(aq) Copper Zinc Lead Magnesium EXPERIMENT 8F-9 Laboratory Manual ACTIVITY SERIES Question 1 a) Which metals from Table 8.1 are considered VERY ACTIVE (react with water and aqueous hydrochloric acid), MODERATELY ACTIVE (react with aqueous hydrochloric acid but not water) or have LOW ACTIVITY (do not react with either)? VERY ACTIVE: MODERATELY ACTIVE: LOW ACTIVITY: b) For the metals that reacted with water and/or aqueous hydrochloric acid, write out an appropriate redox reaction equation and assign oxidation numbers to ALL elements (HINT: Use Equations 8.10 and 8.11 in your laboratory manual in order to guide you). c) From your equations above, identify the species that are the oxidising agents and reducing agents in each redox reaction. Provide an explanation for your choices. (HINT: Think about the oxidation numbers of each element and how they change). EXPERIMENT 8F-10 Laboratory Manual PART TWO ACTIVITY SERIES SINGLE REPLACEMENT REACTIONS OF LOW ACTIVITY METALS Hazardous substances Lead Pb toxic and teratogen Magnesium Mg highly flammable Copper Sulfate CuSO4 harmful and irritant Lead Nitrate Pb(NO3)2 harmful and very toxic Zinc Sulfate ZnSO4 toxic to aquatic organisms PROCEDURE 1 Label three test tubes with, ‘Zn with CuSO4’ , ‘Zn with Pb(NO3)2’ and ‘Zn with MgSO4’, respectively. Place a small piece of zinc in each labelled test tube. 2. In the test tube labelled ‘Zn with CuSO4’, add enough drops of copper sulfate (CuSO4) solution onto the zinc such that it is covered by the solution. Make note of any observations in Table 8.2 such as if there is any gas evolved, if there is a black or grayish layer on the metal (indicating that there is some other type of metal deposited on the zinc) or if the solution has changed colour. If you observe no obvious reaction in the first 30 seconds, wait 10 minutes and then check the test tube again. 3. Repeat Step 3 with lead nitrate (Pb(NO3)2) solution and magnesium sulfate (MgSO4) solution in their respectively labelled test tubes. Make note of any observations in Table 8.2. As before, if you observe no obvious reaction in the first 30 seconds, wait 10 minutes and then check the test tube again. 4. Repeat Steps 1-3 for each remaining metal – copper, lead and magnesium such that you are able to complete Table 8.2. Note that you will need to test each metal against each solution except the solution that contains that same metal. 5. For each row count up how many reactions actually took place. Write this number in the final column of Table 8.2 (You should have a number between 0 and 3 for each column). EXPERIMENT 8F-11 Laboratory Manual ACTIVITY SERIES Table 8.2 CuSO4 Metal ZnSO4 Pb(NO3)2 MgSO4 Total No. Reactions Copper Zinc Lead Magnesium Question 2 a) Why do we not test each metal against a solution containing that metal? What would you observe if you did? b) Using your results complete the following redox equations. If no reaction occurred, then write ‘no reaction’ like in Equation 8.8 of your laboratory manual. Mg(s) + Pb(NO3)2(aq) → Equation 8.13 Pb(s) + MgSO4(aq) → Equation 8.14 c) What is the oxidation number of magnesium on the left and right-hand sides of Equation 8.13? Has magnesium been oxidized or reduced? Explain your answer. d) Explain the outcome of the reaction in Equation 8.14. e) Complete the Metal Activity Series below by considering the numbers you wrote in the final column of Table 8.2 (ignore Metal ‘Pg’ for the moment). > Most Easily Oxidised > Metal Pg > > Least Easily Oxidised EX ...
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Tutor Answer

jorgearf92
School: Carnegie Mellon University

Hi there 😀 , answers in attachment, pdf and doc have the same content.

1.
a) Very active: Magnesium
Moderately active: Zinc, Lead
Low activity: Copper

b)
𝑀𝑔(𝑠) + 2𝐻2 𝑂(𝑙) → 𝑀𝑔(𝑂𝐻)2 (𝑎𝑞) + 𝐻2 (𝑔)
Reactants oxidation numbers: Mg: 0, H: +1, O: -2
Products oxidation numbers: Mg: +2, H(hydroxide): +1, O: -2, H(H2): 0
𝑀𝑔(𝑠) + 2𝐻𝐶𝑙(𝑎𝑞) → 𝑀𝑔𝐶𝑙2 (𝑎𝑞) + 𝐻2 (𝑔)
Reactants oxidation numbers: Mg: 0, H: +1, Cl: -1
Products oxidation numbers: Mg: +2, Cl: -1, H: 0
𝑍𝑛(𝑠) + 2𝐻𝐶𝑙(𝑎𝑞) → 𝑍𝑛𝐶𝑙2 (𝑎𝑞) + 𝐻2 (𝑔)
Reactants oxidation numbers: Zn: 0, H: +1, Cl: -1
Products oxidation numbers: Zn: +2, Cl: -1, H: 0
𝑃𝑏(𝑠) + 2𝐻𝐶𝑙(𝑎𝑞) → 𝑃𝑏𝐶𝑙2 (𝑎𝑞) + 𝐻2 (𝑔)
Reactants oxidation numbers: Pb: 0, H: +1, Cl: -1
Products oxidation numbers: Pb: +2, Cl: -1, H: 0

c)
𝑀𝑔(𝑠) + 2𝐻2 𝑂(𝑙) → 𝑀𝑔(𝑂𝐻)2 (𝑎𝑞) + 𝐻2 (𝑔)
The oxidizing agent is the water because it causes magnesium to lose electrons and the
reducing agent is the magnesium metal because it causes hydrogen to gain electrons.
𝑀𝑔(𝑠) + 2𝐻𝐶𝑙(𝑎𝑞) → 𝑀𝑔𝐶𝑙2 (𝑎𝑞) + 𝐻2 (𝑔)
The oxidizing agent is t...

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