RC2

MDQ2 How is the reactivity of various metals predicted?

TEXT REFERENCE: CHAPTER 11

2.1 conduct practical investigations to compare the reactivity of a variety of metals:

PRACTICAL:

From PS&A p131, carry out PA 3.4 Reactivity of metals

POWERPOINT:

Below (or can be uploaded as 11C RC2 2 Activity Series PPT from https://drive.google.com/open?id=1SABRdKchHyE_0mghbkfyaLerdGsqVrEj )

11C RC2 2 Activity Series PPT.pptx

2.1 conduct practical investigations to compare the reactivity of a variety of metals in:

a) water

Some metals react with water or steam, while others do not. The more active metals, such as sodium or potassium, form hydroxides with water. The less active metals, such as aluminium and zinc, form oxides with steam. Low activity metals, such as copper, do not react with either. (Smith, 2008)

When reactions occur with water, the products are hydrogen gas and the metal hydroxide.

General word equation: metal + water → metal hydroxide + hydrogen gas

Balanced formula equation: 2Na_(s) + 2H₂O_(l) → 2NaOH_(aq) + H_2(g)

With steam, the product is a metal oxide.

General word equation: metal + steam→ metal oxide + hydrogen gas

Balanced formula equation: Zn_(s) + H₂O_(g) → ZnO_(aq) + H_2(g)

VIDEOS:

RC12 Metal and water https://www.youtube.com/watch?v=oQN9mJ5ewBY&list=PLeFSFSJ9WqSAJfb02g1rwALZbMuRTFaf9&index=16 4.12

TASK:

1. Write balanced equations to represent the following reactions:

a) Potassium metal with cold water

b) Sodium metal with cold water

c) Aluminium metal with steam

d) Zinc metal with steam

2. Use the Periodic Table to suggest why potassium and sodium might react with cold water whereas zinc and aluminium do not.

3. Apart from the temperature of the water, what is a difference between the reaction of sodium with water and zinc with water?

15 Metal and Water.doc

2.1 conduct practical investigations to compare the reactivity of a variety of metals in:

b) dilute acid

Most metals react with dilute hydrochloric acid and sulfuric acid to form hydrogen gas. The speed at which the reactions occur varies depending on the reactivity of the metal. Active metals like sodium and calcium react vigorously with dilute acids at room temperature. Low activity metals, such as copper, may need to be heated in order to react with dilute acids. (Smith, 2008)

Acids are substances which in solution produce hydrogen ions H⁺ in solution. It is usually the hydrogen ion from HCl or H₂S0₄ (sulfuric acid) which reacts with the metal.

Li, Na, K (group 1 of Periodic Table) Mg, Ca, Bar (group 2 Periodic Table) Al, Sn, Pb, Fe, Zn all react with dilute acids to form hydrogen gas.

When metals react with acids, electrons are transferred from the metal to the hydrogen ions in solution. This leaves the metal with a positive charge.

General equation: metal + acid → salt + hydrogen gas

Balanced formula equation: Mg_(s) + H₂SO_4(aq) → MgSO_4(aq) + H_2(g)

When less reactive metals react with concentrated acids like sulfuric or nitric acid, they can reduce the anion as well as or instead of the hydrogen cation. This produces a different set of products. Water may form instead of hydrogen gas, as well as either sulfur dioxide or nitrogen dioxide (or other NOx) depending on the metal used and the concentration of the acid.

VIDEOS:

RC13 Metal and acid https://www.youtube.com/watch?v=Woq3juJBVVo&list=PLeFSFSJ9WqSAJfb02g1rwALZbMuRTFaf9&index=15 4.59

TASK:

1. Write balanced equations to represent the following reactions:

a) Calcium metal with dilute hydrochloric acid

b) Iron metal with dilute sulfuric acid

c) Zinc metal with dilute nitric acid

d) Magnesium metal and dilute phosphoric acid

e) Tin metal and dilute hydrochloric acid

f) Copper metal with concentrated nitric acid

13 Metal and Acid.doc

2.1 conduct practical investigations to compare the reactivity of a variety of metals in:

c) oxygen

All oxides formed from the reaction between a metal and oxygen are ionic compounds.

Li, Na, K, Ca, Ba react rapidly at room temperature.

Mg, Al, Fe, Zn react slowly at room temperature but burn vigorously if heated in air or pure oxygen.

Sn, Pb, Cu react slowly and only if heated.

Metals which burn in oxygen form crystalline white solids and have none of the physical properties of the original metal (lustre, strength, malleability, and conductivity). When metals slowly react with oxygen at room temp they lose their lustrous appearance. Some such as Al and Zn become coated with a dull layer of tightly adhering oxide which prevents further reactions. The metals have been oxidised.

General word equation: metal + oxygen → metal oxide

Balanced formula equation 4Li_(s) + O_2(g) → 2Li₂O_(s)

Corrosion

Corrosion of a metal involves oxidation of metal atoms to metal ions so that the metal structure is broken down. Corrosion is a chemical process involving the reaction of a metal with other chemicals in its environment, often oxygen and/or water. Corrosion reactions are oxidation-reduction reactions that involve electron transfer from metals to other substances. (Alliband, 2003)

Rusting of iron on the surface of iron or steel is the most common type of corrosion. The iron atoms become part of a layer of hydrated iron(III) oxide. The formula for rust is Fe₂O₃.xH₂O where x can be from 0.5 to 2 and colour from orange-brown to black colour. The layer of rust is porous and easily flakes off, leaving a new surface of iron atoms to be rusted. The outer rust layer is porous allowing oxygen and water to penetrate and convert the iron layer below the surface to rust. (Alliband, 2003)

Aluminium, although a more reactive metal than iron, reacts less readily than the iron. This is because aluminium reacts with oxygen to form an inactive and impervious coating on its surface. This film is inert, tenacious and reforms immediately if removed by abrasive action. This makes aluminium metal, despite its high activity, a useful construction metal that doesn’t need painting or protective coatings. (Alliband, 2003)

Metals that corrode very slowly in dry air can undergo rapid corrosion in sea water where ions are plentiful and move freely. These reactions require the movement of ions in an electrolyte between an anode and a cathode and so salty sea water facilitates corrosion. (Alliband, 2003)

VIDEOS:

RC14 Corrosion https://www.youtube.com/watch?v=m42fAQjSbB0&list=PLeFSFSJ9WqSAJfb02g1rwALZbMuRTFaf9&index=14 4.55

TASK:

1. Some metals react with oxygen until all the metal has reacted, while other metals react to form an oxide coating that protects the metal from further reaction. (Smith, 2008)

a) Name one metal that undergoes complete reaction with oxygen (corrosion).

b) Name two metals that form an oxide coating that makes them resistant to further reaction.

2. Write balanced equations to represent the following reactions:

a) Barium metal with oxygen

b) Lithium metal with oxygen

c) Warmed zinc with oxygen

14 Metal and Oxygen.doc

VIDEOS:

Reactions involving metals https://www.youtube.com/watch?v=16sO_odtXSA&list=PLeFSFSJ9WqSCyQ8K25M7JD5yAPZVre-VH&index=4 7.48

2.1 conduct practical investigations to compare the reactivity of a variety of metals with:

d) other metal ions in solution

What happens if the results of our previous experiments do not allow us to separate the metals that react equally with oxygen, water and acid?

In this case, we can use displacement reactions.

When metals react with other metals in solution, there is a transfer of electrons.

When electrons are lost, we call it OXIDATION (LEO— Lost Electrons = Oxidation)

When electrons are gained, we call it REDUCTION (REG —Reduction = Electron Gain)

OIL RIG (Oxidation Is Loss, Reduction Is Gain)

All these reactions involve the metal losing electrons to become metal ions, or metal ions gaining electrons to form elemental metal. The activity series then lists the metals in order of decreasing ease of losing electrons.

Eg. Lithium Li → Li⁺ + e^-

Lithium readily loses an electron to a hydrogen ion as it is a very active metal. It would also lose an electron to a copper ion as it is more active than copper.

VIDEOS:

RC15 Displacement https://www.youtube.com/watch?v=eIvljmjbI04&list=PLeFSFSJ9WqSAJfb02g1rwALZbMuRTFaf9&index=13 4.55

TASK:

11C RC2 7 Displacement WS.doc

REVIEW:

11.1 Review p359
Key Qs p359

2.2a construct a metal activity series using the data obtained and compare this series with that obtained from standard secondary-sourced information

PRACTICAL:

From PS&A p133, carry out PA 3.5 Activity series of metals

If we were to carry out these reactions with each metal, we would be able to build up an order from the least reactive metals to the more reactive ones, see table below (Woollett, et al., 2018, p. 358). The placement of metals in an order of activity has been justified on the basis of their reactions with different substances. More reactive metals react more violently or vigorously than others.

VIDEOS:

RC16 Metal Activity Series https://www.youtube.com/watch?v=SKFGcHKZ7YI&list=PLeFSFSJ9WqSAJfb02g1rwALZbMuRTFaf9&index=12 6.04

TASK:

a) What is the colourless gas being formed?

b) Use the results in the table to arrange the metals into three groups—highly reactive, moderately reactive, not reactive.

c) If there is more than one metal in each of the categories moderately reactive, not reactive, suggest an investigation you could do to distinguish between the reactivities of these metals.

REVIEW:

Complete:

11.2 Review p365
Key Qs p365

2.3 analyse patterns in metal activity on the periodic table and explain why they correlate with:

a) ionisation energy

The reactivity of metals is related to the movement of electrons. In fact, most of chemistry is related to what’s happening to those electrons! The easier it is to take an electron away from an atom, the more easily it will lose electrons (be oxidised). If we could measure the amount of energy needed to remove 1 electron from an atom, we could use this to compare how reactive different metals are.

The amount of energy needed to remove 1 electron from a stable atom is called the FIRST IONISATION ENERGY. The first ionisation energy can be used determine the relative reactivity of metals.

The reactivity of metals increases as their ionisation energy _ _ _ _ _ _ _ _ _ .

In other words, the lower the ionisation energy the _ _ _ _ _ _ it is to remove an electron and hence the _ _ _ _ active the metal is.

The reactivity of metals increases as their ionisation energy decreases (with a few exceptions!). When metals react (with water, the H+ of acids, oxygen and ions of another metal), they are losing an electron (or more than one). Hence there is a strong link between metal activity and ionisation energy.

VIEW Videos:

RC17 Ionisation Energy Trends https://www.youtube.com/watch?v=DZxd5B6Zfc0&list=PLeFSFSJ9WqSAJfb02g1rwALZbMuRTFaf9&index=22 4.53

2.3 analyse patterns in metal activity on the periodic table and explain why they correlate with:

b) atomic radius

Atomic radius is a measure of the distance between the nucleus and the outer shell electrons. Atomic radius increases down a group as an extra electron shell is added. Atomic radius decreases across a period due to the increasing size of the attractive charge from the nucleus which has the effect of drawing the outer shell electrons closer to the nucleus.

For the active metals in Groups I and II, atomic radius increases down the group, as does reactivity.

As the valence shell electron(s) get further away from the nucleus, the force of electrostatic attraction exerted by the positively charged nucleus and the valence shell electrons decreases, hence electrons are more easily lost, hence the increase in activity.

VIEW Videos:

RC18 Atomic Radius Trends https://www.youtube.com/watch?v=NryIfNdh-Tg&list=PLeFSFSJ9WqSAJfb02g1rwALZbMuRTFaf9&index=21 5.24

TASK 2.3.1

Use atomic radius to explain why potassium is more reactive than both sodium and calcium. Write in 7- to 13-word sentences.

2.3 analyse patterns in metal activity on the periodic table and explain why they correlate with:

c) electronegativity

Electronegativity is a measure of the tendency of an atom of an element to attract electrons. Electronegativity increases across a period as the nuclear charge increases and the elements lose their metallic character. The electronegativity decreases down the group, as the atomic radius increases.

We can summarise the key properties linking metal activity to the trends observed in the Periodic Table as follows (Woollett, et al., 2018, p. 366).

There is an inverse relationship between electronegativity and metal activity as these are opposite effects. Metal activity is about the ability to lose electrons while electronegativity is about the ability of an element to attract electrons.

VIDEOS:

RC19 Electronegativity trends https://www.youtube.com/watch?v=-Wb9uoZ3ib8&list=PLeFSFSJ9WqSAJfb02g1rwALZbMuRTFaf9&index=20 6.17

TASK 2.3.2

REVIEW:

11.3 Review p368
Key Qs p368

CHAPTER REVIEW:

Chapter 11 Review p369-70

TEXT REFERENCE: CHAPTER 12

2.4 apply the definitions of oxidation and reduction in terms of electron transfer to a range of reduction and oxidation (redox) reactions

Oxidation Is Loss of electrons (OIL)

Reduction Is Gain of electrons (RIG)

For a species to be reduced, it must gain electrons from a species which has been oxidised. Hence these processes happen in pairs and we can discuss REDuction Oxidation (or Redox) reactions.

PRACTICAL:

Text p372 Observing the Corrosion of Iron
Blue Bottle experiment
Photolysis of silver chloride

11C RC2 4 Blue Bottle Prac.docx

11C RC2 4 Redox Photolysis Prac.docx

2.5 construct relevant half-equations and balanced overall equations to represent a range of redox reactions

POWERPOINT:

11C RC2 4 Redox Equations PPT (loaded below, or available from https://drive.google.com/open?id=1jU8ow-7_oE8XCS2NiNbKBl4xn1fmWrDK)

VIDEOS:

Metals and electron transfer https://www.youtube.com/watch?v=cMgrb2SYSsc&index=1&list=PLeFSFSJ9WqSCyQ8K25M7JD5yAPZVre-VH 5.28
RC20 Redox Reactions https://www.youtube.com/watch?v=MNbxM2g8enk&list=PLeFSFSJ9WqSAJfb02g1rwALZbMuRTFaf9&index=19 8.30

11C RC2 4 Redox Equations PPT.pptx

TASK 2.5.1

1. Select two pairs of metals from the list below and write an electron transfer reaction for each pair. Identify which metal would be displaced from a solution of its ions.

Sodium

Potassium

Copper

Iron

Magnesium

Zinc

2. Describe the characteristics of a displacement reaction.

3. Copper is a low activity metal. Why can it displace silver from a solution, but not zinc?

4. [2004 HSC Q19] An experiment was performed to investigate the relative activity of metals. The beaker initially contained 250.0 mL of 0.050 mol L⁻¹ copper sulfate solution and then a zinc rod was placed into it.

After several hours the dark blue colour of the solution had become lighter and a red-brown deposit had formed on the piece of zinc metal.

(a) Account for the changes observed. Provide a balanced oxidation–reduction equation in your answer. (3 marks)

(b) The red-brown deposit was removed from the piece of zinc metal and dried. It was found to weigh 0.325 g. Calculate the concentration of copper(II) ions remaining in the solution. (3 marks)

TASK 2.5.2

Complete PS&A p118 WS 3.8 Balanced species neutral and ionic equations

REVIEW:

12.1 Review p381
Key Qs p381

2.6 apply the definitions of oxidation and reduction in terms of oxidation numbers to a range of reduction and oxidation (redox) reactions

Oxidation Is Loss of electrons (OIL)

Reduction Is Gain of electrons (RIG)

The oxidation state is a measure of the degree of oxidation of a chemical species. Oxidation numbers are assigned to elements using these rules:

Rule 1: The oxidation number of an element in its free (uncombined) state is zero — for example, Al(s) or Zn(s). This is also true for elements found in nature as diatomic (two-atom) elements (Have No Fear Of Ice Cold Beer), and for sulfur, found as S8.
Rule 2: The sum of all oxidation numbers in a neutral compound is zero. The sum of all oxidation numbers in a polyatomic (many-atom) ion is equal to the charge on the ion. (This rule often allows chemists to calculate the oxidation number of an atom that may have multiple oxidation states, if the other atoms in the ion have known oxidation numbers.)
Rule 3: The oxidation number of a monatomic (one-atom) ion is the same as the charge on the ion, for example Na+ = +1, S2- = -2
Rule 4: The oxidation number of an alkali metal (Group 1 family) in a compound is +1; the oxidation number of an alkaline earth metal (Group 2 family) in a compound is +2.
Rule 5: The oxidation state of hydrogen in a compound is usually +1. If the hydrogen is part of a binary metal hydride (compound of hydrogen and some metal), then the oxidation state of hydrogen is –1.
Rule 6: The oxidation number of oxygen in a compound is usually –2. If, however, the oxygen is in a class of compounds called peroxides (for example, hydrogen peroxide), then the oxygen has an oxidation number of –1. If the oxygen is bonded to fluorine, the number is +1.
Rule 7: The oxidation number of fluorine is always –1. Chlorine, bromine, and iodine usually have an oxidation number of –1, unless they’re in combination with an oxygen or fluorine.

These rules give you another way to define oxidation and reduction — in terms of oxidation numbers. For example, consider the zinc reaction below, which shows oxidation by the loss of electrons. Notice that the zinc metal (the reactant) has an oxidation number of zero (rule 1), and the zinc cation (the product) has an oxidation number of +2 (rule 2).

A substance is oxidised when there’s an increase in its oxidation number.

Reduction works the same way. Consider the copper reaction below. The copper is going from an oxidation number of +2 to zero.

A substance is reduced when there’s a decrease in its oxidation number.

VIDEOS:

Oxidation states https://www.youtube.com/watch?v=5A515UMQyZ4 15.25

TASK 2.6.1

1. In which of the following reactions does the metal atom show the greatest change in oxidation state? (10, Q12)

A) MnO₄⁻ to Mn²⁺

B) MnO₂ to Mn(OH)₃

C) PbO₂ to PbSO₄

D) VO₂⁺ to VO²⁺

2. One test used for random breath testing in NSW involved crystals of potassium dichromate reacting with ethanol. In this reaction the orange dichromate ion, Cr₂O₇²⁻, changes to the green chromium ion, Cr³⁺. Which statement is true for this reaction? (09, Q9)

A) Chromium has lost electrons and reached a lower oxidation state.

B) Chromium has lost electrons and reached a higher oxidation state.

C) Chromium has gained electrons and reached a lower oxidation state.

D) Chromium has gained electrons and reached a higher oxidation state.

3. Why do oxidation and reduction occur together?

TASK 2.6.2

Complete PS&A p113 WS 3.5 Metals and their cations
Complete:

www.chemistry.wustl.edu/~coursedev/Online%20tutorials/Plink/Redox/redoxset.htm

REVIEW:

12.2 Review p388-9
Key Qs p389

2.7 conduct investigations to measure and compare the reduction potential of galvanic half-cells

A Galvanic Cell is a cell capable of producing an electric current from a redox reaction that occurs within it. Galvanic cells allow the conversion of chemical energy into electrical energy.

In galvanic cells, the cathode (reduction electrode) is positive; the anode (oxidation electrode) is negative.

The purpose is to separate the oxidation and reduction half-cells in order to use the electricity which has been created in a spontaneous reaction.

Galvanic cells use the same principle used in displacement reactions – that more reactive metals will transfer electrons to less reactive metals in order to displace them. However, when that reaction occurs directly, we cannot use the energy it generates. Instead, galvanic cells separate the two metals with a circuit through which the electrons will travel. By placing something in the middle (such as a light globe) we can harness the electricity generated. A simple galvanic cell is shown below this text box.

A key point to begin with: it is important that the electrode and its electrolyte are the same metal species, as otherwise a displacement reaction would occur within this beaker.

Now, what will happen in this situation? Copper ions will be reduced and will solidify on the copper electrode. At the same time, zinc atoms will be oxidised and will be released into solution. This process will cause electrons to flow from the zinc electrode to the copper electrode, through the external circuit. Using this knowledge, we can write half-equations for this cell.

The two electrodes are termed the ‘anode’ and the ‘cathode’. The anode is defined as the site of oxidation, and the cathode is the site of reduction. As a result, we can determine that electrons will flow from the anode to the cathode. Also, each side of the cell (one electrode and its electrolyte) is called a half-cell.

Now, what is a salt bridge? It is a physical bridge between the electrolyte solutions containing an ionic solution, usually potassium nitrate. This is necessary to allow for the ions lost or gained by each solution. If there is no salt bridge, the electrolytes will become charged and the current will not flow. Why choose potassium nitrate? To ensure that the ions flowing into each beaker do not precipitate – potassium and nitrate ions are always soluble.

PRACTICAL: Galvanic cell

Aim: To construct a galvanic cell.

Equipment:

safety goggles
beakers 100mL x 2
voltmeter
copper electrode
zinc electrode
copper sulfate solution 0.1M 50mL
zinc sulfate solution 0.1M 50mL
leads x2
filter paper
potassium sulfate solution 0.1M 50mL

Risk Assessment:

View the diagram and assess the risks

Method (see diagram):

Place copper electrode in copper sulfate solution in beaker 1. Connect wire from electrode to the voltmeter.
Place zinc electrode in zinc sulfate in beaker 2. Connect wire from electrode to the voltmeter.
Soak folded filter paper in potassium sulfate solution and fold over to connect the contents of the beakers.
Observe.

Results:

Anode: Zn(s) → Zn2+ + 2e– oxidation

Cathode: Cu2+ + 2e–→ Cu(s) reduction

VIDEOS:

You start at the anode https://www.youtube.com/watch?v=-bxJXt_69yM 3.28
RC21 Galvanic Cells https://www.youtube.com/watch?v=fe_DZ3mngBM&list=PLeFSFSJ9WqSAJfb02g1rwALZbMuRTFaf9&index=18 6.45
Galvanic cells https://www.youtube.com/watch?v=Ljn2bLy69SI&index=19&list=PLeFSFSJ9WqSA1rv6SqTeF0Uz_4ObE3bE2 20.39
Galvanic cell explained https://www.youtube.com/watch?v=YrTNQzwfs9o&index=57&list=PL49FAB756081270BA 14.45

TASK 2.7.1

The diagram below shows the construction of a galvanic cell from two half cells connected via a conductor and a salt bridge. Electrons are produced at the anode and flow through the conductor through the voltmeter to the cathode.

a) Write the two half-equations and the net ionic equation for the galvanic cell below.

b) Write 7- to 13-word sentences to answer the following: What is the purpose of the salt bridge? Why does it usually contain potassium nitrate solution?

TASK 2.7.2

Answer the following questions

TASK 2.7.3

Complete PS&A p115 WS 3.6 Generating electricity

REVIEW:

12.3 Review p395
Key Qs p395

2.8 predict the reaction of metals in solutions using the table of standard reduction potentials

The key concept underlying the function of galvanic cells is displacement reaction. This occurs between two metals which have different reactivity. In order to determine the reactivity of a metal, the standard reduction potentials on the HSC chemistry data sheet can be used (see table below this text box).

In this table, the most reactive metals are at the top and the least reactive are at the bottom.

The displacement reaction is simple: a more reactive species will displace less reactive species in solution. That means if a more reactive metal is placed in a solution containing ions of a less reactive metal, the metals will ‘swap places’ – the more reactive will move into solution and the less reactive will solidify.

Consider the following example where a piece of zinc is placed into a solution of copper sulfate. Zinc atoms will lose electrons to become ions and copper ions will gain electrons to become neutral atoms. These reactions are shown below this text box.

The zinc atoms are “oxidised” (lose electrons) and the copper atoms are “reduced” (gain electrons). This can be remembered with the acronym OIL RIG.

The reference, the standard hydrogen electrode, consists of:

Platinum foil coated in fine platinum powder dipping into a solution containing 1 mol.L^-1 hydrogen ions.
Hydrogen gas of 100 kPa pressure and 25°C temperature bubbling over the electrode.

When the standard hydrogen half-cell forms one half of a galvanic cell with another half-cell, then the potential difference of the whole cell is a measure of the standard electrode potential of the second half-cell, considering hydrogen to have an E^Θ value of 0.00V.

Once we have set up a second electrode in series with the hydrogen electrode, we can measure the voltage. By assuming hydrogen to be zero, a table of relative values can be produced. For each of these reactions, we can identify a particular oxidation and a reduction. We can represent these processes with simple half-equations.

2H⁺_(aq) + 2e^- → H_2(g) E^Ө = 0.00V Reduction half-equation

Mg_(s) → Mg²⁺_(aq) + 2e^- E^Ө = 2.36V Oxidation half-equation

2H⁺_(aq) + Mg_(s) → Mg²⁺_(aq) + H_2(g) E^Ө = 2.36V Balanced equation