Module 3: Reactive Chemistry

Outcomes

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A student:

› designs and evaluates investigations in order to obtain primary and secondary data and information CH11/12-2

› conducts investigations to collect valid and reliable primary and secondary data and information CH11/12-3

› selects and processes appropriate qualitative and quantitative data and information using a range of appropriate media CH11/12-4

› explores the many different types of chemical reactions, in particular the reactivity of metals, and the factors that affect the rate of chemical reactions CH11-10

Content FocusAll chemical reactions involve the creation of new substances and associated energy transformations, which are commonly observable as changes in the temperature of the surroundings and/or the emission of light. These reactions are harnessed and controlled by chemists to produce substances that lead to the development of useful products.

Chemicals can react at many different speeds and in many different ways, yet they basically involve the breaking and making of chemical bonds. Students study how chemicals react, the changes in matter and energy that take place during these reactions, and how these chemical reactions and changes relate to the chemicals that are used in everyday life.Working ScientificallyIn this module, students focus on designing, evaluating and conducting investigations to obtain and process data in the most appropriate manner in relation to chemical reactions. Students should be provided with opportunities to engage with all the Working Scientifically skills throughout the course.

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Chemical Reactions

Students:

● investigate a variety of reactions to identify possible indicators of a chemical change

Undertake Investigation 10.1 Chemistry in Focus p224

The following are indicators of chemical changes:

Change in Temperature.Change in Colour.Noticeable Odour (after reaction has begun)Formation of a Precipitate.Formation of Bubbles.A solid disappears

● use modelling to demonstrate

- the rearrangement of atoms to form new substances

- the conservation of atoms in a chemical reaction (ACSCH042, ACSCH080)

Investigation 10.2 Chemistry in Focus p225

● conduct investigations to predict and identify the products of a range of reactions, for

example:

- synthesis- decomposition- combustion- precipitation- acid/base reactions- acid/carbonate reactions (ACSCH042, ACSCH080)

1. synthesis

Synthesis reactions are reactions in which two or more substances combine to

form a new substance.

For example: burning of magnesium

2Mg(s) + O2(g) → 2MgO(s)

Discuss reactivity of elements on the Periodic table.

Production of Nylon – teacher demonstration

2. Decomposition

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Types of decomposition reactions:

1. Thermal Decomposition: Thermal decomposition occurs when a substance is

heated and breaks down into 2 or more new substances.

For example: The thermal decomposition of copper carbonate.

heatCuCO3 CuO + CO2

Investigation 10.3 Part A - Chemistry in Focus

2. Electrolysis

Electrolysis is the process in which an electric current is passed through a liquid or solid to bring about a chemical reaction.

For example: The electrolysis of water.

In pure water at the negatively charged cathode, a reduction reaction takes place, with electrons (e−) from the cathode being given to hydrogen cations to form hydrogen gas.

If bromothymol blue is added in this demonstration a yellow colour is observed at the anode due to the formation of H1+ and a blue colour is observed at the cathode due to the formation of OH1- .

Anode 2 H2O → O2 + 4 H1+ + 4 e1-

Cathode 2 H2O + 2 e1- → H2 + 2 OH1-

Teacher demonstration using Hoffman VoltameterStudents preview this video for homework the go over it and demonstrate process in class.https://www.youtube.com/watch?v=rb_ol8BRo_U

3. Decomposition by light

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Substances can be decomposed by sunlight. Silver chloride, when left in sunlight, decomposes to silver and chlorine gas.

light2AgCl(s) 2Ag(s) + Cl2(g)

Investigation 10.3 Part B - Chemistry in Focus

4. Combustion

Combustion reactions involve “burning” in air and are chemical reactions with

oxygen.

Lighting the gas in a Bunsen burner is a combustion reaction.

For example: burning of magnesium

2Mg(s) + O2(g) → 2MgO(s) Teacher demonstration

5. Precipitation

Precipitation reactions are those that when 2 solutions are mixed a solid forms.

When ionic substances mix, the solubility of the product determine if a

precipitate forms.

If the ionic bond is stronger that the interaction with water molecules then the

precipitate forms.

The formation of precipitates can be predicted by applying the solubility rules.

For example Lead ions in solution will form a yellow precipitate when mixed

with chloride ions.

Pb2+(aq) + 2Cl-

(aq) → PbCl2(s)

Investigation 10.4 p233 Chemistry in Focus

6. Acid/base reactions

Refer to previous work with neutralisation reactions

7. acid/carbonate reactions (ACSCH042, ACSCH080)

Refer to previous work with carbonate reactions

Check Your Understanding 10.5 & 10.6

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● investigate the chemical processes that occur when Aboriginal and Torres Strait

Islander Peoples detoxify poisonous food items

Students read and summarise Section 10.9, p243 of Chemistry In Focus

● construct balanced equations to represent chemical reactions

Ongoing throughout the course

Predicting Reactions of Metals

Inquiry question: How is the reactivity of various metals predicted?

Students:

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

– water

– dilute acid (ACSCH032, ACSCH037)

– oxygen

– other metal ions in solution

Reactions of Metals with Water

Most metals do not react with cold water or do so very slowly. However, Group 1

metals do. These metals react according to the following general equation:

Metal + Water Salt (hydroxide) + hydrogen

For example:

Word equation

Sodium + water sodium hydroxide + hydrogen

Chemical equation:

2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g)

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The reaction of Iron with water

On the reaction of iron and water is very slow, but on heating the rate of reaction

increases.

If we take 3 moles of iron and 2 moles of water, as a result of the interaction we

will get iron dioxide and hydrogen in gaseous form. The reaction between iron and

water has the equation:

3Fe + 2H₂O Fe₃O₄ + 2H₂

The reaction of metals with dilute acids

Some metals react rapidly with dilute acids, others react slowly and some have no

reaction.

Metals react with dilute acids according to the general reaction equation:

metal + acid metal salt + hydrogen

For example:

Magnesium reacts with sulfuric acid:

Word equation

Magnesium + sulfuric acid magnesium sulfate + hydrogen

Chemical equation

Mg(s) + H2SO4(aq) MgSO4(aq) + H2(g)

Investigation 11.2 Chemistry in Focus p251

The reaction of metals with oxygen

Iron swiftly oxidizes, or in other words it rusts in the presence of moisture. The re-

action is:

4Fe + 3O₂ + 6H₂O  4Fe(OH)₃

Since, sodium is a very reactive metal, it tends to react with oxygen to form sodium

oxide but this is an unstable compound and soon reacts with hydrogen to form

sodium hydroxide.

4Na + O2 → 2Na2O (s)

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The overall reaction of sodium in air (oxygen and water) is:

2Na(s) + 2H2O (l) → 2NaOH (aq) + H2 (g)

With the formation of sodium oxide being an intermediate phase.

The most common example would be the corrosion of iron (rust).

iron + oxygen iron(III) oxide

4Fe(s) + 3O2(g) 2Fe2O3(s)

Investigation 10.1 Chemistry in Focus p250

Reactivity of metals with other metal ions in solution.

Metal ions can displace electrons from other metals. For example:Zn(s) + Cu2+

(aq) Zn2+(aq) + Cu(s)

This is known as a metal displacement reaction.Metal displacement reactions along with the reactions with water, oxygen and dilute acids allow metals to be place into an activity series:

K Ba Ca Na Mg Al Zn Fe Ni Sn Pb Cu Ag Hg Pt Au

Strongest Reducing agent Strongest Oxidising agent

Note:An oxidising agent (or oxidant) causes another substance to be oxidised. Therefore, the oxidant is itself reduced.

Conversely:A reducing agent (or reductant) causes another substance to be reduced. Therefore, the reductant is itself oxidised.

Explain the displacement of metals from solution in terms of transfer of electrons.

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More active metals will displace less active metal ions from solution in an oxidation-reduction reaction.

When an active metal is placed in a solution containing ions of a less active metal, the active metal displaces the less active metal from solution. This occurs because a more active metal atom loses one or more electrons and becomes a positive ion. The electrons lost are transferred to the ions of the less active metal, resulting in them becoming metal atoms.

This is called a Metal Displacement Reaction

For example, if an iron nail is placed in a solution of blue copper (II) salt, some of the iron nails dissolves.

At the same time, the blue colour of Cu2+ ions disappears and a dark copper coating appears on the nail surface.

The overall reaction is:

The electrons lost by iron atoms undergoing oxidation are used to reduce copper (II) ions to copper atoms. Oxidation–reduction reactions (also called redox reactions) involve transfer of electrons.

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● construct a metal activity series using the data obtained from practical investigations

and compare this series with that obtained from standard secondary-sourced

information (ACSCH103)

Investigation11.3 Chemistry in Focus p254

Check Your Understanding 11.1, 11.2 & 11.3

Questions 1, 2, 7, 8, 9 & 10

● analyse patterns in metal activity on the periodic table and explain why they correlate

with, for example:

- ionisation energy (ACSCH045)

- atomic radius (ACSCH007)

- electronegativity (ACSCH057)

Ionisation Energy

Recap definition of ionisation energy

The reactivity of metals generally increases as their ionisation energy

decreases.

Atomic Radius

Overall, reactivity increases with an increase in atomic radius.

Going down groups 1 and 2 atomic radius and reactivity increase.

Electronegativity

Electronegativity increases from left to right across a period and decreases

down a group.

Metal reactivity is the opposite of this, it decreases from left to right across a

period and increases down a group.

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● apply the definitions of oxidation and reduction in terms of electron transfer and

oxidation numbers to a range of reduction and oxidation (redox) reactions

Previously covered – need to introduce oxidation

The oxidation number (state) of an element in a molecule or ion is the charge the atom of

that element would carry if the molecule or ion were completely ionic.

The sum of the oxidation numbers of all the elements in a species (molecule or

polyatomic ion) must be equal to the net charge on the species.

If the permanganate ion (MnO4-) were completely ionic, it would be Mnz+(O2-)4. To give a

net charge of –1, z would have to be +7. Hence, the oxidation number of manganese in

permanganate is +7.

Oxidation numbers help balance equations by helping to decide whether oxidation or

reduction occurs.

Oxidation numbers are assigned in an order of priorities in accordance with the following

set of rules:

1. A substance present in the elemental state is assigned an oxidation number of zero,

regardless of the formula of the molecule of the element, i.e. Hydrogen is diatomic,

phosphorus exists as P4 and sulfur as S8. All have an oxidation number of zero.

2. The oxidation number of a monatomic ion is simply the charge on the ion.

3. Hydrogen has an oxidation number of +1 except for the metal hydrides where it is

–1, e.g. sodium hydride, NaH or Calcium hydride, CaH2.

4. Oxygen has an oxidation number of –2 unless prior application of rules 2 and 3

dictates that it have a different value. In barium peroxide (BaO2) and hydrogen

peroxide (H2O2), oxygen has an oxidation number of –1. Generally oxygen is –2

except in peroxides when it is –1.

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Elements with Multiple Oxidation States.

Many elements may exist in different oxidation states in different compounds. For

example chlorine has the oxidation state of 0 as a free element but chlorine can exist in

five other oxidation states:

Acid Formula Oxidation state of chlorine

hydrochloric HCl -1hypochlorous HClO +1chlorous HClO2 +3chloric HClO3 +5perchloric HClO4 +7

When a metal atom undergoes a loss of electrons (oxidation), there is an increase in

the oxidation number of the metal from 0 to n.

When a metal reacts with dilute acid and releases hydrogen, the metal undergoes

oxidation (loss of electrons) while the hydrogen ions in the acid undergo reduction

(gain of electrons).

Example:  

Mg + 2H+ → Mg2+ + H2

Mg → Mg2+ + 2e- oxidation0 +2 oxidation state

2H+ + 2e- → H2 reduction+1 0 oxidation state

Magnesium changes from oxidation state 0 to 2. This is an increase, thus this is oxidation.Hydrogen changes oxidation state from +1 (in H+) to 0 (in the element H2). This is a decrease, thus this is reduction.  

Check Your Understanding 11.6 & 11.7 p267

Questions 4, 5 & 7

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● conduct investigations to measure and compare the reduction potential of galvanic

half-cells

● construct relevant half-equations and balanced overall equations to represent a range

of redox reactions

A galvanic cell is a device constructed so that a reductant and oxidant are physically separated, but connected by an external circuit made of a conductor (to carry electrons) and a salt bridge (to carry charged ions in solution). A galvanic cell is thus composed of two half-cells, a reductant half-cell and an oxidant half-cell. This arrangement ensures that electrons cannot go directly from the reductant to the oxidant, but they will move through the external circuit

Oxidation–reduction reactions normally take place by direct transfer of electrons between the reductant and the oxidant. For example, if zinc metal is placed in a solution of blue copper(II) ions, the blue colour fades as the zinc goes into solution (as colourless ions) and the copper metal comes out of solution as atoms.

The parts where electrons flow out of or into half-cells are electrical conductors called electrodes. Some galvanic cells use inert platinum or graphite electrodes.

Any solution containing ions is called an electrolyte. Electricity flows through electrolytes by the movement of charged ions, not electrons.

Electricity flows in electrodes (metals or graphite), or through connecting wires, by the movement of electrons. Electrons do not move through water or water solutions containing ions.

The anode is the electrode where oxidation occurs. In the galvanic cell example, this is the zinc electrode. The following reaction occurs here.

The cathode is the electrode where reduction occurs. In the galvanic cell example, this is the surface of the copper electrode where electrons are available for the following reaction to occur (resulting in a coating of copper).

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The salt bridge could be filter paper soaked in a conducting solution such as potassium nitrate solution.

Potassium ions (K+) and (NO3-) ions do not form insoluble precipitates with other

ions. The salt bridge allows the movement of ions between the half-cells. This prevents the buildup of positive charge in the zinc half-cell as negative electrons leave and the buildup of negative charge in the copper half-cell as negative electrons arrive.Positive and negative ions moving through the salt bridge keep a balance of negative and positive charge in each half-cell.  

This cell can be represented by the cell equation:

Zn(s)│Zn2+(aq)║Cu2+

(aq) │Cu(s)

Where the single vertical line represents a phase change and the double vertical line represents the salt bridge.

The electrons move through the conductor from the reductant half-cell to the oxidant

half-cell.

The energy of the moving electrons is electrical energy that can be used to turn an

electric motor, produce heat or light energy in a light globe. 

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A galvanic cell makes electrical energy available from chemical energy.

 

The larger the galvanic cell, the more chemical energy is stored and the more

electrical energy can be obtained from that cell.

Investigation 12.1 p275

Worked example 12.2 for Cells with gases

● predict the reaction of metals in solutions using the table of standard reduction

potentials

● predict the spontaneity of redox reactions using the value of cell potentials

(ACSCH079, ACSCH080)

Standard Electrode Potentials

Potential, Voltage and EMF

Electrical potential is a measure of the energy that a device can deliver.

Scientists use another term: the electromotive force or EMF of a galvanic cell is

the potential difference (voltage) across the electrodes of the cell when a negligibly

small current is being drawn. It is the maximum voltage that the cell can deliver.

Strictly speaking, for comparative purposes we should talk about the EMFs of

galvanic cells rather than just voltages.

The terms ‘potential difference’, ‘voltage’ and ‘EMF’ more or less

interchangeably.

The Standard Hydrogen Electrode

For reference purposes we tabulate voltages of electrodes relative to one particular

electrode, the standard hydrogen electrode.

It consists of a piece of platinum metal immersed in a 1.000molL-1 solution of

hydrogen ions (hydrochloric acid) and through which hydrogen gas is bubbled at a

pressure of 100.0kPa.

Platinum is used because it is a good electrical conductor and is so inert that it will not

take part in the reaction. The half reaction associated with this electrode is:

2H+(aq) + 2e- → H2(g)

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Measuring Standard Electrode Potentials

Electrode potentials are measured relative to the standard hydrogen electrode.

As cell voltage depends upon electrolyte concentration and gas pressure (if any gas is

present), we use the term ‘standard electrode potential’ for measurements that are

made under what are called standard state conditions.

The standard state for chemical measurements such as electrode potentials is

solute(s) present at a concentration of 1.000mlL-1 and any gas present to be a

pressure of 100.0kPa.

The standard electrode potential ͼ ꝋ, of an electrode is the potential of that electrode in

its standard state relative to the standard hydrogen electrode.

Standard electrode potentials are sometimes called standard reduction potentials or

standard redox potentials.

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Standard Electrode Potentials and Reactivity of Metals

A similar order of activity can be derived from a table of standard electrode potentials.

It has been established that if the standard voltage of a reaction is positive, then the

reaction goes as written; that is, the reaction is spontaneous.

When comparing two reduction half reactions, the one with the algebraically larger ͼ ꝋ goes as written and drives the other one in the reverse direction.

For metals, reactivity means the ease with which the metal can be oxidised (to the

positive ion); that is, the ease with which the oxidation half reaction can occur. Hence

the more negative the electrode potential the more reactive the metal.

Magnesium (ͼ ꝋ = -2.37V) is more reactive than Fe (ͼ ꝋ = -0.45V)

Check Your Understanding 12.4, 12.5 & 12.6 Questions 6, 8 & 9

Worksheets on the calculations of cell EMF.

Rates of ReactionsInquiry question: What affects the rate of a chemical reaction?

Factors Influencing the Rate of a ReactionFor most reactions, the rate increases as the temperature increases. For some reactions a 10oC increase can double the rate of reaction.Sometimes the rate of a reaction is increased by the presence of a substance that is not even involved in the stoichiometric equation for the reaction.Another example is the reaction of hydrogen peroxide with acidified iodide solution.

H2O2(aq) + 2I-(aq) + 2H+

(aq) → I2(aq) + 2H2O(l)

The rate of this reaction is greatly increased by adding sodium molybdite solution, Na2MoO4.A substance that increases the rate of a reaction without undergoing permanent chemical change in the reaction is called a catalyst.

In summary, the factors that influence the rate of a homogeneous reaction are:concentration of the reactants (in solution or in the gas phase)nature and concentration of any catalysttemperature

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Heterogeneous Reactions

Many reactions, called heterogeneous reactions, occur at the interface between two phases, such as between a gas and a solid or between a solution and a solid. Some common heterogeneous reactions are:

reaction of zinc metal with hydrochloric acid (to form hydrogen gas and zinc chloride)decomposition of hydrogen peroxide in solution (to form oxygen gas and water), occurring on the surface of various solids such as manganese dioxidereaction of hydrochloric acid with marble chips (calcium carbonate)

Rates of Heterogeneous Reactionsthe rates of heterogeneous reactions depend upon the three factors already listed for homogeneous reactions. However there are two additional factors:the state of division of the solid involved;the rate of stirring (or whether or not the mixture is stirred at all).

Investigation 13.3 demonstrated the effect of surface area. The reaction in that investigation would have been faster if the mixture had been stirred continuously to keep the crushed limestone dispersed throughout the solution instead of letting it settle to the bottom of the beaker.

Some Industrially Important Heterogeneous ReactionsHeterogeneous reactions are very important in industry. Examples include:

synthesis of ammonia (from N2 and H2) using an iron catalyst (for making fertilisers, nitric acid and explosives)hydrogenation of vegetable oils on finely divided nickel metal (to form semi-solid margarines)‘cracking’ (breaking) of high molecular weight compounds from crude oil to form lower molecular weight ones (for use as petrol); this occurs in the gas phase using certain solids as catalysts.

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investigate the role of activation energy, collisions and molecular orientation in

collision theory

explain a change in reaction rate using collision theory

Activation EnergyActivation energy is interpreted as an energy barrier between reactants and products the

higher this barrier, the higher the activation energy and the harder it is for reactants to

get over it and from products. Therefore, the reaction is slower (at a given temperature).

In order for a chemical reaction to take place, the reactants must collide. The collision

between the molecules in a chemical reaction provides the kinetic energy needed to

break the necessary bonds so that new bonds can be formed.

The minimum energy which must be available to a chemical system with potential

reactants to result in a chemical reaction.

Exothermic reactions transfer energy to the surroundings. 

Endothermic reactions take in energy from the surroundings.

Exothermic Reactions -

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Endothermic Reactions

Collision Theory of Reaction RatesThe collision theory of reaction rates proposes that, for chemical reaction to occur, the reactant particles must:

Collidehave more than a certain minimum amount of kinetic energybe correctly orientated

Quantitatively:

rate off reaction = {number of collisions per unit volume per unit time}

x {fraction of the collisions that involve more than the minimum energy}x {fraction of the molecules correctly orientated}

The number of collisions per unit volume per unit time can be calculated from the kinetic theory of gases. This theory proposes that:

gases consist of very small particles that are well separated from one anotherthese particles are in continuous random motionthe intermolecular forces are extremely small (because the particles are so far apart).

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At 300K only a very small fraction of the molecules in the sample has kinetic energy greater than the activation energy, E1 (the very small blue shaded area in Figure 13.7). Therefore the reaction rate is quite small. At 500K a much greater fraction of the molecules (both shaded areas) has kinetic energy greater than E1. This means that the reaction rate increases quite markedly with temperature.

The fraction of the molecules correctly orientated is illustrated in Figure 13.8 for the reaction of nitrogen dioxide with the carbon monoxide.

CO(g) + NO2(g) → CO2(g) + NO(g)

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How Catalysts Work

Catalysts usually work by providing a pathway of lower activation energy.

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Students:

conduct a practical investigation, using appropriate tools (including digital

technologies), to collect data, analyse and report on how the rate of a chemical reaction

can be affected by a range of factors, including but not limited to:

– temperature

– surface area of reactant(s)

– concentration of reactant(s)

– catalysts

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