Chemistry

Reactivity Series of Metals

A common property of metals is their tendency to form positive ions (cations) when reacting with nonmetals. The reactivity of a metal is measured by:

• How readily it forms cations
• Time taken to react with air, water, or acids
• Violence/exothermic nature of the reaction

The reactivity of common metals is shown below, with hydrogen and carbon as references.

Reactions with Water

Elemental metals react with water to form metal hydroxides and hydrogen gas:

• Potassium: Reacts violently, hydrogen gas may ignite due to exothermic reaction.
• Copper: Does not react with water.
• Nonmetals normally do not react with water, except exceptions like chlorine gas, which forms hydrochloric acid and hypochlorous acid.

Tip: The reactivity series can be remembered with the acronym: "Please Stop Calling Me A Lazy Zebra, Instead Try Learning How Copper Saves Gold".

Trial by Fire – Flame Tests for Metals

Flame tests are used to determine the presence of certain metals or metal ions based on flame colour:

• Potassium: Lilac flame (invisible through cobalt blue glass)
• Calcium: Brick red flame (appears light green through blue glass)
• Lithium: Bright red flame (invisible through green glass)
• Copper (I): Bluish-green flame
• Copper (II): Green flame
• Copper (III): Blue-green flame

Thermal Decomposition and Reactivity of Metals

Reduction, Oxidation, and Redox Reactions

Redox reactions: Reactions involving the transfer of electrons between elements or compounds.

Net Ionic Equations

Definition: Net ionic equations show only the elements or ions directly involved in a chemical reaction, excluding spectator ions (ions that remain unchanged in aqueous solution).

Example:

Reaction: NaCl + AgNO₃ → NaNO₃ + AgCl

All compounds can be broken into ions except AgCl (solid).

Spectator ions: Na⁺ and NO₃^- do not participate in the reaction, so they are cancelled out.

Net Ionic Equation:

Ag⁺ (aq) + Cl^- (aq) → AgCl (s)

Each ion in a redox reaction has a unique oxidation number (O.N.), which may differ from the valency. The oxidation number represents the total electrons an atom must gain or lose to form bonds in that reaction.

Extraction of Zinc from Zinc Blende

Zinc blende: The compound zinc sulphide (ZnS) from which pure zinc is extracted. Zinc has a low boiling point, so gaseous zinc can escape easily if heated excessively.

Steps in Extraction:

Extraction of Iron using Haematite

Haematite (Fe₂O₃) is a type of iron oxide commonly found in nature, from which pure iron is extracted through a series of processes involving a blast furnace.

In preparation, raw materials such as haematite ore, coke, and limestone are added into the blast furnace from the top, while hot air is introduced from below.

Electrolysis

Electrolysis can be indirectly interpreted to mean “the splitting of ionic compounds using electricity.”

Electrolysis is the process of separating ionic compounds using electricity. For soluble ionic compounds, they are dissolved in water to create an aqueous solution, such as sodium chloride. For insoluble ionic compounds, they are melted down to create a molten solution, like lead bromide.

The solution is called the electrolyte.

Next, two conductors of electricity, typically made of metal or carbon (graphite) called electrodes are placed into the solution. The positively charged conductor is called the anode, and the negatively charged conductor is called the cathode.

Finally, the electrodes are connected by a wire to a power source such as a battery to facilitate the flow of an electric current.

When the current flows through the solution, the positively charged ions, or cations, will move towards the negatively charged cathode, accumulating near it. Meanwhile, the negatively charged anions will move toward the positively charged anode. The negatively charged anions give up their borrowed valence electrons and become neutral atoms once more. The electrons travel through the wire and are accepted by the cations near the cathode, which, too, become neutral atoms.

In the case of lead bromide, the lead cations become lead atoms and sink down, while the bromide ions become bromine atoms and escape the solution as a gas.

Electrolysis – Detailed Ion Reactions

The Cl^- ions, being negative, are attracted to the positive anode, and give up their extra electron in their valence shell and form covalent bonds with one another, producing chlorine gas (Cl₂).

The H⁺ ions, being positive, are attracted to the negative cathode, and accept the extra electron from the chlorine gas, forming covalent bonds with one another and producing hydrogen gas (H₂).

Both can be observed as the formation of bubbles around the electrodes.

HCl (hydrochloric acid) is not produced because H⁺ and Cl^- ions are separated during the process.

Half-Equations:

• Oxidation: 2Cl^- → Cl₂ + 2e^-
• Reduction: 2H⁺ + 2e^- → H₂

Sodium hydroxide is less electronegative than chlorine and hydrogen, hence, it is not electrolysed.

Environmental Impact (Criterion D)

The electrolysis of brine has a massive environmental impact, as it requires a large voltage of electricity, which can only be efficiently produced through the burning of fossil fuels, which releases excess greenhouse gases and pollutants into the environment, contributing to air pollution and global warming. Furthermore, mercury is used later down the process when refining the sodium hydroxide, which, if it leaches into the natural environment or builds up inside human beings, can have devastating consequences.

Alloying, Steelmaking, and Types of Steel

What is steel?

Steel is an alloy of iron, carbon, and other metals that is used because of its strength in comparison to pure iron. It is produced through the process of steelmaking, where impurities are removed from molten iron and desirable components are added in required ratios.

Separation of sulphur:

Powdered magnesium is added to molten steel, where it reacts with sulphur to form magnesium sulfide (Mg + S → MgS), which forms on the surface of the steel as a slag and is skimmed off.

Removal of carbon:

Oxygen gas is introduced to the molten iron mixture, where it reacts with carbon to form carbon monoxide (CO), and escapes and removes itself as a gas.

Carbon is removed to ensure no excess. Too much carbon in steel results in brittleness.

Removal of other impurities:

Phosphorus and silicon impurities react with the oxygen too to form oxides. These oxides react with quicklime (CaO) to form a slag on the surface of the molten iron which is removed.

This is Aloy, not an alloy.

Types of Steel:

Carbon steel: 90% of steel produced is carbon steel, which can either be low, medium, or high carbon steel depending on the percentage of carbon in the alloy.

Low-carbon steel is used to make wire, medium-carbon for machinery like gears and shafts, and high-carbon for tools like knives.

Alloy steel: Alloy steels contain large amounts of other metals like manganese and copper, and are used to make things like generators and transformers.

Stainless Steel

This steel contains a minimum of 10% chromium, which makes it resistant to rusting (hence, no rust “stains”).

• Austenitic:
• Ferritic:
• Martensitic:

Electroplating and Electrorefining

Electroplating refers to the process of plating one metal onto another through the process of electrolysis. It can be used for a number of reasons, such as aesthetics (gold- or silver-plated jewellery), utility (astronaut visors used to be gold-plated to protect against the Sun’s radiation), or both (car wheel rims are plated for both reasons).

Process of Electroplating:

• Electrolyte: Solution containing metal to be plated onto the cathode.
• Anode: A strip or chunk of the pure metal to be plated onto the cathode to replenish the metal ions in the electrolyte.
• Cathode: The substance on top of which the metal is to be plated onto.

Using the example of copper plating:

The anode in this diagram is a strip of pure copper, the electrolyte is copper sulfate solution, and the cathode is made of silver.

When the current is turned on, the anode becomes positively charged, and the cathode becomes negatively charged.

Electroplating – Detailed Ion Movements

The current passing through the electrolyte splits the copper sulfate molecule (CuSO₄) into its constituent ions (Cu²⁺, SO₄^2-).

The positive copper ions are attracted to the negative silver cathode and move towards it. They undergo reduction at the cathode and settle on top of it, forming a layer of copper on top.

Cu²⁺ + 2e^- → Cu

The negative sulfate ions are attracted to the positive anode and move towards it. The copper ions in the copper anode undergo oxidation and give up their electrons, before bonding with the sulfate ions and dissolving in the solution, wearing down the copper anode.

Cu → Cu²⁺ + 2e^-

As such, a layer of copper is gradually built up on top of the silver cathode while the copper anode is worn away to replenish the copper sulfate electrolyte.

Therefore, the copper ions indirectly move from the anode to the cathode through the electrolyte solution, while the electrons move from the anode to the cathode through the wire.

Criterion D

There are several environmental concerns associated with the electroplating of chrome, a.k.a chromium. The process produces a harmful air pollutant and human carcinogen, which must be treated accordingly before being released into the environment– a process that is often very costly.

Factors affecting electroplating (courtesy of Iya):

• Voltage level of current: An increase in the electric potential difference results in an increase in voltage, and by extent, the speed at which electrons travel through the wire increases and the speed at which electroplating occurs increases.
• Temperature and chemical concentration of the electrolyte: An increase in the temperature of the electrolyte or the concentration of the electrolyte results in an increase in the speed of electroplating. An increase in the temperature causes an increase in the average kinetic energy of the particles, causing them to collide with one another faster and increase the rate of reaction. Furthermore, an increase in the concentration means more ions are plated per second, increasing the rate of electroplating.
• Duration for which current is supplied: If the current is supplied for a longer period of time, then the plating on the cathode grows thicker as more ions are given the chance to be reduced and plated onto the metal.
• Distance between cathode and anode: The closer the electrodes are, the faster ions can travel and the metal gets plated faster.

Electrorefining

Electrorefining is the process of refining or purifying a metal, usually copper, through the process of electrolysis. Electrorefining is typically the final stage of a metal purification stage, occurring after the metal ore is refined, converted to a metal oxide, and reduced to an impure or crude form of the metal.

Process of Electrorefining:

• Electrolyte: Solution containing the metal to be extracted, typically a molten metal salt of that very same metal. The electrolyte is typically acidified by adding a few drops of sulfuric acid to facilitate smoother movement of ions.
• Anode: A strip or chunk of the impure metal.
• Cathode: A strip or chunk of the pure metal.

Using the example of copper refining:

The anode in this diagram is a strip of impure copper. The cathode is a strip of pure copper. The electrolyte is an acidified copper sulfate solution.

When the current is turned on, the anode becomes positively charged, and the cathode becomes negatively charged.

The ions in the copper sulfate solution become delocalized and are split into its constituent ions, copper (Cu²⁺) and sulfate (SO₄^2-).

The positive copper ions are attracted to the negative cathode, where they are reduced and form a layer of pure copper on top.

Electrorefining – Anode Ions and Anode Mud

The negative sulfate ions are attracted to the positive anode. The copper ions are oxidized and lose their electrons, dissolving into the solution and bonding with the available sulfate ions to replenish the lost copper. Other metal impurities that are easily oxidized may also dissolve into the solution, however, they are not easily reduced, unlike copper, and instead of being plated onto the cathode, fall to the bottom of the container, forming an anode mud or anode sludge.

Criterion D: The value of the metals derived from the anode mud, such as gold, silver, and platinum, make up for the cost of electrorefining.

Corrosion

Corrosion is a process through which a pure metal is converted into a chemically stable oxide, thus becoming an undesirable substance, due to chemical or electrochemical reactions with the environment.

In the electrochemical theory of corrosion, metals have an anodic part and a cathodic part. As a result, the metal atoms in the anodic part of the metal become oxidized and lose one of their electrons, which moves to the cathodic part of the metal. The electrons at the cathodic part react with the compounds in the atmosphere, forming ions. These ions then react with the metal ions in the anodic region, forming chemical oxides and causing corrosion.

Thus, the metal ions left behind in the anodic part of the metal react with ions in the surroundings to corrode.

This pdf explains it really well: Module-2-Dr.SMP-Corrosion-and-Metal-Finishing-Final.pdf

Rusting is a special type of corrosion that occurs only in iron, or iron alloys such as steel. Hence, it is inappropriate to say a metal has “rusted” when the metal in question is not iron or an iron alloy.

Rusting only occurs upon exposure to both air and water. Exposure to only one of the two will not result in rusting.

Prevention of Corrosion

Corrosion can be prevented through two types of methods: barrier methods and sacrificial methods.

Barrier Methods

Barrier methods involve preventing iron from coming in contact with oxygen and water in the first place. This includes, but is not limited to:

• Painting the metal: Paint creates a barrier between the surface of the iron and any moisture outside, preventing rusting.
  Pros: Cheap, has aesthetic applications too.
  Cons: Doesn’t work with water-based paints, doesn’t work if scratched or peeled off.
• Oiling or greasing the metal: Applying a layer of oil or grease on top of the metal can accomplish the above, too. Oil is hydrophobic, hence, water won’t penetrate through the grease layer and corrode the iron. However, this only works for metal parts that are always moving, like a bicycle chain, as it acts as a lubricant and stays on the metal instead of flaking off like paint.
• Electroplating the metal: This is the process of using electrolysis to coat the reactive metal, like iron, with another metal. This could either be one that is less reactive, but forms a tough and durable layer on top, like nickel, or one that is extremely reactive and “sacrifices” itself to protect the iron, like zinc.

Sacrificial Methods

Sacrificial methods involve adding a more reactive metal like aluminum or zinc to the iron that will “sacrifice” itself to protect the iron (bro really thinks he’s the main character 💀💀). Thus, the more reactive metal is oxidized and is corroded in place of the iron.

Furthermore, connecting the metal to a more reactive metal means that the two different metals themselves become cathode and anode respectively. The more reactive metal supplies electrons to the less reactive metal, making it a cathode and protecting it from corrosion.

Galvanizing is a method that uses both sacrificial and barrier methods. It is the process of applying a protective zinc coating over steel or iron to prevent corrosion. Not only does the layer of zinc prevent moisture from reaching the iron in the first place, zinc is higher up on the reactivity series than iron. Thus, even if the zinc coating is scratched or damaged, it will react more readily with the moisture than iron, becoming a “sacrificial anode”.

Cathodic protection is also a method of protecting metals, particularly pipelines submerged underground and exposed to corrosive elements such as moisture in the soil, from corrosion. When it comes to metals in an electrochemical cell, the anodic part of the metal will corrode, while the cathodic part of the metal will not. Thus, by converting the metal into a cathode, corrosion can be prevented.

Impressed current cathodic protection involves connecting the anodic metal to an external power source such as a battery. This supplies a flow of electrons to the metal, making the rest of the metal a cathode and protecting it from external corrosion.

The Electrochemical Cell

An electrochemical cell is a device that generates electrical energy from the chemical reactions occurring in it, or vice-versa: it uses the electrical energy it receives from a power source to facilitate chemical reactions in itself.

Cells which execute the former are called galvanic or voltaic cells.

Cells which engage in the latter are called electrolytic cells.

Parts of an Electrolytic Cell:

• Anode: The positively-charged electrode submerged in the electrolytic solution, which attracts negative ions, or anions, and is the site where oxidation occurs.
• Cathode: The negatively-charged electrode which attracts positive ions, or cations, and is the site where reduction occurs.
• Electrolyte solution: The liquid which conducts electricity in which the electrodes are submerged.
• Salt bridge (only available for galvanic/voltaic cells): As electrolysis occurs and electrons travel from the anode to the cathode, charges begin to accumulate in the electrolytes. For example, in the anode half cell, the zinc ions are oxidized, lose their electrons, and become Zn2+ ions, which slowly become more and more concentrated within the solution as time passes. Meanwhile, in the cathode half cell, the Cu2+ ions in the solution are reduced, gain electrons, and build up on the surface of the cathode, leaving behind a large concentration of negative SO4 ions in the solution.

Thus, the electrolyte in the anode half cell becomes positively charged, and the electrolyte in the cathode half cell becomes negatively charged.

As a result, despite electrons flowing from the anode to the cathode, the flow of electrons stops because negatively-charged electrons will not be attracted towards the negatively-charged electrolyte of the cathode half-cell, as charges repel.

Thus, the salt bridge, which contains an inert cation and anion, helps fix this problem by ensuring electrical neutrality. The cations in the salt bridge move to the electrolyte in the cathode half-cell, reducing the negative charge. Likewise, the anions in the salt bridge move to the anode half-cell and reduce the positive charge. Thus, the flow of electrons resumes.

These ions are inert and do not react with those in the electrolyte, thus not impacting the function of the electrolytic cell.

Voltmeter: Measures the voltage of the electrical energy generated from the chemical reactions.

IMPORTANT NOTE

In electrolytic cells, the anode is positive and the cathode is negative. However, in galvanic cells, the anode is negative and the cathode is positive. In galvanic cells, the anode is the less reactive element, and it is considered to be negatively charged as electrons move away from the anode to the cathode, and like charges repel each other, so negative repels negative, hence the anode is considered negative. Meanwhile, the cathode is the more reactive element, and is considered positive because it attracts electrons, and opposite charges attract, so positive attracts negative, hence the cathode is considered to be positive.

Galvanic/Voltaic Cells:

A cell that uses chemical energy to drive the flow of electrons– essentially, the stark opposite of an electrolytic cell.

Fuel Cells:

Fuel cells are cells that convert chemical energy (such as fuel) into electricity, making them a type of galvanic or voltaic cells. They differ from batteries on the grounds that they require a continuous source of fuel from the surrounding air, whereas batteries already contain fuel. They can supply fuel for longer as a result as long as a constant supply of air is maintained, whereas batteries will eventually cease to function as their fuel is drained and will need to be replaced, generating waste.

The typical principle behind the functioning of a fuel cell is that the fuel enters the cell and is oxidized, creating a potential difference (a difference of charge between the electrodes) across the cell. This difference in charge drives the flow of electrons, creating electricity.

The most common type is a hydrogen-oxygen fuel cell, which combines hydrogen fuel with oxygen to produce water and generate electricity in the process. In these types of fuel cells, the electrolyte is potassium hydroxide, with the electrodes themselves made of porous carbon which contains a catalyst to speed up reactions.

The hydrogen enters through the anode compartment, and the oxygen through the cathode compartment. The water and heat exits through an outlet in the cathode compartment.

First, the hydrogen is oxidized at the anode. The anode initially has a positive charge, which makes the positive hydrogen ions migrate to the cathode, which has an initial negative charge. This leaves behind the electrons in the anode, which gives the anode a negative charge. The concentration of H+ ions in the cathode gives the cathode a positive charge.

The electrons move through a wire, generating a current, before reaching the cathode.

Half-reaction: H2 → 2H+ + 2e-

At the cathode, the hydrogen and the electrons react with oxygen to produce water.

Half-reaction: O2 + 4H+ + 4e- → 2H2O

Combining both half-reactions and cancelling out the spectator ions, we are left with the net ionic equation of the reaction taking place in the fuel cell:

H2 + O2 + 4H+ + 4e- → 2H2O + 2H+ +2e-

H2 + O2 + 2H+ + 2e- → 2H2O

2H2 + O2 → 2H2O

Criterion D– Real-life Applications:

Since they only require a supply of air to function and don’t produce as much waste, hydrogen fuel cells have been used on spacecraft and space probes for decades.

They have also been installed as a source of backup power in places like hospitals and power plants. Waste treatment plants also use fuel cell technology to make use of the methane gas produced from decomposing garbage.

Using hydrogen-oxygen fuel cells in place of fossil fuels can have several benefits. They can be produced readily, as they only require hydrogen fuel and oxygen as the reactants. Furthermore, it doesn’t produce pollutants such as carbon dioxide as waste, only producing water. Additionally, it has a longer lifespan than the average battery, as well as less polluting to dispose of.

However, manufacturing hydrogen fuel requires energy which often comes from burning fossil fuels, negating any positive impact using fuel cells may have had. Furthermore, being a gas, hydrogen is harder and less efficient to store, and is explosive when reacting with air, making it dangerous as well.

Aluminum Oxide (Al2O3)

A.K.A. alpha-aluminum, alumina, alundum, or aloxide. Inorganic compound.

Reaction with sodium hydroxide (NaOH):

Produces sodium aluminate and water, like so:

Al2O3 + 2NaOH → 2NaAlO2 + H2O

Reaction with sulfuric acid (H2SO4):

A notable fact about metal oxides is their typically basic nature, however, aluminum oxide is amphoteric, meaning it is both acidic and basic. In this reaction, it acts like a base:

Al2O3 + 3H2SO4 → Al2(SO4)3 + 3H2O

Reaction with hydrochloric acid (HCl):

Reacts with heated dilute hydrochloric acid to produce aluminum chloride solution:

Al2O3 + 6HCl → 2AlCl3 + 3H2O

Applications:

• Used in sunscreen, as well as cosmetics such as lipstick, blush, and nail polish.
• A good catalyst that helps speed up chemical reactions.
• Purifies water.
• Abrasive found in sandpaper.
• Used in sodium vapour lamps.

Batteries

Batteries are devices that accept, store, and release electricity on-demand. Energy is stored as chemical potential energy and converted into electrical energy through redox reactions.

Dry Cells:

Dry cells are types of batteries with an electrolyte in the form of a thick paste in place of a liquid. The most common type of dry cell is a zinc-carbon cell, better known as a Leclanche cell. The electrolyte in this case is ammonium chloride jelly, obtained by combining ammonium chloride with starch and flour.

Electrons flow from the negative zinc casing, which acts as the anode, to the positive carbon electrode, however, it is the manganese dioxide that is reduced, NOT the carbon electrode.

Li-ion Batteries:

Li-ion batteries are a type of rechargeable battery which uses lithium ions, whose reduction is reversible, to store energy. When discharging, electrons move from the anode to the cathode, however when charging, they move from the cathode to the anode.

The cathode is lithium iron phosphate, or LiFePO4, while the anode is graphite, or C6.

When charging, the electric current from the charger reduces the lithium atom of the cathode, causing the ion to migrate through the electrolyte full of lithium ions to the anode and reunite with its lost electron, only to bind with the graphite to form lithiated graphite, or LiC6. When the battery is discharging, this reaction is reversed– the lithiated graphite is oxidized and the lithium ion bonds with the iron phosphate to form lithium iron phosphate once again.

System, Surroundings, and Reactions

In thermodynamics, the system refers to the substances involved in the reaction only, namely the reactants and products. The surroundings refers to anything not directly involved in the reaction, such as the container in which the reaction occurs. During a chemical reaction, energy is transferred between the system and the surroundings.

E is the internal energy of a system and the sum of all its kinetic and potential energies. In a chemical reaction, there is a change within E, which can be illustrated through the following equation:

EΔ = E(final) - E(initial)

Where E(final) is the final internal energy of the system, and E(initial) is the internal energy of the system before the chemical reaction takes place. ΔE represents the total change in the internal energy of the system.

If ΔE is negative, it means that energy is being lost from the system into the surroundings. However, if ΔE is positive, it means that energy is being gained by the system from the surroundings.

Systems can either be open, closed, or isolated.

Endothermic Reactions

Endothermic reactions are ones where energy is transferred from the surroundings to the system.

A good example of this reaction is that between ammonium nitrate and water. This is commonly seen in hand coolers, where both compounds are mixed. The following reaction is endothermic and heat is absorbed, creating a cooling effect.

This reaction is reversible, too! So yippee for reusable hand coolers!

Exothermic Reactions

Exothermic reactions are ones where energy is transferred from the system to the surroundings.

If E < 0, the reaction is exothermic.
If E > 0, the reaction is endothermic.

Type of System

Energy transfers aren’t exclusive only to heat energy– even an insulated system can have energy transferred in and out of it through mechanical work done.

Thus, it is impossible for an isolated system to exist in real life, making it a purely hypothetical system.

Examples:

Open system: An open bottle filled with water.

• Mass (water) can be transferred in and out of the system by pouring it out or pouring more in.
• Energy (heat) can be transferred in and out of the system by feeling a warm sensation if you hold the bottle and the water is at a higher temperature.

Closed system: A closed bottle filled with water.

• Mass (water) cannot be transferred in and out of the system by pouring it out or pouring more in.
• Energy (heat) can be transferred in and out of the system by feeling a warm sensation if you hold the bottle and the water is at a higher temperature.

Activation energy is the minimum energy required for a reaction to take place.

Enthalpy and Heat Change

Enthalpy is the sum of a system’s internal energy and the product of the pressure exerted by it and its volume, and is most commonly written as a change in enthalpy, or:

ΔH = ΔU + (ΔP x ΔV)

Where ΔH = the change in enthalpy
ΔU = the change in internal energy
ΔP = the change in pressure
ΔV = the change in volume

Take, for example, a system which contains gas. If heat is added to the system, the gas will:

• Expand to fill the container, increasing in volume.
• Exert more pressure on the walls of the container.
• Have an increase in its internal energy.

Hence the formula above, which considers all three factors.

If pressure/conditions is/are kept constant, the change in enthalpy is equal to the heat evolved (meaning heat absorbed or released) by the system, or:

ΔH = q

Where ΔH = change in enthalpy
q = heat change

As a chemical reaction occurs, heat is evolved by the system, substances will increase or decrease in temperature. As a result, heat will flow from the substance to other objects surrounding it or vice-versa till a state of thermal equilibrium is achieved. If pressure remains constant, the heat change is an indicator of the change in enthalpy.

For example, if a fuel is burned in the air and 100 kiloJoules of heat are released, then the enthalpy of the system is reduced by 100 kiloJoules. Thus, the change in enthalpy for the system was -100 kJ, or ΔH = -100 kJ.

If the value of ΔH is negative, then energy is released from the system into the surroundings, and the reaction is exothermic.

If the value of ΔH is positive, then energy is absorbed by the system from the surroundings, and the reaction is endothermic.

H refers to standard enthalpy change, or the change in enthalpy under standard conditions.

Similarly, Hc is the standard enthalpy change of combustion. It is essentially a measure of the change in enthalpy of a combustion reaction, which is the amount of heat energy released by combusting a certain amount of the substance as calculated using a copper calorimeter, divided by the moles of the substance burnt.

Hc = Q / moles of substance burnt

To find the Q of a combustion reaction, consider a calorimeter (notes on this later on). The temperature change of water is used in the formula Q = mcΔT. The mass of water is calculated using the volume of water given and the density of water, which is set at 1g/cm³.

Finally, the specific heat capacity of the water is set at 4.18.

Heat change refers to the heat energy in Joules released or absorbed by 1 gram of a substance when its temperature is changed by 1 degree Celsius. Essentially, it’s the amount of heat absorbed or released by the substance by its specific heat capacity, and to calculate it, the following formula is used:

q = mcΔT

Where q = the amount of heat required to change the temperature of 1 gram of a substance by 1 degree Celsius.
m = the mass of the sample of substance
c = the specific heat capacity of the substance
ΔT = the final temperature of the substance - the initial temperature of the substance

Alternatively, to calculate the specific heat capacity of a substance, the formula can be tweaked as so:

c = q / mΔT

Enthalpy of Combustion

The enthalpy of combustion is the change in enthalpy (ΔH) or the amount of heat produced when 1 mole of a substance combusts (reacts vigorously with oxygen) under fixed conditions.

The heat released by the substance is equal to the negative enthalpy of the system, and is given by the following formula:

ΔH_C=-(mcΔTn)

where n = the number of moles of the substance used in combustion.

n can be calculated through the following formula:

n = Given mass / molar mass

Bond Enthalpy

This is essentially the difference of the energy required to break the bonds of reactants minus the energy released in the formation of the bonds of the products, or:

E_{required to break bonds} - E_{released when bonds are formed} = Bond Enthalpy of that reaction.

The energy required to break bonds will vary depending on the elements bonded to each other, or even the nature of the bond – is it a single bond, double bond, or triple bond?

To determine the nature of the bonds correctly, draw out a diagram first to avoid getting confused! For polyatomic ions that are used multiple times, identify a common bond and then multiply by how many times it is repeated. For example, for ammonium (NH4+):

There are 4 NH bonds, so the enthalpy can be calculated as 4(NH). If the ammonium ion is taken twice (NH4+)2, then you just double the number of NH bonds, so 8(NH).

The Rate of Reaction

A.k.a. the reaction rate. It’s the speed at which a chemical reaction proceeds, and is measured by the increase in concentration of products per unit of time and the decrease in the concentration of reactants per unit of time.

Factors Affecting Reaction Rate

• Temperature: If the temperature of a substance is increased, then the average kinetic energy of the particles of the substances involved in the reaction increases. As a result, the rate of collisions increases and the reaction rate increases.
• Concentration: The greater the concentration of each substance, the more particles of each substance are present within the system, and therefore the rate of collisions increases along with the reaction rate.
• Surface Area: Increasing the surface area of the reactants increases the area of the substances available to react with, thereby increasing the rate of collisions and the rate of reaction as a whole.

Catalysts and Inhibitors

A catalyst is a substance that participates in a reaction without being used up, unlike the reactants, and serves to speed up the rate of reaction by lowering the activation energy of the reaction. It accomplishes this by providing an alternative pathway through which the same products can be formed but with a smaller activation energy.

Inhibitors are the opposite of catalysts. Though they, too, are not used up in the reaction, they serve to slow down the rate of reaction instead of speeding it up.

Real-life applications: Medicines often work this way by inhibiting faulty, damaged, or harmful enzymes in the body from functioning.

Speaking of them, enzymes are biological catalysts which serve to increase the rate of biological reactions. Within the human body, the temperature is far too low for biological reactions crucial to life to occur at the desired rate. Increasing the body’s temperature could be downright fatal, hence, the body uses enzymes to speed up the rate of reaction without increasing the temperature.

This aids in processes such as the digestion of food, production of new cells, and even the clotting of blood.

Types of Reactions

Slow reactions are those which take an extremely long time to complete, often years or decades, and as a result have a low rate of reaction.

Moderate reactions are ones which, as the name suggests, take a moderate amount of time to complete (wow what a shocker, could never have guessed that one), ranging from hours to a few minutes, and have a moderate rate of reaction.

Fast reactions are those which are completed in a matter of seconds or less, and thus have a high rate of reaction and an extremely low activation energy. These include the combustion of a match or the petrol in a car’s engine.

Irreversible and Reversible Reactions

Dynamic Equilibrium

In reversible reactions, the reaction where the reactants are converted into the products is called the forward reaction, and the reaction where the products are converted back into the reactants is called the backward reaction.

For example, in the reaction 2HgO (s) ⥦ 2Hg (l) + O2 (g):

• 2HgO → 2Hg (l) + O2 (g) is the forward reaction.
• 2Hg (l) + O2 (g) → 2HgO is the backward reaction.

Both of these reactions occur at different rates or speeds. At the beginning of the reaction, the concentration of the reactants is large, while the concentration of the products is zero. In the case of the above reaction, there is a large amount of mercury oxide (2HgO) present, but no liquid mercury and oxygen gas (2Hg (l) + O2 (g)).

Hence, the reaction rate for the forward reaction is high as the reactants are decomposed into the products, but the reaction rate for the backward reaction is low because there are no products yet to be converted back into the reactants.

As the reaction proceeds, the concentration of products increases and the concentration of reactants decreases. In this case, there is now a larger amount of liquid mercury and oxygen gas, but a smaller amount of mercury oxide.

As a result, the backwards reaction can now commence due to the greater concentration of products. Hence, the reaction rate for the backwards reaction is high, while the reaction rate for the forwards reaction is low.

Over time, the reaction rate for both the forward and backward reactions will even out. Meaning, the decrease in the concentration of reactants and increase in the concentration of products per unit of time will be equal to the decrease in the concentration of products and increase in the concentration of reactants per unit of time.

As a result, the concentrations of the products and reactants will not change, as for every unit of product consumed to form the reactants, a proportional unit of reactants is consumed to form the products. Essentially, both reactions cancel themselves out and there is no net change in the concentration of either reactants or products.

!!IMPORTANT: Note that the concentrations of the products and reactants will remain constant, but this does not mean that they are equal. Or, to put it visually:

• Constant concentration of reactants and products
• Equal concentration of reactants and products

Even at equilibrium, there could be a large concentration of mercury oxide and only a small concentration of liquid mercury and oxygen gas, but as long as these concentrations do not change and reactions are still occurring, then the reaction has achieved a state of dynamic equilibrium.

However, the concentrations of the reactants and products determine the position of equilibrium. It describes the relative concentrations of reactants and products when the reaction is at dynamic equilibrium.

If the equilibrium lies to the right, it means that at equilibrium, the concentration of products is relatively greater than the concentration of reactants.

If the equilibrium lies to the left, it means that at equilibrium, the concentration of reactants is relatively greater than the concentration of products.

Equilibrium can only be reached in a closed system so that neither the mass of the reactants or the products can escape.

The position of equilibrium can be influenced by temperature, pressure, and the presence of catalysts.

Le Chatelier’s Principle:

Le Chatelier’s principle is as follows: “If the equilibrium of a system is disturbed by a change in one or more of the determining factors (as temperature, pressure, or concentration) the system tends to adjust itself to a new equilibrium by counteracting as far as possible the effect of the change.”

Temperature:

Let’s take the sample reversible reaction N2 (g) + 3H2 (g) ⥦ 2NH3 (g). Assuming that the reaction is conducted under constant conditions, the enthalpy of the reaction is equal to the heat evolved by the system, which is: ΔH = -92 kJ

This can be understood to mean that the enthalpy of the forward reaction, or N2 (g) + 3H2 (g) → 2NH3(g), is -92 kJ. As the enthalpy is less than zero, heat is removed from the reaction and released and the forward reaction is exothermic.

Hence, it can be assumed that the backward reaction, or 2NH3 (g) → N2 (g) + 3H2 (g), is endothermic, and an identical amount of heat is added to the system.

Though this is technically incorrect to write in the actual exam, let’s take heat as another reactant consumed as a result of the backward reaction. We get:

2NH3 (g) + heat → N2 (g) + 3H2 (g) (don’t write this in the exam! Or if you do, don’t throw me under the bus, you did this to yourself.)

Now, let’s turn up the heat– literally. Let us increase the temperature of the system. This means that more heat is added to the reactant side of the equation, or:

2NH3 (g) + even more heat → N2 (g) + 3H2 (g)

All this excess heat needs to be used up by the reaction, so it conducts the heat-absorbing reaction, or the endothermic reaction, at a faster rate or speed than the heat-releasing reaction, or the exothermic reaction. This is because we don’t need more heat to be released, we need it to be absorbed.

In the case of this reaction, the endothermic reaction is 2NH3 (g) + heat → N2 (g) + 3H2 (g), and it happens at a faster rate, compared to the exothermic reaction of N2 (g) + 3H2 (g) → 2NH3 (g) + heat.

Thus, equilibrium is achieved in the sense that the concentration of neither the reactants nor the products changes, but the speed at which each reaction occurs does change.

As the backward reaction is occurring at a faster rate compared to the forward reaction, the position of equilibrium has shifted to the right.

Concentration:

If the concentration of products in a system is increased, then the backward reaction is favoured to use up the additional products. This also occurs if the concentration of reactants is decreased, as the backward reaction is favoured to produce more reactants.

If the concentration of reactants in a system is increased, then the forward reaction is favoured to use up the additional reactants. This also occurs if the concentration of products is decreased, as the forward reaction is favoured to produce more products.

Pressure:

If the volume of a system, or the gas in this case, (not the surroundings, which in this case means the container) is decreased, the pressure it exerts on the walls of the container is decreased as the number of gas molecules colliding with the walls of the container is decreased, and vice versa: if the volume of a system (gas) is increased, then the pressure increases as there are more gas molecules colliding with the walls.

Let’s take the example of a reversible reaction resulting in the formation of gaseous compounds: dinitrogen tetroxide reacts to form nitrogen dioxide, or N2O4 ⥦ 2NO2.

The forward reaction is endothermic and absorbs heat to break bonds holding together dinitrogen tetroxide molecules to form nitrogen dioxide. Hence, the backward reaction is exothermic and releases heat as nitrogen dioxide molecules bond together to form dinitrogen tetroxide once again.

Thus, the forward reaction produces a greater volume of gas, and thus results in an increased pressure. The backward reaction counteracts this by producing a smaller volume of gas and thus producing a decreased pressure.

However, if additional pressure is added to the system, then the reaction which produces a decreased pressure is favoured to counteract this disturbance.

Thus, the backward reaction is favoured, and the position of equilibrium shifts to the right.

The opposite occurs if pressure is removed from the system. Then, the reaction which produces an increased pressure is favoured to counteract this.

Thus, the forward reaction is favoured and the position of equilibrium shifts to the left.

The Haber Process:

Despite the air around us being 78% nitrogen, plants can’t use it unless it’s in the soil, typically as ammonia. But with the world’s population rapidly increasing, and the need to feed everyone more important than ever, increasing crop yield for food is more crucial than ever.

Enter German chemist Fritz Haber, whose titular Haber process revolves around the simple reversible reaction of nitrogen, hydrogen, and ammonia:

N2 (g) + 3H2 (g) ⥦ 2NH3 (g)

This reaction is the reason why you and many others have food on their table today.

The Haber process relies on a reversible reaction, meaning ammonia can be decomposed back into nitrogen and hydrogen gas. However, to maximize the output of ammonia, Haber utilized Le Chatelier’s principle by adjusting the pressure and temperature at which the reaction takes place to achieve a state of dynamic equilibrium where the forward reaction is favoured.

Stages:

The reactants of nitrogen and hydrogen gas are taken. Nitrogen is extracted from the atmosphere, where it is in plentiful supply. Hydrogen is taken from natural gases that contain hydrocarbons like methane. Thus, the Haber process uses up 3-5% of the natural gas produced globally.

Both gases are fed into a compressor, where more and more of both reactants are added till the pressure they exert is equal to 200 atm.

The mixture of gases is moved into the reaction vessel, where they are heated till 450℃. The reaction vessel also contains iron catalysts to speed up the reaction. The catalysts speed up both the forward and backward reactions equally, increasing the reaction rate. Otherwise, it would take an extremely long period of time to achieve equilibrium. The gases react to form ammonia.

The ammonia, along with the unreacted gases, is sent to a cooling tank. The ammonia, as it has a low boiling point, is rapidly cooled to a liquid and extracted for use.

The unreacted nitrogen and hydrogen, as they have higher boiling points, remain as gases and are fed back into the compressor to be reacted again. This is done to increase the yield of ammonia and minimize waste.

The Haber Process (continued):

Temperature:

The Haber process is always conducted at a temperature of 450℃. The temperature is sufficiently high that the reaction rate of the particles is high enough for particles to react with one another at a fast enough rate. If the temperature is too low, it would take an extremely long time to reach equilibrium.

However, it’s also low enough to favour the reaction that releases heat, or the exothermic reaction, to compensate for the low temperature.

In this case, the forward reaction, N2 + 3H2 → 2NH3, is exothermic, as it releases heat to form bonds between nitrogen and hydrogen to make ammonia. Thus, it is favoured and the concentration of ammonia produced increases.

Pressure:

The volume of the mixture of gases involved in the Haber process is increased till the pressure of the system reaches 200 atm.

The pressure exerted by a system of gases on the walls of the container is directly proportional to the volume of the gas. Thus, to counteract the high pressure, according to Le Chatelier’s principle, the reaction which produces a smaller volume will be favoured.

In this case, the forward reaction, N2 + 3H2 → 2NH3, produces a smaller volume as 2 molecules of nitrogen gas and three molecules of hydrogen gas combine to form only 2 molecules of ammonia. Thus, it is favoured and the concentration of ammonia produced increases.

Real-life Applications:

• Dyes: Ammonia is used in hair dye to help the dye penetrate more evenly through the hair, resulting in a stronger, long-lasting colour.
• Nitrogen-based fertilizers: As initially mentioned, the Haber process was used to extract ammonia from the atmosphere to use in nitrogen-based fertilizers. Essentially, this made nitrogen more accessible to plants, boosting crop yields and the productivity of agricultural plots of land.
• If the low productivity of crop yields produced in the 1900s was maintained in the 21st century, four times more land would be required to feed the population. Furthermore, the land used for cultivation would cover half of all the land available today save for Antarctica and the Arctic. Instead, thanks to the Haber process, agricultural land only makes up less than 20% of total land area in most countries.
• Explosives: Nitric acid (HNO3) could also be produced from ammonia, and is a key component in explosives.

Drawbacks:

• Costly: Generating a temperature higher than 450℃ is costly due to the amount of fuel needed to be burnt. Thus, it is kept at that temperature to balance costs as well. Furthermore, increasing the pressure exerted by the system beyond 200 atm is not only very costly due to the amount of reactants needed to be extracted, but also unsafe, as the container could burst open and cause damage.
• Could be used for harm: The ammonia generated from the Haber process was used to produce nitric acid, a key component in explosives. The Germans used this to produce explosives in World War II and kill millions.
• Wastage of nitrogen: Plants can only use so much nitrogen to produce an efficient crop yield. As a result, excess nitrogen in the fertilizer runs off from crop yields and builds up in natural environments, causing harm. Furthermore, nitrogen builds up in the biosphere and gets released sparingly into the atmosphere, disrupting the natural nitrogen cycle.
• Eutrophication: Nitrogen runoff from agricultural areas into nearby water bodies creates “dead zones” as algae outcompete fish for oxygen and block light from entering below.
• Nitrous oxide emissions: Nitrous oxide (N2O), commonly known as laughing gas, also has a greenhouse effect on the environment, contributing to global warming. It is commonly produced as a byproduct of the Haber process.

Calorimeters:

Calorimeters are devices which measure the energy absorbed or released as heat during a physical or chemical change. This is used to measure the energy content of food.

These are insulated vessels with a reaction chamber where reactants are placed and sealed in. This chamber is surrounded by a layer of water, and then, a layer of air, to ensure little to no heat escapes into the surroundings. The reaction takes place in the “bomb”, which is the reaction chamber of the calorimeter. The heat is either absorbed or released in the reaction, which will cause the temperature of the surrounding water to decrease or increase respectively.

The change in the temperature of water is proportional to the amount of heat transferred, which is measured in joules (J) of energy.

The greater the calorific value, or the amount of heat released per unit of fuel, of a fuel, the less fuel is consumed overall.