377Chemistry Unit 4 Notes Complete

Liam Mariadas ©

CHEMISTRY: UNIT 4

CONTENTS

Area of Study 1:

Chapter 15: Fast and Slow Chemistry………..…2-12

Chapter 16: Controlling the Yield of Reactions………..12-21

Chapter 17: Equilibria Involving acids and Bases……….………….22-26

Detailed Study Covered: Sulfuric Acid:

Chapter 18: Production of Sulfuric Acid……….27-33

Area of Study 2:

Chapter 23: Fossil Fuels………33-37 Chapter 24: Alternative Energy Sources………37-42 Chapter 25: Energy from Chemical Reactions………..43-46 Chapter 26: Electricity from Chemical Reactions………46-53 Chapter 27: Cells and Batteries……….53-60 Chapter 28: Electrolysis………..60-65

Additional Information about these notes:

This is for the 2012 study design for VCE Chemistry. There is an additional Health and Safety chapter not present in these notes. Images taken from Heinemann Chemistry. Hopefully these notes will equip you with the knowledge needed for the exam.

Area of Study 1: Industrial Chemistry:

Chapter 15: Fast and Slow Chemistry:

Yes, this is the actual name of the chapter- which examines the changes in energy during chemical reactions as well as factors that influence the rate of these reactions.

Keep in mind:

 During chemical reactions, particles (atoms, molecules, ions) collide and undergo change- where atoms rearrange to form new particles.

 An example: Transfer of electrons (Redox) or transfer of protons (Acid/Base.)  However, collisions between particles do not always result in a chemical change

(i.e: when the use of a catalyst is necessary.)

Chemical Energy:

What is Chemical Energy?

 All substances have chemical energy.

 It is the sum of stored (potential) energy and kinetic (movement) energy. These energies may result from such things including:

 Attractions between electrons and protons in the atoms  Repulsion between nuclei

 Repulsion between electrons  Movement of electrons

 Vibration and rotation of and around bonds.

The chemical energy of a substance is often called its heat content or enthalpy, denoted by the symbol H.

When a chemical reaction does take place, atoms of the reactants are rearranged to form products with different chemical energies.

Depending on the relative energies of the reactants and products, two situations arise:

 Exothermic Reactions:

Total chemical energy of products is less than the energy of the reactants. The difference in this energy is usually released as heat energy (not lost.) Thus, exothermic reactions are the release of energy. (i.e: combustion of petrol.)



reactants. Meaning energy is absorbed from a surrounding environment. This absorbance of energy is an

endothermic reaction.

 Energy released/absorbed during a chemical reaction is called the heat of reaction.  The heat of the reaction is equal to the difference in enthalpy of the reactants and

products of a reaction, this it is denoted as ΔH .  ΔH = H (Products) – H (Reactants)

 ΔH, for Psychology students (simple terms), is the change in energy.

 For exothermic reactions, the value of H(Products) is less than H(Reactants) so the values of ΔH are negative.

 For endothermic reactions, the value of H(Products) is more than H(Reactants) so the values of ΔH are positive.

MEMORY NOTE:

Exo- “ex”-“exit sign”-“leave/release”- release of energy.

Endo-“end”- “end with more”-“end>start”-“products > reactants”- must absorbing energy.  Most reactions we encounter are exothermic, such as the burning of fossil fuels,

wood, coal, and even the oxidation of glucose in our bodies.

 Examples of endothermic reactions include condensation polymerization of glucose to starch.

Thermochemical Equations:

 Photosynthesis is another endothermic reaction, changing carbon dioxide and water to glucose and oxygen. The thermochemical equation for this is:

6CO2(g) + 6H2O(l)  C6H12O6(aq)+6O2(g); ΔH= +2803kJmol-

In this reaction 2803kj per mol of energy is absorbed when 6 moles of CO2 reacts with 6 mol H2O to form 1 mol glucose and 6 mol O2.

 Thermochemical equations show the energy released or absorbed during a chemical reaction. (Will be + in endothermic, - in exothermic.)

 Energy is measured in Joules (J) or kilojoules (kj), yet the heat of reaction or ΔH is in kj mol-, saying the energy corresponds to the mol amounts in the equation.

The thermochemical reaction for the oxidation of glucose: (reverse of photosynthesis)

C6H12O6(aq)+6O2(g)  6CO2(g) + 6H2O(l); ΔH= -2803kJmol- Note the negative ΔH value indicating an exothermic reaction taking place (energy released.)

CH3OH (l) + 3/2O2 (g)  CO2 (g)+ 2H2O(g); ΔH= -726kJmol-

Note: If we double the number of moles of methanol, the energy released also doubles: 2CH3OH (l) + 3O2 (g)  2CO2 (g)+ 4H2O(g); ΔH= -1452kJmol-

Keep in mind when writing, doubling the moles of methanol also doubled the moles of every other substance in the reaction.

Keep in mind: Temperature of surroundings will decrease during endothermic reactions as it absorbs energy, likewise during exothermic reactions when energy is released, the temperature of surroundings increases.

Activation Energies: Getting a Reaction Started:

What causes substances to not react with each other instantly?

 Recall what happens to bonds during a chemical reaction:

 Bonds between atoms in the reactants must be broken, requiring energy to be absorbed.

 Bonds between atoms in the products must then be formed, requiring a release of energy.

 In the combustion of methane, the bonds between methane atoms and oxygen atoms are first broken and this causes absorption of energy, then energy is released in forming the bonds of the products- carbon dioxide and water. As combusting methane is an exothermic reaction, more energy is released in forming, than absorbed in breaking.

 The heat of reaction, ΔH, is the net result of the energy absorbed in breaking bonds of reactants, and the energy released in forming them in products.

 Energy changes during the course of the reaction are shown in energy profiles. The top of the curve (activation or transition complex) is the intermediate step in the reaction. This is where bonds in the

reactants are partially broken and bonds in the products are partially formed.

The energy required to break the bonds of reactants so that the reaction can

successfully take place is the activation energy.

This is why natural gas does not

spontaneously combust- it cannot naturally meet the activation energy requirements, and must be lit to do so.

Making Reactions go Faster:

The rate at which reactions occur is of course important in chemical industry. Some reactions are extremely fast whilst others very slow. In industry reactions need to occur rapidly so profits can be made- maximizing reaction rates are a focus in industry.

 To understand how to speed up or slow down a chemical process we should visualize what happens to particles during a reaction and understand energy changes occurring.

Collision Theory:

 For a reaction to take place, particles must collide with each other with sufficient energy to overcome activation energy of the reaction.

 Majority of collisions don’t result in a reaction. Although the frequency of collisions is very high, explosions do not occur, as not all collisions result in a reaction.

 For collisions not to cause a reaction, the energy involved will be less than the energy needed to break the bonds in the reactant (less than the activation energy.)

 Only successful collisions, when the energy involved exceeds the activation energy, allows the chemical reaction to take place.

 Thus the rate of the reaction depends on the number of collisions taking place, as well, primarily, the amount of successful collisions that take place.

Factors that Affect Rates:

The following factors influence the frequency of collisions and the proportion of successful collisions. There are four main ways in which reaction rates can be increased:

 Increasing the surface area of solids.

 Increasing the concentration of reactants in solution (or increase pressure of gaseous reactants which increases the concentration of the gas particles.)

 Increasing the temperature.

 Adding a catalyst.

Increasing the Surface Area of Solids:

 In solids, only particles at the surface can be involved in reactions.

 If you crush a solid into smaller parts, there are more particles present at the surface to react (by making them into smaller parts we increase the surface area.)  A greater number of exposed particles means a higher frequency of collisions,

meaning a higher proportion of successful collisions, increasing the reaction rate.  Note: Surface area changes do not affect the energy present in the reaction. Increasing the concentration of the Reactants: (Also increasing Pressure of Gases):

 With more particles moving randomly in a given volume of solution, the frequency of the collisions increase and so more successful collisions occur.

 With gases, increasing the pressure of gases raises the concentration of the gas molecules, causing more frequent collisions also.

 Note: Changes in concentration does not affect the energy in the reaction. Increasing the Temperature:

 As temperature increases the average speed and average kinetic energy of the particles increases also. They have increased energy.

 More particles have enough energy to overcome the activation energy, thus more particles can actually react in successful collisions to form reactions.

 Particle’s speed increases, which increases the frequency at which collisions occur.  Thus as temperature increases, so too does the rate of the reaction.

Extending Collision Theory:

 This is to do with the point made above with temperature.

 Particles have a wide spread of kinetic energies, so increasing the temperature will increase the AVERAGE kinetic energy of particles.

 Although there will still be many particles with widespread energies, it increases the proportion of particles with higher energies, as shown below.

 Therefore, more particles at higher temperatures will also exceed the activation energy, increasing the likelihood of collisions forming reactions due to a higher proportion of particles able to overcome the activation energy.

The shaded area under the graph to the right represents the number of particles able to

successfully react due to their ability to overcome the activation energy. EA represents the

activation energy point.

As temperature increases this shaded region becomes larger. More particles are able to be involved in SUCCESSFUL collisions.

The factor of increasing particle energy so more react (as act. energy is met), has a greater impact on the reaction rate then the simplicity of increasing the frequency of collisions.

 Also, the orientation and angle at which particles collide have a factor in whether successful reactions occur. Below shows an unfavourable and favourable

orientation.

Catalyst:

 Many reactions occur more rapidly in the presence of particular elements or compounds. These substances are known as catalysts.

 Chemical industry uses catalysts extensively as without them, many reactions will be too slow and products cannot be obtained in an economical rate.

 Many catalysts are found by trial and error, and when found chemists may look for similar substances that may speed up the reaction rate even more.

Some are shown:

There are two types of catalyst:

 Homogenous Catalysts:

Are in the same state as reactants and products. (i.e: Chlorine gas atoms are catalysts on forming ozone gas to oxygen gas.)

 Heterogenous Catalysts:

Are in different states from the reactants and products. (These catalysts are preferred in industry as they can be easily separated from the product.)

How do Catalysts work?:

As previously studied, the process of adsorption is the forming of bonds between one particle (usually solid) and another particle.

 Catalyst work in the same way, where a solid catalyst is used in industry.

 Generally they use the largest possible surface area of the catalyst, as this will allow more reactant molecules to adsorb onto the catalyst for a faster reaction.

One type of reaction is the Haber process, where nitrogen (gas) molecules undergo hydrogenation (react with H2) using an iron catalyst.

Nitrogen and hydrogen molecules both adsorb onto the iron surface (by forming intermolecular bonds.)

As they do so, the bonds within the molecules of N2 and H2 break (or weaken) as a result of bonds forming with the iron catalyst.

This allows the normal reaction of H2 + N2 to occur, as they combine to form NH3, at a faster rate, as usually this reaction requires 3000 degrees temp. Note: This is an exothermic reaction releasing energy.

How do catalysts increase the rate of reaction?

 Catalyst speed up the rate of the reaction as they lower the activation energy needed for the reaction to occur, by providing an alternative reaction pathway.

 As activation energy is lowered, more particles can overcome the activation energy to result in more successful collisions, increasing the rate of reaction.

 Catalysts lower the activation energy but don’t change ΔH.

What has this chapter covered?

Chemical energy, energy changes during reactions, ΔH (enthalpy/heat of reaction), exothermic and endothermic reactions, thermochemical equations, activation energies, energy profiles, collision theory, factors affecting reaction rates, surface area changes, temperature changes, concentration changes, extended collision theory, catalysts.

Chapter 16: Controlling the Yield of Reactions:

Some reactions occur when not all the product is formed. This fact has consequences on industrial companies – as large amounts of unreacted materials are costly and wasteful.

Thus the profitability of the company depends on the reaction yield- that is the extent of the conversion of reactions to products (how far the reaction will go).

Take this example reaction:

N2 + 3H2  2NH3

One would expect that 2 mol NH3 is formed from 1 mol N2 and 3 mol H2, however, one might find much less than 2 mol of NH3 forming. The reaction has seemed to “Stop”-before the reactants are completely consumed- as it is incomplete, and remains as such. The stage at which the quantities of reactants and products in the reaction remain unchanged during a reaction is called chemical equilibrium.

Chemical Equilibrium: The point in a chemical reaction where the rate of the forward reaction is the same as the rate of the back reaction, so that the quantities of reactants and products in the mixture remain unchanged.

These questions are vital to consider in an effort to maximize the yield in a reaction, so that industrial processes are more efficient:

 How can the amount of product from a reaction that reaches equilibrium be increased?

Why are some reactions Incomplete?: Reversible Reactions:

Some physical and chemical changes can be reversed.

 Water (ice form) melting if placed in a drink and then freezing when in the freezer can be noted by this equation: (The double arrows indicate reversible reaction)

H2O (s)↔ H2O (l)  It is made from the forward reaction of water melting:

H2O (s) H2O (l)  And the back reaction of water freezing:

H2O (l)  H2O (s)

Likewise, the reaction of hydrated copper sulfate (CuSO4.5H2O) when heated is

reversible, as it forms a white precipitate (CuSO4) and water- But, when the water formed reacts with the precipitate, it re-forms the original reactant (hydrated copper sulfate):

 CuSO4.5H2O (s) ↔ CuSO4 (s) + 5H2O (l) In the same way this reaction is reversible.

Equilibrium Explained:

Chemists have shown that forward and reverse reaction occur simultaneously. Take the above example with nitrogen gas and hydrogen gas. The following occurs:

1. Nitrogen gas and hydrogen gas will react immediately to form ammonia.  N2 + 3H2  2NH3

2. As the forward reaction proceeds, the concentrations of nitrogen and hydrogen decreases- so the rate of reaction decreases (less reactants), so

the rate at which ammonia is produced decreases- the quantity of ammonia present increases.

3. As more ammonia is formed, ammonia will break down to reform N2 and H2, thus the rate of the back reaction will increase (as the concentration of ammonia increases).

4. Eventually the forward and back reactions proceed at the same rate. At this rate ammonia is formed at the exact same rate as it breaks down. The concentrations of NH3, N2 and H2 will remain constant.

 N2 + 3H2 ↔ 2NH3

The reaction has reached equilibrium (note reversible arrows now) and the forward and back reaction rate do not change.

The reaction is described as dynamic, as although the forward rate is the same as the back rate, they both are still occurring simultaneously. During dynamic equilibrium: (step 4):

 Amounts and concentrations of substances remain constant.

 The total gas pressure is constant (if gases are involved.)

 The temperature is constant.

 The reaction is incomplete (All substances are present in the equilibrium mixture.)

How far do Equilibrium reactions go?:

 Different reactions proceed to different extents.

Heinemann uses the example of HCL and H2O proceeding more into the reaction than ethanoic acid and H2O, as seen by the electrical conductivity of the resultant solution, in which H+ and Cl+ ions are formed far more from HCL+ H2O, than the ion CH3COO- which is formed in CH3COOH + H2O. (HCL/H2O conducts more, proceeds more into reaction.)

The Equilibrium Law: Consider the reaction:

Now consider the table of data and concentrations (the reaction is at 400 degrees):

Note: [x] denotes the concentration of the substance.

As can be seen there is no relationship between the initial concentrations of N2,H2 and NH3 in any trial mixture (A.B,C or D)- particularly, NH3 is will not be calculated with a moles formula (using equation), as the reaction does not go to absolute completion.

Also, the is not consistent, as it does not reflect the mole ratio of the reaction.

Disregard this, but notice

has indices corresponding to the mole ratio of the equation. This, in the table, gives a constant value (almost, anyway) for each mixture.

K =

“K” is known as the equilibrium constant. “A and B” are the concentrations of the products, and “C and D” are the concentrations of the reactants. “a,b,c,d” are the mole ratios in front of each substance-according to the balanced equation representing it.

 This concentration fraction is known as the reaction quotient for the reaction.

 Only when the reaction is at equilibrium does the reaction quotient have a constant value, equal to K. Until this point it will not be equal to K.

 Different types of chemical reactions have different values for K.

 For any given type of reaction that is the same, the reaction quotient will give about the same value for “K”, if at a fixed temperature.

(i.e: If HCL + NaOH are reacted, but at different concentrations each time, the quotient will still give a relatively consistent value for all HCL + NaOH reactions.)

From the above example, the equilibrium constant “K” for the reaction: N2 + 3H2  2NH3 at 400 degrees is 0.052.

 The units for “K” can vary according to the equation. It requires index laws. For example: (Where M is the concentration in mol L of each substance):

The denominator is virtually M^4, and with index laws, indices in fraction subtract from the numerator leaving M^-2. This means the unit for “K” is , or: .





Gives us an indication of the extent of the reaction- how far the forward reaction proceeds before equilibrium is established.

 For values of K between and , a large amount of both reactants and products will be present at equilibrium.

 For values of K that are very large, above , the equilibrium mixture consists mostly of products with relatively small amounts of reactants.

 For values of K that are very small, below , the equilibrium mixture consists mainly of reactants and relatively small amounts of products.

If not many reactants are present in the reaction mixture it means the forward reaction has proceeded relatively far, and vice versa if there is more reactant in the equilibrium mixture.

Effect of Temperature on Equilibria:

Only temperature effects the value of the equilibrium constant, K.

The affects of temperature change on an equilibrium depends on whether the reaction is exothermic or endothermic. As temperature increases:

 For exothermic reactions, the amounts of products decreases and so the value of K decreases.

 For endothermic reactions, the amount of products increases and so the value of K increases.

Since the value of K depends on the temperature and the mol ratios, both the temperature and the equation should be noted or stated when an equilibrium constant is involved. This means, if the temperature is increased on a reaction that is exothermic, the BACK reaction of this will increase (as it will be endothermic), and vice versa.

Changing the Equilibrium Position of a Reaction:

The composition of the equilibrium mixtures is important to industrial chemists (that is, the more product in the mixture the better.) Ways to maximize the yield for the desired products is sought after.

The equilibrium position (amounts of reactants and products) of a reaction may be changed by:

 Adding or removing a reactant or product.

 Changing the pressure by changing the volume (if equilibria involves gases.)

 Dilution (if equilibria is in solution.)

 Changing the temperature.

Adding an extra Reactant or Product:

For example, take a vessel containing an equilibrium mixture represented by:

N2(g) + 3H2(g) ↔ 2NH3(g)

If more nitrogen was added to the vessel without changing anything else, the mixture that was once in equilibrium will momentarily no longer be. The following occurs as the

composition of the mixture adjusts to return to equilibrium:

 Forward rate of reaction increases due to higher concentration of nitrogen (reactant), which means more ammonia is produced.

 As the concentration of ammonia then increases, the rate of the back reaction also increases to reform N2 and H2.

 Eventually the forward & back rate become equal again- new equilibrium formed. When the new equilibrium is formed the concentrations of all substances have changed.

 Addition of a reactant: Increases the concentration of the product (as more can be formed) as there is more forward reaction. There will also be a higher

concentration of reactants at the new equilibrium than the old one even though reactants get used up during the forward reaction. The other reactant that wasn’t added in addition should decrease in concentration in the next equilibrium.

 Addition of a product: Increases the concentration of the reactant (as more can be formed) as there is more back reaction. There will also be a higher concentration of products at the new equilibrium than the old one even though products get used up during the back reaction.

Removing a product or reactant from the original equilibrium mixture will have the opposite affect on the concentrations of the substances at the new equilibrium point.

 In summary, this principal can be used:

“If an equilibrium system is subjected to change, the system will adjust itself to partially oppose the effect of the said change.”

Changing the Pressure:

By changing the Volume: For gases, the pressure increasing in an equilibrium mixture can be caused by decreasing the volume of the container in which the gas is kept (whilst maintaining a constant temperature.)

So how does pressure affect the equilibrium position? Take this example:

On the reactants side, there are 3 particles of gas, on the product side there are 2 particles. This means in the forward reaction there will be a reduction in pressure (as it turns 3 particles into 2 particles, decreases number of particles, decreases pressure). This causes a net back reaction to occur (to increase pressure, changes the 2 particles to 3) to adjust the pressure to normal- this continues until adjustment has been made.

Therefore, the adjustment of pressure depends on:

 The number of particles on each side of the reaction.

 If the reactant side has more particles in it, the forward reaction will decrease the pressure, the back reaction will increase the pressure. (see above for explanation).  If the container volume is decreased, the pressure increases, so the equilibrium

mixture will undergo more back or forward reaction (whichever one decreases the pressure), to adjust the system. This depends on the particle number on each side.

 Increasing the volume (decreasing pressure) goes more to side with more particles.

(I.e: For N2 + 3H2  2NH3, a net back reaction will occur).

 Decreasing the volume (increasing pressure) goes more to side with less particles.

(I.e: For N2 + 3H2  2NH3, a net forward reaction will occur.)

By Adding an Inert Gas:

If an unreactive gas is added (such as helium, neon or argon) the pressure increases because there is more particles- but none of the original substances are affected since the

gas is inert. The concentrations of the substances are not affected. The system therefore stays at equilibrium with no net forward or back reaction.

Dilution:

 When you dilute a solution, you decrease the number of particles PER VOLUME.

 Therefore, when you dilute something, the reaction will favor the side which creates more number of particles.

 This is similar to how changing the volume works (it depends on the number of particles on each side of the reaction.)

Therefore for:

If we dilute the solution, a net back reaction will occur (this creates more particles.)

Changing the temperature:

As temperature Increases:

 For exothermic reactions, the amounts of products decreases and so the value of K decreases.

 For endothermic reactions, the amount of products increases and so the value of K increases.

As temperature increases:

 A net back reaction (less products) for exothermic reaction occurs.

 A net forward reaction (more products) for endothermic reactions occurs. The Influence of a Catalyst:

 Therefore, Catalysts don’t change the position of an equilibrium, and therefore have no effect on the equilibrium yields of a reaction.

 However, catalysts may greatly increase the rate at which equilibrium is attained. Do All Reactions Reach Equilibrium?

 The simple answer is: No.

Reactions can sometimes continue until they are complete (not plateau to equilibrium), If:

 Reactions that produce products (such as gases) escape from the reaction mixture as they are formed. Continual loss of the product drives the reaction forward. (If a product escapes, the mixture adjusts to produce more (forward reaction), which will continue and continue until no reaction is left- completion.)

 Reactions that form equilibrium in which only minute quantities of reactants are present. (i.e: HCL and water.)

Chapter 17: Equilibria Involving Acids and Bases:

 As we know, acids are proton (H+) donators, bases are proton (H+) acceptors.

 Some substances may behave as both acids and bases depending on the conditions, these substances are amphiprotic. (i.e: Water.)

Ionization constant of Water:

Water can undergo self ionization in this reaction:

This means at equilibrium:

However, in aqueous solutions water’s concentration is usually constant at 56M, and so far more abundant than any other substance- we therefore don’t include H2O in

calculations. Thus:

 Where Kw is the ionization constant specific to water. It is (K * [H2O] where H2O is constant.)

 It (Kw) applies to both pure water and all aqueous solutions.

 In pure water (at 25 degrees celcius) chemists have found the concentration of both H3O+ and OH- to be about the same at: M.

 Thus, the value of Kw at 25 degrees celcius is:

= _{* =}

Note: The concentrations of ions in water is very, very small- there is much more water than ions in pure water. However, the ions are still present so water can conduct electricity (only very slightly!)

Acidic and Basic Solutions:

In solutions of acidic substances, any acid in a reaction with water produces H3O+ ions. This also occurs when water self ionizes.

 Thus, in acidic conditions, the concentration of H3O+ will be greater than

 The concentration of OH- will also be less than This is because the Kw ([OH-]*[H3O+]) is constant.

In Summary, at 25 degrees:

 In pure water and neutral solutions,

 In acidic solutions:

 In basic solutions:

The higher the concentration of H3O+ in the solution, the more acidic the solution is. In strong basic solutions the concentration of H3O+ is low (around ) for example in NaOH. In strongly acidic solutions, like HCL, this can rise to 10M.

Indicators that change colour at specific levels of H3O+ concentrations can help measure and determine the acidity of a solution.

pH: A Convenient way to Measure Acidity:

A pH scale has been used to measure the range of acidity in a solution. pH is defined as:

Alternatively:

 This means at 25 degrees pure water, as [H3O+]= ,the pH of the solution is 7.

 For acidic solution, the pH is < 7, as the more acidic a solution is, the lower the pH.

 For basic solutions, the pH is > 7. As the more basic a solution is, the higher the pH.

Keep in mind, pH is only a measure of the [H3O+] concentration, so in finding the pH you must obtain the H3O+ concentration value for this.

How is pH affected by temperature?

The pH of pure water is only exactly 7 (other times it is very close), at 25 degrees Celsius. But what if a solution is not at 25 degrees Celsius as we have considered?

It works similarly to how equilibrium worked:

 An increase in temperature favors the endothermic reaction.

 A decrease in temperature favors the exothermic reaction.

Therefore, lets take:

It is an endothermic reaction, if we increase the temperature, it’s forward reaction is favored. This means the concentrations of H3O+ and OH- increase (as more product is formed). Thus, according to previous formula, the Kw of the solution increases. This means the pH of the solution decreases (more H3O+).

 Likewise, for endothermic reactions, a decrease in temperature will cause an increase in pH.

Acidity Constants:

Most acid-base reactions in water can be considered as equilibrium reactions. Let HCL and water be an example:

From knowledge, we know that in the above reaction Cl- in the conjugate base of HCL, as it has lost a proton. Likewise, HCL is the conjugate acid of Cl-. (As a note, just reverse the reaction to see the gaining or losing affect of H+ to work out conjugates.)

The acidity constant, , can be calculated as such: (Water is left out as it is constant.)

The acidity constants of some acids are shown below:

The table means that as 25 degrees, the Ka value for Hydrochloric acid is . This means most of the HCL in solution has converted to H3O+ and Cl-. This high value is why HCL is a strong acid.

In contrast, ethanoic acid has a very low Ka value at 25 degrees, . The equilibrium

position has/favours the reactants, so there is less product. This is why CH3COOH is a weak acid.

The acidity constant is thus used as a measure of acids strength.

Buffers: Using Equilibrium to Resist Change:

 Solutions that can absorb the addition of acids and bases with little change in pH.

 Made most of the time by mixing a weak acid and a salt of it’s conjugate base. For example, a buffer can contain ethanoic acid and sodium ethanoate. The equilibrium mixture will contain CH3COOH, H3O+ and CH3COO-.

The important aspect of this solution and of buffers is that it contains significant amounts of BOTH the weak acid and it’s conjugate base. If a strong acid was used- this wouldn’t be the case as it would ionize completely.

 So, if a strong acid is added into a buffer, the pH will decrease but by much less than expected. In this, equilibrium works the same as it usually does.

 If a strong acid is added, more H3O+ is present so the equilibrium will try to oppose it (use up H3O+)- in this case it favours the back reaction to make more CH3COOH.

 Likewise, adding a base consumed H3O+, so the equilibrium will try to oppose this if a base is added and form more H3O+.

Buffers are important for delicate environmental and living systems where it is important to maintain stable conditions, also in labs where surroundings of an experiment should be kept constant.

pH in the Body:

Many reactions in the body are acid-base. The pH must be controlled within a narrow range so the body can survive.

 Acidosis- Lethal drop in pH in the body.  Alkalosis- lethal rise in pH in the body.

Natural buffers prevent this from happening and controls the pH levels.

Chapter 18: Production of Sulfuric Acid:

Sulfuric acid is used extensively in the production of fertilizers (such as superphosphate). It is also used in the manufacturing of paper, dyes, drugs,

detergents, petroleum refining and is an electrolyte in car batteries.

Superphosphate:

 Phosphorus is needed for plant growth, but must be added since Australian soils are deficient in this. In manufacturing superphosphate, insoluble calcium

phosphate (Ca3(PO4)2) contained in rock phosphate is converted to a soluble form that plants can absorb.

 Sulfuric acid is reacted with rock phosphate, over several weeks superphosphate is formed.

 Another fertilizer it also produces is ammonium sulfate:

Other Uses:

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The Contact Process:

Raw Materials for Making Sulfuric Acid:

Sulfuric Acid can be manufactured in a process beginning with the starting substance, Sulfur dioxide. The sulfur dioxide is then oxidized to form sulfur trioxide, and then follows the conversion to the acid- sulfuric acid.

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The sulfur dioxide used to produce sulfuric acid is obtained from two sources:

 Combustion of sulfur recovered from natural gas and crued oils.

 Sulfur dioxide formed during the smelting of sulfur ores of copper, zinc or lead.

 Mining of underground deposits of elemental sulfur by a method known as the Frasch process but this isn’t used much in Australia (other two are more available.)

The sources used in Australia are environmentally friendly as they use by-products from other industries, limiting the amount of SO2 emitted into the atmosphere.

Less fossil fuels are being used to obtain sulfur for use in sulfuric acid, with most of the sulfur dioxide recovered from smelters.

STEP 1: Contact Process: BURNING SULFUR:

 When sulfur is used as a raw material for making sulfuric acid, the first stage of this process is spraying molten sulfur (liquid) under pressure into a furnace.

 This sulfur then is heated and burns to produce sulfur dioxide. (Combustion)  Combustion is rapid due to the sulfur spray being of high surface area.

 Temperatures up to 1000 degrees celcius may be reached.

 Sulfur dioxide gas formed is then cooled for the next stage in the process. STEP 2: Catalytic Oxidation of Sulfur Dioxide:

 Sulfur Dioxide gas is then oxidized in this next step to produce sulfur trioxide gas by using oxygen, and using Vanadium (V) oxide (V2O5) as a catalyst.

 This step is performed in a reaction vessel (known as a converter).

 Sulfur dioxide is mixed with air and passed through trays containing loosely packed porous pellets of catalyst (catalyst beds of Vanadium).

 The air enters from the top (see diagram) supplying oxygen, passing over each catalyst bed level in succession with the Sulfur dioxide gas.

 As the reaction is exothermic, it is neccessary to cool the gas mixture as it passes from one tray to another so a desired reaction temperature is maintained. (Exothermic causes it to be a hotter environment- not always ideal.)

 The temperature is maintained at around 400-500 degrees, with a pressure of 1 atmosphere. Almost all Sulfur dioxide is converted to Sulfur trioxide in this process.

Note that knowledge of equilibrium and reaction rate play an important role in this:

The reaction yield of SO3 will increase as:

 The temperature decreases. The reaction is exothermic, then the temperature decreases the mixture will want to release more heat to maintain the

 As pressure increases. As the pressure increases the mixture will favour the side that produces less particles (to maintain pressure), this means the forward reaction is favoured (see equation above, 3 -> 2).

 Any reactants are added. (SO2 or O2) The rate of the reaction will be faster if:

 Temperature increases.

 Pressure increases.

 A catalyst is used.

Conflict arises- rate increases with higher temperatures, but yield lowers with higher temperatures. A catalyst is used so that at low temepratures there is a higher yield, but the rate is still very fast.

 Vanadium Oxide (V2O5) is used as a catalyst- (economically viable and not poisoned- not neccessarily most efficient but best for these reasons.)

 Temperatures are maintined by cooling the reactant mixture as they pass from catalytic bed tray to another.

 Increasing the yield can also be done by using atmospheric air in the form of oxygen gas (this is also cheaper.) Also, as the ideal pressure is atmospheric, high pressure systems usually are not needed.

Step 3: ABSORPTION OF SULFUR TRIOXIDE:

Sulfur trioxide can react with water to form sulfuric acid.

However, this method is not used as there is a lot of heat involved- it creates a mist of the acid which is difficult to collect. Instead:

Sulfur trioxide is passed through concentrated sulfuric acid in an absorption tower. The reaction is regarded as occurring in two steps.

1. The SO3 gas dissolves almost completely in the H2SO4 to form a liquid (oleum.)

2. Oleum obtained from the abruption tower is then mixed with water to produce sulfuric acid. This is a dilution process.

Waste Management:

Environmental Management: (SO2 is a pollutant)

 Use/recycle SO2 from other industries. This solves a waste problem.

 Double absorption. Unreacted SO2 recycled from absorption tower back into converter to try to react. Increases conversion of SO2 to SO3 from 98% to 99.6%.

 Improved catalyst. More expensive catalyst helps conversion (less SO2 pollute) Waste Heat Produced used to Reduce Energy Costs:

 Since all the reactions are exothermic, this heat energy can be used to boil water to produce electricity for the plant itself.

 Energy can be sold to other industries, for use as heat or to generate electricity. Health and Safety:

Respiratory Irritants:

 Cause: SO2, SO3, H2SO4 mist.

 Precautions: Maximise conversion of SO2 to SO3 or Ventilation and air filtering.

Corrosiveness:

 Precautions: Storage, highly regulated handling and transport. Acid Rain:

 Causes: SO2, SO3

 Precautions: Limit emissions. Acid Spills:

 Precautions: Contain within earth, clay or sand. Neutralize with base (CaCO3 limestone or sodium carbonate.)

Dilution of H2SO4 with Water:

 Only add acid to water- stir. Reaction is exothermic, releases a lot of heat. General:

 Wear protective clothing: Safety glasses, lab coat, gloves.

 Use fume hood (for ventilation).

Area of Study 2: Supplying & Using Energy:

Chapter 23: Fossil Fuels:

Energy Sources Today:

How much energy do we use?

 Energy is measured in units called the joule, J.

 Larger amounts usually expressed as kilo jouled (kj, 1000 joules).

An individual in a modern society uses approximately 1000 MJ a day. This figure is about 100 times more than what the body actually requires- the bulk of our energy is used for transport, heating and domestic purposes. Globally, around 4 * 10 ^20 J is used per day.

Consumption for different sources:

Meeting our Energy Needs:

Dominant in our society are the fossil fuels:

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We are seeking alternative fuel sources, due to their negative effects on the environment:

 Release into atmosphere (SO2 causing acid rain, CO2 causing greenhouse effect.)

 In such a high demand, but non renewable.

Now we look to alternate sources (solar, wind, hydro, nuclear) to combat this problem.

Energy Converters:

 Anything that transforms energy from one type to another is an energy converter.

 For Example, a car engine converting chemical to thermal. Others are shown below

There are always losses in energy through transformations. This is usually as heat energy. Therefore energy converters may not always be 100% efficient. Efficiency is desirable. This would mean 2J of chemical energy is needed to create 1J electrical in a fuel cell. (50%)

Fossil Fuels:

 Sustainability is the need to meet short term needs, considering long term impact.

 Fossil fuels are unfortunately non renewable, they are used at a greater rate than what it is replenished. They are from finite sources of energy.

 Fossil fuel energy are essentially trapped solar energy, the stored chemical energy in the bonds of these fuels were transformed from photosynthesis, from energy from the sun.

 Formed from decaying vegetation. Plants lose their hydrogen and oxygen content, and so their carbon content increases over time. Becomes peat, then brown coal, then black coal. Also contains small amounts of nitrogen and sulphur.

 When coal is burnt, energy is used to vaporize water (reduces the amount of energy released upon combustion.) Thus mining black coal is worthwhile economically despite being buried further underground.

Coal is burnt in a coal fired power station.

 Chemical energy from the coal is released as thermal energy upon combustion.  Thermal energy produced is used to heat water, making a steam vapour.

 This thermal energy in this steam/vapour then turns a turbine, forming mechanical energy. This then generates a current, forming electrical energy.

 This process is about 35% efficient, with 50% of the lost energy occurring in the cooling tower where heat is lost in steam.

Overall, coal is not a clean burning fuel. A lot of ash, water vapour and sulphur dioxide are produced, as well as the primary product of Carbon Dioxide.

Oil and gas produce less pollutants, however coal has the largest reserves remaining in the world- (followed by oil, then natural gas.) It’s use is estimated to increase over time. It can be converted to oil via the process “flash pyrolysis”, where it is pulverized, heated to 600 degrees and then reacted with hydrogen.

 Advantages of Coal: Cheap (economically viable), and can be used as a baseload power source. (Can provide most of the energy demands for a large population.)

 Disadvantages of Coal: Releases greenhouse gases into the air, non renewable, inefficient (process if only 35% conversion from chemical to electrical energy.)

Coal: Energy Transformations Summary:

 Chemical energy in bonding of hydrocarbons  Coal is combusted

 Thermal Energy in temperature of steam from combustion  Steam passes turbine

 Electrical energy in the movement of electrons  Electricity used by public. Crude Oil:

 Also known as petrolium. It’s simply a mixture of alkanes. (the relative amounts of which vary.)

 Crude oil undergoes fractional distillation (seperation of several components of the mixture based on their boiling points. (See Unit 3)

 Oil is heated to about 400 degrees, enters the fractional tower, and the vapour begins to rise. The temperature in the tower decreases the higher up in the tower.

 Horizontal trays with bubble caps impere the rise of the vapour, forcing it to bubble through condensed liquid in the trays. These vapours then condense into trays containing condensed liquids with similar boiling points to their own. These fractions are then collected.

 The higher up the alkanes, the lower the boiling points- i.e lighter alkanes.

 Typically the fractions are: (Where Cx denotes the number of carbons)  Refinery gas (LPG, feedstock) C1 to C4,

 Gasoline (petrol) C5 to C6,

 Naphtha (cracking to petrol feedstock) C6 to C10,  Gas oil (diesel, cracking to petrol) C14 to C20 and  Bitumen (lubricating oil, fuel oil) larger than C20. Keep note that these will vary with location.

The heavier fractions may undergo several seperations:

 “Catalytic cracking” is the process in which heavier alkanes are broken into lighter alkanes, alkenes, and hydrogen. A Zeolite catalyst is used for this process.

 This is done because the lighter alkanes are more useful (but crude oil contains larger ones usually.)

 The main uses of crude oil products are fuels such as petrol, keosene, diesel and liquidified petroleum gas, and as raw materials for the manufacturing of products such as plastics and pharmaceuticals.

Natural Gas:

 Mainly refers to methane, but also small amounts of hydrocarbons such as ethane and propane.

 Popular fuel for heating and cooking purposes, may replace petrol and diesel as oil prices increase.

 In a gas-fired power station, thermal energy in hot gases from combustion expand air to spin the turbine, forming mechanical energy.

 The gas is simply burnt in air to produce CO2 and H2O, so less pollutants than coal.

 Natural gas provides for about 20% of the world’s energy needs. Natural Gas: Energy Transformations Summary:

 Chemical energy in bonds  combusted, creates a vapour

 Thermal energy of vapour  turns turbine

 Mechanical energy of turbine  creates a current

 Electrical energy due to current (electrons moving)  Electricity used by public.

Chapter 24: Alternative Energy Sources:

Nuclear Fission:

 A process where a fissile isotope (usually Uranium-235) is bombarded by a neutron and then splits up into two smaller nuclei (and releases a large amount of energy.)  This is called a nuclear reaction, and new elements are formed from the original.  Neutrons are also released from these reactions and go on to bombard other

nuclei, causing a chain reaction to occur (think of all that energy!)

 The Uranium -235 to start with is taken from Uranium ores. However, these ores contain 99.3% U238, and only 0.7% U235, so the ores must be refined for 3% U235.

 Uranium 238 can also be exploited by fast-breed reactors, decaying into Plutonium 239 to generate further energy.

 Uranium 235 nuclear reactions is the primary source of nucleur energy however, and occur in nuclear reactor power plants.

 In a nuclear reactor powerplant, the chain reaction occurs in a reaction vessel.

 Control rods and heavy water (deuterium oxide, 2H2O) moderate the rate of reaction in the vessel to a safe level. The reaction is shown:

 A small amount of the reactants nuclear energy in the bonds, are converted to kinetic energy in the products, and this then becomes thermal energy in the water.

 This then transfers to the boiler, creating steam which turns a turbine (mechanical energy) and this creates a current (electrical energy.)

 1kg Unranium produces the same amount of energy as 2500 tonnes of coal. Now we look to the environmental impact involving nuclear reactors:

 Produces no gaseous emissions, but products are radioactive and pose waste problems that a long term.

 Uranium and Plutonium may be recycled for further use.

 Non renewable source of energy (Uranium is finite.)

 Nuclear waste is stored for 5 years so short lived isotopes/waste can decay and cool down. Long lived isotopes must be stored safely for thousands of years. This is done by burying in sealed tanks or deep mines in geologically stable regions,

stockpiling in air conditioned warehouses, dumping in ocean trenches or sealing in special types of glass.

 The other concern with nuclear power is the risk of release of radioactive substances (by accident or sabotage, terroist attack).

 Fears of Uranium being stolen for weapons (unlikely).

 Advantages of Nuclear Energy: Relatively cheap, can be used as a baseload power source (power a whole population), high amount of energy per unit mass of fuel, no greenhouse gases released.

 Disadvantages of Nuclear Energy: Non renewable, waste is extremely tocix, potential terroist threats, nuclear meltdown potentially catastrophic.

Nuclear Energy: Energy Transformations Summary:

 Nuclear energy in the nuclear bonds of protons and neutrons  Fission & steam

 Thermal energy in the temperature of steam passes through turbine

 Mechanical energy in the motion of the turbine  Generates a current

 Electrical energy in the movement of electrons  Electricity used by public Nucleaur Fusion:

 The opposite to fission, two or more nuclei are smashed together to form a larger one. This is the source of energy in the sun (when hydrogen fuses with helium).  Fusion of one kilogram of fuel yields more than fission, may reach the same energy

as 10000 tonnes of coal.

 However, great difficulty in sustaining a temperature in a reactor for fusion to occur. Temperature of 100,000,000 degrees needed as well as a place to contain.

Bioethanol, Biodiesel and Biogas:

 The fermentation of glucose (with a catalyst) from plant crops produces bioethanol, which can be blended with fuel to lower the demand and use of petroleum (less pollutants etc.)

 Biodiesel is derived from vegetable oils and animal fats, hydrolysed to fatty acids and esterified.

 Waste plant and animal material can be converted into syngas (carbon monoxide and hydrogen) or biogas (carbon dioxide and methane) which can generate heat/electricity.

Solar:

 Solar energy can be used directly.

 Solar efficient building designs reduce heating, cooling and lighting costs.  Solar water heating systems provide hot water, for homes and for pools.

 However, solar cells are very expensive, and inability to be used efficiently when in the dark- either at night or in cloudy weather. There is also a need for large areas of land due to relatively poor energy conversion efficiency.

 Solar energy is renewable. Utilizes energy from the sun (which always provides energy.) Solar energy is stored for use as electricity.

Solar: Energy Transformations for STORING solar energy:

 Radiant energy in light waves from sun  Current made in solar panel.

 Electrical energy (movement of electrons) in solar panel  Recharges batteries

 Chemical energy of the reactants in the battery made.

The chemical energy is stored, now, more transformations occur for it’s use (below). Solar: Energy Transformations for the USE of solar energy:

 Chemical energy of the reactants in batteries  Discharge of battery

 Electrical energy of movement of electrons Electricity for the home

 Advantages of solar: Renewable source of energy, no greenhouse gas emissions.

 Disadvantages of solar: Expensive, inefficient (only about 1% conversion), cant be used as a baseload power source (for an entire population.)

Hydroelectricity:

 Firstly, Solar energy evaporates this water transferring it to higher altitudes. The water thus obtains more Gravitational potential energy.

 Then, as the water falls, it is converted into mechanical energy as it spins a turbine, which generates a current for electrical energy.

 Hydroelectricity supplies about 25% of the world’s and 10% of Australia’s electricity.

 Developement of hydroelectric stations are restricted, due to the limited number of suitable sites and concern about it’s environmental impact. This is because storage dams are often requires to establish a good flow of the water at all seasons, but lead to further environmental concerns (loss of natural habitat.)  Hydroelectricity is renewable.

Hydroelectricity: Energy Transformations Summary:

 Gravitational Potential energy in high altitude of water Water falls, turns turbine

 Mechanical energy of turbines movement  Generates a current

 Electrical Energy from movement of electrons (Current)  Electricity used by public

 Advantages of Hydroelectricity: No greenhouse gases, no toxic waste, renewable.

 Disadvantages of Hydroelectricity: Building dams results in destruction of natural habitat.

Wind Power:

 Can be used to generate electricity through wind turbines.

 Average wind speeds exceeding 5ms are suitable, but the turbines are only as reliable as the wind.

 The turbines are highly visible and loud, a number of them are needed to match the energy production of a coal-fired plant. (Wind turbines make less energy)

Tidal Power:

 Uses the movement of water caused by the moon to generate electicity.

 Sites are limited, as a large difference in water height between tides are required.

Geothermal Power:

 In volcanic regions, heat from underground rocks can reach underground water.

 This can cause steam to rise to the surface, which can be harnessed to spin turbines to create electricity.

 This is the principal of geothermal energy, however Geothermal power plants are again restricted to suitable sites- but are reliable.

Chapter 25: Energy from Chemical Reactions:

Thermochemical Equations:

 Knowledge of thermochemical equations (covered in chapter 1) is needed for this section.

 In addition to what was covered in chapter 1, is it important to know that for example, the above reaction means 2 moles of C8H18 produces 10108kjmol-, 1 mol will produce half of this: 5054kjmol-.

The Connection between Energy and Temperature Change:

 The specific heat capacity of a substance is how much energy it will take to increase the temperature of the substance by 1 degree Celsius.

 The higher the heat capacity, the more the energy will store heat as thermal energy in it’s bonds, meaning it will heat up at a much slower rate.

 Measuring the energy when a substance is heated:

Where E is energy, m is mass, c is specific heat capacity and T is temperature. Measuring the Heat Released during a Reaction:

Bomb and Solution Calorimetry:

 Enthalpy is measured directly from an instrument known as a calorimeter.

Bomb Calorimeter:

 Used for reactions involving gaseous reactants or products. (Cant be done underwater.)

 The vessel is made to withstand high pressures during reaction.

Note what each component is used for:

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Solution Calorimeter:

 Used for aqueous solutions.

 Insulated to reduce heat loss or gain, to or from the environment. In reference to both Calorimeters:

 When a reaction takes place in a calorimeter, the heat chance causes a rise or fall in the temperature of the contents of the calorimeter. (Which is measured.)

 Before the Calorimeter can be used however, we must determine how much energy is required to change the temperature within the calorimeter by 1 degree.

 The calibration factor can be found by:

 Supplying a known current (I) (and a known voltage(V)) and measuring the time (t) in which this occurs. Energy (E) is calculated from this as: (in joules)

This is the total energy change due to the supplied current.

 Also, measuring the temperature of the calorimeter, then heating the calorimeter with a known amount of thermal energy and measuring the temperature again.

 Then, calculating the Calibration Factor (CF): (In Joules per degrees Celcius)

We now get the energy change per 1 degree Celcius. (Calibration Factor)

With the calibration factor, we can now measure the temperature change of the actual reaction. It is used to determine what energy is responsible for this temperature change.

We can get the energy released during the ACTUAL reaction by reusing the Calibration factor (CF) formula, only with the temperature change of the actual reaction- not due to heating the calorimeter. (CF * of reaction = E of reaction)

Heat of Combustion:

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chemical formula). There the energy released when they burn is measured per gram or per liter rather than per mol.

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Heat of Combustion for mixed solids, mixed liquids, and pure substances:

For the above table, volumes of gases are measured at SLC with H2O and CO2 as products.

Chapter 26: Electricity from Chemical Reactions:

Remember the Daniel Cell from Unit 2? This requires a Zinc anode(in ZnNO3 solution) and Copper cathode (In Copper (II) solution). Refer to the diagram below:

 A current passes through the circuit to the globe. This part of the cell is the external circuit. The globe converts electrical energy to light and heat energy.

 The current flows because of the chemical reaction taking place in the cell (of which there is little indication of happening initially.)

In the above example, if left long enough we observe the following:

 The zinc metal corrodes.

 The copper metal becomes covered with a furry brown-black deposit.

 The blue copper (II) sulfate solution loses some of it’s colour. These changes provide evidence of a chemical reaction.

 There is also a current flowing from the zinc electrode, through the wire, to the copper electrode.

 Current flows only if a salt bridge is present. (Made of some type of salt).

These findings help us decide what is happening inside the cell.

 The reaction in the cell is redox, electrons are being produced and consumed.  The zinc electrode is eaten away, forming zinc ions in solution.

 The oxidation of the zinc metal releases electrons, these flow through the wire to the copper electrode.

 Electrons are accepted by copper ions in the solution when ions collide with the copper electrode.

 Copper atoms are insoluble and deposit on the electrode producing brown-black coating.

Purpose of the Salt Bridge:

 The salt bridge is an essential component of the cell, It allows charges to balance. Without it, the cell will be polarized and electrons will accumulate on one half of

the cell. This would prevent any further current passing through, however the salt bridge stops this.

 The salt bridge contains ions that can migrate to either half cell so that a buildup of electrons (and charge) is prevented. The cation of the salt goes to the Cathode, and the anion goes to the anode.

Below is a summary of the processes occurring in the cell:

 The overall reaction in the cell is found by adding the two half equations that occur in each cell.

Half Cells:

In a half cell, an electrode is in a solution. The species present in each half cell (the elctrode and solution) form a conjugate redox pair.

 Generally, the metal in the conjugate redox pair is used as the electrode. The other is used in the solution.

 If there isnt a metal, graphite or platinum is used as the electrode.

 When a gas is involved as one of the conjugates, a special “gas electrode” is used.

 Sometimes spectator ions are present (not involved in the reaction.)  The oxidation reaction is always at the anode. (Copper is the oxidant)