Electroplating: Videos & Practice Problems

Electroplating involves using an electrical current to deposit metal cations onto an electrode. The electrical current, measured in amperes (A), is the flow of electrons between electrodes at a rate of one coulomb per second. Electrochemical stoichiometry calculations in electrochemical cells require converting the given time to seconds, using the current to find the charge, and then applying Faraday's constant (96,485 coulombs per mole of electrons) to determine moles of electrons. A mole-to-electron comparison using balanced chemical equation coefficients allows conversion to moles of the plated element. From there, moles can be converted into various units such as grams or molecules. When starting with mass, the process involves going from grams to moles using atomic mass, then to moles of electrons, charge, and finally, using the current to find the time for the electroplating process.

Electroplating Process

Electroplating is a process that utilizes electrical current to deposit metal cations onto a metal electrode. In this context, electrical current refers to the flow of electrons moving from the anode to the cathode within a closed circuit. The concept of rate is crucial here, as it represents the amount of charge transferred over a specific time period. The standard unit of electrical current is the ampere (A), where 1 ampere is defined as the flow of 1 coulomb of charge per second. This relationship is essential for performing calculations related to electroplating, as it allows for the conversion of current into charge over time, facilitating a deeper understanding of the electroplating process.

Electroplating Example

To determine the electrical current produced when a charge of \(4.14 \times 10^3\) coulombs passes through a wire for 15 minutes, we start by recalling that electrical current (I) is measured in amperes (amps), where 1 ampere is defined as 1 coulomb per second.

First, we need to convert the time from minutes to seconds. Since 1 minute equals 60 seconds, we can calculate the total time in seconds for 15 minutes:

\[15 \text{ minutes} \times 60 \text{ seconds/minute} = 900 \text{ seconds}\]

Now that we have the charge in coulombs and the time in seconds, we can calculate the current using the formula:

\[I = \frac{Q}{t}\]

where \(I\) is the current in amperes, \(Q\) is the charge in coulombs, and \(t\) is the time in seconds. Plugging in the values:

\[I = \frac{4.14 \times 10^3 \text{ coulombs}}{900 \text{ seconds}} \approx 4.6 \text{ amps}\]

Thus, the electrical current produced is approximately 4.6 amps.

Electrochemical Stoichiometric Chart (Time)

Electrochemical stoichiometry involves calculations related to electrochemical cells, particularly focusing on the relationship between current, charge, and time. When given a time duration in units such as hours, days, or seconds, it is essential to convert all time measurements into seconds for consistency. This conversion allows for the effective use of current, which is measured in amperes (A) and defined as coulombs per second (C/s).

To find the charge (Q), the formula used is:

Q = I \cdot t

where I is the current in amperes and t is the time in seconds. The charge is expressed in coulombs (C). Once the charge is determined, Faraday's constant, which is approximately 96,045 C/mol, can be applied to convert charge into moles of electrons (n_e):

n_e = \frac{Q}{F}

where F is Faraday's constant. This step is crucial as it allows the transition from charge to the number of moles of electrons involved in the reaction.

Next, to relate the moles of electrons to the moles of a specific substance, a mole-to-electron comparison is necessary. This involves using the coefficients from the balanced chemical equation. For instance, if the equation indicates that 1 mole of a particular element corresponds to the transfer of 3 moles of electrons, the relationship can be expressed as:

1 mole of element = 3 moles of electrons

With the moles of the substance known, various conversions can be performed to express the quantity in different units, such as grams, molecules, ions, or kilograms. This systematic approach is essential when starting with time units, ensuring accurate calculations in electrochemical stoichiometry.

Electrochemical Stoichiometric Chart (Time) Example

Gold can be plated from a solution containing gold(III) ions through a specific electrochemical reaction. The half-reaction indicates that for every mole of gold(III) ion, three moles of electrons are required to produce one mole of solid gold. To determine the mass of gold plated by a current of 6.8 amps over a duration of 41 minutes, we first need to convert the time from minutes to seconds. Since there are 60 seconds in a minute, 41 minutes is equivalent to:

41 minutes × 60 seconds/minute = 2460 seconds

Next, we relate the current (in amps) to charge (in coulombs). One ampere is defined as one coulomb per second, so a current of 6.8 amps translates to:

6.8 amps = 6.8 coulombs/second

To find the total charge (Q) over the 2460 seconds, we multiply the current by the time:

Q = 6.8 coulombs/second × 2460 seconds = 16728 coulombs

Now, we convert the total charge to moles of electrons using Faraday's constant, which states that 1 mole of electrons corresponds to 96,485 coulombs. Thus, the number of moles of electrons (n) can be calculated as follows:

n = Q / Faraday's constant = 16728 coulombs / 96485 coulombs/mole ≈ 0.173 moles of electrons

According to the stoichiometry of the half-reaction, 3 moles of electrons are needed to produce 1 mole of gold. Therefore, the moles of gold produced can be calculated by dividing the moles of electrons by 3:

moles of gold = 0.173 moles of electrons / 3 ≈ 0.0577 moles of gold

To find the mass of gold, we use the atomic mass of gold, which is approximately 196.967 grams/mole. The mass (m) can be calculated as:

m = moles of gold × atomic mass of gold = 0.0577 moles × 196.967 grams/mole ≈ 11.373 grams

Considering significant figures, the final mass of gold plated is rounded to:

11 grams

This calculation illustrates the relationship between current, time, and the electrochemical plating of gold, emphasizing the importance of stoichiometry and Faraday's laws in electrochemistry.

Electrochemical Stoichiometric Chart (Mass)

When dealing with electrochemical reactions, particularly when the initial mass for a half-reaction is provided, the mass version of the stoichiometric chart becomes a valuable tool for determining the time required for the reaction. This approach is particularly useful when the given amount is expressed in grams.

To begin, you convert the grams of the substance into moles using the atomic mass of the element. This conversion is essential as it allows for a more straightforward comparison between the amount of substance and the number of electrons involved in the reaction.

Next, to find the moles of electrons transferred, a crucial step involves using the coefficients from the balanced chemical equation. This step is often referred to as a mole-to-electrons comparison, where you relate the moles of the given element to the moles of electrons involved in the half-reaction.

Once the moles of electrons are determined, Faraday's constant (approximately 96485 C/mol) can be employed to calculate the total charge associated with the reaction. The formula used here is:

Charge (Q) = moles of electrons × Faraday's constant

With the charge known, you can then utilize the current (I) in the system to find the time (t) required for the reaction to occur. The relationship between charge, current, and time is given by the equation:

Q = I × t

From this, time can be calculated as:

t = Q / I

This systematic approach allows for the determination of time based on the initial mass of the reactant, making it a powerful method in electrochemistry.

Electrochemical Stoichiometric Chart (Mass) Example

To determine the time required to plate out 42.1 grams of nickel using a current of 3.08 amps, we start by converting grams of nickel to moles. The molar mass of nickel (Ni) is approximately 58.693 g/mol. Thus, the conversion from grams to moles is done as follows:\[\text{Moles of Ni} = \frac{42.1 \text{ g}}{58.693 \text{ g/mol}} \approx 0.717 \text{ moles of Ni}\]Next, we need to find the moles of electrons involved in the reaction. The balanced equation indicates that 1 mole of nickel solid requires 2 moles of electrons. Therefore, the moles of electrons can be calculated as:\[\text{Moles of electrons} = 0.717 \text{ moles of Ni} \times 2 \approx 1.434 \text{ moles of electrons}\]Using Faraday's constant, which is approximately 96,485 coulombs per mole of electrons, we can find the total charge (Q) required for the reaction:\[Q = \text{Moles of electrons} \times \text{Faraday's constant} = 1.434 \text{ moles} \times 96,485 \text{ C/mol} \approx 138,000 \text{ C}\]Now, we can relate the charge to current (I) to find the time (t) in seconds. The relationship is given by the formula:\[Q = I \times t\]Rearranging this gives:\[t = \frac{Q}{I} = \frac{138,000 \text{ C}}{3.08 \text{ A}} \approx 44,835 \text{ seconds}\]To convert seconds into hours, we use the following conversions:\[\text{Minutes} = \frac{44,835 \text{ seconds}}{60} \approx 747.25 \text{ minutes}\]\[\text{Hours} = \frac{747.25 \text{ minutes}}{60} \approx 12.45 \text{ hours}\]Rounding to three significant figures, the final time required to plate out 42.1 grams of nickel is approximately:\[\text{Time} \approx 12.5 \text{ hours}\] This calculation illustrates the process of converting mass to moles, determining the charge required, and using current to find the time needed for electroplating nickel.

Cu²⁺ is reduced to Cu(s) at an electrode. If a current of 1.25 A is passed for 72 hours, what mass of copper is deposited at the electrode? (MW of Cu: 63.55 g/mol)

• A. 91.5 g
• B. 55.8 g
• C. 83.1 g
• D. 110 g

Answer: (D)

A solution of Mn⁺⁵ is used to plate out Mn in an electrochemical cell. If a total of 1.13 g of Mn is plated out in a total time of 1600 seconds, what was the electrical current used? (MW of Mn is 54.94 g/mol)

• A. 3.5 A
• B. 6.2 A
• C. 1,8 A
• D. 7.7 A

Answer: (B)

Here’s what students ask on this topic:

What is electroplating?

Electroplating is a process used to coat the surface of an object with a thin layer of metal through an electrochemical reaction. Essentially, it involves submerging the object you want to plate, known as the substrate, into a solution containing the plating metal, such as gold, silver, or copper. This solution is known as an electrolyte.

The object to be plated is connected to the negative terminal of a power supply, becoming the cathode. A piece of the plating metal is connected to the positive terminal, becoming the anode. When electricity is applied, the metal ions from the electrolyte solution are attracted to the negatively charged substrate. They deposit onto the surface, forming a new metal coating.

Electroplating is used for various purposes, including corrosion resistance, aesthetic enhancement, increasing surface hardness, and improving electrical conductivity. It's a common technique in manufacturing, jewelry making, and electronics.

How does electroplating work?

Electroplating is a process that uses electrical current to coat an electrically conductive object with a thin layer of metal. Here's a simplified explanation of how it works:

1. Preparation: The object to be plated, known as the substrate, is thoroughly cleaned to remove any dirt, grease, or oxides that might interfere with the plating process.
2. Setup: The substrate is then placed into an electrolyte solution, which contains the metal ions of the metal you want to plate with. The substrate acts as the cathode (negative electrode), and an anode (positive electrode) made of the plating metal is also placed in the solution.
3. Electroplating: When an electrical current is applied, the metal ions in the solution are attracted to the negatively charged substrate. They gain electrons (reduction) and deposit onto the substrate's surface as a thin layer.
4. Result: The thickness of the metal coating can be controlled by adjusting the time the substrate is left in the solution and the intensity of the electrical current.

Through electroplating, objects can be coated for various purposes, including corrosion resistance, aesthetic appeal, increased strength, or to improve electrical conductivity.

What is electroplating?

Electroplating is a process that uses electrical current to coat an electrically conductive object with a thin layer of metal. This is achieved by immersing the object to be plated, known as the substrate, along with a piece of the metal to be deposited, into an electrolyte solution. The substrate acts as the cathode (negative electrode), and the metal acts as the anode (positive electrode).

When an electrical current is applied, positively charged metal ions in the solution are attracted to the negatively charged substrate. These ions undergo a reduction reaction, changing from their ionic state to a solid state, and deposit onto the surface of the substrate, creating a new layer of metal. This process can be used for various purposes, such as improving corrosion resistance, enhancing appearance, increasing thickness, or to impart desired surface properties to the substrate.

What is electroplating used for?

Electroplating is a process used to coat the surface of a metal or other conductive material with a thin layer of another metal. This is achieved by using an electrical current to reduce dissolved metal cations so that they form a coherent metal coating on an electrode. The primary purposes of electroplating are to improve a material's properties, such as corrosion resistance, abrasion resistance, aesthetic qualities, and electrical conductivity. For example, chrome plating is often used for its shiny, reflective finish and rust resistance, while gold plating is used in electronics for its excellent conductivity and resistance to corrosion. Additionally, electroplating can be used to build up thickness on undersized parts, create uniformity in surface texture, or to coat objects for decorative purposes.