Electrical Conductivity Measurement Discussion Examples | Relationship Between Electrolyte Concentration and Conductivity

What Is Electrical Conductivity?

Electrical conductivity is a quantity that represents how easily a solution conducts electricity.
Water itself conducts very little electricity, but when electrolytes such as sodium chloride or hydrochloric acid dissolve, ions are produced in the solution.
These ions move in an electric field, allowing current to flow and increasing electrical conductivity.

Electrical conductivity increases as the number of ions in the solution increases and as the ions can move more rapidly.
Therefore, concentration, degree of ionization, type of ion, temperature, solvent viscosity, and other factors affect conductivity.

Main Items to Include in the Results

In the results of electrical conductivity measurements, organize the type of sample, concentration, measurement temperature, electrical conductivity, molar conductivity, cell constant, and whether calibration was performed using a standard solution.
When measurements are performed at different concentrations, plotting concentration against electrical conductivity makes the discussion easier to write.

Main Items to Include in the Results

• Type of electrolyte measured
• Solution concentration
• Measurement temperature
• Electrical conductivity
• Cell constant
• Whether calibration with a standard solution was performed
• Mean value when multiple measurements were performed
• Molar conductivity
• Concentration-electrical conductivity graph
• Concentration-molar conductivity graph
• Comparison between strong and weak electrolytes
• Comparison with literature values
• Error rate
• Condition of electrode cleaning and presence or absence of bubbles

Reference Experimental Values and Analysis Examples for Electrical Conductivity Measurement

Here, reference experimental values are organized for measuring electrical conductivity while changing the concentration of an aqueous electrolyte solution and determining conductance, cell constant, electrical conductivity, and molar conductivity.
Differences between strong and weak electrolytes, concentration dependence, temperature effects, and measurement errors are presented in a form that is easy to discuss in a report.

Aqueous electrolyte solutions conduct electricity because ions in the solution move.
In general, as the electrolyte concentration increases, the number of ions in the solution increases, so electrical conductivity becomes larger.
However, as the concentration increases, interactions between ions become stronger, and molar conductivity may decrease.

Reference Experimental Conditions

Relationship Between Conductance and Electrical Conductivity

Conductance G is the reciprocal of electrical resistance R.

G = 1 / R

In actual measurements, the distance between the electrodes and the electrode area affect the result, so the cell constant K is used to determine the electrical conductivity κ.

κ = K × G

Here, κ is electrical conductivity, K is the cell constant, and G is conductance.
For example, if the cell constant is 1.00 cm⁻¹ and the conductance is 1.25 mS,

κ = 1.00 × 1.25 = 1.25 mS/cm

If the cell constant is not set correctly, all measured electrical conductivity values will be shifted.

Example of Determining the Cell Constant

The following is an example of determining the cell constant using the electrical conductivity of a standard KCl solution.

The cell constant is determined by dividing the electrical conductivity of the standard solution by the measured conductance.

K = κ / G

If the standard electrical conductivity is 1.41 mS/cm and the measured conductance is 1.40 mS,

K = 1.41 ÷ 1.40 = 1.01 cm⁻¹

Example Measurements of NaCl Concentration and Electrical Conductivity

The following is a reference example in which the concentration of an NaCl solution was changed and the conductance and electrical conductivity were measured.
The cell constant is assumed to be 1.00 cm⁻¹.

NaCl is a strong electrolyte and is almost completely ionized in water, so electrical conductivity increases as the concentration becomes higher.
However, at high concentrations, interactions between ions become stronger, and the relationship may gradually deviate from simple proportionality.

Example Calculation of Molar Conductivity

Molar conductivity Λ_m is determined by dividing electrical conductivity by molar concentration.

Λ_m = κ / c

When κ is expressed in S cm⁻¹ and c in mol cm⁻³, the unit of Λ_m is S cm²/mol.

For example, if the electrical conductivity of 0.010 mol/L NaCl is 1.15 mS/cm,

κ = 1.15 mS/cm = 0.00115 S/cm

c = 0.010 mol/L = 1.0 × 10⁻⁵ mol/cm³

Λ_m = 0.00115 ÷ 1.0 × 10⁻⁵ = 115 S cm²/mol

Concentration Dependence of Molar Conductivity of NaCl Solutions

Electrical conductivity increases with concentration, but molar conductivity decreases as the concentration becomes higher.
This is considered to occur because interactions between ions become stronger as the concentration increases, making it more difficult for ions to move freely.

Comparison of Strong and Weak Electrolytes

Strong electrolytes are almost completely ionized in water, whereas weak electrolytes are only partially ionized.
Therefore, differences in electrical conductivity and molar conductivity appear even at the same concentration.

Even at the same concentration of 0.010 mol/L, HCl shows very high electrical conductivity.
This is because H⁺ has high mobility.
In contrast, acetic acid is a weak electrolyte with a low degree of ionization, so its electrical conductivity is low even at the same concentration.

Changes in Molar Conductivity of a Weak Electrolyte With Dilution

In weak electrolytes, dilution may greatly increase molar conductivity because the degree of ionization increases.

For a weak electrolyte such as acetic acid, lowering the concentration shifts the ionization equilibrium toward ionization, increasing the number of ions per mole and therefore increasing the molar conductivity.

Changes in Electrical Conductivity With Temperature

Electrical conductivity is affected by temperature.
As temperature increases, the viscosity of the solution decreases and ions can move more easily, so electrical conductivity generally increases.

In electrical conductivity measurements, the temperature must be kept constant or temperature correction must be performed.
If data measured at different temperatures are compared directly, the effect of temperature differences may be mistaken for the effect of concentration differences.

Simple Example of Temperature Correction

The following is an example of correcting a value measured at 30°C to 25°C, assuming a temperature coefficient of electrical conductivity of 2.0%/°C.

κ₂₅ = κ_t ÷ {1 + α(t − 25)}

Here, if κ_t = 1.27 mS/cm, t = 30°C, and α = 0.020 /°C,

κ₂₅ = 1.27 ÷ {1 + 0.020 × (30 − 25)}

κ₂₅ = 1.27 ÷ 1.10 = 1.15 mS/cm

With this correction, the value measured at 30°C can be compared more easily as a value referenced to 25°C.

Comparison of Pure Water, Tap Water, and Electrolyte Solutions

Sucrose does not produce ions when dissolved in water, so even at the same concentration of 0.010 mol/L, its electrical conductivity does not become large as it does for NaCl.

Example of Variation in Measured Values

In electrical conductivity measurements, measured values may vary because of the condition of electrode cleaning, temperature, bubbles, and the mixing state of the solution.

If only the fourth measurement is high, possible causes include a bubble adhering to the electrode, residue from the previous sample, or an increase in temperature.

Effect of Insufficient Electrode Cleaning

In low-concentration samples and pure water, even slight contamination greatly affects the measured value.
It is important to wash the electrode after each measurement and record the value only after it has stabilized.

Example Calibration Curve of Concentration and Electrical Conductivity

In a low-concentration range, NaCl concentration and electrical conductivity show an approximately linear relationship.
This relationship can be used to estimate the concentration of an unknown sample.

Suppose approximation over the range of 0.001–0.010 mol/L gives the following equation.

κ = 115 × c + 0.010

Here, κ is in mS/cm and c is in mol/L.
If the electrical conductivity of an unknown sample is 0.585 mS/cm,

c = (0.585 − 0.010) ÷ 115 = 0.0050 mol/L

Therefore, the NaCl concentration of the unknown sample can be estimated to be approximately 0.0050 mol/L.

Example Estimation of Unknown Samples

If an unknown sample lies outside the calibration-curve range, as with Unknown Sample D, simply extrapolating to determine the concentration may produce a large error.
The sample should be diluted so that it falls within the calibration-curve range and then recalculated.

Example of How to Write the Results

The concentration of NaCl aqueous solution was varied over the range of 0.001–0.100 mol/L, and the electrical conductivity was measured.
As a result, electrical conductivity increased as the concentration increased.
This is because NaCl is almost completely ionized into Na⁺ and Cl⁻ in water, and the number of ions in the solution increased with increasing concentration.

On the other hand, molar conductivity decreased as the concentration increased.
It was 125 S cm²/mol at 0.001 mol/L but decreased to 102 S cm²/mol at 0.100 mol/L.
This is considered to have occurred because interactions between ions become stronger at high concentrations and interfere with the movement of individual ions.

When 0.010 mol/L NaCl, KCl, HCl, and acetic acid were compared, HCl showed the highest electrical conductivity, while acetic acid showed a low value.
HCl is considered to have shown high electrical conductivity because H⁺ has high mobility.
In contrast, acetic acid is a weak electrolyte and is only partially ionized in water, so the number of ions was small even at the same concentration, resulting in low electrical conductivity.

Points for Connecting the Results to the Discussion

In a discussion of electrical conductivity measurements, it is important to explain not only the relationship between concentration and electrical conductivity but also the number of ions, ionic mobility, degree of ionization, temperature, cell constant, and measurement errors in relation to one another.

• Can you explain the difference between conductance and electrical conductivity?
• Can you calculate electrical conductivity using the cell constant?
• Can you explain why electrical conductivity increases as concentration increases in relation to the increase in the number of ions?
• Can you discuss why molar conductivity decreases at high concentrations in relation to interactions between ions?
• Can you explain the difference between strong and weak electrolytes from the degree of ionization and electrical conductivity?
• Can you relate the high mobility of H⁺ and OH⁻ to the high electrical conductivity of substances such as HCl?
• Can you discuss why electrical conductivity increases as temperature rises in relation to viscosity and ionic mobility?
• Can you explain the differences among pure water, tap water, and nonelectrolyte solutions from the amount of ions present?
• Can you explain insufficient electrode cleaning, bubbles, unstable temperature, and deviation of the cell constant as sources of error?
• When estimating the concentration of an unknown sample, can you evaluate it within the calibration-curve range?

Example Discussion

In this experiment, the concentration of aqueous electrolyte solutions was varied and the electrical conductivity was measured.
In the NaCl solution, as the concentration increased from 0.001 mol/L to 0.100 mol/L, the electrical conductivity increased from 0.125 mS/cm to 10.2 mS/cm.
This is because NaCl is almost completely ionized in water and the number of Na⁺ and Cl⁻ ions increased as the concentration increased.

On the other hand, molar conductivity decreased as the concentration increased.
Electrical conductivity indicates how easily the solution as a whole conducts electricity, whereas molar conductivity represents the contribution to conductivity per mole of electrolyte.
At high concentrations, interactions between ions become stronger and ion movement is hindered, so the conductivity per mole is considered to have decreased.

When compared at the same concentration of 0.010 mol/L, HCl showed higher electrical conductivity than NaCl.
This is because H⁺ has very high mobility and a high ability to carry electric current.
In addition, acetic acid showed much lower electrical conductivity than NaCl.
Acetic acid is a weak electrolyte and is only partially ionized in aqueous solution, so the number of ions in the solution was small and the electrical conductivity was low.

The effect of temperature is also important.
The electrical conductivity of 0.010 mol/L NaCl solution was 0.93 mS/cm at 15°C but 1.39 mS/cm at 35°C.
As the temperature increases, the viscosity of the solution decreases and ions can move more easily, so electrical conductivity is considered to increase.
Therefore, when comparing values measured at different temperatures, the temperature must be kept constant or temperature correction such as conversion to 25°C must be performed.

Possible sources of error include insufficient electrode cleaning, residue from the previous sample, bubbles on the electrode surface, unstable temperature, and deviation of the cell constant.
Particularly in low-concentration solutions and pure water, even slight ionic contamination can greatly change electrical conductivity.
It is important to thoroughly wash the electrode with pure water after each measurement and record the value only after it stabilizes.

Summary

In electrical conductivity measurements, conductance can be corrected using the cell constant to determine electrical conductivity, and dividing it by concentration gives molar conductivity.
Electrical conductivity increases with electrolyte concentration, whereas molar conductivity may decrease at high concentrations.

This reference example covered the cell constant, NaCl concentration and electrical conductivity, molar conductivity, comparison of strong and weak electrolytes, the dilution effect of acetic acid, temperature dependence, temperature correction, comparison of pure water, tap water, and nonelectrolytes, measurement errors, and estimation of unknown-sample concentrations.
In a report, it is useful to discuss the results in relation to the number of ions, ionic mobility, degree of ionization, temperature, and instrument conditions.

Relationship Between Electrolyte Concentration and Electrical Conductivity

As the electrolyte concentration increases, the number of ions present in the solution increases, so electrical conductivity generally becomes larger.
In a low-concentration range, electrolyte concentration and electrical conductivity may show an approximately proportional relationship.
This is because when the concentration doubles, the number of ions carrying electric charge also approximately doubles.

However, at high concentrations, interactions between ions become larger and ion movement may be hindered.
Therefore, even when the concentration increases, electrical conductivity may not increase in simple proportion.

Discussion of a Concentration-Electrical Conductivity Graph

A graph with concentration on the horizontal axis and electrical conductivity on the vertical axis makes it possible to examine the effect of concentration on conductivity.
In the low-concentration range of a strong electrolyte, conductivity may increase almost linearly as the concentration increases.
On the other hand, in the high-concentration range, interactions between ions and changes in solution viscosity may cause deviation from linearity.

If some points on the graph deviate greatly, possible causes include errors in concentration preparation, electrode contamination, temperature changes, bubbles, and reading the value before it had stabilized.

Discussion of Strong Electrolytes

A strong electrolyte is an electrolyte that is almost completely ionized in water.
Representative examples include sodium chloride, hydrochloric acid, and sodium hydroxide.
Because strong electrolytes produce many ions in solution, they tend to have higher electrical conductivity than weak electrolytes at the same concentration.

However, even for strong electrolytes, interactions between ions can no longer be ignored at high concentrations.
Therefore, while conductivity increases as the concentration rises, molar conductivity may show a decreasing trend.

Discussion of Weak Electrolytes

A weak electrolyte is an electrolyte that is only partially ionized in water.
Representative examples include acetic acid and aqueous ammonia.
Because weak electrolytes contain fewer ions in solution than strong electrolytes even at the same concentration, their electrical conductivity tends to be smaller.

In weak electrolytes, the degree of ionization may increase upon dilution.
Therefore, lowering the concentration may greatly increase the conductivity per unit amount of substance, that is, the molar conductivity.

Comparison Between Strong and Weak Electrolytes

When strong and weak electrolytes are compared, strong electrolytes often have higher electrical conductivity even at the same concentration.
This is because strong electrolytes are almost completely ionized and produce many ions, whereas weak electrolytes are only partially ionized.

However, ionic mobility differs depending on the type of ion, so differences in conductivity cannot be explained solely by the degree of ionization.
For example, hydrogen ions and hydroxide ions show particularly high mobility, affecting the conductivity of acids and bases.

What Is Molar Conductivity?

Molar conductivity is a quantity representing the extent to which 1 mol of electrolyte contributes to electrical conduction.
It is often determined by dividing electrical conductivity by concentration and is used when comparing solutions with different concentrations.

Molar conductivity Λ_m = κ / c

Here, κ is electrical conductivity and c is electrolyte concentration.
Because the unit depends on the laboratory manual and the system of units used, the units of concentration and conductivity must be made consistent during calculations.

Why Molar Conductivity Decreases as Concentration Increases

For strong electrolytes, electrical conductivity itself increases as the concentration increases, but molar conductivity may decrease.
This is because the distance between ions becomes smaller as the concentration increases and electrical interactions between ions become stronger.
When ion movement is hindered, the conducting ability per mole decreases.

In contrast, dilution increases the distance between ions and reduces their influence on one another.
Therefore, as the solution approaches infinite dilution, molar conductivity becomes larger.

Changes in Conductivity Caused by Dilution

When a solution is diluted, the number of ions per unit volume decreases, so electrical conductivity κ generally becomes smaller.
However, molar conductivity may increase upon dilution.
This is because interactions between ions weaken and ions can move more easily.

In weak electrolytes, dilution increases the degree of ionization, so an even greater increase in molar conductivity may be observed.
Therefore, electrical conductivity and molar conductivity must be considered separately.

Effect of Ionic Mobility

Electrical conductivity is affected not only by the number of ions but also by how rapidly the ions can move.
The ease with which an ion moves is called ionic mobility.
Mobility varies depending on ion size, degree of hydration, charge, solvent viscosity, and other factors.

Even at the same concentration, an electrolyte containing ions with high mobility may show higher electrical conductivity.
Hydrogen ions and hydroxide ions in particular show high mobility in water, so they are important when considering the conductivity of acids and bases.

Effect of Temperature on Electrical Conductivity

Electrical conductivity changes greatly with temperature.
In general, as temperature rises, the viscosity of the solution decreases and ions move more rapidly.
Therefore, in many aqueous electrolyte solutions, electrical conductivity increases as the temperature rises.

Temperature control is extremely important in electrical conductivity measurements.
If the measurement temperature differs, the conductivity changes even for solutions of the same concentration.
When comparing with literature values, the temperature conditions must also be matched.

Discussion of Conductometric Titration

Electrical conductivity measurement is also used in titration experiments.
In conductometric titration, the ionic species and ion concentrations in the solution change as titration proceeds, so electrical conductivity also changes.
The equivalence point may be determined from the point at which the slope of the graph changes.

For example, when a strong acid is titrated with a strong base, many highly mobile hydrogen ions are present before titration.
When the base is added, the hydrogen ions are neutralized and the conductivity decreases.
After the equivalence point, hydroxide ions increase and the conductivity rises again.

What Is the Cell Constant?

To measure electrical conductivity, the effects of the distance between electrodes and the electrode area must be corrected.
The value specific to this electrode arrangement is the cell constant.
In a conductivity meter, the cell constant is calibrated using a standard solution and then used to determine the electrical conductivity of the sample.

If the cell constant is incorrect, all measured values may be systematically shifted.
Therefore, calibration using a standard solution before measurement is important.

Error Caused by Electrode Contamination

If the electrodes of the conductivity meter are contaminated, the flow of current at the electrode surface is hindered and measurement error occurs.
If residue from the previous sample, precipitates, or oil adheres to the electrodes, the surface condition of the electrodes changes.
Particularly in low-concentration solutions, even slight contamination may greatly affect the measured value.

Error Caused by Bubbles

If a bubble adheres to the electrode surface, the area of contact between the electrode and the solution becomes smaller.
As a result, current flows less easily and the electrical conductivity may be measured as lower than the actual value.
Bubbles may form when the electrode is immersed in the solution or during stirring.

In conductivity measurements, it is important to check that no bubbles are attached to the electrode surface and to read the value only after it has stabilized.

Error Caused by Concentration Preparation

In experiments examining the relationship between electrolyte concentration and electrical conductivity, errors in concentration preparation directly affect the results.
If errors occur when using volumetric flasks, volumetric pipettes, or electronic balances, the actual concentration deviates from the set value.
If the concentration deviates, the measured electrical conductivity also changes.

Particularly in dilute solutions, slight contamination or dilution errors may have a large effect.
Electrolytes remaining in pure water or on laboratory glassware can also become non-negligible sources of error in low-concentration measurements.

Effects of Pure Water and Solvent

In electrical conductivity measurements, the conductivity of pure water or the solvent itself may not always be negligible.
Particularly in low-concentration solutions, the conductivity originating from the sample is small, so the relative effects of trace ions in pure water and dissolved carbon dioxide become larger.
Blank correction and confirmation of pure-water quality are important.

Discussion When the Measured Value Does Not Stabilize

If the conductivity-meter reading does not stabilize, possible causes include unstable temperature, bubbles on the electrode surface, insufficient mixing of the sample, electrode contamination, or improper immersion of the electrode.
If the value is read before it stabilizes, variation occurs in the results.

When the Measured Value Is Larger Than the Literature Value

If the experimentally determined electrical conductivity is larger than the literature value, possible causes include a higher measurement temperature, preparation of the sample at a higher concentration, contamination of laboratory glassware or pure water with electrolyte, or deviation of the cell constant.
Particularly in low-concentration solutions, insufficient cleaning of laboratory glassware or residue from the previous sample may have a large effect.

When the Measured Value Is Smaller Than the Literature Value

If the experimentally determined electrical conductivity is smaller than the literature value, possible causes include a lower measurement temperature, preparation of the sample at a lower concentration, bubbles on the electrode, electrode contamination, or calibration error in the cell constant.
If a bubble is present on the electrode surface, the contact area with the solution becomes smaller and the conductivity may be measured as too low.

Calculation of Error Rate

When experimentally determined electrical conductivity or molar conductivity is compared with a literature value, the error rate can be calculated to express the difference quantitatively.
The error rate is obtained by dividing the difference between the experimental and literature values by the literature value and expressing it as a percentage.

Error rate (%) = |Experimental value − Literature value| ÷ Literature value × 100

Simply showing the error rate is not sufficient as a discussion.
Explain which factors, such as temperature, concentration preparation, cell constant, electrode contamination, bubbles, and impurities in pure water, shifted the measured value and in which direction.

When the Results Can Be Considered Good

Electrical conductivity measurement results can be considered good when repeated measurements under the same conditions agree well and theoretically reasonable trends are observed, such as increasing conductivity with increasing concentration.
It is also important that the difference between strong and weak electrolytes can be explained and that the results do not greatly contradict literature values.

Example Discussion When the Experiment Did Not Go Well

If electrical conductivity measurement does not go well, the causes are considered from results such as variation in measured values, an unnatural relationship with concentration, large deviation from literature values, unstable readings, or large deviations in the low-concentration range.
Organizing the causes separately into temperature control, concentration preparation, electrode cleaning, bubbles, cell constant, and pure-water quality makes the discussion easier.

How to Write Points for Improvement

In a discussion of electrical conductivity measurement, including points for improvement in addition to sources of error makes the report easier to organize.
Improvements are easier to write when divided into temperature control, concentration preparation, electrode operation, calibration, and reading of measured values.

Improvements to Temperature Control

• Keep the measurement temperature constant
• Measure after the sample has reached thermal equilibrium
• Record the measurement temperature
• Match the temperature conditions when comparing with literature values
• Check whether a temperature-correction function is available

Improvements to Concentration Preparation

• Prepare standard solutions accurately
• Use volumetric flasks and volumetric pipettes correctly
• Perform dilution operations carefully
• Do not leave residue from the previous sample on laboratory glassware
• Check the conductivity of pure water in low-concentration measurements

Improvements to Electrode and Measurement Operations

• Calibrate the cell constant using a standard solution
• Wash the electrode thoroughly
• Remove bubbles from the electrode surface
• Mix the sample uniformly
• Read the value only after it stabilizes
• Perform multiple measurements and take the average

Difference Between a Superficial Discussion and a Good Discussion

In a discussion of electrical conductivity measurement, simply writing that “conductivity increased as the concentration increased” or “the strong electrolyte had higher conductivity” results in a superficial discussion.
Relating the number of ions, degree of ionization, mobility, interactions between ions, temperature, and measurement errors produces a more persuasive discussion.

Examples of Expressions That Can Be Used in Reports

The following expressions can be used when writing the results and discussion of electrical conductivity measurements.
Adjust the necessary parts according to your own experimental results.

• Electrical conductivity arises because ions in the solution carry electric charge.
• As the electrolyte concentration increases, the number of ions per unit volume increases, so electrical conductivity becomes larger.
• Strong electrolytes are almost completely ionized, so their electrical conductivity tends to be higher than that of weak electrolytes at the same concentration.
• Because weak electrolytes have a low degree of ionization, the number of ions in the solution is small and conductivity becomes lower.
• At high concentrations, interactions between ions may interfere with ion movement.
• Molar conductivity represents the conducting ability per mole of electrolyte.
• Dilution may weaken interactions between ions and increase molar conductivity.
• As temperature rises, the viscosity of the solution decreases and ionic mobility increases, so conductivity tends to increase.
• Bubbles on the electrode surface reduce the contact area between the electrode and solution and cause conductivity to be measured as low.
• Calibration error in the cell constant becomes a systematic error that shifts all measured values in the same direction.

Points to Check When Discussing Electrical Conductivity Measurements

Checking the following points before writing the report makes the discussion easier to write.

• Have you stated the measurement temperature?
• Have you written the electrolyte concentration correctly?
• Have you distinguished between electrical conductivity and molar conductivity?
• Have you prepared a concentration-electrical conductivity graph?
• Have you explained the reason conductivity increases with increasing concentration in terms of the number of ions?
• Have you explained the difference between strong and weak electrolytes from the degree of ionization?
• Have you discussed deviation at high concentrations in terms of interactions between ions?
• Have you considered the effect of temperature?
• Have you considered the cell constant and calibration using a standard solution?
• Have you considered electrode contamination and bubbles as sources of error?
• Have you considered the effects of pure water and contamination at low concentrations?
• Do the points for improvement correspond to the sources of error?

Summary

Electrical conductivity is a physical quantity that arises because ions in a solution carry electric charge.
As electrolyte concentration increases, the number of ions per unit volume increases, so electrical conductivity generally becomes larger.
Strong electrolytes are almost completely ionized and therefore tend to show higher conductivity than weak electrolytes at the same concentration.

However, at high concentrations, interactions between ions become stronger and may interfere with ion movement.
Therefore, the relationship between concentration and conductivity may deviate from simple proportionality.
In addition, molar conductivity may decrease as concentration increases and is related to the increased ease of ion movement upon dilution.

In a report, do not simply write that “conductivity increases as concentration increases.”
Discuss the results in relation to the number of ions, degree of ionization, ionic mobility, interactions between ions, temperature, cell constant, and electrode condition.
Because electrical conductivity is sensitive to temperature and electrode condition, careful temperature control, calibration using a standard solution, electrode cleaning, and bubble removal are important.