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Inside Collection (Course):

Course by: Mary McHale. E-mail the author

# Bonding 07

Module by: Mary McHale. E-mail the author

## Objective

• To test various compounds and determine their conductivity and bonding.
• To understand how electronegativity can predict bond type.
• To learn the relationship between bonding and conductivity.

• Pre-Lab (10%)
• Lab Report Form (80%)
• TA Points (10%)

## Background Information

A chemical bond is a link between atoms that results from the mutual attraction of their nuclei for electrons. Bonding occurs in order to lower the total potential energy of each atom or ion.  Throughout nature, changes that decrease potential energy are favored.

The main types of bonds that we will be covering are ionic bonds, covalent bonds, and metallic bonds.  An ionic bond is the chemical bond that results from the electrostatic attraction between positive (cations) and negative (anions) ions. The ionic relationship is a “give and take” relationship. One ion donates or “gives” electrons, while the other ion receives or “takes” electrons.

A covalent bond is a chemical bond resulting from the sharing of electrons between two atoms.  There are two main types of covalent bonds.  The first being non-polar covalent bonds.  These are bonds in which the bonding electrons are shared equally by the united atoms-with a balanced electrical charge.  Polar covalent bonds are covalent bonds in which the united atoms have an unequal attraction for the shared electrons.

The role of electrons in bonding has been well-studied. The ability of an atom or element to attract electrons to itself is known as the element’s electronegativity. A scale was first calculated by the Nobel laureate Linus Pauling and is commonly called the Pauling electronegativity scale. The actual electronegativity values aren’t as important as how they compare to a different element. In Part I of today’s experiment, you will compare electronegativity values to predict the type of bond that will exist between two elements.

In the solution state, ionic compounds dissociate to give a separation of charge. The separation of charge allows for the flow of electrons through solution. The flow of electrons is classified as conductivity. A strong electrolyte is a compound that when dissolved in water will completely ionize or dissociate into ions.  That is, the compound exists in water only as individual ions, and there are no intact molecules at all.  This solution conducts electricity well. A weak electrolyte is a compound that when dissolved in water only partially ionizes or dissociates into ions.  That is, the compound exists in water as a mixture of individual ions and intact molecules.  This solution conducts electricity weakly. A nonelectrolyte is a compound that when dissolved in water does not ionize or dissociate into ions at all.  In water, this compound exists entirely as intact molecules.  The solution does not conduct electricity at all. By measuring the conductivity of a dissolved compound, we can classify it as a nonelectrolyte, weak electrolyte, or strong electrolyte and determine its ability to dissociate into ions. There are four common compounds that you will encounter in today’s lab.

ACIDS are molecular compounds which ionize (turn into ions) in water.  The cation that is formed is always H+H+ size 12{H rSup { size 8{+{}} } } {}.  Therefore, in the formulas for simple acids, H is always the first element listed. Some acids are strong electrolytes and some acids are weak electrolytes.  There are no acids which are nonelectrolytes because by definition an acid is a H+H+ size 12{H rSup { size 8{+{}} } } {} donor.

BASES can be molecular compounds or ionic compounds.  Some bases are soluble and some are not.  The soluble bases ionize or dissociate into ions in water, and the anion formed is always OHOH size 12{"OH" rSup { size 8{ - {}} } } {}.  The ionic bases have hydroxide ( OHOH size 12{"OH" rSup { size 8{ - {}} } } {} ) as the anion.  If they are soluble, the ions simply separate (dissociate) in the water.  All of the ionic bases which are soluble are strong electrolytes.

SALTS are ionic compounds which are not acids or bases.  In other words, the cation is not hydrogen and the anion is not hydroxide.  Some salts are soluble in water and some are not.  All of the salts which are soluble are relatively strong electrolytes.

NONELECTROLYTES are compounds which dissolve in water but do not ionize or dissociate into ions.  These would be molecular compounds other than the acids or bases already discussed.

## Experimental Procedure

Caution:Acids and bases are corrosive and can cause burns.

### Part I. Predicting bond type through electronegativity differences.

Using the electronegativity table provided in the lab manual, predict the type of bond that each of the following compounds will have by the following process:

• Find the electronegativity for each element or ion in compound using electronegativity table provided.
• Subtract the electronegativites (using absolute value).
• If values are between:

4.0-1.7---Ionic bond-50-100% ionic

1.7-0.3---Polar Covalent bond-5-50% ionic

0.3-0.0---Non-Polar Covalent-0-5% ionic

Determine the type of bonding in the following compounds: KCl, CO, CaBr2CaBr2 size 12{"CaBr" rSub { size 8{2} } } {}, SiH4SiH4 size 12{"SiH" rSub { size 8{4} } } {}, MgS.

### Chemicals

• tap water
• 0.1 M hydrochloric acid, HCl
• 0.1 M acetic acid, HC2H3O2HC2H3O2 size 12{"HC" rSub { size 8{2} } H rSub { size 8{3} } O rSub { size 8{2} } } {}
• 0.1 M sulfuric acid, H2SO4H2SO4 size 12{H rSub { size 8{2} } "SO" rSub { size 8{4} } } {}
• 0.1 M sodium hydroxide, NaOH
• 0.1 M ammonia, NH3NH3 size 12{"NH" rSub { size 8{3} } } {}
• 0.1 M sodium acetate, NaC2H3O2NaC2H3O2 size 12{"NaC" rSub { size 8{2} } H rSub { size 8{3} } O rSub { size 8{2} } } {}
• 0.1 M sodium chloride, NaCl
• 0.1 M ammonium acetate, NH4C2H3O2NH4C2H3O2 size 12{"NH" rSub { size 8{4} } C rSub { size 8{2} } H rSub { size 8{3} } O rSub { size 8{2} } } {}
• 0.1 M ammonium chloride, NH4ClNH4Cl size 12{"NH" rSub { size 8{4} } "Cl"} {}
• methanol, CH3OHCH3OH size 12{"CH" rSub { size 8{3} } "OH"} {}
• ethanol, C2H5OHC2H5OH size 12{C rSub { size 8{2} } H rSub { size 8{5} } "OH"} {}
• sucrose solution, C12H22O11C12H22O11 size 12{C rSub { size 8{"12"} } H rSub { size 8{"22"} } O rSub { size 8{"11"} } } {}

In today’s lab, you will be using a MicroLab conductivity probe to determine how well electrons flow through a given solution. First, you will need to calibrate the probe with a non-electrolyte (distilled water) and a very strong electrolyte. To quantify how well a solution conducts, we will assign numerical values to the conductance probe. A non-conducting solution will have a conductance value of 0, a poor conducting solution will have a reading of 0 to 1,000, and good conductors will have readings of 3,000 up.

### Instructions for MicroLab Conductivity Experiment

Open the MicroLab Program by clicking on the Shortcut to MicroLab.exe tab on the desktop.

On the “Choose an Experiment Type” Tab, enter a name for the experiment, and then double click on the MicroLab Experiment icon

Click “Add Sensor”, Choose sensor = Conductivity Probe

Choose an input, click on the red box that corresponds to the port that your conductivity sensor is connected to. Choose 20,000 microseconds

“Choose a Sensor”, click radial button that says Conductivity Probe. Click next.

Click “Perform New Calibration”

Click “Add Calibration Point” place the conductivity probe in the non-conductive standard solution, while swirling wait until the value is constant and then enter 0.0 into the “Actual Value” box in MicroLab and hit “ok”.

Again, Click “Add Calibration Point” place the conductivity probe in the conducting standard solution, while swirling wait until the value is constant and then enter 1020 into the “Actual Value” box in MicroLab and hit “ok”. Repeat for 3860 as the Actual Value.

Under Curve Fit Choices , click on “First order (linear)” and then “Accept and Save this Calibration”, when prompted to “Enter the units for this calibration”, leave as is and click ok, save as your name-experiment-date. Click finish.

In the sensor area, left click on the conductivity icon and drag it to the Y-axis over “data source two”, also click and drag to column B on the spreadsheet and also click and drag to the digital display window.

When ready to obtain data, click start.

This is very important: Be sure to thoroughly since the probe with DI water between every use.

Beginning with the tap water, measure the conductance of each of the following solutions. Using the information provided in the lab manual, classify each solution as a non-, weak, or strong electrolyte. For those solutions that are electrolytes, record the ions present in solution.

### Chemicals

• calcium carbonate powder - shake once
• 1 M HCl - stopper it
• 1 M HC2H3O2HC2H3O2 size 12{"HC" rSub { size 8{2} } H rSub { size 8{3} } O rSub { size 8{2} } } {}
• 0.5 M H2SO4 H2SO4 size 12{H rSub { size 8{2} } "SO" rSub { size 8{4} } } {} begin time
• Test tube gas collection apparatus - end at 20mL

Measure 2 g of powdered calcium carbonate ( CaCO3CaCO3 size 12{"CaCO" rSub { size 8{3} } } {}) onto a piece of weigh paper. Obtain 30 mL of 1 M HCl in a graduated cylinder. Pour the acid into the test tube apparatus. Add the calcium carbonate to the acid and QUICKLY stopper the tube to begin collecting gas. Record the time it takes to collect 20 mL of gas. The acid may react very fast with the CaCO3CaCO3 size 12{"CaCO" rSub { size 8{3} } } {} generating the gas very rapidly. Clean out the test tube apparatus and repeat the experiment using 1 M HC2H3O2HC2H3O2 size 12{"HC" rSub { size 8{2} } H rSub { size 8{3} } O rSub { size 8{2} } } {} and 0.5 M H2SO4H2SO4 size 12{H rSub { size 8{2} } "SO" rSub { size 8{4} } } {}.

### Chemicals

• 0.01 M calcium hydroxide, Ca(OH)2Ca(OH)2 size 12{"Ca" $$"OH"$$ rSub { size 8{2} } } {}
• Plastic straws

Obtain ~20 mL of saturated calcium hydroxide solution. Make sure it is clear and colorless. Place the conductivity probe in the solution and begin monitoring it conductivity. With your straw, slowly exhale into the solution. Note any observations in the solution and the conductivity.

## (Total 10 Points)

Hopefully here for the Pre-Lab

Name(Print then sign): ___________________________________________________

Lab Day: ___________________Section: ________TA__________________________

This assignment must be completed individually and turned in to your TA at the beginning of lab. You will not be allowed to begin the lab until you have completed this assignment.

### Part I. Bonding of chemicals in solution

1. Write out the formulas of the following acids:

• phosphoric ____________________

• perchloric ____________________

• nitric ____________________

• sulfuric __________________

• hydrochloric ____________________

• acetic ____________________

1. Write out the formulas of the following bases:

• calcium hydroxide ____________________

• potassium hydroxide ____________________

• sodium hydroxide ____________________

• ammonia ____________________

1. Write out the formulas of the following salts:

• potassium chromate ____________________

• potassium sulfate ____________________

• copper(II) nitrate ____________________

• calcium carbonate ____________________

• potassium iodide ____________________

## Report 5: Bonding 07

Hopefully here for the Report Form

Note: In preparing this report you are free to use references and consult with others. However, you may not copy from other students’ work (including your laboratory partner) or misrepresent your own data (see honor code).

Name(Print then sign): ___________________________________________________

Lab Day: ___________________Section: ________TA__________________________

### Part I. Predicting bond type through electronegativity differences.

 Chemical Formula Electroneg (1) Electroneg (2) Diff Electroneg Type of bond KCl CO CaBr 2 CaBr 2 size 12{"CaBr" rSub { size 8{2} } } {} SiH 4 SiH 4 size 12{"SiH" rSub { size 8{4} } } {} MgS

### Part II. Weak and strong electrolytes

 Solution Tested Numerical Output Electrolyte Strength Ions Present 0.1 M HCl 0.1 M HC2H3O2HC2H3O2 size 12{"HC" rSub { size 8{2} } H rSub { size 8{3} } O rSub { size 8{2} } } {} 0.1 M H2SO4H2SO4 size 12{H rSub { size 8{2} } "SO" rSub { size 8{4} } } {} 0.1 M NaOH 0.1 M NH3NH3 size 12{"NH" rSub { size 8{3} } } {} 0.1 M NaC2H3O2NaC2H3O2 size 12{"NaC" rSub { size 8{2} } H rSub { size 8{3} } O rSub { size 8{2} } } {} 0.1 M NaCl 0.1 M NH4C2H3O2NH4C2H3O2 size 12{"NH" rSub { size 8{4} } C rSub { size 8{2} } H rSub { size 8{3} } O rSub { size 8{2} } } {} 0.1 M NH4ClNH4Cl size 12{"NH" rSub { size 8{4} } "Cl"} {} CH 3 OH CH 3 OH size 12{"CH" rSub { size 8{3} } "OH"} {} C 2 H 5 OH C 2 H 5 OH size 12{C rSub { size 8{2} } H rSub { size 8{5} } "OH"} {} Sucrose Tap water

1. Why do we use deionized water instead of tap water when making solutions for conductivity measurements?

### Part III. Electrolyte strength and reaction rate

2. Time to collect 20 mL of gas using 1 M HCl _______________________. Write the reaction of HCl with CaCO3CaCO3 size 12{"CaCO" rSub { size 8{3} } } {}.

3. Time to collect 20 mL of gas using 1 M HC2H3O2HC2H3O2 size 12{"HC" rSub { size 8{2} } H rSub { size 8{3} } O rSub { size 8{2} } } {}_______________________. Write the reaction of HC2H3O2HC2H3O2 size 12{"HC" rSub { size 8{2} } H rSub { size 8{3} } O rSub { size 8{2} } } {} with CaCO3CaCO3 size 12{"CaCO" rSub { size 8{3} } } {}.

4. Time to collect 20 mL of gas using 0.5 M H2SO4H2SO4 size 12{H rSub { size 8{2} } "SO" rSub { size 8{4} } } {}_________________________.Write the reaction of H2SO4H2SO4 size 12{H rSub { size 8{2} } "SO" rSub { size 8{4} } } {} with CaCO3CaCO3 size 12{"CaCO" rSub { size 8{3} } } {}.

5. Why does it take different lengths of time to collect 20 mL of gas?

6. Based on the time it took to collect 20 mL of gas, rank the acids in the order of increasing strength.

7. Why did we use 0.5 M H2SO4H2SO4 size 12{H rSub { size 8{2} } "SO" rSub { size 8{4} } } {} instead of 1.0 M H2SO4H2SO4 size 12{H rSub { size 8{2} } "SO" rSub { size 8{4} } } {}?

### Part IV. Chemical reactions

8. Write the chemical reaction for calcium hydroxide with your exhaled breath.

9. Write your observations for the reaction that took place (i.e. appearance, conductivity, etc.)

10. When in separate solutions, aqueous ammonia, NH3NH3 size 12{"NH" rSub { size 8{3} } } {}(aq) and acetic acid HC2H3O2HC2H3O2 size 12{"HC" rSub { size 8{2} } H rSub { size 8{3} } O rSub { size 8{2} } } {} conduct electricity equally well. However, when the two solutions are mixed a substantial increase in electrical conductivity is observed. Explain.

11. Separately, ammonium sulfate and barium hydroxide solutions are very good conductors. When the two solutions are mixed a substantial decrease in conductivity is observed. Rationalize this.

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