general chemistry lab manual

GENERAL CHEMISTRY

Laboratory Manual

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General Chemistry - Lab Manual Page: 1

Introduction

The Lab Manual serves as your text for the lab portion of this course (CHEM 111 – General

Chemistry). You must carefully read through the experiment to be performed. Keep the lab manual

with you during the lab hours.

Objective

Learn chemistry the hands-on way, with the general chemistry laboratory experiments.

Table of contents

S.No. Particulars Page

1 Laboratory safety rules

2 Introduction – Qualitative analysis

3 Analysis of acid radicals (Diluted HCl group)

4 Analysis of acid radicals (Con. H2So4 group)

5 Analysis of acid radicals (miscellaneous group)

6 Analysis of base radicals

7 Analysis of unknown sample (salt)

8 Predicting shapes of molecules using molecular models

9 Measurement of Density

10 Measurement of melting points and boiling points


Laboratory Safety Rules

Laboratories are interesting and potentially dangerous place to work. Laboratory Safety is a very important aspect of science. Without it, experimentation could result in very serious injury. To reduce the risks involved with experimentation, there are certain procedures that we should all follow.

“Safety is the first priority in our lab”

1. Wear lab coat. A lab coat should be worn during laboratory experiments. It provides protection at all times.

2. Dress properly. Long hair and loose clothing (such as ties and ghutra) are a hazard in the laboratory. They may either catch fire or be chemically contaminated. Long hair must be tied back, baggy clothing must be secured.

3. Wear shoes. Shoes must completely cover the foot. Sandals are prohibited because of the hazard from the chemical spills.

4. Wear eye protector. Wear appropriate eye protection (goggles) at all times in the laboratory and in any area where chemicals are stored or handled. Such protection will protect you against chemical splashes.

5. Do not wear contact lenses (even with safety goggles). Contact lenses prevent rinsing chemical splashes from the eye. Vapors in the laboratory (eg. HCl) dissolve in the liquids covering the eye and concentrate behind the lenses. ‘Soft’ lenses are especially bad as chemicals dissolve in the lenses themselves and are released over several hours.

6. Behave responsibly. The dangers of spilled acids, chemicals and broken glassware created by thoughtless actions are not tolerated.

7. Do not smoke. Smoking is not just an obvious fire hazard; it also draws chemicals in laboratory air (both vapors and dust) into the lungs.

8. Do not eat or drink. Do not eat food, drink beverages, or chew gum in the laboratory. Do not use laboratory glassware as containers for food or beverages. Wash your hands thoroughly with soap and water when leaving the laboratory.

9. Avoid breathing fumes of any kind. To test the smell of a vapor, collect some in a cupped hand. Obtain your instructor’s written permission before you smell any chemical. Never smell a chemical reaction while it is occurring. Work in a hood if there is the possibility that noxious vapors may be produced.

10. Be prepared for your work in the laboratory. Carefully read the experiment before coming to the laboratory. An unprepared student is a hazard to everyone in the room.


11. KEEP YOUR WORK AREA NEAT AND CLEAN. Clean up spills and broken glass immediately. Clean up your work space; wipe all surfaces and put away all chemicals and equipment at the end of the laboratory period.

12. Never work alone in the lab. There must be at least one other person present in the lab. In addition, an instructor should be quickly available.

13. Do not perform any unauthorized experiments. Perform only those experiments authorized by your teacher. This includes using only the quantities instructed, no more. Carefully follow all instructions, both written and oral. Consult your instructor if you have any doubts about instructions and laboratory manuals.

14. Be careful while heating liquids. Add boiling chips to avoid ‘bumping’. Flammable liquids such as ethers, hydrocarbons, alcohols, acetone, and carbon disulfide must never be heated over direct flame.

15. Careful about spatters. Spatters are common in general chemistry lab. Test tubes being heated or containing reaction mixtures should never be pointed at anyone. If you observe this practice in a neighbor, speak to him or the instructor.

16. Always pour acids into water when mixing. Otherwise the acid can spatter, often quite violently.

17. Never use mouth suction in filling pipettes with chemical reagents. Always use a suction device.

18. Do not force a rubber stopper onto glass tubing or thermometers. Lubricate the tubing and the stopper with glycerol or water. Use paper or cloth toweling to protect your hands. Grasp the glass close to the stopper.

19. Dispose of all chemical waste properly. Dispose of excess liquid reagents by flushing small quantities down the sink. Consult the instructor about large quantities. Dispose of solid in crocks. Never return unused chemicals to their original container.

20. Report any accident (spill, breakage, etc.) or injury (cut, burn, etc.) to your instructor immediately. Do not panic.

21. In case of fire or accident, call the instructor at once. Note the location of fire extinguisher and safety showers now so that you can use them if needed.

- Wet towels can be used to smother small fires.

- In case of chemical spill on your body or clothing, wash the affected area with large quantities of running water. Remove clothing that has been wet by chemicals to prevent further reaction with the skin.

- If a chemical should splash in your eye(s), immediately flush with running water for at least 20 minutes.

- Except for superficial injuries, you will be required to get medical treatment for cuts, burns or fume inhalation.


22. Most importantly, think about what you are doing. Plan ahead. If you give no thought to what you are doing, you predispose yourself to an accident.

23. Report if you have any diagnosed allergies or other special medical needs to your instructor.

QUALITATIVE ANALYSIS

The object of qualitative analysis is the identification of the constituents of a substance. In the

qualitative analysis procedure, the chemical constituents and properties of an unknown substance

are determined by systematically reacting the unknown with a number of different reagents or

chemicals.

A visible chemical reaction that can be utilized in qualitative analysis may involve:

- Formation of precipitate (ppt.)

- Change in color

- Evolution of gas or vapor

QUALITATIVE ANALYSIS OF SIMPLE SALT

In this course, we are going to analysis the chemical constituents of simple SALTS.

A salt is made of two parts: one from the acid and the other from the base.

Eg. HCl + NaOH ----------- NaCl + H2O

- The part derived from the acid is known as the acid radical (Eg. Cl-)

- The part derived from the base is known as base radical (Eg. Na+)

ACID RADICALS

Acid radicals to be studied in this course can be divided according to their behavior towards dilute

hydrochloric acid (HCl) and hot concentrated sulphuric acid (H2SO4) in three groups.

1) Diluted HCl group: Acidic radicals evolving gases or vapors with dil. HCl.

Eg. Carbonates [(CO3)2-]

Bicarbonates [(HCO3)-]

Nitrites

2) Concentrated H2SO4 group: Acidic radicals evolving gases or vapors with hot conc.

H2SO4.

Eg. Chlorides (Cl-)


Nitrates [(NO3)-]

3) Miscellaneous group: Acidic radicals unaffected by either acids (HCl and H2SO4).

Eg. Sulphate [(SO4)2-]

BASE RADICALS

Base radicals to be studied in this course are sodium (Na+) and magnesium (Mg2+).

SALT

Acid radicals

Dil. HCl Group

1. Carbonates (CO32-)

2. Bicarbonates (HCO3-)

3. Nitrites (NO2-)

Con. H2SO4 group1. Chlorides (Cl-)

2. Nitrates (NO3-)

Miscellaneous group 1. Sulphates (SO42-)

Base radicals1. Sodium (Na+)

2. Magnesium (Mg2+)


Experiment: 1 Date:

ANALYSIS OF ACID RADICALS (DILUTED HCl GROUP)

The dil. HCl group contains salts derived from unstable acids which decompose giving gases which

can be identified by odor, color, etc.

1) Carbonates (CO3) 2-

Pure carbonic acid does not exist, and its solution in water rapidly decomposes into CO2 and H2O.

They exist as salts of carbonic acids (H2CO3). All carbonates are insoluble in water except those of

the alkali metals (Na, K) and (NH4)+.

Example for the salt of CO32- = Sodium carbonate (NaCO3)

Main test:

1) Dil HCl + Solid

- Effervescence (emit small bubbles of gas) and evolution of CO2 gas that is colorless and

odorless gas; detected by the turbidity of limewater.

- If the carbonate is soluble in water, carry the following reactions (the confirmatory tests).

Na2CO3 + 2HCl ----------- 2NaCl + CO2 ↑ + H2O

Confirmatory tests:

2) Barium chloride (BaCl2) solution + Salt solution

- White precipitate is formed of Ba2CO3.

- On cold, the precipitate is soluble in mineral acids (eg. HCl, H2SO4).

BaCl2 + Na2CO3 ----------- BaCO3 ↓ + 2 NaCl

3) Silver nitrate (AgNO3) solution + Salt solution

- White precipitate is formed of Ag2CO3.

- On cold, the precipitate is soluble in NH4OH and HNO3.

2AgNo3 + Na2CO3 ----------- Ag2CO3 ↓ + 2 NaNO3


4) Magnesium sulphate (MgSO4) solution + Salt solution

- On cold, a white precipitate is formed of MgCO3.

MgSO4 + Na2CO3 ----------- MgCO3 ↓ + Na2SO4

2) Bicarbonates (HCO3) -

Also known as acid carbonates or hydrogen carbonates. All bicarbonates are soluble in water (not

all metals form bicarbonates).

Example for salt of HCO3-: Sodiumbicarbonate (NaHCO3)

Main test:

1) Dil. HCl + Solid

- The same results as in case of carbonates (Effervescence (emit small bubbles of gas)

and evolution of CO2 gas that is colorless and odorless gas; detected by the turbidity of

limewater).

NaHCO3 + HCl ----------- NaCl + CO2 ↑ + H2O

Confirmatory test:

2) Magnesium sulphate (MgSO4) solution + Salt solution

- When MgSO4 solution added to a cold salt solution (NaHCO3), no precipitate is formed,

because Mg(HCO3)2 is soluble in water.

- When the solution is then boiled, a white precipitate is formed of MgCO3.

MgSO4 + 2NaHCO3 ----------- Mg(HCO3)2 + Na2SO4

Mg(HCO3)2 + Heat ---------- MgCO3 ↓ + H2O + CO2

3) Nitrites (NO2) -

Nitrous acid is only known in the form of its salts. All nitrites are soluble in water, except AgNO2.

Example for the salt of NO2-: Sodium nitrite (NaNO2).


Main test:

1) Dil. HCl + Salt (Solid)

- Brown fumes of NO2 evolved (resulting from the decomposition of HNO2), and the solution

of acquires a pale blue color.

NaNO2 + HCl ----------→ HNO2 + NaCl

3HNO2 ----------→ HNO3 + 2NO + H2O

2NO + O2 (air) ----------→ 2NO2 (Brown)

Confirmatory test:

2) Salt solution (concentrated) + Silver nitrite (AgNO3) solution

- White precipitate is formed, which is soluble in boiling water.

NaNO2 + AgNO3 ----------→ NaNO3 + AgNO2 ↓

3) Salt solution + One drop of KMnO4 + 2 drops of Dil. H2SO4

- The deep violet color of permanganate disappear at warm temperature, because nitrites are

reducing substances.

5NaNO2 + 2KMnO4 + 3H2SO4 ----------→ 5NaNO3 + K2SO4 + 2MnSO4 +3H2O

4) Salt solution + Ferrous sulphate (FeSO4) solution + Few drops of H2SO4.

- The solution becomes brown color, because nitrites are oxidizing substances.

2NaNO2 + 2FeSO4 + 3H2SO4 ----------→ 2NO + Fe2(SO4)3 + 2NaHSO4 + 2H2O

2NO + O2 (air) ----------→ 2NO2 (Brown)

5) Salt solution + Bromine water

- Bromine water is decolorized, because nitrites are reducing substances.

NaNO2 + Br2 + H2O ----------→ NaNO3 + 2HBr

6) To the salt solution, add a crystal of KI and a starch paper, and then acidify with dil. H2SO4.

- A deep blue color appears on the paper due to the liberation of iodine.

2NaNO2 + 2KI + 2H2SO4 ----------→ I2 + 2NO + 2H2O + K2SO4 + Na2SO4


1) Carbonates (CO3) 2-

S.No. Experiment Observation Results

1Salt (solid) + Dil. HCl

Effervescence (emission of small

bubbles of gas).

Presence of Carbonates

or Bicarbonates.

2 Salt solution + BaCl2

solution

White precipitate is formed (on

cold).

Presence of

Carbonates.

3 Salt solution + AgNO3

solution


cold).

Presence of

Carbonates.

4 Salt solution + MgSO4

solution


cold).

Presence of

Carbonates.

Results:

I have performed the above experiments. From the observation, I confirmed that the acid radical

present in the given sample is ………………………………………………………………………………

2) Bicarbonates (CO3) 2-


1 Salt (solid) + Dil. HCl Effervescence (emission of small

bubbles of gas) and evolution of

colorless gas detected by the

turbidity of lime water.

Presence of Carbonates

or Bicarbonates.

2 Salt solution + MgSO4

solution

NO precipitate is formed (on cold).

When solution is boiled, a white

precipitate is formed.

Presence of

Bicarbonates.

Results:


present in the given sample is ………………………………………………………………………………


3) Nitrites (NO2)-


1 Salt (solid) + Dil. HCl Brown fume evolved and

the solution acquires a

pale blue color.

Presence of Nitrites.

2 Salt solution (Concentrated) +

AgNO3 solution

White precipitate is

formed, soluble in boiling

water.


3 Salt solution + One drop of

KMnO4 + 2 drops of dil. H2SO4.

Warm this mixture.

Deep violet color of

permanganate

disappears.


4 Salt solution + FeSO4 solution

+ Few drops of dil. H2SO4

Brown color is formed Presence of Nitrites.

5 Salt solution + Bromine water Bromine water also

decolorized.


6 Salt solution + Crystal of KI +

Starch paper (immerse in

solution) + Few drops of dil.

H2SO4

A deep blue color appears

on the paper.


Results:

I have performed the above experiments.

From the observation, I confirmed that the acid radical present in the given sample is

………………………………………………………………………………………………………..


Experiment: 2 Date:

ANALYSIS OF ACID RADICALS (CON. H2SO4 GROUP)

The Conc. H2SO4 group contains Cl- and NO3-. Their corresponding acids are set free when the

salts are treated with a less volatile acid eg. H2SO4.

1) Chlorides (Cl - )

They are salts of hydrochloric acid (HCl). Nearly all chlorides are soluble in water except AgCl,

Hg2Cl, and PbCl2 (PbCl2 is soluble in hot water).

Example for the salt of Cl-: Sodium chloride (NaCl).

Main test:

1) Salt (solid) + Conc. H2SO4

- Evolution of colorless gas (HCl gas) that fumes in air, and forms white clouds (NH4Cl) with a

wetted paper by NH4OH.

Confirmatory tests

2) Salt solution + Silver nitrate (AgNO3) solution

- White cloudy precipitate of AgCl is formed, which is soluble in ammonia but insoluble in hot

dil. HNO3. The precipitate turns gradually form violet to black on exposure to light.

NaCl + AgNO3 ----------→ AgCl + NaNO3

AgCl + NH4OH ----------→ Ag(NH3)2Cl (soluble silver diamino-chloride complex)

3) Salt solution + Lead acetate (Pb(C2H3O2)2) solution

- A white precipitate of PbCl2 is formed, which is soluble in hot water but reprecipitate on

cooling.


2) Nitrates (NO3- )

All nitrates are soluble in water. Mercury and nitrates yield a precipitate of the basic salt on

treatment with water, but are soluble in dil. HNO3.

Example for the salt of NO3-: Sodium nitrate (NaNO3).

Main test:

1) Salt (solid) + Conc. H2SO4 + Heat

- Reddish brown vapors of NO2 accompanied with vapors of HNO3 are evolved, which

fumes in air.

2NaNO3 + H2SO4 ----------→ 2HNO3 + Na2SO4

4HNO3 ----------→ 4NO2 + O2 + 2H2O

Confirmatory tests

2) Salt (solid) + Conc. H2SO4 + Copper turnings + Heat

- The same result as the above case but the fumes increase and the solution is colored blue

due to the formation of copper nitrate. In this case copper reacts with liberated nitric acid

giving off NO2 and NO that with oxygen of the air gives also NO2.

3) Brown ring test

Salt solution + 2 to 4 crystals of FeSO4(green crystals) then shake well + add 1 to 2 ml of H2SO4

slowly down the side of the test tube.

- We observe a brown ring is formed at the interface between two solutions.

- The ring is due to the formation of [Fe(NO)SO4] compound. On shaking or warming the

mixture, the ring disappears.

- Note: If Pb or Ba nitrates is used, a white precipitate of Pb or Ba sulphates is formed on

adding the ferrous sulphate.


1 ) Chlorides (Cl - )


1 Salt (solid) + Conc. H2SO4 Evolution of colorless gas. Presence of Chlorides.

2 Salt solution + AgNO3 solution White curdy precipitate is

formed. The precipitate

gradually forms violet to

black on exposure to light.

Presence of Chlorides.

3 Salt solution + Pb(C2H3O2)2

solution


formed.

Presence of Chlorides.

Results:


present in the given sample is …………………………………………………………………………….

2) Nitrates (NO3) -


1 Salt (solid) + Conc. H2SO4 Evolution of reddish

brown vapors.

Presence of Nitrates.

2 Salt (solid) + Conc. H2SO4 +

Copper turnings + Heat

Evolution of reddish

brown vapors (higher than

the previous experiment).

The solution colored blue.

Presence of Nitrates.

3 Salt solution + 2 to 4 crystals of

FeSO4 (shake) + add 1 to 2 ml

of H2SO4 slowly down the side

of the tube.

Brown ring is formed. Presence of Nitrates.

Results:


present in the given sample is …………………………………………………………………………….


Experiment: 3 Date:

ANALYSIS OF ACID RADICALS (MISCILLANEOUS GROUP)

The miscellaneous group includes the remaining acid radicals (such as SO42- and PO4

3-) which are

not apparently affected by dil. HCl and Conc. H2SO4.

Example for the salt of SO42-: Sodium sulphate (Na2SO4).

1) Sulphates (SO4) 2-

Most sulphates are soluble in water, CaSO4 is difficulty soluble in water, SrSO4, PbSO4, BaSO4 are

insoluble in water.

Confirmatory tests

1) Salt solution + Barium chloride (BaCl2) solution

- White precipitate of barium sulphate is formed which is insoluble in hot dil. HNO3.

2) Salt solution + Lead acetate solution

- White precipitate of lead sulphate is formed. The precipitate is soluble in ammonium

acetate, or NaOH. Also, the precipitate is appreciable soluble in hot dil. HNO3.

3) Salt solution + Silver nitrate (AgNO3) solution

- White precipitate is formed from concentrated solutions only.


1) Sulphates (SO4) 2-


1 Salt solution + BaCl2 solution White precipitate is

formed.

Presence of Sulphates.

2 Salt solution + Lead acetate

solution


formed.


3 Salt solution + AgNO3 solution White precipitate is

formed.


Results:

I have performed the above experiments.

From the observation, I confirmed that the acid radical present in the given sample is

…………………………………………………………………………………………………..


Experiment: 4 Date:

ANALYSIS OF BASE RADICALS

This group contains magnesium (Mg2+) and sodium (Na+).

1) Magnesium (Mg 2+ )

Example for the salt of Mg2+: Magnesium sulphate (MgSO4).

Confirmatory tests

1) Salt solution + Ammonium carbonate ((NH4)2CO3) solution

- White precipitate of magnesium carbonate, MgCO3.Mg(OH)2.5H2O is formed on boiling or

long standing. The precipitate is soluble in ammonium salt solution.

2) Salt solution + Ammonia solution

- Partial precipitation of white gelatinous magnesium hydroxide, Mg(OH)2 is produced. The

hydroxide is soluble in ammonium salt solution and very sparingly soluble in water.

MgSO4 + 2NH3 + 2H2O ----------→ Mg(OH)2 + (NH4)2SO4

3) Salt solution + Sodium hydroxide (NaOH) solution

- White precipitate of magnesium hydroxide is formed (insoluble in excess of the reagent).

The precipitate is soluble in ammonium salt solution.

4) Salt solution + Sodium phosphate (Na3PO4) solution

- White crystalline precipitate of magnesium ammonium phosphate, Mg(NH4)PO4.6H2O is

formed in the presence of ammonium chloride (to prevent the precipitation of magnesium

hydroxide) and ammonia solution.

- The precipitate is sparingly soluble in water, but is soluble in acetic acid and in mineral

acids (eg. HCl, H2SO4).


- Scratching the side of the test tube with a glass rod accelerate precipitation from the

solutions.

2) Sodium (Na + )

- Example for the salt of Na+: Sodium chloride.

1) Salt solution + Sodium cobaltinate solution

- No precipitate is formed with sodium salts (distinction from potassium salts).

2) Flame test

1. Imparts a golden yellow color to Bunsen flame.


1) Magnesium (Mg 2+ )


1 Salt solution + (NH4)2CO3

solution. Boil the mixture.

White precipitate is formed. Presence of Magnesium.

2 Salt solution + Ammonia

solution

White gelatinous precipitate

(partial precipitation) is

formed.

Presence of Magnesium.

3 Salt solution + NaOH

solution

White precipitate is formed. Presence of Magnesium.

4 Salt solution + Sodium

phosphate solution +

Ammonium chloride +

ammonia solution

White crystalline precipitate is

formed.

Presence of Magnesium.

Result:

I have performed the above experiments. From the observation, I confirmed that the base radical

present in the given sample is ……………………………………………………………………………

2) Sodium (Na + )


1 Salt solution + Sodium

cobaltinate solution

NO precipitate is formed. Presence of Sodium.

2 Flame test Imparts a golden yellow color

to Bunsen flame.

Presence of Sodium.

Result:

I have performed the above experiments. From the observation, I confirmed that the base radical

present in the given sample is ……………………………………………………………………………


Experiment: 5 Date:

ANALYSIS OF UNKNOWN SAMPLE

The salt (unknown sample) contains both acid radicals and base radicals. Carryout the following

experiments in the given order.

Part 1: Qualitative Analysis of Acid Radical

Part 2: Qualitative Analysis of Base Radical

PART 1: Qualitative Analysis of Acid Radical

Diluted HCl group: Carbonates (CO3)2-, Bicarbonates (HCO3)-, and Nitrites (NO2)-.

Concentrated H2SO4 group: Chlorides (Cl-) and Nitrates (NO3-).

Miscellaneous group: Sulphates (SO42-).

Experiment Observation Conclusion

Salt (Solid)

+ Dil. HCl

Emission of small bubbles of gas (colorless).

Gas is CO2. The radical may be carbonate (CO3)2- or bicarbonate (HCO3)-.

(Perform confirmatory test-1).

Reddish brown gas evolved and the solution becomes blue.

Gas is NO2. The radical may be nitrites (NO2)-.


NO emission of gas. Absence of ‘diluted HCl group’.

(Perform next experiment).

Salt (Solid)

+ Con. H2SO4

Emission of colorless gas that forms white clouds with a wetted paper with NH4OH.

Gas is HCl. The radical may be chloride (Cl-).


Evolution of reddish brown gas after heat.

Gas is NO2. The radical may be nitrates (NO3

-).


NO emission of gas. Presence of ‘Miscellaneous group’. The radical may be sulphates (SO4

2-).



Confirmatory test-1 (for Carbonates and Bicarbonates)


* Salt solution + BaCl2

* Salt solution + AgNO3

* Salt solution + MgSO4

White ppt. is formed on cold condition.

Presence of Carbonates (CO3)2-.

* Salt solution + MgSO4 White ppt. is formed after heating.

Presence of Bicarbonates (HCO3)-.

Confirmatory test-2 (for Nitrites)


* Salt solution + AgNO3 White ppt. that turns to yellow and black.

Presence of Nitrites (NO2)-.

* Salt solution + One drop of KMnO4 + 2 drops of dil. H2SO4. Warm this mixture

Deep violet color of permanganate disappears.

Presence of Nitrites (NO2)-.

* Salt solution + FeSO4 solution + Few drops of dil. H2SO4.

Brown color. Presence of Nitrites (NO2)-.

Confirmatory test-3 (for Chloride)


* Salt solution + AgNO3 White curdy ppt. that turns to violet and black.

Presence of chloride (Cl-).

* Salt solution + Lead acetate White ppt. Presence of chloride (Cl-).

Confirmatory test-4 (for Nitrates)


* Salt soln. + 2 to 4 crystals of FeSO4 (shake) + add H2SO4 slowly down the side of the tube.

Brown ring. Presence of nitrates (NO3-).

* Salt soln. + Cu rurnings + Con. H2SO4

Reddish brown fumes and the solution become blue.

Presence of nitrates (NO3-).


Confirmatory test-5 (for Sulphates)


* Salt solution + BaCl2

* Salt soln. + Lead acetate

* Salt solution + AgNO3

White ppt. Presence of Sulphates (SO42-).

PART 2: Qualitative Analysis of Base Radical

This group contains magnesium (Mg2+) and sodium (Na+).


* Salt solution + Ammonia

*Salt solution + NaOH

* Salt solution + (NH4)2CO3

White or gelatinous ppt. Presence of Magnesium (Mg2+).

NO ppt. Presence of Sodium (Na+).



Result

The given sample is (base radical + Acid radical) …………………………………………………………


Experiment: 6 Date:

PREDICTING SHAPES OF MOLECULES USING MOLECULAR MODELS

Objectives

1. To draw Lewis structures of the given molecules.

2. To make models of molecular structures using VSEPR model.

3. To describe and name the molecular geometry (shapes) of each model.

Introduction

The chemical and physical properties of molecules and ions are directly related to the geometry of

the species. The spatial arrangement of atoms and electrons in a chemical species will affect the

function and reactivity of the species. Understanding the geometrical structure of chemical species

gives insight into studying chemical reactivity and reaction rates.

Valence Shell Electron Pair Repulsion (VSEPR) Theory states that bonds and lone pairs are

regions of high electron density in an atom that repel each other until they get as far apart as

possible. This effect determines the atom’s three-dimensional geometry and bond angles. See

table 1 for more details. In order to get correct answers from VSEPR theory, you must first have a

correct Lewis structure. Lewis structures help us to organize the valence electrons in a structure as

bonding and lone pairs on each of the atoms. Knowing how many electron pairs are arranged

around a central atom allows us to predict the geometry of the molecule. In this lab, you will create

several models of molecules using a Ball-and-stick models which are often used to demonstrate

molecular shape (painted wooden balls represent atoms, and wooden sticks represent bonds).

Table 1. Molecular Geometry based on VSEPR Model

Number of atoms bound to central atom

Number of lone pairs of electron around

central atomGeometry Bond angles

2 0 Linear 180°

3 0 Trigonal Planar 120°

4 0 Tetrahedral 109.5°

5 0 Trigonal Bipyramidal 120° (equatorial) & 90° (axial)

6 0 Octahedral 90°


2 1 Bent (or) ‘V’-shaped 104.5°

2 2 Bent (tetrahedral) 104.5°

3 1 Trigonal Pyramidal 107.5°

VSEPR diagrams

1) Linear 2) Trigonal Planar 3) Tetrahedral 4) Trigonal Bipyramidal

5) Octahedral 6) Bent 7) Bent 8) Trigonal pyramidal

Procedure

1. For each of the molecules, determine the number of lone pairs and bonded pairs around

the central atom by drawing the Lewis Dot Structure.

2. Use the model kit to build the structure. Arrange the atoms to maximize the distance

between all the electron pairs. Remember that lone pair electrons cause more repulsion

than electrons between atoms.

Use yellow (or any colour available) balls for hydrogens.

Use black balls for other atoms that desire an octet.

Use short sticks for non-bonded electron lone pairs.

Use long sticks for single bonds.

3. Describe the structure you have created as linear, V-shaped, trigonal planar, trigonal

pyramidal tetrahedral, square planer, trigonal bipyramidal, or octahedral.

4. Estimate the angle between the atoms attached to the central atom.


.... ••

:

Observations (Complete the following table for the indicated species)

Molecule Lewis Structure

# of atoms bonded to

central atom

# of lone pairs on central atom

VSEPR Diagram

BeCl2

BF3

NH3

SF6

H2O

PCl5

CH4


Questions & Answers

1. What does VSEPR stand for?

VSEPR stands for Valence Shell Electron Pair Repulsion.

2. Why it is important to know the shape of molecule?

Explaining the shapes of molecules is important in understanding how molecules react.

3. When you made molecules, it was important to fill all of the open holes. What would an

.....open hole in a molecule represent?

An area of electron density may be:

2. a lone (nonbonding) pair

3. a single bond

4. a double bond

5. a triple bond

4. List three of the models you built that were molecules.

Linear:

V-shaped:

Trigonal Planar:

5. Draw a Lewis Dot Diagram for ammonia. Using your Lewis dot diagram, explain why your

….molecular model for ammonia looked like it did. Use the terms “paired electrons” and

….“unpaired electrons” in your answer.

See textbook (Chemistry by Zumdahl, 8th edition)

6. What effect does the presence of lone-pair electrons have on the bond angles in a

….molecule?

The lone-pair electrons in a molecule reduce the bond angles.


Experiment: 7 Date:

MEASUREMENT OF MELTING POINT AND BIOLING POINT

Objectives

To determine the melting temperature of a solid, and melting temperature of a liquid.


Experiment: 8 Date:

MEASUREMENT OF DENSITY

Objectives

To determine the density of unknown samples (solid and liquid) at room temperature.

By the end of this experiment, the student should be able to understand:

- The relationship between mass, volume and density.

- The relationship between experimental measurements, precision and significant figures.

- Use of laboratory equipments, such as electric balance, graduated cylinder, pipette and

Vernier Caliper.

Introduction

Substances (i.e. molecules) have physical properties of mass, volume and density. Density is

defined as the ratio of the mass divided by the volume. Typically, densities are reported in units of

g/cm3 (for solids) or g/ml (for liquids). Please note that 1 ml = 1 cm3.

Density = Mass / Volume (or) d = m/V

Measurement of density

1. Density of a solid is determined by weighing a sample of the solid and then measuring the

volume it occupies.

The volume of a solid can be determined by 2 methods.

- By measuring the dimensions of the solid and calculating its volume.

Example: A solid cylinder has a length of 3.270 cm and a diameter of 1.345 cm.

Volume of the solid cylinder = l (length) x x d2 (diameter) / 4

= 3.270 x 3.141 x 1.3452 / 4

= 18.358 / 4

= 4.645 cm3

- By measuring the volume of liquid that is displaced by the solid.

Example: A lump of solid is submerged in water in a 50 ml graduated cylinder. The

water level was 25.5 ml before adding the solid; and 30.0 ml after adding the solid.


The volume of the solid = 30.0 – 25.5

= 4.5 ml (or) 4.5 cm3

2. Density of a liquid is determined by weighing a known volume of the liquid. A

VOLUMETRIC FLASK is a glass container which holds a defined volume of liquid. The

mass of the liquid is determined by weighing the volumetric flask first empty and then fill of

the liquid.

Weight of liquid = Weight of empty flask – Weight of flask with liquid

3. The density of solid or liquid can be calculated from the formula: Density = Mass / Volume.

4. Because the volume (as well as the density) of a substance changes with its temperature, it

is necessary to keep a record of the temperature at which densities are measured.

Note:

1. With any equipment (balance, pipette, etc.), properly

reading the scale or mark is essential, if

the full precision of the device is to be obtained

(precision is defined as the extent to which a given

set of measurements of the same sample agree

with their mean. Example. 20.16 ± 0.01).

2. For aqueous solutions, the meniscus (the concave

surface of the liquid in the glass) should be read

at the bottom of the curve.

Also, the reading should be performed with the

eye at the level of the meniscus.


This reads 43.0 ml (Not 43.5)

Write it as: 43.0 ± 0.02

Eye level

It reads

20.15 ml or 20.16 ml or 20.17 ml

Write it as:20.16 ± 0.01

3. The volume of a solid with cylindrical shape is calculated from its length and diameter,

which can be measured with the help of Vernier Caliper.


Procedure

1. Determining the density of a cylindrical solid.

a. Obtain an unknown solid and vernier calipers from your instructor.

b. Weigh the solid on the electric balance.

c. Measure the solid dimension (length and diameter), using the calipers.

d. Calculate the volume of the solid (V = l x x d2 / 4; = 3.141).

e. Calculate the density of the unknown solid (Density = Mass / Volume).

f. Repeat the above procedure once again. Calculate the average density.

g. Use the thermometer to measure the room temperature and record it.

Trial 1 Trial 2

Mass of solid* ± ±

Diameter* ± ±

Length* ± ±

Volume of solid

(V = l x 3.141 x d2 / 4)

Density of solid

(Density = Mass / Volume)

Average density

= (Density 1 + Density 2) / 2

Room temperature* ±

*The precision is required for mass, diameter, length and temperature.


2. Determining the density of an irregular solid.

a. Obtain an unknown solid of irregular shape.

b. Weigh the solid on the electric balance.

c. Pour distilled water into a 100 ml graduated cylinder to approximately 50 ml mark. Note

down the level of the water accurately (V1). Always read the bottom of the meniscus.

d. Hold the cylinder (with distilled water) at 45° angle and slowly slide the unknown solid into

the water. Try to avoid having air bubbles trapped by the solid. Note down the level of the

water accurately (V2).

e. Calculate the volume of the solid (V = V2 – V1).

f. Calculate the density of the unknown solid (Density = Mass / Volume).

g. Repeat the above procedure once again. Calculate the average density.

Trial 1 Trial 2

Mass of solid* ± ±

Volume of water without solid (V1)* ± ±

Volume of water with solid (V2)* ± ±

Volume of solid

(V = V2 – V1)

Density of solid


Average density

= (Density 1 + Density 2) / 2


*The precision is required for measured mass, volume and temperature.


3. Determining the density of liquid.

a. Obtain an unknown liquid from your instructor.

b. Clean the 10 ml volumetric flask with acetone in a fume hood and leave it for five minutes to

dry.

c. Weigh the dry volumetric flask (with stopper). Consider it as m1.

d. Fill the volumetric flask with the unknown liquid exactly up to the mark.

e. Weigh the flask (with unknown liquid and stopper). Consider it as m2.

f. Calculate the mass of the liquid (Mass= m2 – m1).

g. Calculate the density of the unknown solid (Density = Mass / Volume).

h. Repeat the above procedure once again. Calculate the average density.

Observation

Mass of flask + stopper (m1)* ±

Mass of flask + stopper + Liquid (m2)* ±

Mass of liquid (m2 – m1)

Volume of liquid* ±

Density of Liquid



*The precision is required for measured mass, volume and temperature.


general chemistry lab manual

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