Types of Bonding in Solids Lab
Ionic, Covalent, and Metallic Bonding
(Chapter 6)
Name: __________________
Period: _________________
Date: ___________________
Objective
To compare melting points, solubility, and electrical conductivity of ionic, covalent (molecular), and metallic solids.
To relate these macroscopic, observable properties to the microscopic behavior of the atoms/ions in each bonding type.
To identify an unknown solid based on these properties.
Background Information
Chemical bonds are the forces which bind atoms together. We will examine 3 types of bonding in solids:
(1) Ionic-results from electrical attractions between + (metal) and –(nonmetal) ions. Electrons are “stolen,” not shared.
(2) Covalent-occurs in molecular compounds, and results from the sharing of electrons between nonmetal atoms.
(3) Metallic-results from attraction between atoms & “seas of electrons” that don’t belong to any single metal atom
Pre-Lab Questions
1. Sketch what an ionic bond between the elements sodium and chlorine would look like.
2. Sketch what a covalent bond between one carbon atom and four hydrogen atoms would look like.
3. Sketch what the “sea of electrons” might look like for a metal such as copper.
Safety
1.
2.
3.
4.
5.
Wear ___________________ and apron at all times in the laboratory.
Tie back __________ hair.
Never leave Bunsen burner ________________ unattended.
Never aim opening of a test tube at a person while _________________.
Allow glassware to ____________ before washing.
Procedure
Part 1-Solid Conductivity
1. Locate the solids your teacher has already prepared for you in spot plates at the lab station labeled Part 1.
2. Use the key provided to identify solids #1-7 in Table 1.
3. Use the conductivity meter as demonstrated to determine whether each solid conducts electricity.
4. Note which solids do and do not conduct electricity in Table 1.
5. After each test, rinse the conductivity meter with distilled water (squirt bottle) & carefully dry the electrodes.
6. Return all supplies in clean and dry condition for other students to use.
Parts 2 & 3 Solubility & Aqueous Conductivity
1. Locate Lab Station Parts 2 & 3 where your teacher has already mixed each solid with distilled water for you.
2. Use the key provided to identify the solids/solutions #1-7 in Table 1.
3. For Part 2, Solubility, note which solids dissolved in water and which did not in Table 1.
4. For Part 3, submerge meter in liquid portion of each to determine which solutions conduct electricity.
5. After each test, rinse the conductivity meter with distilled water (squirt bottle) & carefully dry electrodes.
6. Return all supplies in clean and dry condition for other students to use.
7. Note which solutions do and do not conduct electricity in Table 1.
Part 4-Melting Point
1. Label seven medium/large test tubes with the chemical names shown in Table 1 below.
2. Obtain one small piece of copper and one small piece of aluminum, and place them in separate, labeled test tubes.
3. Use a pea-sized amount of the remaining solid chemicals, and place them in their labeled test tubes.
SAFETY IS VERY IMPORTANT IN THE NEXT 4 STEPS-READ #4-7 CAREFULLY BEFORE BEGINNING!
4. When heating each solid in the Bunsen burner flame, NEVER AIM THE OPENING OF THE TEST TUBE AT A PERSON
and IMMEDIATELY REMOVE the test tube from the flame when the solid inside it begins to melt.
5. If a solid does not melt after thirty seconds, remove it from the flame and label its melting point as “High”.
6. Place the hot test tubes in an empty glass beaker to cool before cleaning/disposing.
7. After they have cooled, dispose of the sucrose and lactose test tubes in the broken glassware container.
8. The remaining solids can be disposed of in the trash can when they have cooled. Clean all test tubes.
Table 1: Comparison of Ionic, Covalent, and Metallic Bonding in Solids.
#
Solid
Type of
Bonding
1
copper
Metallic
2
sucrose
Covalent
3
sodium chloride
Ionic
4
lactose
Covalent
5
aluminum
Metallic
6
calcium chloride
Ionic
7
unknown
Observations
(1) Solid
Conductivity
(yes/no)
(2) Soluble in
Water?
(yes/no)
(3) Aqueous
Conductivity
(yes/no)
(4) Melting
Point
(Low/High)
Post-Lab Questions (Answer the following questions below or on a separate sheet of paper stapled to this handout.)
1. Ionic solids have low/high (pick one) melting points, are/are not (pick one) soluble in water, do/do not (pick one)
conduct electricity in their solid forms, and do/do not (pick one) conduct electricity when dissolved in water.
2. Covalent solids have low/high (pick one) melting points, are/are not (pick one) soluble in water, do/do not (pick one)
conduct electricity in their solid forms, and do/do not (pick one) conduct electricity when dissolved in water.
3. Metallic solids have low/high (pick one) melting points, are/are not (pick one) soluble in water, do/do not (pick one)
conduct electricity in their solid forms, and do/do not (pick one) conduct electricity when dissolved in water.
4. Identify your unknown from solids 1-6 above. Justify your answer with evidence from your experiment.
5. Why was it important to use distilled water instead of tap water to rinse the conductivity meters?
6. Sketch what an ionic bond between one calcium ion and two chloride (chlorine) ions might look like (hint-see
example #1 we did together in class).
7. What is the octet rule, and how is useful in predicting chemical bonding?
8. When cations & anions join, they form what kind of chemical bond?
(a.) ionic
(b.) hydrogen
(c.) metallic
(d.) covalent
9. Some of the molecules found in the human body are NH2CH2COOH (glycine), C6H12O6 (glucose), and CH3(CH2)16COOH
(stearic acid). The bonds they form are
(a.) nuclear
(b.) metallic
(c.) ionic
(d.)covalent
10. What type of bond do all of the following molecules have in common: H2, Cl2, NH3 (ammonia), and CH4 (methane)?
(a.) covalent
(b.) ionic
(c.) metallic
(d.) polar
11. What type of force holds ions together in crystals such as calcium fluoride?
(a.) electrostatic
(b.) magnetic
(c.) gravitational
(d.) nuclear
12. The reason salt crystals, such as KCl, hold together so well is because the cations are strongly attracted to
(a.) neighboring cations (b.) protons in the neighboring nucleus (c.) free electrons in crystals (d.) neighboring anions.