Week 2, Lab 1: 9/15: Atoms



Week 4, Section 1: 9/28: Types of Bonds

1) Polar

2) Ionic

3) Metallic

4) Bonding Triangle

Reading: 8.7 p- 354-357; 9.5 p405-409; 3.2 p 81-82; 3.5 p 88-99

New Group Formation

Introduce yourselves to the group. Tell everyone three things about yourself (such as favorite place to travel, what you did last summer, your favorite movie, your favorite meal to fix, what you like to do on the weekends, what you aim to be when you “grow up,” etc.) Exchange information such as e-mail/phone/address.

Then share responses to the following questions:

1. What do you perceive to be your responsibilities to the group (list these on the board or on paper)

2. What are the responsibilities the group has to each member

3. Describe the advantages of working in a group or as a team

4. Describe the disadvantages of working in a group or as a team

Draft a Group Covenant based on your discussion of the questions above which outlines each group members' responsibilities to the group and the group responsibilities to each member. (3-5 sentences.) Write/print out a copy of the covenant and have every group member sign it and give a copy to your GSI. You may want to keep a copy for yourselves.

Group assignments for today: A Recorder; B Skeptic; C Leader

1) Bonds, Polar Bonds (In conjunction with Powerpoint whole section discussion/demonstration)

For a molecule to be polar it must:

a) have at least 1 polar bond or 1 lone pair and

b) polar bonds or LP must be arranged in space so that the dipoles do not cancel out

Fill in each box with:

# valence electrons

Lewis structure

Molecular structure

Dipole arrow

|HCl |CO2 |NH3 |SF4 |

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|C2H2 |BF3 |CHCl3 |ICl3 |

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|XeF2 |XeF4 |SF6 |BrF5 |

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2) Bonds, Ionic Bonds

Does the difference in electronegativity between atoms ever get so large that the more electronegative atom takes the electrons away from the less electronegative atom?

The result is an ionic bond: the atoms involved undergo a transfer of one (or more electrons) to produce two charged species: a positive ion (cation) and a negative ion (anion). Remember from studying periodic trends, you can predict which elements form cations and which form anions and what their charges are.

There are also many polyatomic ions. They have covalent bonds between the atoms of the ion, but form ionic complexes with other ions. For example NaNO3 has an ionic bond between Na+ and NO3-, and covalent bonding between nitrogen and oxygen.

Table 1: Polyatomic cations and anions

|Cation | |Anions | | | |

|NH4+ |Ammonium |OH- |hydroxide |PO4-2 |phosphate |

|H3O+ |Hydronium |CN- |cyanide |MnO4-2 |permanganate |

| | |CO3-2 |carbonate |CrO4-2 |chromate |

| | |CH3COO- |acetate |Cr2O7- |dichromate |

| | |SO4-2 |sulfate |ClO4- |perchlorate |

| | |SO3-2 |sulfite |ClO3- |chlorate |

| | |NO3- |nitrate |ClO2- |chlorite |

| | |NO2- |nitrite |ClO- |hypochlorite |

You should commit the names, molecular formula, and charges of the ions in Table1 above to memory by next week!

Question: If ionic bonds are created when atoms undergo a transfer of electrons to form cations and anions, how do they form bonds?

Data/Fact Gathering: There is a force, described by Coulomb’s Law which exists between charges.

The force is attractive between positive charges and negative charges and repulsive between multiple positive charges, and multiple negative charges. The magnitude of the forces increases as the magnitude of the charges increases. The force decreases with the distance between the charges.

[pic]

A bond distance is set when the electron clouds of the atoms overlap and the attractive forces between the oppositely charged ions is balanced by the repulsive forces between the two nuclei and the electrons.

Ions form an ordered solid as shown in Figure 1 so that symmetry causes the attractive and repulsive forces to balance each other and the energy to be minimized.

Figure 1: a) A model of the distribution of charges in an ionic compound; b) Ionic compound upon deformation

a) [pic] b) [pic]

1) Ionic compounds tend to have very high melting point. Why is this?

2) Ionic materials are generally brittle, that is they break suddenly with little deformation before breaking. Why is this?

3) Bonds, Metallic Bonds

Metals are not modeled like covalent or ionic bonds. In metals none of the atoms exert a strong pull on the electron, the atoms become positive ion cores surrounded by a sea of electrons. The electrons act like a glue holding the ions together. These ions still create a lattice structure.

Figure 2: a) Model for metallic bonding; b) metal after deformation

a) [pic] b) [pic]

Metals are very good conductors because the free electrons can move through a circuit causing a current. When a metal is deformed, the metallic crystal structure is able to accommodate greater deformation than an ionic crystal structure. This property of metals is called ductility.

Question: So how do you know whether a bond is covalent, ionic, or metallic?

Data/Fact Gathering

Table 2: Electronegativity, EN parameters and melting points for selected compounds[1]

|Compound |EN |EN |___ |ΔEN |Melting point (oC) |Type of bonding |

| |1st atom |2nd atom |EN | | | |

|CsF(s) |0.66 |4.19 |2.42 |3.53 |682 |ionic |

|NaCl(s) |0.87 |2.87 |1.87 |2.00 |801 |ionic |

|NaI(s) |0.87 |2.36 |1.62 |1.49 |661 |ionic |

| | | | | | | |

|Cs(s) |0.66 |0.66 |0.66 |0 |28 |metallic |

|Na(s) |0.87 |0.87 |0.87 |0 |98 |metallic |

|CuZn(s; brass) |1.8 |1.8 |1.7 |0.2 |932 |metallic |

| | | | | | | |

|F2(g) |4.19 |4.19 |4.19 |0 |-220 |covalent |

|O2(g) |3.61 |3.61 |3.61 |0 |-218 |covalent |

|HI(g) |2.30 |2.36 |2.33 |0.06 |-51 |covalent |

|C(s; diamond) |2.54 |2.54 |2.54 |0 |>3000 |covalent |

| | | | | | | |

|CH4(g) |2.54 |2.30 | | |-182 | |

|GaAs(s) |1.76 |2.21 | | |1238 | |

|Si(s) |1.92 |1.92 |1.92 |0 |1410 | |

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EN is the average electronegativity of the two elements

ΔEN is the absolute difference in electronegativity between the two elements

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3) Based on the data in Table 2, what combination of ΔEN and EN values leads to metallic bonding?

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4) Based on the data in Table 2, what combination of ΔEN and EN values leads to ionic bonding?

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5) Based on the data in Table 2, what combination of ΔEN and EN values leads to covalent bonding?

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6) Fill in the ΔEN and the EN values for methane (CH4) in Table 1. Classify the C-H bonds in methane as metallic, ionic, or covalent bonding? Explain.

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7) Fill in the ΔEN and the EN values for GaAs in Table 1. Is it possible to classify GaAs as metallic, ionic, or covalent bonding? Explain.

Model 1: Bond-type Triangle

Many compounds have properties that are intermediate between the three bond types. Si, for example is known as a semiconductor; a compound with properties which are intermediate between metallic and covalent.

The bond-type triangle is a model that enables one to predict the properties of a compound based on the electronegativities of the elements that comprise the compound. The data for CsF, F2, and Cs from Table 1 have been used to generate three points at the corners of the bond-type triangle shown in Figure 3.

Figure 3: Bond-type Triangle

[pic]

8) Place points on the chart for sodium chloride (ionic bonding), sodium (a metal), hydrogen iodide (covalent bonding), and Si (a semiconductor.)

The bond-type triangle can be divided into regions which indicate the predominant type of bonding present in compounds. The dividing lines between regions are not absolute, but they give a general idea of the boundaries between different types of bonding.

9) Associate the regions (A, B, C) with bond types (metallic, covalent, ionic). The “SM” region is sometimes called semimetals.

A

B

C

10) Quartz, SiO2 is a very high melting, hard solid. Place a point for SiO2 on the bond-type triangle. What type of bonding would you predict to be predominant in quartz?

11) For each of the following compounds, place a point on the bond-type triangle. Classify each compound as metallic, covalent, ionic, semimetal.

a) CO2

b) NH3

c) BaO

d) SO2

e) AlSb

f) GaAs

g) CdLi

h) BaBr2

i) ZnO

j) NaH

12) Give the type of bonding in each of the following: CO3-2; BaCO3(s); NaClO4(s)

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[1] Moog, R. S.; Farrell, J. J. “Chemistry A Guided Inquiry” 2nd edition John Wiley & Sons 2002, p 131-135.

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q1, q2 -are the magnitudes of the two charges

r -is the distance between them

k -is a constant

Cs

F2

CsF

SM

B

C

A

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