Unit 2: Compound Structure and Properties
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Essential Knowledge
◼ Electronegativity values for the representative elements increase going from left to right across a period and decrease going down a group. These trends can be understood qualitatively through the electronic structure of the atoms, the shell model, and Coulomb’s law. (SPQ-1.A.1) Go to Section
◼ Valence electrons shared between atoms of similar electronegativity constitute a nonpolar covalent bond. For example, bonds between carbon and hydrogen are effectively nonpolar even though carbon is slightly more electronegative than hydrogen. (SPQ-1.A.1) Go to Section
◼ Valence electrons shared between atoms of unequal electronegativity constitute a polar covalent bond.
a) The atom with the higher electronegativity will develop a partial negative charge relative to the other atom in the bond.
b) In single bonds, greater differences in electronegativity lead to greater bond dipoles.
c) All polar bonds have some ionic character, and the difference between ionic and covalent bonding is not distinct but rather a continuum. (SAP-3.A.3) Go to Section
◼ The difference in electronegativity is not the only factor in determining if a bond should be designated as ionic or covalent. Generally, bonds between a metal and nonmetal are ionic, and bonds between two nonmetals are covalent. Examination of the properties of a compound is the best way to characterize the type of bonding. (SAP-3.A.4) Go to Section
◼ In a metallic solid, the valence electrons from the metal atoms are considered to be delocalized and not associated with any individual atom. (SAP-3.A.5) Go to Section
◼ Electronegativity values for the representative elements increase going from left to right across a period and decrease going down a group. These trends can be understood qualitatively through the electronic structure of the atoms, the shell model, and Coulomb’s law. (SPQ-1.A.1)
Rank the following set of elements in order of increasing electronegativity: Se, Br, As, Ga
Electronegativity is a measure of an atom's ability to attract electrons towards itself in a chemical bond.
General trends in electronegativity:
Increases across a period (left to right): Atoms with more protons have a stronger pull on electrons.
Decreases down a group: Increasing atomic size means outer electrons are farther from the nucleus, reducing attraction.
Analyzing the Elements
Let's locate these elements on the periodic table:
Se (Selenium) is in Period 4, Group 16.
Br (Bromine) is in Period 4, Group 17.
As (Arsenic) is in Period 4, Group 15.
Ga (Gallium) is in Period 4, Group 13.
Ranking based on the periodic trends:
Ga (Gallium) has the lowest electronegativity as it's furthest to the left in the period.
As (Arsenic) is to the right of Ga, so it has a higher electronegativity.
Se (Selenium) is to the right of As, so it has a higher electronegativity than both Ga and As.
Br (Bromine) is to the right of Se, making it the most electronegative element in the group.
Final Order
Therefore, the elements in order of increasing electronegativity are: Ga < As < Se < Br
Remember: While these trends are generally reliable, there can be exceptions, especially with transition metals. However, for the main group elements in this question, the trend holds true.
Try it yourself …
The table below gives information about the dipole moment for the H–Br bond. Do you predict that the dipole moment for the H–Cl bond should be less than 0.82 D or greater than 0.82 D? Justify your answer based on periodic trends in electronegativity.
| Bond | Dipole moment (D) |
|---|---|
| H--Cl | ? |
| H--Br | 0.82 |
Check your answer here.
◼ Valence electrons shared between atoms of similar electronegativity constitute a nonpolar covalent bond. For example, bonds between carbon and hydrogen are effectively nonpolar even though carbon is slightly more electronegative than hydrogen. (SPQ-1.A.1)
◼ Valence electrons shared between atoms of unequal electronegativity constitute a polar covalent bond.
a) The atom with the higher electronegativity will develop a partial negative charge relative to the other atom in the bond.
b) In single bonds, greater differences in electronegativity lead to greater bond dipoles.
c) All polar bonds have some ionic character, and the difference between ionic and covalent bonding is not distinct but rather a continuum. (SAP-3.A.3)
Place the following bonds in order of increasing polarity. C-H, C-F, C-C, C-O, Ca-C
A polar bond is a covalent bond between two atoms with different electronegativities. The more electronegative atom pulls the shared electrons closer to itself, creating a partial negative charge on that atom and a partial positive charge on the other. The greater the difference in electronegativity, the more polar the bond.
Analyzing the Bonds
To rank the bonds by increasing polarity, we need to compare the electronegativity differences between the atoms in each bond.
C-H: Carbon and hydrogen have similar electronegativities, so this bond is essentially nonpolar.
C-F: Fluorine is the most electronegative element, so this bond is highly polar.
C-C: Both atoms are carbon, so there is no electronegativity difference, making the bond nonpolar.
C-O: Oxygen is more electronegative than carbon, making this bond polar.
Ca-C: Calcium is a metal and carbon is a nonmetal. This is an ionic bond, not a covalent bond. However, if we were to consider it as a covalent bond for comparison purposes, the electronegativity difference would be quite large, making it highly polar.
Ranking the Bonds
Based on the electronegativity differences:
C-C: Nonpolar
C-H: Slightly polar
C-O: Polar
C-F: Highly polar
Ca-C: Ionic (or highly polar if considered covalent)
Therefore, the order of increasing polarity is: C-C < C-H < C-O < C-F < Ca-C
Try it yourself …
Arrange each of the following bonds in order of increasing polarity: C—P, P—F, and C—Cl.
Check your answer here.
◼ The difference in electronegativity is not the only factor in determining if a bond should be designated as ionic or covalent. Generally, bonds between a metal and nonmetal are ionic, and bonds between two nonmetals are covalent. Examination of the properties of a compound is the best way to characterize the type of bonding. (SAP-3.A.4)
Classify the substances the follows into either: (1) nonpolar covalent, (2) polar covalent, (3) ionic, or (4) metallic
H2, NaF, ZnCl2, NO, CuZn, NCl3, CH4, Al
1. $\text{H}_2$ (Hydrogen gas):
Bond Type: Covalent
Electronegativity Difference: The two hydrogen atoms have the same electronegativity (2.1).
Classification: Nonpolar Covalent
Since the two atoms are identical, the electrons are shared equally, resulting in a nonpolar covalent bond.
2. NaF (Sodium fluoride):
Bond Type: Ionic
Electronegativity Difference: Sodium (0.9) and fluorine (4.0) have a large difference in electronegativity (~3.1).
Classification: Ionic
The large difference in electronegativity causes the sodium atom to lose an electron, forming $\text{Na}^+$, and the fluorine atom to gain an electron, forming $\text{F}^-$, leading to an ionic bond.
3. ZnCl2\text{ZnCl}_2ZnCl2 (Zinc chloride):
Bond Type: Covalent (but with significant ionic character)
Electronegativity Difference: Zinc (1.6) and chlorine (3.0) have a moderate difference in electronegativity (~1.4).
Classification: Polar Covalent
While the bond has some ionic character, the difference in electronegativity suggests a polar covalent bond.
4. NO (Nitric oxide):
Bond Type: Covalent
Electronegativity Difference: Nitrogen (3.0) and oxygen (3.5) have a small difference in electronegativity (~0.5).
Classification: Polar Covalent
The slight difference in electronegativity between nitrogen and oxygen results in a polar covalent bond.
5. CuZn (Brass - an alloy of copper and zinc):
Bond Type: Metallic
Electronegativity Difference: Both copper and zinc are metals.
Classification: Metallic
In metallic bonds, atoms share a "sea of electrons" that are free to move, which is characteristic of metallic bonding.
6. $\text{NCl}_3$ (Nitrogen trichloride):
Bond Type: Covalent
Electronegativity Difference: Nitrogen (3.0) and chlorine (3.0) have a small difference in electronegativity (~0.5).
Classification: Polar Covalent
Even though the difference in electronegativity is small, NCl3\text{NCl}_3NCl3 has a polar covalent bond due to the asymmetric distribution of electrons.
7. $\text{CH}_4$ (Methane):
Bond Type: Covalent
Electronegativity Difference: Carbon (2.5) and hydrogen (2.1) have a very small difference in electronegativity (~0.4).
Classification: Nonpolar Covalent
The small difference in electronegativity and symmetrical distribution of bonds make CH4\text{CH}_4CH4 nonpolar covalent.
8. Al (Aluminum):
Bond Type: Metallic
Electronegativity Difference: Aluminum is a metal.
Classification: Metallic
Aluminum atoms are held together by metallic bonds, characterized by a "sea of electrons."
Summary of Classifications:
$\text{H}_2$: Nonpolar Covalent
NaF: Ionic
$\text{ZnCl}_2$: Polar Covalent
NO: Polar Covalent
CuZn: Metallic
$\text{NCl}_3$: Polar Covalent
$\text{CH}_4$: Nonpolar Covalent
Al: Metallic
Try it yourself …
Here is a situation in which examination of the properties of substance is the best way to characterize the type of bonding present in that substance.
(A) Tin (Sn) is classified as a (pick one): metal non-metal
(B) Chlorine (Cl) is classified as a (pick one): metal non-metal
(C) What type of bonding, ionic or covalent, would you predict for a compound that contains tin (Sn) and chlorine (Cl)?
(D) Consider the properties of the substances SnCl2 and SnCl4 listed in the table below. This information should help you to predict the type of bonding that is present in each substance.
| SnCl2 | SnCl4 | |
|---|---|---|
| Appearance | white crystalline solid | colorless liquid |
| Melting Point | 247°C | - 34°C |
| Boiling Point | 623°C | 114°C |
| Type of Bonding |
◼ In a metallic solid, the valence electrons from the metal atoms are considered to be delocalized and not associated with any individual atom. (SAP-3.A.5)
(Multiple Choice) What does it mean that valence electrons in a metal are delocalized?
(A). The valence electrons move between atoms in shared orbitals.
(B). The valence electrons move from outer orbitals to inner orbitals of each atom.
(C). The valence electrons move from inner orbitals to outer orbitals of each atom.
(D). The valence electrons move out of the orbitals and go into the air.
The answer is choice A. The valence electrons move between atoms in shared orbitals. The result of millions of metal atoms crowding together so that molecular orbitals become combined most likely is the formation of localized valence electrons where the electrons are called the "electron sea". This is a characteristic of compounds that are held together by metallic bonding.
Try it yourself …
Consider a sample of pure copper (Cu) metal. Copper is known to form a metallic solid where the valence electrons are delocalized.
Part A . Describe the nature of bonding in the copper metal. How do the delocalized electrons contribute to the properties of copper, such as electrical conductivity and malleability?
Part B. Predict how the delocalization of electrons would affect the melting point of copper. Compare this with a molecular solid (like ice) and explain the difference in melting points based on the nature of bonding.
Part C. Explain why copper is a good conductor of electricity at room temperature. How would this ability change as the temperature increases?