Chemical bonding: flashcards, quiz and summary

Flashcards

1. 1Why do atoms form chemical bonds at all?

   Atoms bond by transferring, sharing, or pooling electrons to reach a lower-energy, more stable arrangement, often a full outer shell of 8 valence electrons.

2. 2How does an ionic bond form, and why does the compound form a lattice?

   A metal loses electrons to become a cation and a nonmetal gains them to become an anion. The opposite charges attract, and the ions pack into a crystal lattice rather than forming separate molecules.

3. 3How does electronegativity difference determine the type of bond?

   A large difference (about 1.7 or more, depending on the textbook) gives an ionic bond. A small difference (under about 0.4) gives a nonpolar covalent bond. Differences in between give polar covalent bonds, where electrons are shared unequally.

4. 4Why can a molecule with polar bonds be nonpolar?

   If the shape is symmetric, the bond dipoles point in opposite directions and cancel. CO2 is linear, so its dipoles cancel and it is nonpolar, even though each C=O bond is polar.

5. 5Why do metals conduct electricity and still bend without breaking?

   Metal atoms release valence electrons into a shared sea around a lattice of positive ions. The mobile electrons carry charge and heat, and the layers of ions can slide without breaking the bonding.

6. 6How does VSEPR predict the shape of a molecule?

   Electron groups around the central atom spread as far apart as possible. Multiple bonds count as one group, and lone pairs push harder than bonding pairs, which changes the angle but is not named in the shape.

7. 7Why does water boil at 100 °C while H2S boils near -60 °C, even though H2S is larger?

   Water has hydrogen bonding, which is much stronger than the dipole-dipole and London dispersion forces in H2S. More energy is needed to separate water molecules, so its boiling point is much higher.

8. 8Why does boiling water not break the O-H bonds inside the molecules?

   Boiling overcomes the intermolecular forces between molecules. The covalent O-H bonds inside each water molecule stay intact, so the water is still H2O afterward.

Practice quiz

Pick an answer to check it.

1. 1Which pair of elements is most likely to form an ionic bond?

   1. Hydrogen and nitrogen
   2. Carbon and oxygen
   3. Chlorine and chlorine
   4. Sodium and chlorine
   Show the answer

   D. Sodium and chlorine Sodium is a metal and chlorine is a nonmetal with a large electronegativity difference, so sodium transfers an electron to chlorine and the ions attract. Carbon and oxygen are both nonmetals, so they share electrons covalently.

2. 2CO2 has polar C=O bonds, but the molecule is nonpolar. What explains this?

   1. The bonds are actually nonpolar
   2. The molecule is linear, so the bond dipoles cancel
   3. London dispersion forces cancel the polarity
   4. Carbon has no lone pairs, so the molecule has no charge
   Show the answer

   B. The molecule is linear, so the bond dipoles cancel CO2 is linear, so its two polar bond dipoles point in opposite directions and cancel. Polar bonds do not guarantee a polar molecule; the shape must be checked.

3. 3Which molecule has a bent shape with about 104.5° bond angles?

   1. NH3
   2. CH4
   3. BF3
   4. H2O
   Show the answer

   D. H2O H2O has 4 electron groups around oxygen, 2 of them lone pairs, which gives a bent shape of about 104.5°. NH3 has one lone pair and is trigonal pyramidal, and CH4 has no lone pairs and is tetrahedral.

4. 4Why does NH3 have stronger intermolecular attractions than CH4 of similar size?

   1. CH4 has double bonds that attract neighboring molecules
   2. NH3 has a higher molar mass and therefore only London dispersion forces
   3. NH3 has hydrogen bonded directly to nitrogen, which allows hydrogen bonding
   4. CH4 has more London dispersion forces than NH3
   Show the answer

   C. NH3 has hydrogen bonded directly to nitrogen, which allows hydrogen bonding In NH3, hydrogen is bonded directly to nitrogen, so hydrogen bonding can occur with lone pairs on other molecules. CH4 has its hydrogens bonded to carbon, so it cannot hydrogen bond and only has London dispersion forces.

5. 5Why are solid ionic compounds poor conductors of electricity, while molten ionic compounds conduct well?

   1. In the solid, the ions are fixed in the lattice and cannot move to carry charge
   2. Solid ionic compounds have no ions at all
   3. Ionic compounds conduct only when they are gases
   4. Melting creates free electrons from the metal atoms
   Show the answer

   A. In the solid, the ions are fixed in the lattice and cannot move to carry charge In a solid ionic compound, the ions are locked in place in the crystal lattice, so charge cannot flow. When melted or dissolved, the ions are free to move and carry charge.

6. 6Which bonding situation is described by metal atoms releasing valence electrons into a shared sea around positive ions?

   1. Metallic bonding
   2. Polar covalent bonding
   3. Hydrogen bonding
   4. Ionic bonding
   Show the answer

   A. Metallic bonding In metallic bonding, metal atoms release valence electrons that move freely around a lattice of positive ions. This mobile electron sea explains conductivity and malleability.

Test traps

What teachers like to test on this, and the mistakes students make.

• Confusing boiling with breaking bonds: boiling water overcomes the forces between molecules, but the O-H covalent bonds inside each molecule remain intact.
• Assuming polar bonds always make a polar molecule. Students often forget to check the shape, and CO2 is the classic example of symmetric cancellation.
• Saying hydrogen bonding happens with any hydrogen. It only occurs when H is bonded directly to N, O, or F, so CH4 cannot form hydrogen bonds.
• Calling an ionic solid a conductor. Solid ionic compounds do not conduct because the ions cannot move; they conduct when melted or dissolved.
• Putting hydrogen in the center of a Lewis structure. Hydrogen is never the central atom and only needs 2 electrons.
• Forgetting that lone pairs count as electron groups for VSEPR, so NH3 is trigonal pyramidal and H2O is bent even though lone pairs are not named in the shape.

Summary

• Atoms bond by transferring, sharing, or pooling electrons so that they reach a lower-energy, more stable arrangement, often with 8 valence electrons.
• Ionic bonds form between a metal and a nonmetal when electrons transfer, and the resulting oppositely charged ions attract and pack into a crystal lattice.
• Covalent bonds form when nonmetals share pairs of electrons, and unequal sharing creates polar bonds with partial charges.
• Bond polarity does not always make a molecule polar, because the molecular shape can cancel the bond dipoles, as in linear CO2.
• Metallic bonds are a sea of delocalized electrons around positive ions, which explains why metals conduct and can be bent or drawn into wires.
• VSEPR theory predicts molecular shapes from the number of electron groups around the central atom, with lone pairs pushing harder than bonding pairs.
• Intermolecular forces are weaker than bonds, and their strength (hydrogen bonding, dipole-dipole, London dispersion) controls melting and boiling points.

Key terms

Ionic bond

A bond formed when a metal loses electrons to become a positive cation and a nonmetal gains them to become a negative anion, and the opposite charges attract.

Covalent bond

A bond in which two atoms, usually nonmetals, share one or more pairs of electrons.

Polar covalent bond

A covalent bond in which electrons are shared unequally, so the more electronegative atom has a partial negative charge (δ-) and the other has a partial positive charge (δ+).

Octet rule

The tendency of atoms to gain, lose, or share electrons to reach 8 valence electrons (hydrogen aims for 2).

VSEPR

Valence shell electron pair repulsion. Electron groups around a central atom spread as far apart as possible, which determines the molecular shape.

London dispersion force

A weak intermolecular attraction caused by brief, shifting dipoles. It is present in all molecules and is the only force in nonpolar molecules; it gets stronger in larger molecules.

Hydrogen bonding

An extra-strong dipole-dipole attraction that occurs when a hydrogen bonded directly to N, O, or F is attracted to a lone pair on N, O, or F in another molecule.

Metallic bond

The attraction between a lattice of positive metal ions and the shared sea of valence electrons released by the metal atoms.

How to study it

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Questions