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Molecular compounds

Molecular compounds. SO FAR. We’ve seen how ions combine to form ionic compounds Electrons are gained or lost so that the atom can form an ion Opposite charges attract Tend to form crystal structures as salts Yet, this is not the only way that compounds are formed. COVALENT BONDS.

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Molecular compounds

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  1. Molecular compounds

  2. SO FAR . . . • We’ve seen how ions combine to form ionic compounds • Electrons are gained or lost so that the atom can form an ion • Opposite charges attract • Tend to form crystal structures as salts • Yet, this is not the only way that compounds are formed

  3. COVALENT BONDS • Sometimes there is not enough energy to fully remove an e- • In this case, the atoms still want to be like the noble gases. • The only way to achieve this is by SHARING electrons. • Molecular compound: a compound formed when two or more atoms combine by sharing electrons.

  4. Covalent bond • Covalent bond: a bond in which e- are shared between two atoms. • Single bond: a bond in which two atoms share one pair of electrons between them.

  5. Single Bond • Abbreviated as: H-H H H

  6. Distance of Bond • The hydrogen atoms have a definite distance • If the two hydrogen’s get too close together then the electrons repel one another

  7. Electronegativity • Electronegativity: the tendency of an atom to attract bonding electrons to itself when it bonds with another atom.

  8. Electronegativity (Handout) • Electronegativity refers to how likely an element will form an ionic bond. • F is the most electronegative element, thus it is assigned a 4.0. • Electronegativity decreases as it goes down a group and increases as it goes across a period.

  9. Nonpolar Covalent Bond • Nonpolar covalent bond: a bond in which the electronegativities of two atoms are equal. • For instance, H-H bond. • In a nonpolar covalent compound, the e- are shared evenly

  10. Polar Covalent Bond • Polar covalent bond: a bond in which two atoms form a covalent bond, but one atoms attracts electrons more strongly than the other atom.

  11. Electronegativity Difference • How do we determine if a compound is covalent or ionic? • Electronegativity difference: the difference between two elements electronegativities.

  12. Electronegativity Difference • What is the electronegativity difference between HBr and classify it as nonpolar covalent, polar covalent, or ionic?

  13. Electronegativity Difference • What is the electronegativity difference between AlF3 and classify it as nonpolar covalent, polar covalent, or ionic? • When doing compounds with multiple bonds, you look at all bonds

  14. Dipole • Dipole: a molecule in which one end has a partial positive charge and the other end has a partial negative charge. • Polarity: when one end of the molecule is more electronegative than the other

  15. Dipole • For instance, HF • Which element is more electronegative? • F • So, we say that H is more positive and F is more negative.

  16. Dipole • Dipole moment: the electronegativity difference. • Dipole moments only occur in polar covalent compounds. • Remember ionic compounds don’t share e-, so there is no partial charge. • So, HF the dipole moment is 1.8. • We write a dipole molecule as • HΔ+-FΔ- • Δ = a small change in charge.

  17. SPECIAL CASE: HYDROGEN BONDS • Water has some special properties because of polarity: • Surface tension • High specific heat (we will discuss this later in the quarter) • Density

  18. PRACTICE PROBLEMS • For the following, find the electronegativity difference. • Is the compound, ionic, polar covalent or nonpolar covalent? • If the compound is covalent, what is the dipole moment?

  19. PRACTICE PROBLEMS • CO2 • RbCl • N2O3 • FeN • AgCl • ZrO • PO4 • NaCl • H2 • CdCl2 • SnO • LiF

  20. ANSWERS • Polar covalent, 1.0 • Ionic, 2.2 • Polar covalent, 0.5 • Polar covalent, 1.2 • Polar covalent, 1.1 • Ionic, 2.1 • Polar covalent, 1.4 • Ionic, 2.1 • Nonpolar covalent, 0.0 • Polar covalent, 1.3 • Polar covalent, 1.7 • Ionic, 3.0

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