Ionic Bonds, Covalent Bonds & Molecular Structure

1. Ionic Bonds,Covalent Bonds & Molecular Structure

2. Chemical Bonding • What is chemical bonding? • It is a strong attraction or force which holds atoms or ions together in a chemical compound. • Why do atoms form bonds? • Octet Rule says that atoms want a full valence shell of 8 e- • It is the valence e- which are responsible for chemical bonds • So by reacting, they may fulfill the octet rule • But most importantly, they form stable compounds!

3. •You know that metals tend to have low IE values and so lose electrons fairly readily to form cations • •You also know that nonmetals tend to have more negative EA values and so attract electrons fairly readily to form anions • So what happens when a metal atom collides with a nonmetal atom? Ionic Bonds

4. •The metal atom with its small IE gives an electron (or more) to the nonmetal with its negative EA • •The cation and anion have achieved a Noble Gas electron configuration • And the cation and anion are held together by electrostatic forces (opposite charges attract). This is the ionic bond. Ionic Bonds

5. Ionic Bonds • Ionic Bonds: a chemical bond between ions of opposite charge (classically, a metal cation bonded to a nonmetal anion). • Electrons are transferred from the metal to the nonmetal.

6. •In an ionic solid like NaCl, you can’t separate out individual Na-Cl ionic bonds, instead it is a 3-D network of Na+ and Cl- ions which are interconnected. • • This network is the crystal lattice. • •Let’s delve in deeper! Ionic Bonds

7. •How exactly does sodium metal combine with chlorine gas to produce sodium chloride? •F irst, write the overall equation. •Although the rxn occurs simultaneously, we can break the overall rxn into 5 distinct steps. •These steps are called the Born-Haber Cycle Born-Haber Cycle

8. As sodium metal is a solid, we must first convert it to the gaseous state: • Na(s) -> Na(g) • This is the heat of sublimation for sodium, ΔHsub. It is also called the heat of formation of Na(g), or ΔHf. Energy is always required in this step as the gas state is higher energy. • We will call this ΔH1 Step 1: Convert Metal to Gas

9. As chlorine gas is a diatomic element, we must break the Cl-Cl bond to form Cl(g) atoms: • 1/2 Cl2(g) →Cl(g) • This is the heat of formation of Cl(g), orΔHf, OR we may use 1/2 the Cl-Cl bond energy (the energy required to break a Cl-Cl bond), D(Cl-Cl). Breaking bonds ALWAYS takes energy. • We will call this ΔH2 Step 2: Convert Cl2 to Cl Atoms

10. Now we ionize the sodium gaseous atom: • Na(g) → Na+(g) + e- • This is simply the IE1 for sodium. This requiresenergy. • We will call this ΔH3 Step 3: IE1 for Na(g)

12. Now we ionize the chlorine gaseous atom: • Cl(g) + e- →Cl-(g) • This is simply the EA1 for chlorine. This releasesenergy. (the first step so far to release energy) • We will call this DH4 Step 4: EA1 for Cl(g)

14. Now we form the ionic solid sodium chloride from the gaseous ions: • Cl-(g) + Na+(g) →NaCl(s) • This step releases energy, as bonds are formed. • Energy is always released when bonds are formed. • We will call this DH5 Step 5: -Lattice Energy

15. The energy released in Step 5 is the negative or reverse of what we call the Lattice Energy of an ionic compound. • Lattice Energy is abbreviated LE, U, or ΔHLatt. • Lattice Energy is DEFINED as the energy REQUIRED to separate a mole of a solid ionic compound into its gaseous ions. Step 5: -Lattice Energy

17. What do you notice about the 5 steps? Ionic Compound Formation

19. What factors affect the Lattice Energy? • This is Physics! • Charge on ions • Distance between ions (size of ions) Lattice Energy Factors

20. How would you draw the Born-Haber Cycle for MgCl2? Ionic Compound Formation

22. Ionic Compounds • The attractive force between full opposite charges is very strong. • So ionic bonds are very strong. • Therefore, ionic compounds have very high melting points, boiling points, and high lattice energies. • NaCl melts at 804°C • LE are in thousands of kJ/mol, so it takes a lot of energy to break apart a solid ionic crystal. (But it does happen, does salt dissolve in water?)

23. Covalent Bonding • Covalent Bonding and Covalent Compounds: A bond where atoms share electrons. • Remember that it is difficult for atoms to gain or lose 3 or more electrons. • So many atoms share electrons in order to have 8 valence electrons. • It’s like sharing a room with someone, it’s both your room, but you’re sharing. • When atoms share one or more electrons, a covalent bond is formed because both nuclei are attracted to the shared electrons. • Compounds which contain covalent bonds are called covalent compounds or molecular compounds or molecules.

24. Covalent Bonding • The two atoms share one or more electrons; with the shared electrons having a high probability of being found between the two nuclei. • The above figure represents the hydrogen molecule, where 2 electrons are shared equally between the two atoms. • The 7 diatomic elemental molecules share electrons equally just as the above figure shows. • These covalent bonds where the electrons are shared equally are also called nonpolar bonds or nonpolar covalent bonds.

25. Covalent Bonding • But as many of us know from personal experience, not everyone shares equally! • Just as there are greedy people, some atoms are more electron-greedy than others and take more than their fair share of electrons. • When two atoms share electrons unequally, a polar covalent bond results. • The atom which has a stronger attraction for electrons will pull the shared electrons towards its nuclei. • Thus, the unequally shared electrons will tend to be closer to the electron-greedy atom.

26. Covalent Bonding • In the figure below, the HF molecule would look like the left picture if the two atoms shared two electrons equally. • But F is extremely electron-greedy so it pulls the sharedelectrons towards its nuclei as in the picture on the right. • Thus, the HF molecule is polar covalent.

27. Covalent Bonding • As electrons have a negative charge, the electron-greedy atom will have what we call a partial charge. • It doesn’t have a full negative charge as it is still sharing electrons, but as it has more than its fair share, it has a slight or partial negative charge. • If the one atom has a partial negative charge, what type of charge do you think the other atom has? • How do we show these partial charges?

28. Covalent Bonding • We show these partial charges like this:δ+for a partial positive charge, andδ-for a partial negative charge. • The HF molecule can be drawn to show these partial charges (Note that we show the bond between the H and F atoms with a horizontal line between them.):

29. Water and Polar Covalent Bonding • Water has two polar covalent bonds. • Oxygen has a partial negative charge, while both hydrogens have a partial positive charge. • As there is a positive end and a negative end of the bond, there is a charge separation or a dipole. • So water has 2 polar covalent bonds and 2 bond dipoles.

30. Water and Polar Covalent Bonding • In water, these 2 bond dipoles make water a very polar molecule. • This is why water has some very special properties including its high freezing and boiling points. • It is also why ionic compounds tend to be water soluble. • Life on our planet would be very different (and might be nonexistent) if water were not a polar molecule.

31. Water and Polar Covalent Bonding • However, there are many nonpolar molecules with polar covalent bonds. • Sometimes, due to the shape or geometry of the molecule, bond dipoles cancel out. • Carbon dioxide is an example of a nonpolar molecule with polar covalent bonds. • If the bond dipoles cancel, the molecule is nonpolar with polar covalent bonds. • If the bond dipoles do not cancel, the molecule is polar with polar covalent bonds.

32. The Bonding Continuum Here’s a picture of covalent bonding, polar covalent bonding, and ionic bonding. In (a), electrons are shared equally as in H2. In (b), electrons are shared unequally as in HF. In (c), electrons have been transferred from one atom to another, as in NaCl. Notice that a polar covalent bond is between a covalent bond and an ionic bond, so sometimes we say that a polar covalent bond has partial ionic character.

33. Covalent Bond Lengths & Strengths • Let’s go back to the hydrogen molecule! • You can see why the H2 molecule forms as the attractive forces between the nuclei and shared electrons overcome the repulsive forces. • But there is an optimum distance between the 2 nuclei where the bond between the 2 nuclei and electrons is greatest. • This optimum distance is called thebond length.

34. Covalent Bond Lengths & Strengths • In the following figure, you can see what happens if the H atoms get too close or too far apart: the bond is unstable and the molecule falls apart!

35. Covalent Bond Lengths & Strengths

36. Covalent Bond Lengths & Strengths • So the bond length is the optimal distance between 2 atoms. • As this is the distance where the bond strength is greatest, it is also the distance at which the most energy is required to break the bond. • The energy required to break a covalent bond is called the bond energy or the bond dissociation energy. • Every type of bond has its own characteristic bond energy, D, but it always takes energy to break a bond. • But this means that energy is released when a bond forms! • There are Tables of bond energies for many different types of bonds and you’ll use them later.

37. Covalent Bond Lengths & Strengths • As a rough idea of how strong covalent bonds are, bond energies range from about 100 kJ/mol to over 600 kJ/mol. • By comparison, Lattice Energies for ionic compounds were thousands of kJ/mol! • Here are a few things to point out: • Double and triple bonds are shorter and have higher bond energies than single bonds. • There may be many bonds in a molecule, so it may take a lot of energy to break ALL of the bonds.

38. Covalent Bond Lengths & Strengths

39. Covalent Bond Lengths & Strengths

40. Properties of Ionic & Molecular Cmpds • You’ve seen how ionic and molecular compounds form. • Do they have different properties?

41. Properties of Compounds

42. Predicting Bond Types: Electronegativity • Ionic and molecular compounds have different properties. • How can you predict whether a compound is ionic or covalent? • Several ways: • Nonmetal bonded to nonmetal equals covalent (But is it polar or nonpolar?) • Metal bonded to nonmetal equals ionic • Use electronegativity values

43. Predicting Bond Types: Electronegativity • As was mentioned earlier, some atoms are electron-greedy or are very strongly attracted to electrons in a bond. • Electronegativity is a measure of the ability of an atom in a molecule to attract electrons towards itself within a chemical bond. • Chemist Linus Pauling developed the electronegativity scale for the elements. • Fluorine is the most electronegative element, with its electronegativity set at 4.0. • As F is the most electronegative element, the electronegativity of the elements increases going across from left to right across a period, and it would decrease going down a group. • If F is the most electronegative element, what is the LEAST electronegative element (or most electropositive)?

44. Predicting Bond Types: Electronegativity • Why aren’t the Noble Gases on the Table?

45. Predicting Bond Types: Electronegativity • To determine the type of bond, the difference in electronegativity values must be calculated. • Ex: H-F bond: • Your Turn: Determine the electronegativity difference in the Na-Cl bond. • Your Turn: Determine the electronegativity difference in the C-H bond.

46. Predicting Bond Types: Electronegativity • After the electronegativity difference has been calculated, the bond type can easily be determined. • If the difference is ≥0.0 but ≤ 0.4, the bond is covalent. • If the difference is > 0.4 but < 2.1, the bond is polar covalent. • If the difference is ≥ 2.1, the bond is ionic. • Ex: Determine the bond type in H-Cl. • 1) Cl = 3.0; H = 2.1, so the difference is 3.0 - 2.1 = 0.9 • 2) Thus, the H-Cl bond is polar covalent. • Your Turn: Determine the following bond types: • 1) O-O 2) Na-F 3) Si-N

47. Bond Energies, D, & Enthalpy Changes • We can use Bond energies to approximate the enthalpy change for a reaction. • If we know how much energy it takes to break or make a chemical bond, we can calculate the energy change for a rxn. • Now this is an energy, not an enthalpy, but the difference may be less than 1%, so we usually ignore this and just say that it is a bond enthalpy.

48. Bond Energies & Enthalpy Changes • There are Tables and Tables of bond energies for many different types of bonds. • We can use these to findΔH°rxn if we don’t have the rightΔH°f values or we can’t use Hess’s Law.

49. Bond Energies, D, & Enthalpy Changes

50. Bond Energies & Enthalpy Changes • To use D values to find the enthalpy change for a rxn, we use the following equation: