Section 5.1—Types of Bonds

Why atoms bond Atoms are most stable when the outer shell of electrons is full. Atoms bond to fill this outer shell The outer shell electrons of an atom are called the Valence Electrons. (outermost s or s and p) Most atoms want 8 electrons in their valence. We call this the Octet Rule Common exceptions are Hydrogen and Helium which can only hold 2 electrons.

One way valence shells become full - - - - - - - - - - - - - Na - Cl - - - - - - - - - - - - - - Sodium has 1 electron in it’s valence shell Chlorine has 7 electrons in it’s valence shell Some atoms give electrons away to reveal a full level underneath. Some atoms gain electrons to fill their current valence shell.

One way valence shells become full - - - + - - - - - - - - - - - Na - Cl - - - - - - - - - - - - - - The sodium now is a cation (positive charge) and the chlorine is now an anion (negative charge). These opposite charges are now attracted, THIS FORMS AN IONIC BOND!

Ionic Bonding—Metal + Non-metal Metals have low valence counts. For Example: Group 1 Metals have ONE valence electron. Metals also have low ionization energies. This means it doesn’t take much energy to remove their valence electrons. Therefore, metals tend to lose their electrons and non-metals gain electrons Metals become cations (positively charged) Non-metals become anions (negatively charged) The cation & anion are attracted because of their charges—forming an ionic bond – ionic means formed from ions!!!

The Ionic Crystal – take note of how the ion charges force the atoms into a repetitive structure called a Crystal!

Covalent Bonding - between Non-metals When two non-metals bond, they share electrons Non-metals that share electrons evenly form non-polar covalent bonds Non-metals that share electrons un-evenly form polar covalent bonds

Metallic Bonding – Metal to Metal Example: The Gold Atoms in a bar of Gold The metals do not (as previously mentioned) hold onto their electrons well. The valence electrons in metallic bonding are free to move throughout the structure - like a sea of electrons. This makes the atom appear to have a positive center.

Bond type affects properties The type of bonding affects the properties of the substance.

Melting/Boiling Points Ionic bonds tend to have very high melting/boiling points as it’s hard to pull apart those electrostatic attractions They’re found as solids under normal conditions

Ionic vs. Covalent

Melting/Boiling Points Polar covalent bonds have the next highest melting/boiling points Most are solids or liquids under normal conditions

Melting/Boiling Points Non-polar covalent bonds have lower melting/boiling points Most are found as liquids or gases

Solubility in Water Ionic & polar covalent compounds tend to be soluble in water

Solubility in Water Non-polar & metallic compounds tend to be insoluble

How does soap work?

Conductivity of Electricity In order to conduct electricity, charge must be able to move or flow Metallic bonds have free-moving electrons—they can conduct electricity in solid and liquid state Ionic bonds have free-floating ions when dissolved in water or when they are molten (liquid) form that allow them conduct electricity Covalent bonds are NOT formed from charges and therefore cannot conduct electricity in any situation

Electrical Conductivity of solutions No Light Dim Light Bright Light Bright Light Question: Do all of these dissolve in water? If so, do all water soluble compounds conduct a current? Why or Why Not?

Distribution of electron density in H2. Figure 9.13 Distribution of electron density in H2. At some distance (bond length), attractions balance repulsions. Electron density is high around and between the nuclei.

Figure 9.12 Covalent bond formation in H2.

Bonding Pairs and Lone Pairs Atoms share electrons to achieve a full outer level of electrons. The shared electrons are called a shared pair or bonding pair. The shared pair is represented as a pair of dots or a line: H •• or H–H An outer-level electron pair that is not involved in bonding is called a lone pair, or unshared pair. •• F F–F or

Properties of a Covalent Bond The bond order is the number of electron pairs being shared by a given pair of atoms. A single bond consists of one bonding pair and has a bond order of 1. The bond energy (BE) is the energy needed to overcome the attraction between the nuclei and the shared electrons. The stronger the bond the higher the bond energy. The bond length is the distance between the nuclei of the bonded atoms.

Trends in bond order, energy, and length For a given pair of atoms, a higher bond order results in a shorter bond length and higher bond energy. For a given pair of atoms, a shorter bond is a stronger bond. Bond length increases down a group in the periodic table and decreases across the period. Bond energy shows the opposite trend.

Table 9.2 Average Bond Energies (kJ/mol) and Bond Lengths (pm)

Table 9.3 The Relation of Bond Order, Bond Length, and Bond Energy

Bond length and covalent radius. Figure 9.14 Bond length and covalent radius. Internuclear distance (bond length) Covalent radius 72 pm Internuclear distance (bond length) Covalent radius 114 pm Internuclear distance (bond length) Covalent radius 100 pm Internuclear distance (bond length) Covalent radius 133 pm

Sample Problem 9.2 Comparing Bond Length and Bond Strength PROBLEM: Using the periodic table, but not Tables 9.2 or 9.3, rank the bonds in each set in order of decreasing bond length and decreasing bond strength: (a) S–F, S–Br, S–Cl (b) C=O, C–O, CΞO PLAN: S is singly bonded to three different halogen atoms, so the bond order is the same. Bond length increases and bond strength decreases as the atomic radius of the halogen increases. The same two atoms are bonded in each case, but the bond orders differ. Bond strength increases and bond length decreases as bond order increases.

Sample Problem 9.2 SOLUTION: (a) Atomic size increases going down a group, so F < Cl < Br. Bond length: S–Br > S–Cl > S–F Bond strength: S–F > S–Cl > S–Br (b) By ranking the bond orders, we get Bond length: C–O > C=O > CΞO Bond strength: CΞO > C=O > C–O