Unit 08: Bond Types

Valence Electrons Reminders

  • The e- is the outermost energy level of an atom.

  • The e- involved in bonding.

  • The number of valence electrons- determines how many bonds the atom will form.

  • They determine what type and charge ion is formed by an element.

Octet Rule Reminders

  • In forming compounds, atoms tend to achieve the e- configuration of a Nobel gas.

  • An octet is a set of 8.

  • Electron configurations will end with ns2np6.ns^2np^6.

  • They will lose, gain, or share e- in order to obtain this state.

01: Covalent Bonds

  • Formed between nonmetal atoms

  • Make up molecules

  • Share electrons to form bonds

  • Both atoms will have a full valence shell

  • All covalent bonds are weaker than ionic bonds

  • Compounds with these bonds can be liquids, solids, or gases at room temperature

  • Most have low melting points

  • Most have relatively low boiling points

  • Do not conduct electricity

  • The diatomic elements (BrINClHOF) Form covalent bonds

  • There are two types of covalent bonds:

    • Sigma bonds - these are single bonds between elements

    • Pi bonds - these are multiple bonds between elements

      • Double bonds =

      • Triple bonds (three lines)

  • Remember each dash represents two e-

Types of Covalent Bonds!:

02: Metallic Bonds

  • Metal bonds are arranged in very compact and orderly patterns.

  • The e- are not shared, gained, or lost to form these bonds.

  • Outer energy levels in the metal atoms will overlap.

  • The metal atoms in a metallic solid contribute their valence e- to form a “sea of e-“

  • The e- are not held to any certain atom, they move freely from one atom to the next.

  • These e- are known as delocalized e-.

  • They do not have a specific location in the bond.

  • Think of a metallic bond as a peanut butter sandwich.

    • The two pieces of bread represent the atoms.

    • The peanut butter represents the “sea of e-“

    • The peanut butter hold the sandwich together and it is not specific to either piece of bread

    • The e- hold the bond between the atoms together and they are not specific to either atom

  • When the valence e- leaves to join the sea of e-, it leaves behind a metallic cation (the individual atoms have a positive charge)

  • The bond is the attraction of a metallic cation for delocalized e-

  • Compounds containing metallic bonds have the following properties:

    • Malleable

    • Ductile

    • Durable

    • Fairly high boiling points

    • Melting point vary, but are generally moderate

  • Alloys are a mixture of elements with metallic properties

  • The properties of the alloy are often superior to the properties of their individual parts

  • Alloys are important because they increase the usage of metals

  • Some examples of alloys are steel, pewter, and brass

  • There are two types of alloys:

    • Interstitial

      • Small holes in the metal lattice are filled with smaller atoms

      • Ex: carbon steel - Iron crystals are filled with carbon atoms producing a stronger, harder, and more ductile substance

    • Substitutional

      • Atoms of the original metal are replaced by atoms of similar size of another metal

      • Ex: sterling silver, brass, and pewter

Alloys

03: Ionic Bonds

  • Formed between oppositely charged ions

  • Metals form cations

  • Nonmetals form anions

  • E- are transferred from one atom to another

  • The formation of ionic compounds are always exothermic

  • Ionic bond form crystal lattice structures

  • Ionic bonds are different from metallic bonds, but share some of the same properties

  • They are relatively strong bonds and require large amounts of energy to break apart

  • Compounds will have high melting and boiling points

  • Compounds are hard, brittle solids at room temperature

  • In their solid state, they do not conduct an electric current

  • When they are dissolved in water, the ions dissociate to form ions

  • These are known as electrolytes

  • Electrolytes do conduct electricity very well

04: Bond Types Based on EN Difference

  • Differences in electronegativity (EN) determine whether a bond will be classified as nonpolar covalent, polar covalent, or ionic

  • EN is how strongly atoms attract bonding e- to themselves

  • They higher the EN, the stronger the “pull” that atom has for e-

  • You can look at it like a tug of war for the e-

  • A nonpolar covalent compound is one in which the e- are shared equally or fairly equal in the bond

  • The tug of war is essentially equal

  • The difference in the atoms’ electronegativities will be no greater than 0.4

  • An example of a nonpolar covalent compound would be any of the diatomic molecules since the EN difference is 0

  • Another example would be SeS. Se has an EN of 2.4 and S has an EN of 2.5, giving a difference in EN of 0.1

  • In a polar covalent compound, there will be one atom that has a stronger attraction for the e-

  • Because of this difference in attraction, there will be a slightly positive end and a slightly negative end to the compound

  • In our tug tug of war, one element is pulling more than the other

  • The difference in EN will be 0.4 to 1.5 in a polar covalent compound

  • In a ionic compound, there will be one atom that has a much larger attraction for the e-

  • The tug of war will be extremely one-sided. It would be like the Atlanta Falcons defensive line versus a group of kindergarten students

  • The difference in EN will be 1.6 or greater