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Intro to Bonding: Part 2: Covalent Compounds (Type 3 Binary Compounds)

Intro to Bonding: Part 2: Covalent Compounds (Type 3 Binary Compounds). Covalent Bonds. A bonding force resulting from the sharing of valence electrons. Covalent compounds result when a nonmetal shares electrons with another nonmetal . Example: Carbon dioxide (CO 2 ) Carbon (C)

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Intro to Bonding: Part 2: Covalent Compounds (Type 3 Binary Compounds)

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  1. Intro to Bonding:Part 2:Covalent Compounds(Type 3 Binary Compounds)

  2. Covalent Bonds • A bonding force resulting from the sharing of valence electrons. • Covalent compounds result when a nonmetal shares electrons with another nonmetal. • Example: • Carbon dioxide (CO2) • Carbon (C) is a nonmetal • Oxygen (O) is a nonmetal

  3. Valence Electrons • Valence electrons are the electrons found in the outermost energy level. • The periodic table organizes the number of valence electrons based on the group that the element is located in.

  4. Valence Electrons • Valence electrons tell us how many electrons are available for bonding. • In covalent bonds, electrons are shared. • So the valence electrons are the electrons that are shared in a covalent bond. • Valence electrons can be illustrated using a Lewis Dot Structure.

  5. Lewis Dot Structures • A representation of a molecule showing how valence electrons are arranged among the atoms in the molecule. • Dots are used to represent valence electrons. • The dots are written around the element symbol.

  6. Lewis Dot Structures • Each atom wants to achieve noble gas electron configuration. • To do this, it must complete it’s outermost energy level. • For H and He, this means fulfilling the duetrule: having 2 electrons in the outermost energy level.

  7. Lewis Dot Structures • For atoms that contain more electrons than H and He, there is the octet rule: having 8 electrons in the outermost energy level. • Example: Fluorine gas • Each fluorine atom will share 1 electron to satisfy the octet rule. • By doing this, both fluorine atoms bond covalently and create a covalent compound.

  8. Steps for Writing Lewis Structures • Obtain the sum of the valence electrons from all of the atoms. • Use one pair of electrons to form a bond between each pair of bound atoms. • Use a line (instead of a pair of dots) to indicate each pair of bonding electrons. • Arrange the remaining electrons to satisfy the duet rule for hydrogen and the octet rule for each second-row element.

  9. Single, Double, Triple Bonds • By drawing the Lewis dot structure, you can determine if there is a single, double, or triple bond between atoms. • Single bond • Each atom shares one electron • Double bond • Each atom shares two electrons • Triple bond • Each atom shares three electrons.

  10. Single, Double, Triple Bonds • Bonds are represented by solid lines:

  11. Nonpolar Covalent Bonds • A nonpolar covalent bond is when electrons are shared equally. • An example is Oxygen gas • Each oxygen atom has 6 valence electrons • Each oxygen is going to share 2 valence electrons. • Oxygen is left with 2 lone pairs on each oxygen atom. Thus there is an equal distribution of electrons. • Oxygen gas would be a nonpolar covalent compound.

  12. Polar Covalent Bonds • A polar covalent bond is when electrons are NOT shared equally. • An example is HCl • Hydrogen has 1 valence electron. • Chlorine has 7 valence electrons. • Hydrogen will share it’s one electron so that Chlorine can fulfill the octet rule. • Chlorine will share one of it’s electrons so that Hydrogen can fulfill the duet rule. • Hydrogen is left without any lone pairs. Chlorine has 3 lone pairs. Thus there is an unequal distribution of electrons. • HCl would be a polar covalent compound.

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