Types of chemical bonds
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Types of chemical bonds. Bond : Force that holds groups of two or more atoms together and makes the atoms function as a unit. Example: H-O-H Bond Energy : Energy required to break a bond. Ionic Bond : Attractions between oppositely charged ions.

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Types of chemical bonds

Bond: Force that holds groups of two or more atoms together and makes the atoms function as a unit.

Example: H-O-H

Bond Energy: Energy required to break a bond.

Ionic Bond: Attractions between oppositely charged ions.

Example: Na+ Cl-


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Types of chemical bonds

Ionic Compound: A compound resulting from a positive ion (usually a metal) combining with a negative ion (usually a non-metal).

Example: M+ + X- MX

Covalent Bond: Electrons are shared by nuclei.

Example: H-H

Polar Covalent Bond: Unequal sharing of electrons by nuclei.

Example: H-F

Hydrogen fluoride is an example of a molecule that has bond polarity.


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Lewis structures

Lewis Structure: Representation of a molecule that shows how the valence electrons are arranged among the atoms in the molecule.

Bonding involves the valence electrons of atoms.

Example: Na● H-H


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Lewis structures of elements

  • Dots around elemental symbol

    • Symbolize valence electrons

      • Thus, one must know valence electron configuration


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Lewis Structures of molecules

Single Bond: Two atoms sharing one electron pair.

Example: H2

Double Bond: Two atoms sharing two pairs of electrons.

Example: O2

Triple Bond: Two atoms sharing three pairs of electrons.

Example: N2

Resonance Structures: More than one Lewis Structure can be drawn for a molecule.

Example: O3


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Rules for Lewis structures of molecules

  • Write out valence electrons for each atom

  • Connect lone electrons because lone electrons are destabilizing

    • Become two shared electrons

      • Called a “bond”

  • Check to see if octet rule is satisfied

    • Recall electron configuration resembling noble gas

      • In other words, there must be 8 electrons (bonded or non-bonded) around atom

        • Non-bonded electron-pair

          • Called “lone pair”


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    Let’s do some examples on the board

    • H2

      • Duet rule

    • F2

      • Octet rule

    • O2

    • N2


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    Lewis structures

    Example

    Write the Lewis Structure for the following molecules:

    • H2O

    • CCl4

      • Where does the carbon go & why?

    • PH3

    • H2Se

    • C2H6


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    Lewis structures continued

    • CO2

    • C2H4

    • C2H2

    • SiO2


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    Polyatomic ions

    • If positive charge on ion

      • Take away electron from central species

    • If negative charge on ion

      • Add electron to central species

    • Example:

      • H3O+


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    Your turn

    • NH4+

    • ClO-

    • OH-


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    Resonance structures

    • When structures can be written in more than one way

      • O3

    • Actual molecule is “in-between”

      • Resonance hybrid

    • Another example

      • HCO3-

        • What would its resonance hybrid look like?


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    Practice

    • NO2-

    • NO3-


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    Formal Charge

    • Charge calculated on atom based on Lewis structure

      • Yields best Lewis structure of competitors

    • FC = VE - [LE + ½(BE)]

    • Rules:

    • Sum of all FC’s must equal to charge on species, if any

    • Smaller or zero FC’s on atoms better than large FC’s

    • Negative FC should be on most electronegative species


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    Examples

    • HBr

      • FC on H = 1-[0 + ½ (2)] = 0

      • FC on Br = 7 – [6+ ½ (2)] = 0

      • Net sum of FC’s = charge on ion = 0

    • OH-

      • FC on O = 6 – [6 + ½(2)] = -1

      • FC on H = 1 – [0 + ½(2)] = 0

      • Net sum of FC’s = charge on ion = -1


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    Practice

    • H2O2

    • H3O+


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    Aberrant compounds

    • Odd-electron species

      • NO

      • NO2


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    Aberrant compounds

    • Incomplete octet

      • BH3


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    Aberrant compounds

    • Expanded octet

      • Some central atoms can exceed an octet

    • Third period and higher elements can do this

      • E.g., Al, Si, P, S, Cl, As, Br, Xe, etc.

      • d-orbitals can accommodate extra electrons


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    Examples

    • AsI5

    • XeF2


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    Practice

    • SCl6

    • XeF4


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    Aberrant compounds

    • Write this out:

      • SO42-

    • Can we reduce the formal charges?

      • If so, how?

    • We can also find the average FC

      • Let’s take a look


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    Aberrant compounds

    • Its OK to expand the octet for those atoms that can take it in order to lower FC’s


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    Practice

    • SO32-

    • PO33-

    • SO2

    • SO3

    • H2SeO4


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    Electroneutrality principle

    • Electrons distributed so that charges on atoms are closest to zero

    • If “-” charge present, should be on most electronegative atom

    • (so, “+” charge should be on least electronegative atom)

    • Good for deciding which resonance structure is best

    • Example: OCN-


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    Electronegativity

    Electronegativity: The relative ability of an atom in a molecule to attract shared electrons to itself.

    Example: Fluorine has the highest electronegativity.

    • Similar electronegativities between elements give non-polar covalent bonds (0.0-0.4)

    • Different electronegativities between elements give polar covalent bonds (0.5-1.9)

    • If the difference between the electronegativities of two elements is about 2.0 or greater, the bond is ionic


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    Electronegativity

    Example

    For each of the following pairs of bonds, choose the bond that will be more polar.

    • Al-P vs. Al-N

    • C-O vs. C-S


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    Dipole moment

    • Dipole Moment

      • A molecule that has a center of positive charge and a center of negative charge

        • Will line up on electric field

    • In Debye units

      • 1 D = 3.34 x 10-30 C  m


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    Examples

    • F2

    • CO2

    • H2O

    • NH3

    • BF3

    • CCl4


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    Molecular polarity

    • Net-dipole moment leads to molecular polarity

    • Thus the following two that have net-dipole moments are polar:

      • H2O

      • NH3


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    Molecular structure

    Molecular Structure: or geometric structure refers to the three-dimensional arrangement of the atoms in a molecule.

    Bond Angle: The angle formed between two bonds in a molecule.


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    Molecular structure:VSEPR

    The VSEPR Model: The valence shell electron pair repulsion model is useful for predicting the molecular structures of molecules formed from nonmetals.

    The structure around a given atom is determined by minimizing repulsions between electron pairs.

    The bonding and nonbonding electron pairs (lone pairs) around a given atom are positioned as far apart as possible.


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    Molecular Structure:VSEPR

    Steps for Predicting Molecular Structure Using the VSEPR Model

    1. Draw the Lewis structure for the molecule.

    2. Count the electron pairs and arrange them in the way that minimizes repulsion (that is, put the lone pairs as far apart as possible).

    3. Determine the positions of the atoms from the way the electron pairs are shared.

    4. Determine the name of the molecular structure from the positions of the atoms.


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    Example

    • Br2

    • CO2

    • CF4

    • PF3


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    Your turn

    • NH4+

    • XeF4

    • AsI5

    • SF3 +

    • I3 -


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