Electron-Dot Symbols

Chemical properties of elements depend on the electrons in the highest energy levels of an atom.

            •  Valence shell: highest occupied energy level of an atom

            •  Valence electrons: Electrons in the valence shell

            •  Number of valence electrons for main group elements is shown by position on periodic table.

            •  Valence electrons are shown by dots

           

Example 1: Write the electron dot symbol for C.

 

               

Example 2: Write the electron dot symbol for Mg.

 

 

Example 3: Write the electron dot symbol for O.

 

               

Forming Ions

Ion:  charged particle formed by the gain or loss of electrons

            Example:  Na+, Cl -

 

Why do ions form?

 

Octet Rule: stability associated with having a filled valence shell (8 electrons for most, 2 electrons for H and He)

            Example:  Draw the electron-dot symbol for Ne.

 

                        Note:  Noble gases are particularly unreactive.
Question:  How does F obtain the same number of valence electrons as Ne?

 

Answer:

Examples:

 

            O                    

 

            N

 

            Al

 

            Ca

 

            B

 

            Na

 

Metals tend to lose electrons (form cations).

·        Non-metals tend to gain electrons (form anions).

·        The charge on simple ions from the main group elements can be determined based on the element’s position in periodic table.

 

Where could F get 1 electron?

The tendency to attract electrons is called electronegativity.

            Trend: 

 

The larger the difference in electronegativity between two elements, the more likely it is for them to bond by forming ions.

Attractions between oppositely charged ions result in formation of ionic bonds.

 

Example:

   

 

Use electron dot symbols to show the ionic bond(s) which would form between the following combinations:

 

            Mg and O

 

 

            Mg and Cl

 

 

            Li and O

 

Consider atoms of Calcium and Phosphorous:

1.  Write electron configurations for each atom.

 

 

2.  Write electron dot symbols for each atom.

 

 

3.  Write electron dot symbols for the ions that form from each atom, and then write the electron configuration for the ion.

 

 

4.  Write the electron dot symbol for the ionic compound that would form when the atoms combine.

Write balanced chemical reactions which show how the following elements combine.  Use correct formulas for diatomic and polyatomic molecules.

 

Magnesium and oxygen

 

 

Magnesium and chlorine

 

 

Lithium and oxygen

 

 

Aluminum and oxygen

 

 

Calcium and phosphorous

 

 

Potassium and nitrogen

 

 

 

Covalent Bonding

Question:  What is another way for atoms to fill their valence shell?

Answer:  Share electrons!

·        Elements which are similar in electronegativity tend to share electrons rather than forming ions.

·        They form covalent bonds.

 


Ex:  1) Water (H2O)

  

           

2) Ammonia (NH3)

 

A covalent bond is formed when orbitals from two atoms overlap, so that a pair of electrons can be shared back and forth.

·        A single bond results from the overlap of one pair of orbitals, with one pair of electrons being shared.

·        A double bond results from the overlap of two pairs of orbitals, with two pairs of electrons being shared.

·        A triple bond results from the overlap of three pairs of orbitals, with three pairs of electrons being shared.

 

Lewis formulas can be used to represent the distribution of electrons in covalently bound substances.

·        Shared pairs of electrons are represented by dashes, while non-bonding or unshared electrons are represented by dots.

·        Most covalent compounds fulfill the octet rule – the main group elements share electrons to achieve noble gas configurations.


Formula for Lewis structures: S = N – A

S = number of shared electrons

N= number of electrons needed for all atoms to have noble gas configurations

A= number of electrons available (from atoms or from charge on polyatomic ions)

 

Steps to writing Lewis structures:

1.      Select a symmetrical structure with:

·        the element with least e-affinity in the center

·        oxygen atoms not bonded to each other (except O2, O3, O22-, and O2-)

·        most oxoacids have H bonded to oxygen

·        if more than one central atom is needed, the most symmetrical possible structure is used

2.      Calculate N, A, and S

3.      Place S (#) electrons in the symmetrical structure as bonding pairs (dashed lines), using double or triple bonds only when necessary.

4.      Place additional available (A – S) electrons as non-bonding electrons to fulfill the octet of each element

 

Examples:

CO2

 

 

 

C2H4

 

 

 

H2SO4

 

 

 

CO32-

 

Resonance and formal charge

 


There are three possible ways to write the Lewis structure for the carbonate ion.

 

The double bond is equally likely to occur between carbon and any of the three oxygen atoms. 

·        Each structure is referred to as a resonance structure, with the double arrow used to signify that each contributes to the overall structure.

·        The true structure is an average of the three.

·        Carbon to oxygen single bonds are usually 1.43 A in length, and carbon to oxygen double bonds are usually 1.22 A in length.

·        In carbonate, all three carbon to oxygen bonds are 1.29 A, an intermediate length explained by resonance.

·        Resonance structures can only involve movement of electrons – overall arrangement of the atoms must be the same.

 

Formal charges can be calculated for each atom in a Lewis structure, and can be helpful in determining correct Lewis structures.

·        The best Lewis structure will have formal charges on each atom equal to zero or as near zero as possible.

Example: thionyl chloride SOCl2

 

Possible Lewis structures:


 


Formal charge = number of valence electrons – number of bonds – number of unshared electrons.

 

Guidelines for Lewis structures:

·        Most likely structure has formal charges closest to zero.

·        Negative formal charges are more likely to occur on more electronegative elements.

·        Adjacent atoms should not have the same formal charge.

 

Notes:

·        Formal charges of zero are obtained based on the number of bonds formed.  Group 3 = 3 bonds, group 4 = 4 bonds, group 5 = 3 bonds, group 6 = 2 bonds, group 7 = 1 bond.

·        Formal charges are written inside a circle to differentiate from ionic charges.  Ex. Ε

·        Atoms which have d sublevel electron orbitals available (level 3 and higher) can have more than eight electrons in their Lewis structures.

Exceptions to the octet rule:

1.      Beryllium usually only has four electrons around it.

2.      Group IIIA elements (especially Boron) usually only have six electrons around them.

3.      Some compounds contain an odd number of electrons.  Ex. nitrogen monoxide

4.      Some compounds require that the central atom have more than eight electrons (because of formal charge or because there are extra available electrons).

·        If S (# of shared electrons) is less than number needed to bond all atoms, then S is increased

·        If S must be increased and extra electrons are leftover after filling all valence shells, extra electrons are placed on the central atom.

 

Many of these exceptions involve molecules which are highly reactive.

 


Draw Lewis structures for the following molecules.

 

BeCl2

 

 

 

 

BCl3

 

 

 

 

XeO3

 

 

 

 

ICl4-

 

 

 

 

Covalent bonds can be polar or non-polar

 

Non-polar bonds involve an equal sharing of bonding electrons  -- electron density is distributed evenly around the bond.

 

In polar bonds, one atom tends to attract electrons more than the other (has higher electronegativity) – the sharing is unequal

·        In polar bonds, the atom with the higher electronegativity has a higher electron density around it, symbolized as a partial negative charge. (d-)

·        The atom with lower electronegativity is assigned a partial positive charge. (d+)

·        The separation of charges creates a dipole. The larger the difference in electronegativity, the larger the dipole.


 

 

·        Molecules containing a dipole (polar molecules) line up in an applied electrical field to minimize electrostatic repulsions.

·        Non-polar molecules are unaffected by an applied electrical field.

 

There is a continuous range of bonding character, ranging from highly ionic to non-polar covalent.

·        The larger the difference in electronegativity, the more ionic the character of the bond.

·        The smaller the difference of electronegativity, the more covalent the character of the bond.

Polar covalent bonds are sometimes referred to as having partial ionic character. 

Ex. HCl    DElectroneg = 0.9

·        In water (aqueous state), it ionizes to produce H+ and Cl- ions and acts as a conductor.

·        In the liquid state it is a non-conductor.

Polyatomic ions are held together internally by covalent bonds, and combine with other ions via ionic bonding.

 

 


VSEPR Theory  (Valence shell electron pair repulsion)

 

Valence electrons are arranged in a way which minimizes repulsive forces. 

·        This arrangement is based on the number of areas of electron density surrounding a central atom.

·        Bonding areas (single, double, or triple bonds) are counted as one area of electron density.

·        Pairs of non-bonding electrons are counted as one area of electron density.

·        There may be anywhere from 2-6 areas of electron density surrounding the central atom.

·        The number and types of areas (non-bonding pairs, single/double/triple bonds) help to determine the shape of the molecule and the bond angles observed.

 

 Examples:

CO2

 

 

 

 

Electronic shape:                  Molecular shape:                   Bond angles:

 

 

NO3-

 

 

 

 

Electronic shape:                  Molecular shape:                   Bond angles:

 

 

SO2

 

 

 

 

Electronic shape:                  Molecular shape:                   Bond angles:

 

 

 

CF4

 

 

 

 

Electronic shape:                  Molecular shape:                   Bond angles:

NH3

 

 

 

 

Electronic shape:                  Molecular shape:                   Bond angles:

 

 

H2O

 

 

 

 

Electronic shape:                  Molecular shape:                   Bond angles:

 

 

PF5

 

 

 

 

Electronic shape:                  Molecular shape:                   Bond angles:

 

 

SF4

 

 

 

 

Electronic shape:                  Molecular shape:                   Bond angles:

 

 

ClF3

 

 

 

 

Electronic shape:                  Molecular shape:                   Bond angles:

 

 

XeF2

 

 

 

 

Electronic shape:                  Molecular shape:                   Bond angles:

 

 

SF6

 

 

 

 

Electronic shape:                  Molecular shape:                   Bond angles:

 

 

BrF5

 

 

 

 

Electronic shape:                  Molecular shape:                   Bond angles:

 

 

ICl4-

 

 

 

 

Electronic shape:                  Molecular shape:                   Bond angles: