ionic compounds and metalsgravesscience.zohosites.com/files/unit 6b...exceptions to the octet rule...

Unit 6B Covalent Bonding

Learning Goals

• I can recognize the difference between an ionic, covalent and metallic compound.

• I can explain the difference between ionic, covalent and metallic bonding.

• I can correctly name ionic, metallic and covalent compounds.

Chemical Bonding

• Chemical Bond The force that holds 2 or more atoms

together. • Bonds form by the attraction between the

positive nucleus of one atom and the negative electrons of another atom.

• Only the valence electrons are involved in chemical bonds.

• Chemical bonding results in a more stable situation for each atom in the bond.

Ionic, Metallic & Covalent Compounds

• 3 types of compounds

–Ionic: NaCl, FeCl3, CuCl2

–Metallic: Na, Fe, Cu, brass (alloy of Cu and Zn)

–Covalent:H20, Dextrose-C6H12O6; methane-CH4; NH3

Covalent Bonding

• The chemical bond that results from sharing electrons is a covalent bond.

• A molecule is formed when two or more atoms bond.

• We will use Lewis structures to represent covalent bonding.

Ionic vs. Metallic vs. Covalent Bonding

Explain the difference in bonding between the 3 different types of compounds.

• Ionic bonding-

• Metallic bonding-

• Covalent bonding-

Ionic vs. Metallic vs. Covalent Bonding

• Ionic bonding-

– Valence electrons are transferred forming a cation (loses the e-) and an anion (gains the e-). Generally formed between a metal and a nonmetal.

• Metallic bonding-

– Valence electrons are delocalized and are attracted to all of the metal cations. No electrons are lost or gained. Formed within all elemental metals and alloys.

• Covalent bonding-

– Valence electrons are shared between 2 atoms forming a molecule. Generally formed between 2 nonmetals.

Ionic, Metallic and Covalent Compounds

• Are the following compounds ionic, metallic or covalent? – CaBr2

– Ca

– Bronze- alloy of Cu, Zn and Sn

– HBr

– LiF

– CO2

– LiOH

– Zn

– Zn(NO3)2

Naming Covalent Compounds

• First element in the formula is always name first.

• The second element in the formula is named using its root name and adding the suffix –ide.

• Prefixes are used to indicate the number of atoms of each element that are present in the compound.

• NEVER use “mono” for the first element


In your notes, name the following:

• CO2

• NH3

• N2

• P2O5

• H2

• H2O

• NF3

• CCl4 (that’s an “L”, as in CL)

Naming Acids

• There are two types of acids:

–Binary acids contain hydrogen and only one other element such as HF, HCl, HBr

–Oxyacids contain hydrogen and some polyatomic ion containing oxygen such as HNO3, H2SO4, H3PO4

– The exceptions to the naming system are sulfur and phosphorus.

Naming Binary Acids

• To name a binary acid, we use the prefix “hydro” followed by the base word of the other element in the compound. We then add “-ic” to the end.

• For example:

– HF – Hydrofluoric Acid (fluor comes from fluorine)

– HBr – Hydrobromic Acid (brom comes from bromine)

Try These:

• On your whiteboard, try these:

– Name the following acids:

• HI

• H3N

Write the formulas for:

- Hydrochloric acid

- Hydroselenic acid

Naming Oxyacids

• When naming oxyacids, we NEVER use the prefix hydro.

• We look at the polyatomic ion in the acid:

– If its name ends in –ate, we change the ending to –ic

– If its name ends in –ite, we change the ending to –ous

For example:

HNO3 – NO3 is named “nitrate”, so we change the ending to –ic --- so HNO3 is called “nitric acid”.

H2CO2 – CO2 is named “carbonite”, so we change the ending to –ous --- so H2CO2 is called “carbonous acid”.

Try These

• On your whiteboard:

–Name the following acids:

•HNO2

•HBrO4

Naming Exceptions

• Sulfur and phosphorus follow the same rules for the most part, except they don’t lose their root words.

• For instance:

– H2SO4 – SO4 is named “sulfate”, so by the rules, it should be sulfic acid. However, H2SO4 is actually named “sulfuric acid”.

– H3PO4 does the same thing – its name is “phosphoric acid”.

Mixed Practice

• When given a compound – we have to first determine whether it is ionic or covalent:

– If the first element is a metal IONIC

– If the first element is not a metal Covalent

– If IONIC – Use the ion sheet for naming

– If COVALENT – Use the prefixes

Mixed Practice

• With a partner, name the following compounds:

1. Ba(NO3)2 2. CO

3. PBr3 4. KF

5. CF4 6. P2F5

7. CaCO3 8. LiCl

9. KMnO4 10. N6O7

Mixed Practice

• With a partner, write the formulas for the following compounds:

1. Magnesium Oxide 2. Trisulfur monoxide

3. Copper (I) Sulfate 4. Nitrogen tetrachloride

5. Xenon hexafluoride 6. Sulfur dioxide

7. Lithium sulfide 8. Strontium fluoride

9. Nonaoxygen pentasulfide 10. Sodium carbonate

Covalent Bonding

• There are 7 diatomic molecules:

H2, N2, O2, F2, Cl2, Br2, I2

• The mnemonic device used to remember this is “Have No Fear Of Ice Cold Beer

• Diatomic molecules form when 2 atoms of the same element share electrons to form a covalent bond.

• These 2-atom molecules are more stable than the individual atoms.

• By sharing electrons each atom achieves the stable octet of (8) valence electrons.

Covalent Bonding

• The figure shows two hydrogen atoms forming a hydrogen molecule with a single covalent bond, resulting in a more stable H2 molecule.

Single Covalent Bonds

• Sigma bonds are single covalent bonds.

• Sigma bonds occur when the pair of shared electrons is in an area centered between the two atoms.

Covalent Bonding • Double bonds form when two pairs of

electrons are shared between two atoms.

• Triple bonds form when three pairs of electrons are shared between two atoms.

Multiple Covalent Bonds

• A multiple covalent bond consists of one sigma bond and at least one pi bond.

• The pi bond is formed when parallel orbitals overlap and share electrons.

• Double bonds contain 1 sigma bond and 1 pi bond.

• Triple bonds contain 1 sigma bond and 2 pi bonds.

Lewis Structures

• In a Lewis structure dots or a line are used to symbolize a single covalent bond.

• The halogens—the group 17 elements—have 7 valence electrons and form single covalent bonds with atoms of other non-metals.

Lewis Structures, cont’d.

• Atoms in group 16 can share two electrons and form two covalent bonds.

• Water is formed from one oxygen with two hydrogen atoms covalently bonded to it.


• Atoms in group 15 form three single covalent bonds, such as in ammonia.


• Atoms of group 14 elements form four single covalent bonds, such as in methane.

Drawing Lewis Structures

– Predict the location of certain atoms.

– Determine the number of electrons available for bonding.

– Determine the number of bonding pairs.

– Place the bonding pairs.

– Determine the number of bonding pairs remaining.

– Determine whether the central atom satisfies the octet rule.

Exceptions to the Octet Rule

• Some molecules do not obey the octet rule.

• A small group of molecules might have an odd number of valence electrons.

• NO2 has five valence electrons from nitrogen and 12 from oxygen and cannot form an exact number of electron pairs.

Exceptions to the Octet Rule

• A third group of compounds has central atoms with more than eight valence electrons, called an expanded octet.

• Elements in period 3 or higher have a d-orbital and can form more than four covalent bonds.

Resonance Structures

• Resonance is a condition that occurs when more than one valid Lewis structure can be written for a molecule or ion.

• This figure shows three correct ways to draw the structure for (NO3)-1

.

Hybridization

• To explain how the orbitals of an atom become rearranged when the atom forms covalent bonds, a different model is used:

• Hybridization – The mixing of two or more atomic orbitals of similar energies on the same atom to produce new hybrid atomic orbitals of equal energies.

• Hybrid Orbitals – Orbitals of equal energy produced by the combination of two or more orbitals on the same atom.

Hybridization • To figure out hybridization, you need to remember that there

is 1 S orbital, 3 P orbitals, 5 D orbitals and 7 F orbitals.

• By adding bonded atoms and lone pairs, you get electron domains. Then you can figure out which orbitals fill for each.

Bonded Atoms

Lone Pairs

Electron Domains

Shape Hybridization

2 0 2 Linear sp

3 0 3 Trigonal Planar sp2

4 0 4 Tetrahedral sp3

2 2 4 Bent sp3

3 1 4 Trigonal Pyramidal sp3

5 0 5 Trigonal Bipyramidal sp3d

6 0 6 Octahedral sp3d2

Covalent Bonding

• We can determine whether a covalent bond is polar or nonpolar based upon the ELECTRONEGATIVITY difference between the bonded atoms and the shape the bonded molecule has.

• What is electronegativity?

– The ability of an atom to attract electrons to itself in a chemical bond

Electronegativity

Electonegativity & Bond Character

• A chemical bond is never 100% ionic or 100% covalent. The character of a chemical bond (ionic vs. covalent) depends on the ELECTRONEGATIVITY DIFFERENCE between the bonded atoms.


Example: H2O

Electronegativity of O = 3.44

Electronegativity of H = 2.20

EN Difference 3.44- 2.20 = 1.24

According to table 8.7 this makes

H2O a polar covalent compound.


Example: NaCl

Electronegativity of Cl = 3.16

Electronegativity of Na = 0.93

EN Difference 3.16 – 0.93 = 2.23


NaCl an ionic compound.


Example: CH4

Electronegativity of C = 2.55

Electronegativity of H = 2.20

EN Difference 2.55- 2.20 = 0.35


CH4 a mostly nonpolar covalent compound.

KEY POINT!!!

• The use of electronegativity to predict bond character is just that, a prediction.

• A bond is ionic usually above 2.0 and nonpolar covalent ONLY if the result is exactly 0.

Your Turn

• Use the electronegativity values and table 8.7 to determine if the following compounds are either (in textbook): – mostly ionic

– mostly polar covalent or

– mostly nonpolar covalent

– Then provide the NAME of each compound using our naming rules.

Your Turn KBr

CaO

CO

N2

H2S

MgCl2

NO2

LiF

HBr

CS2

Apply the concept of “like dissolve like”. Only compounds

“like” water will dissolve in water (the ionic and polar

compounds).

Put a star next to the above compounds which you predict

will dissolve in water

Determining Polarity from VSEPR Structures

• After you’ve drawn a correct VSEPR structure – you need to look at the center atom and see if the charge is evenly distributed around the center atom. If it is not, then you have a polar molecule.

Determining Polarity from VSEPR Structures

• Basically –

– Bent and Trigonal Bipyramidal molecules are ALWAYS POLAR

– Tetrahedral, Octahedral and Trigonal Planar and Trigonal Pyramidal molecules are NONPOLAR IF ALL ATTACHED ATOMS ARE THE SAME (otherwise the molecules are polar)

– Linear is almost always nonpolar (there are a slight few exceptions)

Intermolecular Forces


o Intermolecular Forces: attraction between molecules

o Much weaker than chemical bonds

Strength of Forces

• STRONGEST

• Covalent Bonds (400 kcal)

• Hydrogen Bonding (12-16 kcal)

• Dipole-Dipole Interactions (0.5-2 kcal)

• London Dispersion Forces (<1 kcal)

• WEAKEST

• kcal = kilocalorie (unit of energy stored in chemical bonds)

London/Van der Waals

o London Dispersion Forces: o Attraction due to the

constant motion of electrons oWill cause temporary

concentration of charge on one side of an atom/molecule

o Exist between ALL molecules

Dipole

o Dipole-dipole:

o Attraction between polar molecules

The squiggly symbol means a partial charge.

An Analogy

• Dipole-dipole attraction

• - - - - - - -

• The two couples are very happy with each other, but the guy is attracted to the other girl.

Hydrogen

o Hydrogen Bonding:

o Stronger type of dipole-dipole interaction

o Occurs only between molecules with lone pairs on center atom AND hydrogen bonded to F, O, or N

oHydrogen bonding is FON!

Is the molecule polar?

No Yes

London Dispersion Forces Does the molecule have lone pairs on the center

atom, with H bonded to F, O, or N?

Yes No

London Dispersion Forces Dipole-dipole Hydrogen bonding

London Dispersion Forces Dipole-dipole


• To summarize:

– NONPOLAR ATOMS only have London dispersion forces.

– POLAR ATOMS always have dipole-dipole and London dispersion forces and will have a hydrogen bond if there is an H-F, H-O or H-N bond.

Intermolecular Forces Examples

Cl2

• The molecule is nonpolar.

• This molecule has London Dispersion Forces.


• CH3I

• The molecule is polar.

• The central atom does not have lone pairs, and H is not bonded to F, O, or N.

• The molecule has London and Dipole-dipole forces.


• CH2O


• The central atom does not have lone pairs, and H is not bonded to F, O, or N.



• PF3


• The central atom has a lone pair, but does not have H bonded to F, O, or N.



• NH4+

• The molecule is nonpolar…

• So even though H is bonded to N…

• This molecule has only London Dispersion forces.

Intermolecular Forces Examples EXTRA

• H2O


• There is a lone pair on the central atom.

• Hydrogen is bonded to F, O, or N.

• This molecule has London, Dipole-Dipole forces AND hydrogen bonding.

ionic compounds and metalsgravesscience.zohosites.com/files/unit 6b...exceptions to the octet rule...

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