Unit 6: Covalent Bonds • How are covalent compounds named? • How are molecular formulas written? • What do covalent molecules look like? • How are ionic and covalent bonds different?

Chemical Bond • Chemical Bond – Attraction between different atoms – Holds elements together in a compound • Each element has an octet (stability) – Bond type determined by: • Atoms’ electronegativity difference • Type of atoms involved

Compounds • Compounds – Two or more atoms chemically combined – Electrically neutral – Can be ionic, molecular, or metallic

Ionic Bonds • Ionic Bonds – e- donated from one atom to another – Metals and nonmetals • Metal gives up e- to the nonmetal – High EN difference (>1.7) – One atom is MUCH more attracted to electrons than the other

Metallic Bonds

Covalent Bonds

• Metallic Bonds

• Covalent Bonds

– 2 or more metals – Metal cations surrounded by a “sea of electrons”

– e- are “shared” by atoms – 2+ nonmetals – Similar EN values • Difference < 1.7 • More equal sharing

1

Covalent Bonds

Diatomic Elements • Diatomic Elements

• Covalent Bonds – Pairs of electrons are shared

Type of Bond

Pairs Shared

Electrons Shared

Single

1

2

Double

2

4

Triple

3

6

Polyatomic Ions • Polyatomic Ion – Ion made up of more than one atom – Held together by covalent bonds, but bond as ions – Almost always anions • Ex: SO42- - sulfate ion – Names, charges, symbols on back of table

Molecular Compounds • Chemical Formulas – Show kind and number of each type of atom in a compound

• Subscripts – Show there is more than one atom of an element • Fe2O3 – 2 iron and 3 oxygen atoms – If there is only one of that atom, no subscript • H2O– 2 hydrogen and 1 oxygen

– Some elements are never found as a single atom – If they aren’t bonded to another type of atom, they bond to themselves – There are 7: Br I N Cl H O F 1. Hydrogen (H2) 2. Nitrogen (N2) 3. Oxygen (O2) 4. Fluorine (F2) 5. Chlorine (Cl2) 6. Bromine (Br2) 7. Iodine (I2)

Molecular Compounds • Molecular Compounds – Contain covalent bonds – Nonmetals bonded to other nonmetals – Single units called Molecules – Can be two or more elements – CO2 – Can be one element – O2

Molecular Formulas • Molecular Formula – Use prefixes to show number of atoms of each element – Don’t reduce formulas!!

1 2 3 4 5 6 7 8 9 10

Mono Di Tri Tetra Penta Hexa Hepta Octa Nona Deca

2

Naming Molecular Compounds • Naming Molecular Compounds – Prefix + first element then Prefix + second element • Don’t write mono for first element –Ex: CO2 – Carbon Dioxide • For second element, drop ending and add suffix (-ide) –Ex: F – Fluoride • Some vowels are dropped (ao, oo,) with oxygen –Ex: Pentoxide, Monoxide

Lewis Structures • Step 1: Add up valence electrons – For polyatomic ions, include the charge • Ex: 2- charge means add 2 electrons

• Ex: H2O H = 1 valence electron x 2 atoms O = 6 valence electrons 8 valence electrons

Lewis Structures • Step 3: Draw a pair of electrons between the central atom and each bonded atom – A pair of bonding electrons can be represented by a dash between the bonding atoms

H : O : H or

H–O–H

Lewis Structures • Lewis Structures – Show arrangement of e- in a compound – Used to determine molecular shape • Properties are determined by molecular shape

Lewis Structures • Step 2: Determine the central atom – If Carbon is in the compound, it is the central atom – If Carbon is NOT present, the LEAST electronegative element is central – Hydrogen is NEVER the central atom

H

O

H

Lewis Structures • Step 4: Fill in the rest of the electrons to satisfy the octet rule for all elements in the compound – Hydrogen must satisfy the duet rule

.. H : O : H or ..

.. H–O–H ..

3

Lewis Structures • Step 5: Count up all the electrons – Must match count from step 1 – Double check that everything has an octet (duet for hydrogen)

Lewis Structures • What if there are too many electrons in the drawing?  Ex: CO2 • For every 2 extra electrons, in the structure, add another bond – Can be double or triple bonds

–Ex: N2

Lewis Structures • Deficient octet – Boron is stable with six valence e– Ex: BF3

• Expanded Octet – Elements in the 3rd period and after can hold more than 8 valence e– Ex: SF6

VSEPR • VSEPR Theory – Valence – Shell – Electron – Pair – Repulsion – Electron pairs repel each other – Atoms space pairs as far as possible

Lewis Structures • Resonance Structures – Sometimes more than one structure is valid – Draw all possible structures with arrows in between – Actual structure is an average of the resonance structures

• Ex: Ozone (O3)

VSEPR • Molecular Geometry – 3-D arrangement of molecule’s atoms in space – Based on central atom – Shape determined by number and type of electron pairs

4

VSEPR

VSEPR

• Bonding Pair

Lone Pair

– Shared between two atoms

– Angle between bonding pairs – Based on shape and types of electrons – Lone pairs decrease bond angles

• Lone Pair – Unshared electrons – Further apart than bonding pairs

Bonding Pair

• Constituent Group – Atoms/groups bonded to central atom

• Bond Angle

Constituent Groups

VSEPR • To determine the shape:

Bond Polarity • Polar Bonds

1. Draw Lewis Structure 2. Count how many lone pairs/constituent groups surround central atom 3. Compare to the chart

– Intermediate between covalent bonds and ionic bonds – Electrons are unevenly shared

Bond Polarity

Bond Polarity

• Recall Electronegativity – The ability of an atom to attract electrons in a compound • High EN – very attracted to electrons • Low EN – not so attracted to electrons – Values are on the reference sheet

• Bond Polarity – Determined by the difference in electronegativity for the atoms in a bond – Subtract higher EN from lower EN • |EN1 – EN2|

• Compare to the chart: 3.3

1.7

Ionic

Polar Covalent

0.3

0.0

Nonpolar Covalent

5

Bond Polarity

Molecular Polarity • Molecular Polarity

• Dipole – Pull of electrons in polar bond – Bond has slightly negative (δ-) and slightly positive ends (δ+) • Higher EN – δ– Indicated with arrow pointing to more EN element

Molecular Polarity • Molecular Polarity – For BINARY Compounds • Polar when there are lone pairs on central atom –Nonpolar Shapes »Tetrahedral »Linear »Trigonal Planar –Polar Shapes »Trigonal Pyramidal »Bent

Bond Strength • Bond Strength – Energy needed to break bond • More energy – stronger bond – Stronger bonds have shorter length – Single bonds are weakest, triple bonds strongest

– Net dipole for whole molecule – Requires bonds to be polar – Bond polarity is cancelled when molecule is symmetrical • Asymmetrical molecules are polar • Polar bonds do not always mean a polar molecule

Molecular Polarity • Molecular Polarity – For TERNARY Compounds • Generally polar –Asymmetrical due to different EN • Ex: CH3Cl

Polarity • “Like Dissolves Like” – Polar substances can be dissolved by other Polar substances – Nonpolar substances can be dissolved by other Nonpolar substances – Polar substances will NOT dissolve in nonpolar substances

6

Polarity • Miscible – Two liquids are capable of mixing

Polarity • Is density the only reason why your salad dressing separates? – No, the oil layer is immiscible in the water layer

• Immiscible – Liquids will not mix – Ex: Oil and Water

7