unit 12 chemical bonding. definitions chemical bonds force that holds atoms together it’s all...

Unit 12Unit 12

Chemical Bonding

Definitions

Chemical Bonds Force that holds atoms together It’s all about the electrons (e-)

Electrons are attracted to positively charged nucleus of other atom

Types of Chemical Bonds

Ionic BondIonic Bond Bond between metal and nonmetal due to

“electrostatic interactions” Attraction between positively and negatively

charged ions (cations and anions) Electrons are transferred from metal to

nonmetal

Ionic bonds Result from a Transfer of Valence Electrons

+ -


Covalent BondsBonds in which e- are shared Most common type

Shared Electrons Complete Shells

F F


Metallic BondsAtoms are bonded to one another (not to

other elements)Positive ions in a “sea” of negative

charge (e-)

Metallic Bonding

Definitions

Octet rule (Rule of 8) 8 e- in the outer shell very stable

H2 and He want a “duet”

Electron configuration for duet = ns2

Electron configuration for octet = ns2 np6

Examples of Bonding Types

Ionic Bonding: NaCl, K2S

Covalent Bonding H2 , Cl2

Metallic Bonding Cu, Ag

Lewis Dot Diagrams

A Lewis dot diagram depicts an atom as its symbol and its valence electrons. Ex: Carbon

Carbon has four electrons in its valence shell (carbon is in group 14), so we place four dots representing those four valence electrons around the symbol for carbon.

Drawing Lewis Dot Diagrams

Electrons are placed one at a time in a clockwise manner around the symbol in the north, east, south and west positions, only doubling up if there are five or more valence electrons. Same group # = Same Lewis Dot structure

Ex. F, Cl, Br, I, At

Example: Chlorine (7 valence electrons b/c it is in group 17)

Paired and Unpaired Electrons

As we can see from the chlorine example, there are six electrons that are paired up and one that is unpaired.

When it comes to bonding, atoms tend to pair up unpaired electrons.

A bond that forms when one atom gives an unpaired electron to another atom is called an ionic bond.

A bond that forms when atoms share unpaired electrons between each other is called a covalent bond.

Writing Lewis Dots Structures for Ions

Uses either 0 or 8 dots, brackets and a superscript charge designate to ionic charge

Ex.) Li+, Be+2, B+3, C+4, N-3, O-2, F-1

Writing Lewis Dots Structures(Ionic Compounds)

Lewis Dot Diagrams of Ionic Compounds

Ex. 1) NaCl

Ex. 2) MgF2

A substance made up of atoms which are held together by covalent bonds is a covalent compound.They are also called molecules.

Lewis Dot Diagrams for Covalent Compounds

Covalent Compounds and Lewis Dot Diagrams Diagrams show bonds in a covalent compound

and tells us how the atoms will combine Shared e- = bonding e-

Non-shared e- = lone pair e- (a.k.a. non-bonding e-)

Ex. F2

Drawing Electron Dot Diagrams for Molecules

Chemists usually denote a shared pair of electrons as a straight line.

F F Sometimes the nonbonding pair of electrons

are left off of the electron dot diagram for a molecule

Examples

CH H

H

H

CH4

NH3

H H

H

N

Types of Covalent Bonds

Single Bond 2 e- are shared in a bond (1 from each atom)

Double Bond 2 pairs of e- are shared (4 e- total, 2 from

each atom)Triple Bond

3 pairs of e- are shared (6 e- total, 3 from each atom)

Rules for Drawing Lewis Dot Diagrams

1. Add up the total number of valence e- for each atom in the molecule.

Each (-) sign counts as 1 e-, each (+) sign subtracts one e-

2. Write the symbol for the central atom then use one pair of e- to form bonds between the central atom and the remaining atoms.

3. Count the number of e- remaining and distribute according to octet rule (or the “duet” rule for hydrogen)

4. If there are not enough pairs, make sure the most electronegative elements are satisfied. Then, start shifting pairs into double and triple bonds to satisfy the octet rule.

5. If there are extra e-, stick them on the central atom.

Checking Your Work!

But Remember.... The Structure MUST Have: the right number of

atoms for each element, the right number of electrons, the right overall charge, and 8 electrons around each atom (ideally).

Covalent Compounds and Lewis Dot Diagrams

F2 NH3

H2O NH4+

O2 N2

VSEPR: Shapes of Molecules

VSEPR Theory (definition) “Valence Shell Electron Pair Repulsion” Based on idea that e- pairs want to be as far

apart as possibleGives molecule its shape


Electron Pair Any two valence e- around an atom that repel

other e- pairsLone pair e- (unshared/non-bonding pair only on

one atom)Shared e- pair (bonding pair shared between two

atoms) – can be single, double, or triple bonds


Basic Shapes and Bond Angles

Total # e- Pairs

# of Bonded e- Pairs

# of Unshared Pairs

Molecular Shape

Bond Angles

Ex.

1 1 0 Linear 180o F2

2 2 0 Linear 180o BeH2

1 1 Linear 180o

V. VSEPR: Shapes of Molecules

Total # e- Pairs


# of Unshared Pairs

Molecular Shape

Bond Angles

Ex.

3 3 0 Trigonal Planar 120o BF3

2 1 Bent < 120o SO2

1 2 Linear 180o

V. VSEPR: Shapes of Molecules

Total # e- Pairs


# of Unshared Pairs

Molecular Shape

Bond Angles

Ex.

4 4 0 Tetrahedral 109.5o CH4

3 1 Pyramidal < 109.5o NH3

2 2 Bent < 109.5o H2O

1 3 Linear 180o

See handout with all of

the molecular geometry options!!

To determine the electron pair geometry:

1. Draw the Lewis structure.2. Count the number of bonded (X) atoms and non-bonded or lone pairs (E) around the central atom. 3. Based on the total of X + E, assign the electron pair geometry. 4. Multiple bounds count as one bonded atom!

Electron-pair geometry around a central atom

Sum of X + E Shapes

2 linear

3 trigonal planar

4 tetrahedral

5 trigonal bipyramidal

6 octahedral

VSEPR Examples:

What shape would the following compounds have according to VSEPR theory?

CH4

CO2

Bond and Molecule Polarity Polar Bond

Covalent bond in which the electrons are unequally shared Ex. H2O

Non-polar Bond Covalent bond in which the electrons are equally

shared Ex. F2 or CH4

Predicting Bond Polarity Use Electronegativity!! (see next slide)

Predicting Bond Polarity Calculate the difference between the Pauling

electronegativity values for the 2 elements

Type of BondIONIC

(COVALENT)

POLAR NON-POLAR

Types of Atoms1 metal & 1 nonmetal

(ex. NaCl)

(generally)

2 nonmetals

Ex. NH3, H2O

(generally)

2 nonmetals

Ex. CCl4, O2

Electronegativity Difference

≥ 1.7 ≥ 0.4 but < 1.7 ≤ 0.4

0 – 0.4 Non-polar covalent0.4 – 1.7 Polar covalent (more e/n element has greater pull)

1.7 and up Ionic (e- are transferred between atoms)

Polar Molecules

Polar Molecules (dipole) Molecule with separate centers of (+) and (-)

charge In other words, molecules are polar if the pull in

any one direction is not balanced out by an equal & opposite pull in the opposite direction

Polar Bonds and Polar Molecules

Drawing Polar Molecules Positive and Negative regions shown by

“delta”(δ) Ex. CH3Cl

Determining the Polarity of a Molecule

Shape is crucial (determine the VSEPR shape 1st)

All non-polar bonds = nonpolar molecule Polar bonds see if they cancel each other

out If they all cancel = nonpolar molecule If they are unbalanced = polar molecule

Examples: Polar or non-polar?

Determine if the following molecules are polar or nonpolar.

H2S

F2

H2O

Special Types of Bonding

Hydrogen Bonding Force in which a hydrogen atom covalently

bonded to a highly electronegative element (F, O, or N) is simultaneously attracted to a neighboring nonmetal atom

Hydrogen Bonding Elements that undergo H-bonding

Hydrogen bonding is FON! (Fluorine, Oxygen, and Nitrogen)

Effects on Physical Properties H2O is most notable example of H-bonds

Ice forms rigid, open structures Increases volume upon freezing (floats)

Molecules w/ higher molar mass have lower BP than H2O

Special Types of Bonding

Van der Waals (London Dispersion) Forces Intermolecular force between the molecules of a

substance Force of attraction between an instantaneous and

induced dipole Molecules “make these up” (more or less)

Solids Classes of Solids

Molecular Formed by molecules containing covalently bonded atoms

Ionic Formed by cations and anions

Network Covalent Formed by atoms, usually from Group IV A (Group 14)

Metallic Formed by positive ions in a “sea” of electrons

SolidsComparison of Solids

Type of Solid

Hardness Malleability ConductivityMelting Point

Ease of Phase

Change

Molecular(2 nonmetals)

Soft Shatters Insulator LowEasy to convert

to gas

Ionic(1 metal, 1 non)

Hard ShattersConducts if melted or in

waterHigh

Difficult to convert to gas,

solid

Network Covalent

Very Hard ShattersPoor

ConductorHigh

Difficult to convert to gas,

solid

MetallicVaries from soft to hard

Very MalleableGood

conductor as liquid or solid

Usually High

Difficult to convert to gas, Easy to covert

to solid

unit 12 chemical bonding. definitions chemical bonds force that holds atoms together it’s all...

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