Bonding

Core v Valence Electrons

• The core electrons (represented by the noble gas

from the previous row) are those electrons held

within the atom.

– These electrons are not involved in the bonding, but

contribute to the shielding effect

• The valence electrons are those electrons in the

outermost energy level of an atom.

– These electrons are involved in bonding.

– Can be determined by location in periodic table.

Octet Rule

• The most stable atoms will have 8 outer electrons (noble gases).

• The octet rule states that those elements with 8 outer “s” and “p” electrons will be particularly stable.– These valence electrons largely determine

the chemical properties of an element.• Metals have low numbers of valence electrons.

• Nonmetals have high numbers of valence electrons.

• Intramolecular forces are forces of attraction between atoms in a compound (ie. ionic bonding, etc.)

• Intermolecular forces are forces of attraction between compounds in a collection of compounds. (ie. LDFs, dipole-dipole forces, etc.)

Intramolecular Forces vs.

Intermolecular Forces

Bonding Between Atoms

• Ionic Bonds (cations and anions)

• Covalent Bonds (sharing)

• Metallic Bonds (delocalized electron cloud)

• Network Covalent Bonds (diamond)

Ionic bonding

• Electrostatic attraction (in lattice)

• Metals/nonmetals (widely diff electronegativity)

• Electrons transferred

• Polyatomic ion possibly

• High melting points

• Do not conduct as solids

• But conduct when dissolved

or melted

Covalent Bonding

• Share electrons

• Two or more nonmetals

• Similar electronegativities

• Low melting points

• Do not conduct in any phase

• May or may not dissolve in water (depends

upon polarity of compound)

Metallic Bonding

• Metal cations surrounded by their mobile

valence electrons (in lattice)

• Malleable and ductile; shiny

• Conducts electricity in all phases

• Melting points vary (generally lowers as you

move down a group)

• Metallic character increases as you move

down a group (due to shielding effect)

Alloys

• An alloy is a solid solution of two or more metals (could include a nonmetal or a metalloid).

– Made by melting the elements involved, mixing them together, then re-solidifying.

– The properties of the alloy are much better than any of the original components.

• Common alloys

– Brass (Zn and Cu)

– Bronze (Sn and Cu)

Network Covalent Bonding

• Atoms held together in lattice of covalent bonds (like one big molecule)

• Very hard

• Very high mp and bp

• Electrons are localized about

the atoms

• In diamond, all of the carbons are

bonded in a tetrahedral arrangement

Bond Summary

Let’s Look at Electronegativity

Determining Type of Bond

• Simplest – use periodic table

– metal/non-metal should exhibit ionic bond

– Non-metal/non-metal should be covalent bond

– Metals will exhibit metallic bond

• Find electronegativity difference

– Large difference (>1.67) generally means ionic

– Small to medium difference (<1.67) generally means covalent

• Plot electronegativity computations using Bond Triangle

Bond Triangle

• For years, chemists have determined bond type between atoms by using the idea of electronegativity difference.

– Large electronegativity difference = ionic

– Small to medium electronegativity difference = nonpolar or polar covalent

• The bond triangle also uses electronegativity values, but in a slightly different manner

– A comparison is made between the average of electronegativities and the difference in the electronegativites of the two elements involved.

Bond Triangle, cont.

Types elements Ave EN Diff EN Bond type

Metal/nonmetal ~ 1.5 - ~ 2.5 Usually large Ionic

Metal/metal ~ 1 Very low (~0) metallic

Nonmetal/nonmetal Usually > 2.5 Low covalent

• By plotting Ave EN v Diff EN, you get a triangle that is divided into four

areas.

• Area A represents ionic bonding.

• Area B represents metallic bonding

• Area C represents covalent bonding

• Area SM represents semimetals

• This bond triangle is not absolute, but it more closely can predict the bond

type between atoms than just using electronegativity differences.

Ionic

Covalent

Metallic

Covalent bonds and polarity

• Within the bond, one atom is usually more

electronegative than another

– That atom pulls the electrons in the bond to

himself

– Causes a dipole…one side of the bond is more

negative and the other side of the bond is more

positive

– Not an ionic bond…but a very polar bond

• Lewis dot diagrams can predict type of covalent bond between atoms (single, double, or triple) but it cannot imply the shape of the molecule

• Structural formulas will replace a bonded pair of electrons with a stick.

• Then by understanding the VSEPR theory, the shape can be implied on two dimensional paper

Determining Covalent Bonds

Valence-shell electron pair repulsion theory

• The VSEPR theory states that in a small

molecule, the pairs of valence electrons are

arranged as far apart from each other as

possible.

• This repulsion is different for the following

possibilities.

– Unshared-unshared repulsion is the most.

– Shared-shared repulsion is the least.

– Unshared-shared repulsion is midway between

the two.

5 Basic Shapes

• This theory does explain a wide variety of

molecular shapes.

Example Name of Shape Bond Angle

BeF2 Linear 180º

BF3 Trigonal planar 120º

CH4 tetrahedral 109.5º

NH3 Trigonal pyramid (pyramidal) 107º

H2O Bent 104.5º

More complex shapes

Name of Shape Bond Angle

Trigonal bipyramidal 90º & 120º

See-saw 90º & <120º

T-shaped 90º

octahedral 90º

Square pyramidal 90º

Square planar 90º

These shapes arise because certain atoms can

expand their octets (which means they can have more

than 8 electrons around them). The only atoms that

can expand their octet are those who have a “d”

sublevel available to them.

Lewis structures show the arrangement of valence electrons in

covalently bonded molecules and ions. We can use Lewis structures to

predict the properties of the bonds in molecules and ions, such as the

properties of bond length and bond energy, and use these bond

properties to predict physical and chemical properties of molecules

and ions.

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Lewis Dot Diagram steps

1. Count valence electrons

2. If charged, add or subtract the charge (ie. changes the number of electrons)

3. Arrange symbols, make no trains

4. Put in shared pairs

5. Determine remaining electrons

6. Determine who wants octet– Groups 1, 2, 3 do not want eight

– Groups 4-7 want eight (know who can expand octet)

7. Put in unshared electrons

8. Rearrange bonds as needed

Polar vs. Non-polar

• If a molecule contains only nonpolar bonds, it will be

a nonpolar molecule. (seven diatomics, ozone, and diamond)

• If a molecule contains polar bonds it is not

necessarily a polar molecule. The shape of a

molecule and the polarity of its bonds together

determine whether the molecule is polar or

nonpolar.– If the shape of the molecule is symmetrical (which means that there

are no unshared electrons on the central atom), then the molecule will

be nonpolar.

– If the shape of the molecule is asymmetrical (which means that there

is at least one unshared pair of electrons on the central atom), then

the molecule will be polar.

Dipole moment

• The dipole moment measures the polarity of a

molecule (based on shape of molecule)

– The larger the dipole moment, the more polar the

molecule

– The greater the charge at the ends of the dipole

and the greater the distance between the charges,

the greater the value of the dipole moment

Bond Lengths

• In general, single bonds are longer than double

bonds, which are longer than triple bonds.

• How do they determine bond length?

– X-ray diffraction

– Neutron diffraction

– X-ray crystallography

– Microwave spectroscopy

FYI – Computational Chemistry

• Computational chemistry is simply the application of chemical, mathematical and computing skills to the solution of interesting chemical problems.

• It uses computers to generate information such as properties of molecules or simulated experimental results. Some common computer software used for computational chemistry includes:

– Gaussian xx (Gaussian 94 currently)

– GAMESS

– MOPAC

– Spartan

– Sybyl

• Computational chemistry has become a useful way to investigate materials that are too difficult to find or too expensive to purchase.

• It also helps chemists make predictions before running the actual experiments so that they can be better prepared for making observations.

• For instance, you can calculate: – electronic structure determinations

– geometry optimizations

– frequency calculations

– transition structures

– protein calculations, i.e. docking

– electron and charge distributions

– potential energy surfaces

– rate constants for chemical reactions (kinetics)

– thermodynamic calculations- heat of reactions, energy of activation

Hybridization

• One of the theories used to explain molecular geometry (AP expects you to know “bold”)

– sp – linear (2 identical lobes)

– sp2 – trigonal planar, bent (3 identical lobes)

– sp3 – tetrahedral, trigonal pyramid, bent (4 lobes)

– sp3d – trigonal bipyramidal, seesaw, T-shaped, linear (5 identical lobes)

– sp3d2 – octahedral, square pyramidal, square planar (6 identical lobes)

• Sigma bonds overlap end to end

• Pi bonds overlap above and below

πσ verses

• A sigma bond results in a single bond

• A sigma bond and a pi bond results in a

double bond

• A sigma bond and 2 pi bonds results in a triple

bond

• Therefore the picture from the previous slide

represents a double bond

you? tell and can What πσ

Molecular Orbital Theory

• In the molecular orbital theory, as two atomic

orbitals approach each other they create two

molecular orbitals.

– Bonding orbital of lower energy

– Antibonding orbital of higher energy

• Type of covalent bond (single, double, or triple)

can be predicted using the formula

2

## electronsgantibondinelectronsbondingorderbond

−=

Molecular Orbital Diagram

Resonance

• Some Lewis dot diagrams can have a variety of

arrangements

• Ex. Carbonate ion CO32-

• Resonance explains the disparity in bond

lengths

Van der Waals forces (Intermolecular Forces)

• Ion-dipole forces – attraction between an ion and a polar molecule

• Dipole-dipole forces – attraction between

two polar molecules

• Dipole-induced dipole forces – attraction between a polar and nonpolar molecule

• Dispersion forces – attraction between two nonpolar species

Very Important van der Waals force

• Hydrogen bonds (FON) -- attraction between

molecules containing fluorine, oxygen, or

nitrogen and hydrogen…the hydrogen is

electron starved…holds hands with

neighboring F, O, or N

Three famous molecules

• H2O

• NH3

• HF H:F:

::

Intermolecular Forces of Attraction

• A substance will normally exist in the solid,

liquid, or gas phase depending upon the force of

attraction between the particles

– A solid substance will have strong intermolecular

forces of attraction (lots of snuggling)

– A gaseous substance will have weak intermolecular

forces of attraction between the particles.

– A liquid substance has intermediate forces of

attraction between its particles

Three States of Matter