Chemical Bonding

UNIT 4 Chapters 15 & 16

Ionic Bonding

 The bond in ionic compounds

(two ions)



Held together tightly



High melting points

Compounds are formed from chemically bound atoms or ions Substances become more stable through chemical bonding, where 2 or more atoms are joined together by a simultaneous attraction.

Valence electrons are electrons in the highest occupied energy level of an atom ( the last shell).

Bonding involves only electrons

Bonding involves only the valence electrons.

Na Cl

i.

Ionic Bonding ii.

Covalent Bonding iii.

Metallic Bonding

Ionic Bonds occur when the more electronegative element “steals” the electron pair away from the other atom.

The atom that has stolen the electron pair becomes a negative ion (anion) while the “victim” becomes a positive ion (cation). The two atoms are held together by their opposite charges.

Can you predict which atoms will gain electrons and which will loose electrons by looking at the trend in electronegativity?

Increase in Electronegativity

When you consider that for an ionic bond to form there must be a great deal of difference in electronegativity between the atoms, can you predict what two types of atoms allow this to occur?

Metals Non Metals

Ionic Properties

Why do most ionic compounds have similar properties?

We can hypothesis that it is due to the bonds formed between the ions, holding them firmly in a rigid structure

Forming ions

 Na and Cl which one will lose electrons which one will gain electrons  Write out Tin’s electron configuration what will it do??

Lattice energy

 The change in energy that takes place when separated gaseous ions are packed together to form an ionic solid

The anions and cations in an ionic compound are locked in a regular neutrally charged structure, held by the balance of attractive bonds and electrical repulsion.

The component ions in such crystals are arranged in repeating three-dimensional (3-D) patterns.

Alkali metals combine with halogens in 1:1 ratios since alkali metals need to lose 1 e 1 and halogens need to gain 1e 1 .

Alkaline earth metals combine with halogens in 1:2 ratios since alkaline earth metals need to lose 2 e 1 and halogens need to gain 1e 1 .

Lewis Structures can be used to illustrate the formation of ionic bonds.

Be +

2

F F 1-

[

Be

]

2+ F 1-

Write an equation with electron dot diagrams to illustrate the formation of aluminum chloride.

Al +

3

Cl Cl 1-

[

Al

]

3+ 1 Cl Cl 1-

Lewis Structures

 Duet Rule = applies to H and He and states these two atoms are stable with 2 electrons in their outer shell  Octet Rule= elements are most stable with 8 electrons in their outer shell

1.

Most are crystalline in structure 2.

High melting/boiling points 3.

Electrically neutral 4.

Can conduct electricity when melted or in aqueous solution 5.

Hard/ Brittle

LEWIS DOT STRUCTURES: Elements Board Practice Elements #1-20

Covalent Bonding

Br + Br Br Br O + O O O

Covalent Bonding

 Electrons are shared by nuclei  Polar covalent bonds – unequal sharing of electrons

Types of Bonds

1) Single bond – 1 pair of e- are shared - lowest in energy -longest bond length 2) Double bond – 2 pairs of e- are shared 3) Triple bond- 3 pairs of e- are shared

The most common chemical bond results when the nuclei of 2 atoms are attracted to a pair of shared electrons. If the sharing is equal, because the atoms are the same, this is called COVALENT BONDING

H H

Electron pair

H H

 One atom becomes slightly positive the other slightly negative

The force of attraction of an element’s nucleus for electrons is called electronegativity (En). Atoms of different elements have different electronegativities. The higher the En, the stronger the attraction for electron pairs.

Difference in En?

HF

The bonding electrons are on the average closer to the fluorine than to the hydrogen atom.

The movement of the negatively charged electrons away from hydrogen toward fluorine, due to a difference in electronegativity, builds up a partial negative charge on the fluorine and a partial positive charge on the hydrogen.

This is not a complete transfer of an electron from hydrogen to fluorine; it is merely a drifting of electrons toward fluorine.

H

H

When a charge separation of this type is present, the molecule possesses an electric dipole, and the bond is called a POLAR COVALENT BOND , or simply a POLAR BOND .

H

Polar covalent bond (polar bond)

 covalent bond joins two atoms of different elements and the bonding electrons are shared unequally

Resonance

= occurs when more than one valid Lewis structure can be written for a particular molecule Ex. NO 3 -1

Exceptions to the Octet Rule

1) B and Be usually have less than 8 electrons 2) Elements in the 3 rd energy level and above can have more than 8 electrons in their outer shell

Non-polar covalent bond

 bonding electrons are shared equally

1.

Soft and squishy 2.

Low boiling/melting points 3.

Tend to be more flammable 4.

Do not conduct electricity 5.

Usually non-soluble in water

Electrostatic attraction force between the cation and free electrons

.

 Any successful bonding model for metals must account for the typical physical properties of metals: malleability, ductility, and efficient and uniform conduction of heat and electricity in all directions.

 Most metals are durable and have high melting points.

 These facts indicate that the bonding in most metals are

strong

and

nondirectional.

Metal atoms are arranged in very compact and orderly patterns.

i) Body-centered cubic ii) Face-centered cubic iii) Hexagonal close-packed

1.

Can conduct electricity (free electrons) 2.

Malleable (put into shape) 3.

Ductile ( made into wires) 4.

Good conductors of heat and electricity 5.

Metals are usually shiny

Lewis Structures for Molecular Compounds

N N

Bonding capacity is the number of covalent bonds (shared electron pairs) that an atom can form. Covalent molecules often consist of atoms of different elements, with different bonding capacities.

How do we decide on their structural arrangement , when we draw structural formulas?

STEP 1:place the single atom in the center and other atoms around it evenly spaced

H H C H H

STEP 2: count the total # of valence e for all atoms involved in the bonding

Carbon: 1 carbon with 4 valence electrons (1x4) = 4 Hydrogen: 4 hydrogen with 1 valence electrons (4x1) = 4

CH 4 4+4 =8

STEP 3: place the electrons in pairs between the central atom and each non-central atom CH 4 4+4 =8

H H C H H

STEP 4: place the remaining electrons around the non central atom until each has 8 electrons (H atoms have) Ex: AsBr 3 only 2e 5 + (7x3) = 26 Br As Br Br

STEP 5: if electrons remain they are placed in pairs around the central atom Ex: AsBr 3 5 + (7x3) = 26 Br As Br Br

STEP 6: if the central atom is in group 14, 15, 16, 17 or 18, the octet rule must be satisfied by moving electron Ex: SO 2 6 +(6x2) pairs from non-central atoms, creating multiple bonds.

=18

O

S

O

1. Central B(6e ) and Be(4e ) will have less than 8 electrons 2. If the central atom is in energy level 3 or more it may have more than 8 electrons around it (these energy levels can have 18e )

SCl 4 6+(7x4) = 34

Cl Cl S Cl Cl

NOTE: central atom(s) tend to have the highest bonding capacities or/and the lowest En.

Draw Lewis Structures for the following: H 2 O, NF 3 , Cl 2 , SnCl 2 , PCl 5 , SO 3 , BeCl 2 , C 2 H 6 , C 2 H 2 , ClF 3 , CHCl 3, ICl, O 2 , N 2 , SF 6 , CO 2 , BF 3 , C 2 H 4 , O 3 , IF 7

Compounds are arranged in many different shapes

The VSEPR Theory states that because electron pairs repel, molecular shape adjusts so the valence-electron pairs are as far apart as possible.

VSEPR model

 Valence shell electron pair repulsion  Used to predict the geometry of molecules  The structure will minimize electron pair repulsions

LINEAR : the two bonding pairs arrange themselves at 180 °. They are connected in a straight line. Ex. CO 2 , BeF 2 , HCN, CS 2 Groups = 2 Pairs = 0

A molecule that has a tetrahedral shape has all four pairs of electrons bonded Group = 4 Pair = 0 The four electron pairs repel each other forming an angle of 109.5

°

Molecules with a bent shape have four pairs of electrons, but only two pairs are bonding pairs (two are lone pairs). Ex H2O, SO2 The bond angle is 109.5 ° Group = 2 Pairs = 2

A molecule with trigonal planar shape has three bonds all of which lie in the same plane Ex. Boron trifluoride The bond angle are 120 °.

Group = 3 Pairs = 0

A molecule with trigonal pyramidal shape has four pairs of electrons all repelling each other.

Groups = 3 Pairs =1 Ex. ammonia

Intermolecular forces play a key role in determining the physical and chemical properties of covalent compounds.

Van Der Waals consists of 2 possible types of forces: 1.

London Dispersion Forces 2.

Dipole-Dipole Forces

This is the only type of force present in non-polar covalent molecules.

It is the weakest of the intermolecular interactions caused by the motion of the electrons.

The strength of dispersion forces generally increases as the number of electrons in a molecules increases.

Ex. Halogen diatomic molecules.

- This occurs when

polar covalent bonds

are attracted to one another.

- Electrostatic attractions occur between oppositely charged regions. (partially ( –) and partially (+)).

- Dipole interactions are similar to but much weaker than ionic bonds.

Dipolar molecules

 Have a center of positive charge and a center of negative charge. aka: dipole moment Ex. HF

Dipole moment in NH

3

Dipole cancels out in CO

2

This is found in

polar covalent molecules

that have hydrogen that is bonded to a very electrostatic element (N 2 , F 2 , O 2 ) Hydrogen bonds are the strongest of the intermolecular forces.

Hydrogen > dipole-dipole > London Dispersion Bonds interactions forces

Hydrogen bonds are extremely important in determining the properties of water and biological molecules such as proteins.

The water molecule has a

bent

shape (105 °) and is considered to be polar and the universal solvent.

The attraction in water results from the intermolecular hydrogen bonds.

Surface tension

: the inward force, or pull that tends to minimize the surface area of a liquid - this surface tension tends to hold a drop of liquid in a spherical shape The higher the surface tension, the more nearly spherical is the drop of that particular.

Because of hydrogen bonding, water absorbs a large amount of heat as it evaporates or vaporizes.

The hydrogen bonds must be broken before water changes from the liquid to vapor state.

Vapor Pressure

the force exerted due to the gas above the liquid

Boiling Point

: occurs when the temperature at which the vapor pressure of the liquid is just equal to the external pressure.

Boiling leads to evaporation of a liquid. In the case of water, hydrogen bonds break in order for the liquid to vaporize.