Ionic Bonds Ionic compounds High melting points Conduct

Ionic Bonds • Ionic compounds – High melting points – Conduct electricity • When dissolved in water • In molten state – Tend to be soluble in water – Crystallize as sharply defined particles

• Properties of ionic compounds are best explained by assuming the complete transfer of e- from one atom to another – Ex. 2 Na + Cl 2 2 Na. Cl Lewis Diagram is shown as follows: Do problems

• e- are transferred to achieve s 2 p 6 ECN • The Na+ and Cl- ions now have attractive force between them • When the ions are brought close together the force is strong enough to hold them together • This ELECTROSTATIC FORCE is what holds the ions together due to differing charges • This is the IONIC BOND

• Elements can be assigned oxidation #’s in ionic bonding • The oxidation # designates the number of elost or gained by the atom of an element in bonding • Positive values indicate loss of e • Negative values indicate gain of e. Ex. Phosphorous with five outer valence electrons tends to gain 3 e-. This allows P to achieve a stable octet configuration. P oxidation # in ionic bonding is 3 -

• Note: – In the examples of Na and Cl the Cl atom has a higher electronegativity value and ionization value. – This is the reason it is the Cl atom that gains the e • High ionization energy does not allow the Na to pull the e- from the Cl • High electronegativity allows the Cl atom to pull the e- from the Na • Low ionization energy of Na allows for easy removal of e- by Cl

Atom vs. Ion Radii • When atoms transfer e- the size of the ion changes • Positive ions become smaller – A e- is lost – The # protons remains same – Nuclear pull increases on e– Ion holds on to e- harder – Ion gets smaller

• The ion loses an energy level – Na 1 s 22 p 63 s 1 the 3 s 1 e- is transferred – Na new configuration 1 s 2 2 p 6

• Negative ions become larger – Gain an e– # of protons remains the same – Nuclear charge decreases – Ion has more difficult time holding on to e– Ion gets larger

• The extra e- forces the other e- to repel harder against each other to find optimal spacing along the p x, y, z axis's • The e- can only move further apart and out due to repulsion of inner energy level e-

• A tug of war happens between the two opposite charged and attracted ions for the transferred e • The amount of dispersion of the e- between the two ions results in the bond strength • The electronegativity and ionization energies play a vital role • Explain size change with tugs

Covalent Bonds • Atoms that share e- in bonding process • Atoms have very similar electronegativity values • Shared pair of e- constitute a covalent bond • There may be more than one shared pair of e - between the two atoms which constitute multiple bonds between them

• Most covalent bonds are formed between two nonmetal atoms • Covalent compounds typically – Have low melting points – Do not conduct electricity under any condition or state – Are brittle • When 2 or more atoms form covalent bonds, the particle is called a molecule • • • EDD Hybridization of C Bond number

Hybridization of Carbon • e- reconfigure by transferring energy to sp 3 configuration • Do so because only unpaired e- are involved in the covalent bonding process

• Atoms can form multiple numbers of bonds in covalent bonding • The # of e- an atom shares with another atom determines the # of bonds – Single bond • One shared pair – Double bond • Two shared pair – Triple bond • Three shared pair • Which is the strongest of the three?

• Bond Axis – The line joining the nuclei of two bonded atoms in a molecule • Bond Angle – The angle of the bond axes when one atom is bonded to two other atoms • Bond Length – The distance between nuclei along the bond axis • Length is not fixed • bonds act like stiff springs • Atoms vibrate as if the bonds were stretching and shrinking

• Bonds undergo – Bending – Wagging – Rotational vibrations • Causes bond angles and lengths to vary

Nonpolar and Polar Covalent Bonds • e- are shared in covalent bonds • Not all atoms have same electronegativity • It is therefore possible to have different scenarios – Two atoms will share fairly equally the e– Two atoms will not equally share the e-

• NON-POLAR COVALENT BOND – When electronegativity difference between 2 bonding atoms is 0. 00 -0. 20 – e- are shared almost equally – no atom has e- more than other – there is no pos. or neg. side of bond ex

POLAR COVALENT BOND • 2 atoms electronegativity difference is 0. 201. 67 • Atoms do not equally share e • One atom pulls harder • The harder pulling looks neg. and other side of bond looks pos. • Opposite sides (polar)

Coordinate covalent bonds • Occur when only one of the two bonding atoms donates BOTH of the e- in the bonding process • Prominent in the formation of polyatomic ions

Intramolecular forces • Bonds between atoms in a molecule or crystal

Resonance • Used to account for structure of molecules when DOT DIAGRAMS PRODUCED BY Octet Rule fail to give a representation of bonding that is in agreement with experimental evidence • Ex. SO 2 • Diagram shows one coordinate covalent bond and one double bond • Experimental evidence shows both bonds are same strength

• Each bond has characteristics of a single and double bond • Diagram does not show this so scientists wrote two structures as equivalents • Believed the extra e- in resonance structure moved back and forth from one bond location to another • This gave the appearance that the e- resonated between bond locations and was called Resonance • New evidence points to the concept that the e- are “spread” across the locations (remember e- act as waves as well as particles and set up an area of charge)

• This occurs at the bonds areas where resonance is occurring • The concept of resonance is used to help make Dot diagrams more accurate in describing the properties of a substance • Ex.

Polar and Non-polar Molecules • Observed that e- in bonding are not always shared equally • • • Caused polarity in the bonds Causes uneven charge distribution in the molecule Molecule could have pos. & neg. charged ends If happens, molecule referred to as a dipole If there are no polar bonds, not possible to have polar molecule • If polar bonds are present, the molecule may also be polar • Molecular polarity is also dependent on symmetry of molecule (shape)

• Molecules shape dependent on VSEPR Theory • VSEPR – Valence shell electron pair repulsion • Describes the repulsion of e- pairs in the valence shell dependent on how they are or are not bonded • Unshared-unshared > unshared-shared > shared • Ex. CO 2 H 2 O CH 4 NH 3 LASER DISC

Hydrogen Bonding • Occurs when H atoms are bonded to small atoms of high electronegativity • O, N, F H 2 O NH 3 HF • H has only a small share of e- pair forming the bond • H pos. charge on these dipole molecules is great – H acts almost as a bare proton • H is thus attracted to the unbonded pairs of e- on the neighboring molecules and forms a weak bond – – Bond called HYDROGEN BOND H is now bonded to two atoms H bond does not have to be in same plane Strength increases as the electronegativity difference increases

• H bond results in unusual properties for the substances involved – – Raises B. P. Raises M. P. Raises V. P. (vaporization pressure) Allows for surface tension in water ex. • Reason for these changes? – It takes extra energy to separate the formed weaker second H bond to separate the molecules • Also has effect on the crystalline structure of ice – Forces H 2 O molecules into 3 D hexagonal shape with a empty interior – Explains lower density of ice • Biological systems – Protein folding for shape ; holds DNA double helix together

Intermolecular Forces • Attractive forces between molecules • Dipole-dipole attractions

Van der Waals Forces • When dipole molecules come close together an electrostatic force is formed • Weak attractive forces • This attraction allows molecules to bond to each other • The greater the attraction of dipoles(electronegativity) the greater the attraction • Determines molecular substances state at room temp. – Solid, liquid, gas – Molecular mass also helps determine state

• How does CO 2 form dry ice? • N 2 form a liquid? – No dipoles – As two molecules are forced closer together they form temporary dipoles – Dipole is said to be induced • Forces that cause this are known as LONDON FORCES • Van der Waals forces then take effect

How do metals stay together? • Metallic Bonds – Metals are malleable, ductile, have luster and conduct heat and electricity well • What gives them these properties • Metal is a crystal – Every atom is surrounded by 8 or 12 other atoms – Metals typically have 1, 2 or 3 valence e • • • Do not form cov. or ionic bonds Crystals form when atoms lie next to each other E- orbitals overlap Atoms exchange e- ; said to be delocalized Properties determined by # of e- exchanging Electron-sea