# week 3 notes.docx

7/28/2019 # Week 3 Notes.docx

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AVERAGE VALENCE ELECTRON ENERGY (AVEE)

ELECTRONEGATIVITY

Electronegativity () is defined as the relative ability of an atom to attract electrons to itself in a chemical

compound. Elements with high electronegativity tend to acquire electrons in chemical reactions and are

found in the upper right corner of the periodic table (where the ionization energy is highest and electron

affinity is lowest).


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Elements with high electronegativity ( 2.2) have very negative affinities and large ionization potentials,

so they are generally non-metals and electrical insulators that tend to gain electrons in chemical reactions

(i.e., they are oxidants). In contrast, elements with low electronegativity ( 1.8) have electron affinities

that have either positive or small negative values and small ionization potentials, so they are generally

metals and good electrical conductors that tend to lose their valence electrons in chemical reactions (i.e.,

they are reductants).

IONIC BONDS

The reaction of a metal with a non-metal usually produces an ionic compound, that is, the electrons are

transferred from the metal (the reductant) to the non-metal (the oxidant). Ionic bonds are formed when

positively and negatively charged ions are held together by electrostatic forces.

If is negative, it means that energy is released when oppositely charged ions are brought together froman infinite distance to form an isolated ion pair (not crystalline lattice).

As they get closer together ( ), repulsive interactions between the electron clouds of each atombecome stronger than the attractive interactions of the oppositely charged ions. The total energy of the

system is a balance between the attractive and repulsive interactions and it reaches its minimum at ,called the bond distance, the point where the electrostatic repulsions and attractions are exactly balanced.

The energy level at bond distance is called bond energy.


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LATTICE ENERGIES IN IONIC SOLIDS

Ionic compounds are usually rigid, brittle, crystalline surfaces with flat surfaces that intersect atcharacteristic angles. They are not easily deformed and they melt at relatively high temperatures. These

properties result from the regular three-dimensional arrangement of ions in the crystalline lattice and from

the strong electrostatic attractive forces between ions with opposite charges. In such an arrangement each

cation is surrounded by more than one anion and vice versa, so it is more stable than a system consisting of

separate pairs of ions, in which there is only one cation-anion interaction in each pair. may differbetween the gas-phase dimer and the lattice.


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The melting point is the temperature at which the individual ions have enough kinetic energy to overcome

the attractive forces that hold them in place. At the melting point, the ions can move freely, and the

substance becomes a liquid.


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The hardness of ionic materials (resistance to scratching or abrasion) is also related to their lattice energies.

Hardness is directly related to how tightly the ions are held together electrostatically, which is also

reflected in the lattice energy.

Also, the higher the lattice energy, the less soluble a compound is in water.

THE BORN-HABER CYCLE


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Enthalpy of sublimation (

), ionization energies (

) and disassociation energy (

) are all

positive. Electron affinities ( ) can be positive, zero or negative and are not that great inmagnitude. That means that lattice energy () is the most important energy factor in determining thestability of an ionic compound. We can find the lattice energy by using Hesss Law, which relates all

enthalpy changes from the reactions to the enthalpy of formation ():

LEWIS ELECTRON DOT SYMBOLS


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Lewis dot symbols are written as dots around the symbol of an element. The number of dots in the Lewis

dot symbol is the same as the number of valence electrons of the element, which is the same as the last

digit of the elements group number in the periodic table.

Lewis used the unpaireddots to predict the number of bonds that an element will form in a compound.

Atoms tend to lose, gain or share electrons to reach a total of eight valence electrons to completely fill the and orbitals (octet rule).DRAWING LEWIS STRUCTURES

1) Arrange the atoms to show specific connections. The central atom is usually the least electronegative element in the molecule or ion; hydrogen and

halogens are usually terminal.

2)

Determine the total number of valence electrons in the molecule or ion.3) Place a bonding pair of electrons between each pair of adjacent atoms to give a single bond.4) Beginning with the terminal atoms (start with atoms with the highest AVEE), add enough electrons to

each atom to give each atom an octet (two for hydrogen).

5) If any electrons are left, place them on the central atom.6) If the central atom has fewer electrons than an octet, use lone pairs (electron pairs not involved in

covalent bonding) from terminal atoms to form multiple (double or triple) bonds to the central atom to

achieve an octet.

This will not change the number of electrons on the terminal atoms. If the most stable structure leaves the molecule or ion electron-deficient (less than an octet), it is

likely to react with another ion or molecule to gain an additional pair of electrons.

FORMAL CHARGE

Formal charges are used as a bookkeeping method for predicting the most stable Lewis structure for a

compound.

( ) Formal charges do not represent the actual charges on atoms in a molecule or ion. The sum of the formal charges on the atoms within a molecule or an ion must equal the overall charge

on the molecule or ion.

Typically, the structure with the most charges on the atoms closest to zero is the more stable Lewisstructure.

In cases where there are positive or negative formal charges on various atoms, stable structuresgenerally have negative formal charges on the more electronegative atoms and positive formal charges

on the less electronegative atoms.

RESONANCE


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Bond polarity the extent to which it is polar is determined largely by the relative electronegativities ofthe bonded atoms. If the electronegativities of the bonded atoms are not equal, the bond is polarized

toward the more electronegative atom. A bond in which the electronegativity of B () is greater than theelectronegativity of A () is indicated with the partial negative charge on the more electronegativeatom:

POLAR COVALENT BOND ENERGY


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DIPOLE MOMENTS

The asymmetrical charge distribution in a polar substance produces a dipole moment (). The dipolemoment is defined as the product of the partial charge (in Coulombs) on the bonded atoms and thedistance (in meters) between the partial charges (distance of charge separation):

Calculating the partial charge from the equation above and then dividing it by full charge if the substancewere completely ionic will give you the percent ionic character of the substance.


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( )

PROPERTIES OF COVALENT BONDS

In the Lewis bonding model, the number of electron pairs that hold two atoms together is called the bond

order.

When analogous bonds in similar compounds are compared, bond length decreases as bond order

increases. Molecules or ions whose bonding must be described using resonance structures usually have

bond distances that are intermediate between those of single and double bonds.

Triple bonds between like atoms are shorter and stronger than double bonds, which are in turn shorter and

stronger than single bonds. However, bonds of the same order between different atoms show a wide range

of bond energies.


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Bonds between hydrogen and other atoms become weaker as we go down within a column in theperiodic table because as atoms become larger, the region of space in which electrons are shared

becomes smaller.

Bonds between like atoms usually become weaker as we go down a column. As two bonded atomsbecome larger, the region between them occupied by bonding electrons becomes smaller. Exceptions

are NN, OO and FF single bonds, which are unusually weak.

We can estimate the enthalpy for a chemical reaction by adding together the average energies of the bonds

broken in the reactants and the average energies of the bonds formed in the products and then calculating

the difference between the two. If the bonds formed in the products are stronger than those broken in the

reactants, then energy will be released in the reaction ( ): The sign is used because we are adding together average bond energies, hence this approach does

not give exact values for .

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