Chapter 2 - Lewis Structures

27/09/26

compounds

• chemicals that are combinations of two or more different elements in defined proportions

• they do not have overall charges

• they are isolable, which allows chemists to easily study them

what does it mean to be isolable?

it means that they can be isolated from condensed phases (liquid/solid) because they are neutral, allowing them to be easily studied

ions and their subdivisions

• species carrying an overall charge

• can be cations or anions (single positive or negative charged ____)

• ____ that contain a single atom are called simple ____

• ____ that contain multiple atoms are called complex, polyatomic, or molecular ___

• in contrast to compounds, ____ are not isolable due to thier overall charge, and need a counter-___ (an ___ of opposite charge) to form a net neutral compound and become isolable

chemical bonds

these consist of electrons that simultaneously feel the electrostatic forces of attraction for both of the positively charged nuclei of adjacent atoms.

by forming these, the elements achieve a lower overall energy than they would in their atomic form

there are three types: ionic, covalent, and, metallic

metallic bonds

occur when electrons are shared between many atoms simultaneously and are free to “flow” between the atoms. in elemental metals or metallic alloys (mixtures of two or more metals), the most important type of bond is the metallic bond

what decides between which kind of bond is b/w two atoms?

the electronegativity difference between the two atoms (the higher it is, the more likely its ionic)

electronegativity

• the tendency of an atom in a molecule to attract electrons to itself

• the thing that causes polarity in molecules

• it is not a property of the atom, it is the property of an atom within a molecule. it cannot be measured when the atom is outside of a bond, instead is calculated based on other characteristics

electronegativity and size

• typically, the more protons in the nucleus, the more positivity is there to attract more electrons, but in actuality, a high amount of protons means a high amount of electrons already with the atom, that repel any incoming electrons

Electronegativity trends on the periodic table, and a small ranking

the most electronegative atom is fluorine, and then oxygen, then chlorine, then nitrogen and then so on

homonuclear diatomic bonds and electronegativity

in homonuclear diatomic bonds (a bond between two identical atoms, think HOFBRINCL) are an example of non-polar covalent bonds, as polarity stems from electronegativity differences, which wont exist between two of the same atom

all heteronuclear covalent bonds are polar to some degree

polar covalent bonds

in these bonds, electrons are polarized toward the atom with greater electronegativity, which gives rise to a bond dipole

i.e., they are shared unequaly and the atoms have partial charges

ionic compounds (salts)

these contain ions held together by electrostatic forces

these are formed by a redox reaction b/w a metal and non-metal (have a large diff in EN)

covalent bonds

these exist when pair(s) of electrons are shared b/w two atoms

in general, an element will form one of these to achieve a closed shell, with an electron configuration matching a noble gas

lewis structure

a chemical representation that uses lines as bond pairs and pairs of dots for lone pairs

bond order

the number of bonding pairs of electrons b/w two atoms

b.o. = 1 = single bond, etc.

bond length

the distance b/w the two nuclei participating in the bond

bond dissociation energy

the amount of energy required to completely separate two bonded atoms

relationships b/w bond adjectives

• a high bond length usually means a low bond order

• and a high bond order means a higher dissociation energy

note: these are not laws, usually just trends

the Octet Rule

its the goal of any atom to fill its outer shell and have 8 valence electrons

note: hydrogen makes duets, as it only has one 1 shell and needs 2 electrons

note: theres exceptions to this rule

electron deficient species

some group 13 elements can exist with just 6 valence electrons. this does violate the octet rule, but it just happens, and they’re stable

on occasion, an electron deficient species could manage to fill its octet if it had a double bond, but often, this doesn’t happen because the bonded atoms are high EN

hypervalent species

these are species containing elements of the third and later periods of the p-block, and sometimes have more than an octet of electrons

free radicals

species containing upaired electrons

ex: nitric oxide

what is percent ionic character?

% ionic character = ((measured dipole moment of X-Y)/ (calculated dipole moment of X⁺Y^-)) x100

formal charge

represents the difference between the number of valence electrons present in the neutral atom and the number of valence electrons in a particular Lewis structure

formal charge values assume that the bonding electrons are shared equally b/w the bonded atoms

calculating formal charge

1. determine the number of valence electrons in the desired atom as if it was unbonded

2. count the outer shell electrons surrounding the bonded atom in the compound or ion (draw a circle around its lewis, count split bonds as 1e^-)

3. formal charge = (regular valence count) - (what it has in the structure)

4. label the formal charges:

   1. show only non-zero charges

   2. they must be indicated by putting the charge in a circle near the atom that has it

   3. dont show the number if its a singular charge

5. any unmatched formal charges become the overall charge of the atom

how to draw lewis structures

1. put the least EN atom as the central atom

2. count total valence e^-s

3. add/subtract formal charges

4. form the skeleton between the atoms using single bonds

5. add electrons to outer atoms (prioritizing high EN atoms)

6. add remainders to the central atom

7. move lone pairs from outer pairs to form double bonds w/ central to fil its octet (when possible)

   1. make the central hypervalent (if possible) to reduce FC

8. add FC’s

9. check for alternate structures, eliminate the least likely

checks to determine what makes a good lewis structure

1. minimized formal charges

2. valid octets

3. you haven’t made something hypervalent

4. high EN atoms have most electrons??? (if theres a - FC, put it on something EN)

resonance structures

when two or more valid Lewis structures can represent the arrangement of electrons in the same molecule/ion

only electrons in multiple bonds or lone pairs can “move“ between the resonant lewis structures

the arrangement of atoms does not change, and the overall charge doesn’t change either

what’s the best resonance structure? (in comparison to chemically reasonable)

a structure that has minimal formal charges, but keeps negative ones on more electronegative atoms (and vice versa). if appropriate, heavier elements can exceed the octet to minimize formal charge

what is a chemically reasonable resonance structure? (in comparison to best)

where the magnitudes of the FC are small, (maybe not minimized). these represent bonding that is typical of various elements and is consistent with experimental observations

what are isomers?

chemical species that have the same formula, but different arrangements of atoms

isomers vs resonance structures

isomers rearrange atoms

resonance structures rearrange electrons

to identify, look for changes in bonds/lone pairs

or movement of atoms

hypervalence and the octet rule

hypervalence only exists on the central atom. terminal atoms cannot have expanded octets

the octet rule is not exceeded unless it is necessary to reduce formal charges to to make a bond with all the atoms

RULE OF THUMB: when expanding the octet to minimize formal charge, the octet rule should not be exceeded if it results in placing a negative formal charge on an atom of lower electronegativity

formal charges, oxidation states, and partial charges

formal charges: assumes the electrons are shared equally between the two bonded atoms (covalent bonding)

oxidation state: assumes the bonding electrons are given to the more electronegative atom (ionic bonding)

partial charger: most representative of the actual charge distribution in a bond

oxidation state rules

1. the o state for a free element is 0

2. for a simple (monoatomic) ion, the o state is = to the net charge on the ion

3. hydrogen has a o state of +1, and oxygen has an o state of -2 in most compounds.

   1. exceptions: hydrogen in hydrides of active metals (-1)

why does oxygen typically have an o state of 2-?

because its highly EN, and therefore it “wins” all the electrons in its circle, meaning it almost always has 8 electrons to it, and with its usual 6 valence e^-, it always gets 2-

in a similar vein, fluorine always has a - because it brings 7 but always gets 8

edge cases of o states

if two of the same atom are bonded, the have the same EN, so they can split a bond 50/50

and, if two atoms are splitting a bond, they don’t have to share with their other neighbours if they don’t have the same EN

is it possible for an atom to have the same FC and o state?

yes, all homonuclear diatomic are, and there’s more