Ch2: Chemical Bonding & Lewis Structures Flashcards
What are anions
Ions carrying a negative charge
What are the three types of chemical bonds?
Ionic bonds, covalent bonds, and metallic bonds
What are cations?
Ions carrying a positive charge
What are ionic bonds?
oppositely charged ions (positive and negative) are held together by electrostatic forces/attractions between them
- electrons are lost or gained
What are covalent bonds?
held together by the mutual attraction to the nuclei of the atoms through the sharing of electrons
- electrons are shared
- non-metals
What are metallic bonds?
electrons are shared between many atoms all at one time and “flow” between atoms
What is electronegativity?
atoms in a bond/molecule/molecular ion and their tendency to attract electrons towards itself
eg) Fluorine is the most electronegative element (top right) and the element’s electronegativity decreases as you move away from fluorine
- electronegativity values can be used to predict the nature of bonds that form between elements
What are ionic compounds?
also known as SALTS
- contains ions held together by electrostatic forces/attractions
- formed by redox reactions between metals and non-metals (large differences in electronegativities)
ALL ABOUT BALANCING THEIR CHARGES - NET NEUTRAL
How do you identify ionic bonds?
- contains any element from group 1 and 2 and any element from groups 17 and 18
- elements in groups 1 and 2 always lose electrons, carry the positive charge, and never participate in covalent bonds
What are nonpolar covalent bonds?
atoms in the bonds have the exact same electronegativity (found in homonuclear diatomic species)
- electrons are shared equally between the atoms
What are polar covalent bonds?
- electrons are shared UNEQUALLY
- polarized towards the atom with greater electronegativity
- atoms have PARTIAL CHARGES (+ and -)
- elements with the greatest electronegativity differences will have the most polar bonds
- one atom has more electron density than the other
What is a redox reaction?
A reaction in which the metal loses electron(s) and the nonmetal gains electron(s)
LEO says GER
lose electrons - oxidization
gain electrons - reduction
What are the three characteristics to describing a bond?
- bond order
- bond length
- bond dissociation energy
What is bond order?
number of bonds - single, double, or triple bond
What is bond length?
distance between the two nuclei participating in the bond
What is bond dissociation energy? (bond strength?)
amount of energy required to completely separate the two bonded atoms
Describe the relationship between bond length, bond order, and bond dissociation energy.
As bond order increases, bond length decreases, and bond dissociation energy also increases!
What is percent ionic character?
A ratio that helps us determine if the bond is covalent or ionic
Although covalent and ionic bonds exist, no bond is _____________________________________
Although covalent and ionic bonds exist, no bond is TRULY ONE OR THE OTHER!!
What are the two concepts associated with Lewis Structures?
the octet rule and formal charge
What is the octet rule?
an atom that satisfies the octet rule has EIGHT VALENCE ELECTRONS
How do you determine if an atom satisfies the octet rule?
- Count the lone pair electrons
- Count all bonding pair electrons (twice the # of bonds)
- Add numbers of from 1 and 2, if it adds to 8, then you have a full octet!
- if not then you have an incomplete octet or hypervalent (expanded octet)
What are the 4 exceptions to the octet rule?
- Hydrogen - can only have max 2 electrons/only one bond
- Electron deficient species - incomplete octet, elements in group 13 can have just six valence electrons in their valence shell
- Hypervalent - expanded octet, opposite of electron deficient, containing elements in the p-block third row and below that may have more than an octet of electrons
- Free radicals - species containing unpaired electrons (very reactive), an odd # of electrons indicate that you’re dealing with a radical
What is formal charge?
the difference between the # of valence electrons in a neutral atom compared to the # of electrons (lone pairs and bonded pairs) of an atom in a molecule
- formal charges assume that electrons are shared EQUALLY between the bonded atoms
How do you calculate formal charge? (steps)
- determine the # of valence electrons in the neutral unbonded atom
- determine the # of lone pair electrons and the # of bonds
- subtract FC=(1)-(2)
- only label NON-ZERO formal charges with a +/- in a circle
When the bonded atom has _______ electrons in the molecule than the neutral atom, it will have a ____________formal charge
When the bonded atom has MORE electrons in the molecule than the neutral atom, it will have a NEGATIVE formal charge
= too much electron density
When the bonded atom has _______ electrons in the molecule than the neutral atom, it will have a ____________formal charge
When the bonded atom has LESS electrons in the molecule than the neutral atom, it will have a POSITIVE formal charge
= reactive because it wants electron density
How do you calculate formal charge? (formula)
FC = # valence electrons - # lone pair electrons - # bonds
What is the relationship between formal charge and overall charge of the molecule?
the formal charge should be EQUAL to the overall charge of the molecule!
Which of the following contain only ionic bonds?
NaN3
CaF2
H2SO4
NH4Cl
Li3PO4
CaF2 is the only one that contains only ionic bonds
NaN3 contains Na, which only has one electron to lose, so it cannot be bonded to every N… The N must be bonded to each other covalently
CaF2 contains Ca, which can lose two electrons, so it can ionically bond to the two F
H2SO4, NH4Cl, and Li3PO4 all contain polyatomic ions, so they must contain covalent bonds
Where does the overall charge of a molecule come from?
The overall charge of a molecule comes from its formal charge
Terminal halogen atoms (group 17) should have a formal charge of _________
Terminal halogen atoms (group 17) should have a formal charge of ZERO
- full octet - single bond and 3 lone pairs
What are the steps to drawing Lewis structures?
- Count the # of valence electrons available (in each atom)
- Adjust the total to account for overall charges (add or remove e-)
- Lay out atoms so that the central atom is the least electronegative (H and F are NEVER central atoms)
- Draw single bonds between all the terminal atoms to the central atom
- Add in lone pairs starting with the more electronegative atom
- Check formal charges, can they be minimized?
- Form multiple bonds or make the central atom hypervalent to account for all valence electrons
- Check the octets of all atoms
- Count the total # of electrons drawn (should be the same as 1)
- Calculate the final formal charges (should be equal to the overall)
- Check for resonance structures
What are resonance structures?
When two or more valid Lewis structures can represent the arrangement of electrons in the same molecule or molecular ion, each structure is called a resonance structure
- electrons can move around if they are in multiple bonds or lone pairs
- the real structure of a molecule with resonance structures is an average weighed between all of them
What are the two conditions for resonance structures?
- You CANNOT change the arrangement of atoms - keep them attached in the same way
- CANNOT BREAK SINGLE BONDS, but multiple bonds will vary - Formal charge must be equal to the overall charge (or else it indicates that something has moved when it should not have been)
What are the three patterns to identifying resonance structures?
- double bond - single bond - double bond
- double bond - single bond - lone pair
- double bond - single bond - formal charge
What does the best resonance structure look like?
The best resonance structure has:
- minimal formal charges
- negative FC on more electronegative elements
- positive FC on less electronegative elements
- heavier elements can exceed the octet rule to minimize FC
What does a chemically reasonable structure look like?
A chemically reasonable structure has:
- full octets wherever possible
- lighter elements never exceed octet rule
- only heavier central atoms can exceed the octet rule
What is the difference between resonance and rotation?
changing the bonding pattern = resonance structure
changing the perspective = rotating the molecule
resonance does not equal rotation
What are isomers? How are they different from resonance structures?
Isomers = different arrangement of atoms but the same molecular formula
Resonance structures = same arrangement of atoms but different placements of electrons
What does hypervalent mean?
Central atoms (p block 3rd row or below) can exceed the octet rule BUT only if expanding the octet is beneficial in that it:
- Reduces formal charge
- Is needed to form bonds with all terminal atoms
- terminal atoms NEVER expand their octets
What is the difference between chemically reasonable and best for Lewis structures?
Chemically reasonable Lewis structures may have a SMALL formal charge, but BEST Lewis structures have minimized formal charges that are AS CLOSE AS POSSIBLE TO ZERO
- if there MUST be a formal charge, the negative formal charges will be on the most electronegative atoms and the positive formal charges will be on the less electronegative atoms
____________ put a negative formal charge on the central atom
NEVER put a negative formal charge on the central atom
Central atoms can have a ____________ formal charge but NEVER a ______________
Central atoms can have a POSITIVE formal charge but NEVER a NEGATIVE formal charge
What does a negative formal charge indicate?
excess electron density
- too many electrons on the central atom means a lot of negative charge in the middle, but like charges want to repel!!
What is the Rule of Thumb?
When expanding the octet to minimize formal charge, the octet rule should NOT be exceeded (or further exceeded) if it results in placing a negative formal charge on the central atom (atom of lower electronegativity)
What are the six rules of Lewis structures?
- Valence electrons are shown as dots or lines
- All electrons get assigned - either as lone pairs or in bonds
- Atoms are typically surrounded by octets of electrons
- The best Lewis structure has minimum formal charge (zero)
- If non-zero formal charges are necessary, the best structure will have negative formal charge on the most electronegative atom and positive formal charge on the least electronegative atom(s)
- When 2+ Lewis structures are possible, the actual arrangement is an average between the two (same with bond characteristics)
Summarize formal charges, oxidation state, and partial charges
Formal charges
= calculated as (# valence e-) - (lone pairs - bonds)
= assumes that electrons are shared equally (not true because of polar covalent bonds)
Oxidation state
= opposite of formal charges
= assumes that electrons are given to the more electronegative atom
Partial charges
= delta +/-
= not a full sharing of electrons, one will have slightly more/less
= no bond is truly ionic or covalent
= electrons are still being shared, just not equally
= applies to COVALENT bonds only
What are the oxidation rules?
- Oxidation state on a free element (on its own) is ZERO
- Monatomic ions have an oxidation state equal to their NET CHARGE
- Hydrogen has an oxidation state of +1
- Oxygen has an oxidation state of -2 in most compounds
Differentiate overall charge from partial charge
Overall charge = sum of all oxidation states
Partial charge = difference in electronegativity
Can terminal atoms expand their octets?
NEVER!!!
What is the rationale behind negative formal charges being placed on the most electronegative atom(s)
More electronegative atoms have a greater ability to attract electrons - greater ability to take on more negative charge
Why don’t partial charges apply to ionic bonds?
Partial charges ONLY apply to COVALENT bonds, THEY NEVER APPLY TO IONIC BONDS
= ions have full charges so they cannot have partial charges
How would you go about calculating the oxidation numbers of the elements in NaOH?
Na: (?)
O: -2(1) = -2
H: +1(1) = +1
You know that the overall charge must equal 0
So working backwards, Na must have an oxidation number of +1
Na: +1(1) = +1
O: -2(1) = -2
H: +1(1) = +1
= 0
