Reading...
Play button
Play button
Use LEFT and RIGHT arrow keys to navigate between flashcards;
Use UP and DOWN arrow keys to flip the card;
H to show hint;
A reads text to speech;
20 Cards in this Set
- Front
- Back
|
9.1
Lewis Dot Symbols |
> atoms combine in order to achieve a more stable e- config.
> Max stability comes when an atom is isoelectronic (same # of e) w/ a noble gas (He, Ne, Ar, Kr, Xe) > Lewis Dot Symbol consists of the symbol of an element and one dot for each Val e- in an atom of the element. > Except Helium, number of Val e- each atom has is the same as the group number of the element. > elements in the same group have similar outer e- config and hence similar Lewis Dot Symbols. > Transition metals, Lanthanides, and Actinides have incompletely filled inner subshells, can't write Lewis structures for them. |
|
|
9.2
The Ionic Bond |
> elements most likely to form Cations in ionic compounds are Alkali Metals and Alkaline Earth Metals
> elements most likely to form Anions are the Halogens and oxygen. > Ionic Bond is electrostatic force that holds ions together in an ionic compound. |
|
|
9.4
The Covalent Bond |
> A bond which two e- are equally shared by two atoms.
> Covalent Compounds: are compounds that contain only covalent bonds. > shared pair of e- is often represented w/ a single line. (eg, covalent bond for H is: H - H) > In a covalent bond, each e- in a shared pair is attracted to the nuclei of both atoms. > Covalent bonding between many-e- atoms involves only the valence e-. |
|
|
The Covalent Bond (part 2)
|
> only two Val e- participate in formation of (eg) F2
> Other nonbonding e- are called lone pairs > Lone pairs: pairs of Val e- that are not involved in a covalent bond formation. > Lewis Structure represents covalent bonding, where shared e- pairs are shown either as lines or as pairs of dots between two atoms, > lone pairs are shown as pairs of dots on individual atoms. Only Val e- are shown in Lewis Structure. |
|
|
Covalent Bond
- Single, multiple bonds - bond length |
> Single Bond: type of covalent bond, two atoms are held together by one e- pair.
> Multiple Bonds: bonds formed when two atoms share two or more pairs of e- > Double Bond: if two atoms share two pairs of e-, covalent bond is called a double bond. (eg: Carbon Dioxide) > Triple Bond: when two atoms share three pairs of e- (eg: Nitrogen molecule, N2) > w/ the exception of C02, stable molecules containing carbon do not have lone pairs on the carbon atoms. > Bond length is the distance between the nuclei of two covalently bonded atoms in a molecule. > Multiple bonds are shorter than single covalent bonds. |
|
|
Octet Rule
|
> Formulated by Lewis: an atom other than H tends to form bonds until it is surrounded by 8 Val e-
> A covalent bond forms when there are not enough e- for each individual atom to have a complete octet. > Octet rules works mainly for elements in 2nd period of periodic table |
|
|
9.5
Electonegativity - polar covalent bond |
> In a covalent, it is expected that e-'s are shared equally among atoms. Spend same amt of time in vicinity of each atom.
> Polar covalent Bond, or Polar bond: e-'s spend more time in the vicinity of one atom more than another. Unequal sharing of e-, partial e- transfer or shift in e- density. > e-'s gravitate to more e- dense atom. (eg: HF molecule) > HF bond and other polar bonds, thought of as being between (nonpolar) covalent bond, sharing e-'s is exactly equal, and an ionic bond, Xfer of e-'s is nearly complete. > Electronegativity, helps distinguish a nonpolar covalent bond from a polar covalent bond - ability of an atom to attract toward itself the e-'s in a chemical bond. > Elements w/ higher electronegativity have greater tendency to attract e' than do elements w/ low elecnegative. > Elec-Neg is related to e- affinity and ionization E. > Atom that has high e- Aff (tends to pick up e- easily) and a high I.E (doesn't lose e- easily) has high Elec. Neg. |
|
|
Electronegativity (part 2)
|
> Electronegativity increases from left to right across a period (row), as metallic character of elements decreases.
> Elec-Neg decreases w/ increasing Atomic Number, and increasing Metallic character. > Most Elec-Negative elements - halogens, oxygen, N, and sulfur. (upper right hand corner) > Least Elec-Neg elements: Alkali & Alkaline earth metals (lower left corner) |
|
|
Electronegativity
Trends |
> Atom of the less Elec-Neg element give up its e-(s) to the atom of more Elec-Neg element.
> ionic bond joins atom of Metallic element w/ an atom of nonmetallic element > Atoms w/ comparable Elec-Neg tend to form polar covalent bonds, because shift in e- density is small. > No sharp distinction between Polar bond and Ionic bond. > General Rule: Ionic bond forms when the Elec-Neg difference between two bonding atoms is 2.0 or more. > Difference between Elec-Neg & e- Affinity: - Elec-Neg: signifies ability of an atom in a chem bond (with another atom) to attract shared e- - e-Affin: isolated atom's attraction for an additional e-. > Electron Affinity is measurable quality. Elec-Neg is an estimated number. |
|
|
Example 9.2 (p. 384)
Classify the following bonds as ionic, polar covalent, or covalent a) the bond in HCl b) bond in KF c) CC bond in H3CCh3 |
Strategy: use the 2.0 rule of Elec-Neg difference (fig 9.5)
a) Polar covalent. not larger than 2.0. More e- on Cl side. b) Ionic. larger than 2.0 difference c) Covalent, equal sharing. |
|
|
Electronegativity and Oxidation Number
|
> Oxidation number refers to the number of charges an atom would have if e-'s were transferred completely to the more Elec-Neg of the bonded atoms in a molecule.
eg) NH3 molecule. N forms three single bonds w/ H atoms. N is more Elec-Neg than H, e- density will be shifted from H to N. H donates e- to N. N would have total charge -3, H would have +1 Oxidation Number of -3 to N Oxidation Number of -1 to H > Number of e- a molecule gains or loses. |
|
|
9.6
Writing Lewis Structures - basic steps |
1) Write skeletal structure of compound, using chem symbols and placing bonded atoms next to on another
- Least Elec-Neg atom occupies central position (less Val e-) - H and F usually occupy terminal (end) positions in Lewis Structure 2) Count total # of Val e-'s present. - For polyatomic Anions, add # of neg charges to that total. - Polyatomic Cations, subtract # of positive charges from total. 3) Draw single covalent bond between Central atom and each of surrounding atoms. - Complete the octets of atoms bonded to central atom. (Vale shell of H atom is complete w/ 2 e-) - e- not in a covalent bond, must be shown as lone pairs. 4) If central atom has fewer than 8 e-, try adding DBL or Triple bonds. Use lone pairs from surrounding atoms to complete octet of central atom. |
|
|
9.7
Formal Charge and Lewis Structure |
> By comparing the # of e- in an isolated atom w/ # of e- associated w/ the same atom in a Lewis Structure, we can determine the distribution of e- in the molecule and draw the most plausible Lewis Struc.
> In an isolated atom, the # of e- associated w/ the atom is the number of Val e- > An Atom's Formal Charge is the electrical charge difference between the Val e- in an isolated atom and the number of e- assigned to that atom in a Lewis Structure. - Break the bond(s) between the atom and other atom(s) and assign half of the bonding e- to the atom. > Val e- Minus e- assigned to atom (after bonds are broken) |
|
|
Formal Charge
Rules |
1) For molecules, the sum of the charges must add up to Zero, because molecules are electrically neutral species.(eg, Ozone, 03)
2) For Cations, sum of formal charges must equal the positive charge. For Anions, sum of formal charges must equal the negative charge. |
|
|
Form Charge
Guidelines |
> sometimes there are more than 1 acceptable Lewis Structure
> We can select the most plausible Lewis Structure by using formal charges and following the guidelines 1) For molecules, a lewis structure where there are no formal charges is preferable to one in where there is a formal charge present. 2) Lewis Structures w/ large Formal Charges (+2, +3, and/or -2,-3) are less plausible than small Formal charges. 3) Among Lewis structures having similar distributions of formal charges, most plausible structure is one there negative formal charges are placed on the more electronegative atoms. |
|
|
9.8
Concept of Resonance |
> double headed arrow <----> indicates resonance structures.
> Two ore more Lewis structures to represent a particular molecule. > Neither resonance structure adequately represents the actual molecule. It is human invention. RULE: Positions of e- can be rearranged, not the Atoms. |
|
|
9.9
exceptions to the Octet Rule |
> octet rule applies mainly to the second-period elements.
> Exceptions fall into 3 categories: incomplete octet, and off number of e-, or more than 8 Val e- around the central atom. |
|
|
Exceptions to Octet Rule
1) Incomplete Octet |
> some compounds, # of e- surrounding central atom in a stable molecule is fewer than eight
(eg: Beryllium. has two Val e-'s in 2s orbital. NO way to satisfy octet rule for Be in this molecule) > Coordinate Covalent Bond: a covalent bond which one of the atoms donates both e-'s. |
|
|
Exceptions to Octet Rule
2) Odd-Electron Molecules |
> Some molecules contain odd number of e-'s (Nitric oxid, NO. and Nitrogen dioxide NO2)
- because an even # of e- are needed for complete pairing to reach eight, can't satisfy octet rule. - Odd e- Molecules sometimes called radicals. Highly reactive - Have tendency for unpaired e- to form a covalent bond w/ an unpaired e- on another molecule. |
|
|
Exceptions to Octet Rule
3) Expanded Octet Rule |
> Atoms of 2nd period elements can't have more than eight Val e- around the central atom
- but atoms of elements in and beyond the 3rd period of the periodic table form some compounds which more than eight e- surround the central atom. > Some expanded octets, a central atom from the 3rd period and beyond, octet rule satisfied for all atoms, but still Val e- left to place - in this case, the extra e- should be placed as lone pairs on central atom. > Elements in third period also have 3d orbitals that can be used in bonding > These orbitals enable an atom to form an expanded octet. |
