AP Chem Unit 2

2.1 Types of Chemical Bonds

• electronegativity values for the representative elements increase going from left to right across a period and decrease going down a group valence electrons shared between atoms of similar electronegativity constitute a nonpolar covalent bond

• valence electrons shared between atoms of unequal electronegativity constitute a polar covalent bond

  • atoms with the higher electronegativity will develop a partial negative charge relative to the other atom in the bond

  • in single bonds, greater differences in electronegativity lead to greater bond dipoles

    • all polar bonds have some ionic character

      • difference between ionic and covalent bonding is not distinct, but rather continuum

• difference in electronegativity is not only factor in determining if a bond should be considered ionic or covalent

  • bond between metal an nonmetal = ionic

  • bond between two nonmetals = covalent

• examination of the properties of a compound is best way to determine type of bonding

• in metallic solid, valence electrons from metal atoms are considered to be delocalized and not associated with any individual atom

• ionic bonds formed between two ions by transfer of electrons

• polar covalent bonds are unequal sharing of electrons between atoms in a molecule

• nonpolar covalent bonds are equal sharing of electrons between atoms in a molecule

• pure covalent bond has < 0.4 electronegativty difference

• polar covalent bond has between 0.4-1.8 electronegativity difference

• ionic bond has > 1.8 electronegativity difference

2.3 Structure of Ionic Solids

• cations and anions in an ionic crystal are arranged in a systematic, period 3D array that maximizes the attractive forces among cations and anions while minimizing the repulsive forces

Increasing Electronegativity Difference Between Bonding Atoms

• Nonpolar covalent

  • little to no difference

    • 2 of the same atoms

    • C and H

• Polar covalent

  • moderate electronegativity difference

    • 2 different nonmetals

• Ionic

  • large electronegativity difference

    • 1 metal, 1 nonmetal

Ionic Bonds

• between atoms of metals and nonmetals with very different electronegativity

• cation is positive and anion is negative

• bond formed by transfer of electrons

• form crystalline solids

• produce charged ions all states

• conductors

• high melting point

Crystalline solids (salts)

• hard, but brittle

• specific pattern means if it is disrupted they will shift and shatter

• lot of ionic bonds in slat

• takes more energy to break them apart

Lattice Energy

• lattice energy is the energy required/released to form the crystal lattice structure of ionic compounds

• the greater the attractive forces, the more negative the lattice structure energy

  • negative means energy released when it is formed or the energy required to break it

• big charge = big attraction = high potential energy

• small ions = big attraction = high potential energy

2.4 Structure of Metals and Alloys

• metallic bonds form between metals and metals

• metallic bonds happen between metals with similar electronegativity

• sea of electrons is shared between nucleus

• compounds are not made

  • alloys are mixtures of metals

• substitutionally and interstitial alloys are based on the size of the metals atoms

Properties of Metals

• reflective

• high conductivity of electricity and heat

• malleable (ability to be beaten into a flat sheet)

• ductile

Metallic Bonding

• electron cloud around atoms

  • electrons are delocalized

• formed between atoms of metallic elements

• good conductors at all states

• lustrous

• very high melting point

Electron Flow

• when an electrical current is applied to the metal electrons can flow easily from one mental atoms to another in the direction of the current

  • leads to man metal sample's’ ability to conduct electricity

Metals Form Alloys

• metals do not combine with metals, they form alloys which is a solution of a metal in a metal

• alloys are homogeneous mixture of metals made by combining two or more metallic elements

• goal of making an alloy is usually to give greater strength, resistance to corrosion or other desirable properties

2.5 Lewis Diagrams

Covalent Bond Notes

• between nonmetals with similar electronegativity

• bond formed by sharing electrons

• several electrostatic interactions

  • attraction between electrons and nucleus

  • repulsions between electrons

  • repulsions between nuclei

Predicting Structure of Covalent Bonds

• lewis structure for single atoms

• shows only the valence electrons

• place one electron on each of four sides of atom betfore making pairs

Octet Rule

• atoms prefer to have eight valence electrons

• reason that atoms bond and form molecules is to create octet

  • exception: H only required a duet ( 2 electrons )

• lewis structures account for all the electrons and attempt to meet the octet rule

General Reminders

• Hydrogen only has one orbital

• single, double and triple bonds exist

  • triple are the shortest

• there are exceptions to octet rule

Lewis Dot Structures

1. add up the valence electrons

   • polyatomic ions assign an extra electron for each negative charge, subtract electron for each positive charge

2. select central atom and connect all atoms with single bond

   • central atom = Carbon or least electronegative atom

3. complete the octet of all atoms

   • unbonded electrons should be in pairs

4. check electrons

   • if you run out of electrons use multiple bonds

   • if there are extra electrons, put them on central

     • expanded octet

5. if molecule has net charge, place the structure in brackets with the charge in the upper right corner

Notable Rules

• sometimes the ordered list may help

  • HCl

• Central atoms is often written first

• H’s and F’s are terminal (on outside always)

• C’s love to hook together and do not like to be terminal

  • avoid unshared pairs on carbons

• oxygens do not bond together

  • except in O₂ and peroxide molecules

• avoid rings of three or fewer atoms

Expanded Octets

• elements with three or more energy levels have unused orbitals which allow for expanded octets

  • must be on n = 3 of p table or lower

  • only applies on the central atom of a molecule

• maximum number of domains is 6

Polyatomic Ions

• have charge, which means the amount of electrons are effected

• when you complete a Lewis structure of polyatomic ion, place the structure in brackets and place the charge on the upper right corner

Isomers

• molecules with the same chemical formula, but different arrangements of atoms

2.6 Formal Charge and Resonance

Formal Charge

• hypothetical charge the atom would have if we could redistribute electrons in bonds evenly between the atoms

  • subtracting nonbonding electrons, then subtract number of bonds connected to that atom in Lewis structure

• formal charge = number of valence shell electrons (free atoms) * number of lone pair electrons - ½ number of bonding electrons

• sum of formal charge of all atoms in molecule must be zero

  • in an ion should equal charge of ion

• formal charge ≠ actual charge

  • used as bookkeeping procedure

    • no indication of actual charge

• helpful with determining isomers

• lewis structure with formal charge closest to zero

Resonance

• triple bond > double bond > single bond

• if ≥ 2 lewis structures with same atom arrangement can be written for molecule/ion, actual distribution of electron is average shown from various lewis structures

• individual lewis structures "resonance forms”

• actual electronic structure of molecule called resonance hybrid of individual resonance forms

  • doe not fluctuate between resonance forms

  • always the average shown

• dotted line indicates that some models have bond there, but not all

• solid line indicates all models have bond there

Formal Charge Summary

1. all atoms must be counted

2. sum of all formal charges in a neutral molecule must be zero

3. sum of all formal charges in a n ion must equal the charge of the ion

4. smaller the formal charge on an individual, the better

5. when formal charge is needed, negative formal charges should go on the most electronegative atom

   • positive formal charges should go on the least electronegative atom

6. molecular structure in which all formal charges are zero is preferred

   • if must have nonzero formal charges, the arrangement with the smallest non-zero formal charge in preferred

7. lewis structures are preferable wen adjacent formal charges are zero or the opposite sign

8. when choosing from several lewis structures with similar distributions of formal charge, the structure with the negative charges on the more electronegative atoms is preferred

2.7 VSPER and Hybridization

VSPER

• valence shell electron pair repulsion theory

• steps

  1. count electron domains

     • single/double/triple bonds count as 1

  2. predict shape

     • arrangement of electron groups around central atom (electron geometry), exclude lone pairs (molecular geometry)

  3. VSPER chart

     • must memorize

Hybridization and Dipole