AP Chemistry Unit 2

Electronegativity

measure of the ability of an atom to attract shared electrons

Trends of electronegativity

increases across groups —> higher nuclear charge which leads to a stronger attraction between electrons and nucleus

increase going up the periodic table —> smaller distances = stronger attraction

Nonpolar

equal sharing of electrons

Polar

unequal sharing of electrons

No difference in electronegativity (x<0.5) results in a..

nonpolar covalent bond and no bond dipole movement

slight difference in values of electronegativity (unequally) + results in a partial bond dipole movement (0.5-1.7) results in a..

polar covalent bond

Ionic Bonds

metal and non-metal atoms when they lose/gain electrons to form ions

• bonds are stronger when charges are higher

Properties of Ionic Substances

• brittle

• solid at room temp

• form crystals

• high melting and boiling points

• hard

• conduct electricity when dissolved and when molten (not solid)

• good insulators as solid

Why do ionic compounds not conduct electricity at solid?

because they are locked in rigid crystal positions due to the strong electrical forces holding them together

• dissolve it

• melt it

Why are ionic compounds brittle?

external force is applied which shifts the ions to slide near their like charges (pos/pos) —> repulsion between ions occurs which leads to the seperation of layers

Metallic

bond between metal atoms

• can be one type of metal, a pure substance, or for different types of metals (alloys)

• due to multiple metallic cations being attracted to a delocalized sea of valence electrons

IMF is stronger during metallic bonding when there are…

smaller metallic cations and when there are more valence electrons

Properties of Metallic Substances

• shiny (luster)

• malleable and ductile

• conduct heat/electricity

• metallic oxides are basic and ionic

• loses electrons to form cations

Properties of Nonpoalr and Polar covalent molecules (non-metals)

• non-lustrous, various colors

• brittle, hard or soft

• poor conductors

• nonmetallic oxides are acidic and covalent

• form anions by gaining electrons

Covalent Bonds

bonds between two nonmetals when they share valence electrons (polar or nonpolar)

• occur at the lowest energy state—> happens when attraction between nuclei is greatest for shared electrons, but the repulsions between electrons and between the nuclei is the least

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If atoms are too close together in a covalent bond, the nuclei will…

repel off each other

Bond length

distance between 2 nuclei from two bonded atoms in a molecule

• decreases in length if the number of bonds increase between two atoms (distance is increasing)

If the atoms are too far apart, the attraction will..

not be enough to hold them together

Bond energy

energy required when breaking a bond/energy released when a bond is formed

Larger atomic radii ___ the bond length

increase (more electron shells—> longer bond length between nuclei) (two tennis balls compared to two tennis basketballs)

Longer bond length ___ the bond energy

decrease —> the atoms are farther apart, making the electrostatic attraction weaker between them, meaning less energy is needed to break the bond (opposite for shorter bond length)

Bond Order

number of bonds between two atoms

Increasing bond order ____ the bond energy

increase —> more electrons involved and therefore more energy required/more coulombic attraction —> bond length decreases

Ionic compounds - lattice energy

amount of energy released when gaseous atoms come together to form a solid ionic compound (glue that holds ionic crystal tg) (similar to ionization energy)

• measures how strongly the positive and negative ions attract each other

• higher charges (multiplied together absolute value) —> stronger attraction—> more lattice energy

• smaller ions —> stronger attraction —> more lattice energy

Electrons in Metals

• do not stay with one atom

• able to move throughout the entire substance

• cations and electrons are attracted to one another through Coulombic Attraction (holds metal atoms together)

Delocalization

the sharing of electrons or electrical charge across multiple atoms (not confined to a single atom or a specific covalent bond)

Trend of Metals

# of valence electrons determines the amount of electrons in the delocalized sea of electrons

charge on cations + number of electrons increase —> greater attraction

ionic radius decreases —> attraction increases (shorter distance and less electrons)

Alloys

mixtures of metals

Interstitial Alloys

atoms aded to the metal are small and fit between metal atoms in existing holes

Substitutional Alloys

atoms added to the metal have similar radii so they replace the atoms in the lattice

Lewis Dot Diagram

1. count number of valence electrons—> add up all of them

2. Use single bonds to connect atoms —> central atom is the lowest electronegtivity

3. hydrogen and fluorine will never be the central atom

4. add lone pairs of electrons around each terminal atom in order to complete the octets for the terminal atoms

VSEPR Model

predicts the geometries of molecules and polyatomic ions

• shape determined by lone pairs/bonds on the central atom

• determined by areas of electron density

• lone pairs repel more than bonds—> compress the angle between atoms

Determine molecular geometry

1. draw lewis structure of molecule

2. count bonds and lone pairs around central atom (multiple bonds count as one charge cloud) (single electron —> one charge cloud)

3. use electron domains/lone pairs to determine which shape

Potential Energy Graph:

graph of potential energy vs distance between atoms

• helps us find equilibrium bond length, bond energy, and seperation between atoms at which potential energy is lowest

Too far apart for potential energy graph

Atoms do not interact, meaning that the stored energy of interaction (potential energy) is equal to 0 —> not a covalent bond

Stable potential energy:

stable covalent bond, distance between the nuclei of the two atoms is the bond length (74 pcm)

at the ideal bond distance (74pcm) the attraction between each nucleus and shared electrons is balanced by the repulsion

Too close for potential energy:

high positive potential energy because their negative electron clouds push and overlap each other, creating strong repulsive forces that spike the potential energy (unstable)

When determining where to draw a bond (Ex. where to draw Cl-Cl on a Br-Br potential enery graph), think about..

1. atomic radius

2. bond length (if larger atomic radius, then longer bond length)

3. bond energy (if shorter bond length, then more bond energy required to break it)

4. - If the bond has a longer bond length, you draw it closer to the right and reciprocated

5. - If the bond has a higher bond energy, it is more negative poential energy and reciprocated

Bonds with a higher bond order (triple or double) are _____ in bond length

shorter in bond length —> the more shared electrons means a stronger attraction and pulls the electrons closer, creating a smaller distance

Smaller atomic radii means..

A shorter bond length —> results in a higher bond energy because distance is smaller meaning forces are stronger

Which bond is harder to break, shorter or longer?

Shorter because the distance between the electrons is closer, meaning the forces are stronger making it harder to break.

Higher bond order —>

results in a shorter bond length because more electrons means more attraction between those electrons and the nucleus, resulting in smaller radius. —> higher bonding energy because its stronger

2 areas, 0 lone pairs

linear structure 180 degrees

3 areas, 0 lone pairs

trigonal planar 120 degrees

3 areas, 1 lone pair

bent <120 degrees

4 areas, 0 lone pairs

tetrahedral 109.5 degrees

4 areas, 1 lone pair

trigonal pyramidal <109.5 (107.5)

4 areas, 2 lone pairs

bent <109.5

Formal charge

valence electrons - (lone pair electrons + 1/2(bonding electrons)

Hybridization

atomic orbitals on an atom mix to form hybrid orbitals —> mix of orbitals of s and p

Hybridization Remember

2 areas —> sp

3 areas —> sp2

4 areas —> sp3

5 areas —> sp3d

6 areas —> sp3d2

Sigma Bonds

single overlap of the orbitals from a single bond are sigma bonds—> strong bonds

for 5 areas molecular geometry, order is..

trigonal bypyramidal (90-120), seesaw (90 and 120), t shaped (90), and linear (180)

for 6 areas molecular geometry, order is..

octahedral (90), square pyramidal (90), square planar (90)

Pi bonds

double and triple bonds are formed from unhybrdized p orbitals —> contain 1sigma and 1pi bond

• unable to rotate —> geometric isomers

Pi bonds (triple bonds) (bond length + energy) are…

strongest and shortest bonds with the highest energy

Single Bonds (sigma bonds) (bond length + energy) are..

the longest and weakest bonds with the lowest energy

Determining polar/nonpoalr from lewis dot structures

Highly symmetrical geometries (linear, trigonal planar, tetrahedral, octahedral) —> nonpolar

Asymmetrical structures (bent trigonal pyramid) —> polar