Final exam review Flashcards
What is a covalent bond?
A sharing of electrons between two non-metal nuclei
What is an ionic bond?
A transfer of electrons between a metal and a non-metal atom
What is metallic bonding?
Why most metals are solids
What do lewis dot structures represent?
- Valence electrons around an atom
- H follows the duet rule, it forms stable
compounds by sharing 2 electrons - Most main-group atoms follow the octet rule:
they tend to have 8 electrons (bonds + lone
pairs) in their valence shell
How do you quantify a bonds strength?
Ionic bonds - lattice energy
Covalent bonds - bond energies
What is lattice energy?
The delta H when gas phase ions form one mole of an ionic solid
- Always exothermic
- More exothermic lattice = stronger ionic bonds
- Calculated using Hess’ Law
- 3A + B = C, C = 2D so 3A + B = 2D
- H1 H2 (H1 + H2)
What does it mean to use the Born-Haber Cycle?
Using Hess’ Law to calculate lattice energies (weaker ionic bonds)
- k(Q1Q2/r)
- More exothermic LE’s have the elements with
the larger charge (BaTe over LiF), the larger
element (FeCl3 over FeCl2), or are smaller (ScCl3
over ScBr3)
How do you show polar covalent bonds?
- Partial charges
- Dipole; arrow points to more electronegative
atoms - Electrostatic potential map, more red means the
region is more electron-rich and more blue
means the region is more electron-poor
What is electronegativity?
The ability of an atom to attract shared electrons to itself in a chemical bond
- If delta EN is small (< 0.4) the bond is covalent
(example: Cl2)
- If delta EN is medium (0.4 - 2.0) the bond is polar
covalent (example: HCl)
- If delta EN is large (> 2.0) the bond is ionic
(example: NaCl)
What are the characteristics of an ionic compound?
- Large delta EN
- Cation-anion arrays
- Solids at STP
- High melting/boiling points
- Conduct as liquids
What are the characteristics of a covalent compound?
- Small delta EN
- Discrete molecules
- Solid/liquid/gas at STP
- Low melting/boiling points
- Insulating as liquids
What is a dipole moment?
- A measure of positive and negative charge
separation in a molecule - Dipole moment is the vector sum of individual
bond dipoles- EN is calculated; DM is measured by the
current flow when the sample is between the
charged plates of a capacitor - Molecules have a DM if they can be oriented
such that one side has a positive partial
charge and the other has a negative one
- EN is calculated; DM is measured by the
How do you draw a lewis structure?
- Count valence electrons from all atoms
- Connect atoms using single bonds
- Leftover electrons go to multiple bonds and lone
pairs to follow duet (H) and octet rules - The central atom is usually the least
electronegative- Exception: H is never the central atom
- If there is one of a certain type of atom, it is
usually the central atom
What does it mean for a molecule to have resonance?
- > or = 2 valid Lewis structures for one species
- The real structure is an average of the resonance
forms- Resonance structures aren’t in equilibrium
What are the two types of charge?
- Ionic: on entire ion
- Formal: on each atom
What is a formal charge?
- Oxidation state completely assigns electrons in a
bond to the more electronegative element - Formal charge divides electrons equally, ignoring
electronegativity- Formal charge = (valence electrons on neutral
atom) - [non-bonding electrons + (1/2 *
bonding electrons)] - There are 2 non-bonding electrons in 1 lone
pair
- Formal charge = (valence electrons on neutral
How do you solve for formal charge?
- For a molecule: E(FCs) = 0
- For a polyatomic ion: E(FCs) = its charge
- FCs allow us to evaluate resonance structures:
- Good resonance structures have low units of
FC (not ) is best) - Negative FCs should be on the most
electronegative atom; positive FCs should be
on the least EN atom
- Good resonance structures have low units of
What are the octet rule exceptions?
- Second-row elements can’t exceed an octet
since 2d orbitals don’t exist - Third-row and heavier elements often obey the
octet rule, but can use empty d orbitals to
exceed it - Some main group atoms typically hold fewer
than 8 electrons, (example Be (4), B (6) and Al (6))
What do you prioritize when finding the best resonance form of a structure?
Fewer formal charge units > obeys octet rule
What is a bond energy?
The energy required to break the bond of a gas-phase molecule
- Endothermic; it always takes energy to break a
bond
- Compounds with stronger bonds are more
stable and less chemically reactive
- General bond strength trends: X3Y > X=Y > X-Y
- BE’s are easy to obtain for molecules with only
one type of bond
What is average bond energies?
An average of the measured bond energies for a particular bond in several compounds
What can delta H rxn be estimated as?
Delta H rnx = E B.E. (bonds broken) - E B.E. (bonds formed)
- B.E. (bonds broken) - energy consumed
- B.E. (bonds formed) - energy released
What are the average bond lengths?
- Triple (shortest) < double < single (longest)
- Length decreases left to right across a period
- Length increases down a group
- In general, as bonds get longer, they also get
weaker
What is VSEPR theory?
Valence Shell Electron Pair Repulsion
- Predicts the shape of covalent compounds
- Regions of electron density orient around a
central atom as far apart as possible
What is molecular geometry?
- The shape of covalent compounds is described
by regular polygons and Platonic solids
- Trigonal planar
- Square planar
- Tetrahedral
- Octahedral - The central atom is placed in the center
- The corners of the shape are the positions of
surrounding atoms
What are electron groups?
- Molecular geometry is found by counting
electron groups on the central atom - These count as one electron group:
- Lone pairs
- Single bonds
- Multiple bonds
What is the corresponding electron geometry for each electron group?
- 2 groups = linear geometry
- 3 groups = trigonal planar geometry
- 4 groups = tetrahedral geometry
- 5 groups = trigonal bipyramidal geometry
- 6 groups = octahedral geometry
Characteristics of two electron groups
- Linear
- A central atom with two electron groups has
linear electron geometry - Bond angle is 180
- Examples: NO2, BeF2, HCN, SCN, CO
Characteristics of three electron groups
- Trigonal planar
- A central atom orients three electron groups 120
apart and coplanar
- Examples: SO2, SO3, NO3, NO2, BF3
Characteristics of four electron groups
- Tetrahedral
- A central atom puts four electron groups 109.5
apart in a tetrahedral orientation
- Examples: PO4, NH4, CF4, ClO4
Characteristics of five electron groups
- Trigonal bipyramidal
- A central atom with five electron groups adopts
a trigonal bipyramidal orientation - 90, 120, and 180 bond angles
- Examples: PF5, SF4, I3, ClF3
Characteristics of six electron groups
- Octahedral
- A central atom with six electron groups has
octahedral electron geometry - 90 and 180 bond angles
- All positions are equivalent
- Examples: PF6, SF6, XeCl4, BrF5
What are lone pairs?
- Lone pairs have hemispherical electron density
regions - Bonded electrons (not lone pairs) have
cylindrical electron density regions - Lone pairs occupy more space around an atom
than bonded electrons - Since lone pairs take up more space than
bonding electrons, they force atoms closer
together
What is AXE notation?
- A = central atom
- X = outer atoms
- E = lone pairs on central atom
AX2 and AXE
Two electron groups on a central atom, whatever they are, is always linear molecular geometry
Ax2E
- Electron geometry = trigonal planar
- Molecular geometry = bent
AX3E
- Electron geometry = tetrahedral
- Molecular geometry = trigonal pyramidal
AX2E2
- Electron geometry = tetrahedral
- Molecular geometry = bent
AX4E
- Electron geometry = trigonal bipyramidal
- Molecular geometry = see-saw
AX3E2
- Electron geometry = trigonal bipyramidal
- Molecular geometry = T-shaped
AX2E3
- Electron geometry = trigonal bipyramidal
- Molecular geometry = linear
AX5E
- Electron geometry = octahedral
- Molecular geometry = square pyramidal
AX4E2
- Electron geometry = octahedral
- Molecular geometry = square planar
What are the two main bonding theories?
- Valence bond theory
- Atomic orbitals mix to become hybrid
orbitals between two atoms - Molecular orbital theory
- Atomic orbitals mix to become molecular
orbitals that span the entire molecule
What is hybridization?
Mixing atomic orbitals (s, p, d, …) to form a new set of degenerate hybrid (sp, sp2, …) orbitals
- x atomic orbitals mix to x hybrid orbitals
- Hybridization explains the observed shapes of
molecules
Overview of hybridization
- 2 groups, sp hybridization, linear geometry
- 3 groups, sp2 hybridization, trigonal planar
geometry - 4 groups, sp3 hybridization, tetrahedral
geometry - 5 groups, dsp3 hybridization, trigonal
bipyramidal geometry - 6 groups, d2sp3 hybridization, octahedral
geometry
sp3 hybridization
- One s and three p orbitals to give four sp3
orbitals - sp3 orbitals have equivalent shapes and
energies - Each points to a different vertex of an imaginary
tetrahedron
sp3 orbitals
- sp3 orbitals are higher in energy than their
parent s orbitals but lower in energy than p
orbitals - Unlike p orbitals, hybrid orbitals have lobes of
different sizes - The larger lobe overlaps with other orbitals for
bonding
sp2 hybridization
- For three electron groups, one s and two p
orbitals hybridize to three sp2 orbitals - Trigonal planar geometry is obtained by
hybridization of one s and two p orbitals to give
three sp2 orbitals - One 2p orbital remains unhybridized
sp2 orbitals
- Trigonal planar geometry is obtained by
hybridization of one s and two p orbitals to give
three sp2 orbitals - One 2p orbitals remains unhybridized
What are sigma bonds?
- Head-to-head orbital overlap
- Electron density between nuclei
- All orbitals form sigma’s
What are pi bonds?
- Side-to-side orbital overlap
- p orbitals on adjacent atoms overlap
- Electron density above and below the
internuclear axis
What types of bonds are contained within single, double and triple bonds?
- Single bond = one sigma bond
- Double bond = one sigma and one pi bond
- Triple bond = one sigma and two pi bonds
sp hybridization
- sp hybridization is for two electron groups on a
central atom - One s and one p orbital hybridize to two sp
orbitals - Two p orbitals are unchanged
What are the problems with Lewis Theory?
- Often requires > or = 2 structures to describe
bonding (resonance - Doesn’t predict bond strength
- Doesn’t predict bond length
- Doesn’t predict magnetic behavior
- Doesn’t describe odd-electron species well
What can be done with molecular orbitals?
- The wavefunctions of two atomic orbitals can:
- be added, forming a bonding MO
- be subtracted, forming an antibonding MO - x atomic orbitals produce x MOs
Bonding vs. Antibonding MOs
- Electrons in bonding MOs are stabilizing
- Lower energy than atomic orbitals (AOs)
- Electron density is between nuclei
- Electrons in antibonding MOs are destabilizing
- Symbolized by *
- Higher energy than AOs
- Electron density is outside the internuclear
axis - Electrons in antibonding orbitals cancel
stability gained by electrons in bonding
orbitals
How do you use MO diagrams?
- Select correct diagram
- Count the number of valence electrons
- Add valence electrons
- Fill lowest energy MOs first (Aufbau)
- Each MO holds 2 electrons (Pauli Principle)
- Half-fill degenerate orbitals (Hund’s rule)
How do you find bonding order (BO)?
bond order = (bonding electrons - antibonding
electrons) / 2
- BO = 0: the bond is unstable and will not form
- BO > 0: the bond will form
- The larger the BO, the stronger and shorter the
bond
What are the two types of p orbital overlap?
- Head to head in a sigma bond
- Side to side in a pi bond
What are the two types of MO overlaps?
- Bonding - between nuclei
- Antibonding - outside nuclei
What is photoelectron spectroscopy?
- A molecule is bombarded with X-rays to remove
an electron - The electron’s energy is measured, giving the
relative energy levels
Inter vs Intramolecular forces
- Intermolecular - attractions between > or = 2
species
- Hydrogen bonds (4-50 kJ/mol of strength)
- Dipole-dipole (5-20 kJ/mol of strength)
- London force (0.1-5 kJ/mol of strength) - Intramolecular - hold atoms together in a
compound (covalent or ionic bond)
- Ionic bond (400-500 kJ/mol of strength)
- Covalent bond (100-400 kJ/mol of strength)
What are intermolecular forces (IMFs)?
- IMF strength determines phase:
- Solids have lots of strong IMFs
- Liquids have fewer and weaker IMFs
- Gases have none (theoretically)
What are dipole dipole forces?
- Electrostatic attractions between partial charges
of polar molecules - One molecule’s positive pole attracts the
negative pole of another
What is hydrogen bonding?
- A strong dipole-dipole force that occurs in
molecules with H-X bonds - X can only be N, O and F because they are:
- Small, molecules approach closely
- Electronegative, strong partial charges on H
and X
What are dispersion forces (london)?
- Present in all atoms, molecules and ions
- Responsible for liquid and solid phases of
nonpolar molecules - A temporary charge build-up at one side induces
a dipole in its neighbor
What is polarizability?
- How easily a species’ electron cloud distorts
- More polarizable species:
- From stronger dispersion forces
- Have higher molar masses
- Have greater surface areas - In general: the stronger the IMFs, the higher the
boiling point (bp)
What is boiling point?
The temperature where P vapor = P atmosphere
- When a liquid reaches its bp, its particles have
enough energy to break IMFs
- To boil a liquid, energy is needed to break all
IMFs between the particles
- Higher boiling points result from:
- Strong IMFs
- Large molar mass
- Shape
What is ion-dipole forces?
An IMF between an ion and the oppositely charged end of a polar molecule
What is vaporization?
- Vaporization or evaporation - liquid-to-gas phase
change below the liquids boiling point - Endothermic; energy is required to break a
liquid’s IMFs - Heat of vaporization (delta H vap) - energy
required to vaporize 1 mol of liquid at 1 atm
What is dynamic equilibrium?
- In a closed container a liquid’s volume decreases
then remains constant - At equilibrium: evaporation rate = condensation
rate
What is vapor pressure?
The partial pressure of vapor over a liquid or solid at equilibrium
What is a vapor?
The gas phase of a substance that is solid or liquid at 25 C and 1 atm
How does vapor pressure increase with temperature?
- P vap increases exponentially with temperature
- At higher T, more particles have the energy to
enter the gas phase, so P vap increases
What is heat of fusion?
The energy required to melt one mole of a solid
What is heat of sublimation?
The energy required to sublime one mole of a solid
What is the heating curve of water?
- T is measured as a sample is heated
- T is constant during phase changes
- Slope differ; Cs (ice) not equal to Cs (water) not
equal to Cs (steam)
What are phase diagrams?
- Phase diagrams display phases as a function of
temperature and pressure - They are unique to each substance
- Most substances have at least three phases
What is the triple point?
- At the triple point all three phases are in
equilibrium - The triple point of water occurs at 0.01 C and
0.0060373 atm - It can only be observed in sealed containers
under a high vacuum
What is a phase diagram?
- Ask first if the substance is denser as a liquid or
a solid - As a solid; its liquid - solid boundary has a
positive slope
What are supercritical fluids?
- SFs fill their container and diffuse through solids
(like a gas) and dissolve solids (like a liquid) - Beyond the critical point, the meniscus
disappears - Supercritical CO2 is used to remove caffeine
from coffee, nicotine from tobacco, fat from raw
meat, dirt from clothes (dry cleaning), and
pesticides from fruit
What kind of solvent will a solute be dissolved in?
- A solute will dissolve in a solvent if both have
similar IMF types - Solvent molecules will attract the solute particles
at least as well as the solute particles to each
other
What can affect solubility?
- Even molecules with hydrogen bonding become
water-insoluble if the rest of the molecule only
has London forces - The solubility of most salts increases with
temperature - Solubility: the amount of a substance that will
dissolve in a given amount of solvent
Solvent vs solute
- Solvent: the majority component of the solution
- Solute: the minority component of the solution
- Example: mouthwash
- Solvent = water
- Solute = ethanol - Example: absinthe
- Solvent = ethanol
- Solute = water
- Example: mouthwash
Molality vs Molarity
- Both use moles of solute
- Molarity uses solution volume, molality uses
solvent mass - Molarity is temperature-dependent, molality is
not - They are close for dilute aqueous solutions
- molarity (M) = mol solute / L solution
- molality (m) = mol solute / kg solvent
What is colligative properties?
- Solution properties that depend on the number
of solute particles, not their identity - The colligative properties are:
- Vapor pressure lowering
- Boiling-point elevation
- Freezing-point depression
- Osmotic pressure
What is boiling point elevation?
- A nonvolatile solute raises a solvents boiling
point - The solute can be an electrolyte or a
nonelectrolyte
- Delta T b = boiling point elevation
- K b = bp elevation constant
- m solute = solute molality
- T = K * m
What is freezing point depression?
- Any nonvolatile solute raises a solvent’s boiling
point and lowers its freezing point - The solute can be an electrolyte or a
nonelectrolyte - For a given solvent: K f > K b
What is osmosis?
Solvent flow through a membrane from lower to higher concentration
What is osmotic pressure?
The pressure that must be applied to a solution to stop osmosis
What is the van’t Hoff Factor?
- The ratio of moles of particles in solution to
moles of solute - For nonelectrolyte, the van’t Hoff factor is 1
- For an electrolyte, the van’t Hoff factor is the
number of ions formed in solution
What is an ion pair?
- An IMF between a dissolved anion and cation
- An ion-pair acts as a single solute particle
- Ion pairing affects colligative properties by
reducing the number of solute particles - For electrolytes, observed i < expected i
- All electrolytes ion-pair, but less so in dilute solutions
