Quiz 1 - Chapter 11 Flashcards
Polar vs non-polar
Polar: one atom is more electronegative than the other
Non-polar: electrons are shared equally, very similar electronegativity
Density (mass/volume) of gas, liquid and solid?
Least to most….
Gas, liquid, solid
Strength of intermolecular forces of gas, liquid and solid?
Least to most….
Gas, liquid, solid
Why is water an exception to the density rule?
H20(l) is more dense than H2O (s) =ice
*ice floats on water
Crystalline vs Amorphous Solids
Crystalline solid: regular ordered structure
Amorphous Solid: not predictable, no order
Ex) glass vs quartz
○ Glass is not regular or repeating arrangement of atoms
○ Quartz is regular and repeating
What is the energy level of the molecules in solids, liquids and gases?
Least to most energy….
Solid, liquid gas
Intermolecular vs Intramolecular Forces
Intra=within a molecule, shorter and stronger bonds
Inter=between two molecules or more, longer and weaker bonds
*The forces that hold condensed states together
Intermolecular force: Dipole-dipole
- between 2 polar molecules (permanent dipoles)
- Happens when positive end of one permanent dipole is attracted to the negative end of another permanent dipole
- water is a special case
Intermolecular force: Ion-induced dipoles
- ion+non-polar
- induced dipole: starts out as non-polar but once ion interacts it then becomes polarized
Intermolecular force: Dipole-induced dipoles
- polar+non-polar
- Polar compound comes closer to nonplar molecule but electrons push it away onto other side of molecule
Intermolecular force: Dispersions
- always present
- weakest IMF
- based on polarizability
- larger cloud=more dispersion force
- instantaneous dipole of an atom induces instantaneous dipoles on neighbouring atoms which then attract one another
Intermolecular force: Ion dipole
- strongest IMF
- ion+polar
Intermolecular force: Hydrogen bonding
- occurs in polar molecule with H atom bonded to a small electronegative atom F, O, N
- super dipole-dipole force
- Not a BOND, but an INTERMOLECULAR FORCE
- Intermolecular forces span much further distances than bonds
Increasing dispersion force does what to electrons, polarizability, dispersion, surface area, boiling point and energy?
- increase number of electrons
- increase polarizability
- increase dispersion
- increase surface area
- increase boiling point
- more energy to break forces
If you have an isomer (same chemical formula but different structural arrangement and same number of electrons), which two structure, linear or branched arrangements, will have a higher boiling point?
Linear=large area for interaction=high boiling point
Branched=smaller area for interaction=low boiling point
Miscibility
- Like dissolves like (Polar dissolves in other polars, non-polars dissolve in non-polars)
- Liquids with the same polarity tend to mix without separating
- F, O, N has a possibility for hydrogen bonding with water
Increasing hydrogen bonding does what to electrons, polarizability, dispersion, and dipole?
- increasing electrons
- increasing polarizability
- increasing dispersion
- decreasing dipole
Which intermolecular force is strongest? Weakest?
Strongest=ion-dipole
Weakest=dispersion
Which element is the most electronegative in the entire periodic table?
Fluorine (F)
Surface tension
- how hard it is to break the surface of a liquid
- increase surface tension=increasing intermolecular forces
- energy required to increase the surface area by a unit amount
- molecules at the surface have more potential energy than those surrounded inside a liquid
An increasing boiling point does what to number of electrons, molecule arrangements/structures and dipoles?
- increasing electrons
- linear arrangement = higher boiling point
- branched arrangement= lower boiling point
- increasing dipole
What happens to intermolecular forces if surface tension increases?
increasing intermolecular forces
What happens to intermolecular forces if viscosity increases?
increasing intermolecular forces
Is a longer chain or shorter chain tend to be more viscous?
longer chain tends to tangle more, so more viscous
Viscosity
- resistance of a liquid to flow
- stronger IMF increase viscosity
- molecules that are longer can tangle more and tend to be more viscous
Capillary action
- ability of a liquid to flow against gravity up a narrow tube
- a balance of cohesive and adhesive forces
Cohesive vs adhesive forces
Cohesive: attraction between liquid molecules; forms a dome
ex) Hg metal with glass
Adhesive:attraction between liquid molecules and the surface of a tube (allows for upward movement); meniscus is formed
ex) water with glass
Vaporization
liquid to gas (endothermic: low to high energy)
Condensation
gas to liquid (exothermic: high to low energy)
What does an increase in temperature mean for heat and kinetic energy?
heat is transferred to molecules and kinetic energy increases
Volatile vs nonvolatile liquids
Volatile: easily vaporizes, have weak intermolecular forces
Nonvolatile: don’t easily vaporize, have stronger intermolecular forces (H2O)
If a liquid is volatile, what happens to its intermolecular forces?
high volatility=low intermolecular forces
Heat of vaporization (delta vap H)
amount of heat required to vaporize one mole of a liquid to gas
Vapour pressure
- the pressure of a gas in dynamic equilibrium with its liquid
- pressure at which the amount of molecules leaving and entering the container is equal
- increasing vapour pressure lowers IMF
If the vapour pressure increases, what happens to the boiling point and temperature?
increasing boiling point and increasing temperature
If volume of a cylinder is increased, what happens to the pressure and gas molecules?
pressure falls, more gas vaporizes and pressure is restored to equilibrium
If volume of a cylinder is decreased, what happens to the pressure and gas molecules?
pressure rises, more gas condenses and pressure is restored to equilibrium
Boiling point
the temperature where the vapour pressure equals the external pressure
Normal boiling point vs standard boiling point
normal boiling point: temperature where the vapour pressure is 1 atm
standard boiling point: temperature where the vapour pressure is 1 bar
Sublimation
the transition of a substance from solid to gas (endothermic, low to high energy)
Deposition
transition of a substance from gas to solid (exothermic, high to low energy)
Melting or fusion
transition of a substance from a solid to a liquid (endothermic, low to high energy)
Heat of fusion (delta fus H)
amount of heat required to melt 1 mol of solid
Freezing
transition of a substance from a liquid to a solid (exothermic, high to low energy)
Heat capacity
How much heat can be absorbed by the substance
What is the 1st Law of Thermodynamics?
Energy can be neither created nor destroyed, only transformed.
Phase diagrams
- map the state/phase of a substance as a function of pressure and temperature
- have major regions that are labelled with respective state
- lines and curves represent a set of temperatures and pressures where the substances are in equilibrium
Triple point
- represents the temperature and pressure where all 3 states are in equilibrium
- in the center of phase diagram where all three equilibriums of each state meets
Where are the fusion curve, vaporization curve and sublimation curve in the phase diagram?
Fusion curve: equilibrium between solid and liquid
Vaporization curve: equilibrium between liquid and gas
Sublimation curve: equilibrium between solid and gas
What must happen to pressure and temperature in order for a solid to form?
high pressure, low temperature
What must happen to pressure and temperature in order for a gas to form?
low pressure, high temperature
Critical point or supercritical fluid
- have a liquid and a gas at the same time at the highest possible (critical) temperature and (critical) pressure
- located at the top end of the curve in the phase diagram
- a superficial fluid has properties of both a liquid and a gas
- makes it an amazing solvent
Polarity
Difference in electronegativity in a molecule and permanent dipole
What is the distance between two atoms in a crystal? Why can X-ray radiation be used to study crystals?
10^2 pm
X-ray radiation has wavelengths in this range
When x-ray interacts with crystals what are the results?
Crystals are ordered so interference and diffraction patterns result. Diffraction pattern can then be used to relate distances between atoms of crystal.
Light parts= constructed interference
Dark parts=deconstructive interference
Crystalline lattice
Regular arrangement of atoms/ions in a crystalline solid
Unit cell
Small collection of atoms/ions/molecules that can be repeated over and over again to reproduce a 3D structure
Coordination number
Number of atoms touching the central atom
Simple cubic lattice
Coordination number: 6
Atoms per unit cell: (1/8 corners of each atom* 8 corners total=) 1
Body-centred cubic
Coordination number: 8
Atoms per unit cell: 2
Face-centred cubic
Coordination number: 12
Atoms per unit cell: 4
List the cubic cells from least so most efficient in packing.
Simple cubic = 52%
Body-centred = 68%
Face-centred = 74%
Simple cubic edge length
2r
Body-centred cubic edge length
4r/root3
Face-centred cubic edge length
(2root2)r
What are the 2 closest-packed systems?
1) hexagonal closest packing
2) cubic closest packing
Hexagonal closest packing (HCP): alternating layers, packing efficiency and coordination number.
- 2 alternating layers slightly to the side (in between empty spaces)
- ABABAB
- 74% packing efficiency
- coordination number=12
- not simple cubic
Cubic closest packing (CCP): alternating layers and coordination number
- 3 alternating layers slightly to the side (in between empty spaces, top and bottom layers are different orientations)
- ABCABCABC
- coordination number=12
- same as face-centred cubic!
Name the types of crystalline solids.
1) molecular solids
2) ionic solids
3) atomic solids
Crystalline solid: molecular solids
- Composite units are molecules
- Low melting point
- ex) ice
Crystalline solids: ionic solids
- Composite units are formula units (cations and anions)
- High melting points
- Ex) Table salt
Crystalline solids: Atomic solids
- composite units are atoms
- types: nonbonding, metallic and network covalent
Crystalline solids: Atomic solids: Nonbonding
- held together by dispersion force
- low melting points
- Ex) solid xenon
Crystalline solids: Atomic solids: Metallic
- Held together by metallic bonds
- Variable melting points
- Ex) gold
Crystalline solids: Atomic solids: Network covalent
- held together by covalent bonds
- high melting points
- ex) quartz
