Chemistry 2 Exam 1 - Rabson

Exam 1 CH 10/11

Solid

Solid

-particles are very close together

-in a regular pattern=crystalline

-no pattern=amorphous

-definite shape and volume

-fixed shape=does not conform to container=particles don’t move around

Liquid

-more dense than solid

-particles close together

-definite volume

-no definite shape (conforms to a container)

-fills the container only to the volume of liquid present

Gas

-particles are separated over a broad area

-no definite shape or volume (They expand to fill any space)

-conforms to container and fills the entire container

-can undergo changes in pressure=compressible

*smaller container=greater pressure

*larger container=less pressure

Solid-energy

-mostly energies if attraction

-more attraction = less kinetic energy

-stringer attractive forces between particles

Gas-energy

-mostly kinetic energy

-less attraction = more kinetic energy

-energy of motion = particles move around and do not stick together

-lower boiling point

Liquid-energy

-kinetic energy and energy of attraction = particles move around but remain touching

Types of Intermolecular Forces

1. Intra molecular = bonding forces within a molecule or compound

2. Inter molecular = non-bonding forces (weak forces between molecules)

Intramolecular forces

-Intra = strong bonds within

-bonding forces within a molecule to compound

-relatively string between molecules

-larger charges/closer together

-ionic bonds = +/- ions holding each other together

-covalent bonds = shared bonding

Intermolecular forces

-non-bonding forces = can be disrupted without chemical change

- weak forces between molecules

-smaller charges = further apart

-examples =

1. hydrogen bonds (strongest) “of the weak”

2. dipole-dipole

3. London dispersion forces =

van der waals =refers to both dipole and dispersion forces

Phase changes (6 total)

Solid to liquid = melting or fusion

Liquid to solid = freezing

Solid to gas = sublimation

Gas to liquid = condensing

Liquid to gas = vaporising

Gas to solid = deposition

Solid to liquid?

Melting or fusion

Liquid to solid?

Freezing

Solid to gas?

Sublimation

Gas to liquid?

Condensing

Liquid to gas?

Vaporizing or evaporation or boiling

Gas to solid?

Deposition

Exothermic?

Heat/ energy exiting

Endothermic?

Heat/energy entering

*Boiling points increase?

Going down the periodic table

Dispersion (london) forces

-everything has it

-temporary/momentary fluctuations in electron charge = produce instant dipoles/one side negative for a split second

-only attractive forces in non-polar compounds

-present in even non-bonding forces

-increase with number of electrons

-larger molecules have more a attraction to each other =higher boiling and melting points

-electrons “stack up” to one side = it takes more energy to take them out of their phase

Linear (small) molecule

-fewer dispersion forces

Linear (large) molecule

-more dispersion forces

Spherical molecule

-fewer dispersion forces than a linear molecule of the same size

Polar molecules and Dipole-Dipole forces

-exist in all polar molecules

-electron density is not even across molecule = creates positive/negative ends

-attract each other and stay together

-causes polar compounds to have higher melting/boiling points than non-polar compounds

*also, to not mix with non-polar compounds

-similar forces = ion-dipoles and dispersion-dipole (mix of both)

*ion-dipole is the strongest Intermolecular force

-polarizability = ability of a molecule to become polarized

Hydrogen Bond

-especially strong

*dipole-dipole = +/- side intermolecular force

-a hydrogen bond may occur when an H atom in a molecule, bound to a small, highly electronegative atom with lone pairs of electrons, is attracted to the lone pairs in another molecule

*the elements N2, O2, and F2 are strong but small and are extremely electronegative

-One element is the hydrogen bond acceptor and one element is the hydrogen bond donor

Formal dipoles?

=permanent dipole

The molecular basis of surface tension

-surface tension does not work for non-polar, and does not respond to charge

-hydrogen bonding occurs across the surface and below the surface *like a “skin layer” on top of the H2O

-Hydrogen bonding/attraction occurs in three dimensions

-The net vector for attractive forces is downward

-surface tension = the energy required to increase the surface area of a liquid by amount —> very high for water

-surface tension works for polar liquids = water

-surface tension does not work for non-polar liquids = hexane

Viscosity?

= the resistance of a liquid to flow

Viscous liquids?

1. motor oil

2. syrup

*number on motor oil increase with viscosity

Viscosity Info

-unit = poise (P) or kg/m:s

-increases with:

1. Intermolecular forces

2)longer molecules that become engaged (e.g. certain petroleum fractions)

-decreases as temperatures increase *think of hot vs. Cold pancake syrup and the difference in flow of each

Capillary action?

= ability of liquid to flow up a narrow tube

-seen in:

1. taking blood samples

2. liquid going up a straw on its own

3. the meniscus of a graduated cylinder or test tube

Capillary action is due to?

-due to:

1. Cohesive forces, the intermolecular attractive forces

2. Adhesive forces, between the molecules and the tube

*cohesive forces greater than adhesive: convex (rainbow) meniscus effect (as in mercury)

*adhesive forces greater than cohesive: concave (bowl) meniscus effect (as in water)

Vapor pressure as a function of temp and intermolecular forces

-boiling point (external pressure = vapor pressure)

Vapor pressure

-higher the temperature = higher the vapor pressure

-molecules move faster

-more molecules escape from liquid phase than return

-vapor pressure increases

-weaker intermolecular forces result in high vapor pressure at any given temperature

-vapor pressure = external pressure —>boiling=bubbles of gaseous molecules form in the liquid

-at BP: heat energy is used to overcome attractive, intermolecular forces to go from liquid —>gas phase

Vaporization

-all molecules move with thermal energy

-some break free from surface of liquid to enter Vapor phase

*happens more in volatile liquids (tends to be less polar; you can often smell the Vapor)

-increases with:

1. increasing temperature

2. increasing surface area

3. decreasing intermolecular forces

Energetics of vaporization

-vaporization = endothermic (heat enters/absorbed)

*water evaporating cools the skin

*means condensation is exothermic (heat exits/released)

-heart of vaporization is (triangle H vap) *ALWAYS POSITIVE

*condensation energy = (negative triangle H vap)

Vapor pressure

-dynamic equilibrium between liquid and gas phase molecules

-vapor pressure = pressure of this gas

* for small amount, has more effect in a closed container

-Obeys, dynamic equilibrium rules

*if you increase the volume, the Vapor pressure decreases

(Boules gas law)

*when pressure decreases, more liquid vaporises (response to equilibrium aka La Chatelier’s principle)

-reverse situation —> more gas condenses

-at boiling point, vapor pressure = extreme pressure

*”normal boiling point” defined where Vapor pressure = 1 atm

*higher altitudes = lower external pressure —> lower BP

Sublimation?

=the phase changes from solid to gas

-CO2, dry ice

-liquid CO2 only at high pressure (>5 atm)

=freeze-dried foods

-may also sublime out of frozen foods

Sublimation reversed?

=Deposition

Opposite of frozen?

=melting (fusion)

-phase change from solid to liquid

Opposite of fusion (melting)?

=Freezing

Heat of Fusion?

= is ALWAYS POSITIVE because fusion is endothermic (heat enters or you GAIN heat)

Phase diagrams: Regions?

-equilibrium between different phases of a compound

–temperature on x axis, pressure on y axis

-solid-liquid, liquid-gas, solid-gas curves

-increase in pressure —> transition to lowest volume

Phase diagrams: Critical point?

-density of liquid and Vapor are equal at Tc and pc

-difference between liquid and Vapor disappears

-above this point, a “supercritical fluid” exists (mixed gas and liquid properties)

Phase diagrams: Triple point?

-3 phase transition point

-equilibrium between the 3 phases

-at its triple point, CO2 sublimes and deposits, melts and freezes, Vapor sizes and condenses at the same time

*dry ice (solid CO2) does not melt at room temp and pressure

Phase diagram left?

=solid

Phase diagram middle?

=liquid

Phase diagram right?

=gas

Types of Solids?

1. ionic

2. non-bonding

3. metallic

4. network covalent

5. molecular solids

Ionic solid?

=held together by opposite charges (+/-)

=ionic compounds

=high melting and boiling points

=not conductive as solids

=many are H2O soluble —> conduct in solution

Non-bonding solids?

=only dispersion forces

=this only applies to noble gases at extreme low temperatures

Metallic solids?

=metal (element)

=alloy (mix of metals)

=high melting/boiling point

=conductive as solids

=not H2O soluble

=metallic bonds (“sea” of delocalised electrons)

=close-packed structures with high melting points

Network covalent solids?

= long extended molecule (indefinite)

=not soluble

=high melting/boiling point

=might conduct

=held by covalent bonds (carbon)

*carbon:

1. diamond = whole network is covalent

2. graphite = sheets of covalent bonds, only weak dispersion forces between sheets

Molecular solids?

=covalent compound/molecules (mostly compounds)

=fairly low melting/boiling point

=definite formulas (size/weight limit to it)

=covalent molecules held together by intermolecular forces (ice, dry ice)

Allotropes of carbon?

1. diamond

2. graphite

3. grapheme

4. fullerene, C60

Carbon: diamond?

=poor conductor

=not a free moving electron

=hard/rigid substance

Carbon: graphite?

=good conductivity

=ONE free delocalised electron

=soft, layers slide across each other

=pencil

=lube

=electrodes

Carbon: graphene?

=very good conductor

= ONE free delocalised electron moving across the layers easily

=lightest/strongest material replacing silicon in photovoltaic cell

=drug delivery

=electronic/transparent

Carbon: fullerene?

=buckyball

=semiconductor

=surface is sphere, not planar

=electrons CANNOT flow easily (low viscosity)

=lower electron mobility

Solid structure: Crystal Lattice?

-crystal lattice = regular arrangement of atoms or ions in a solid

-unit cell = smallest repeating unit

-types of unit cells:

1. Simple cubic = 52% packing efficiency, 1 atom in unit cell (“primitive”)

2. Body centered cubic (BBC) = 68% packing efficiency, 2 atoms in unit cell

3. Face centered cubic (FFC) = 74% packing efficiency, 4 atoms on unit cell

Simple cubic?

= one atom

-8 corners x 1/8 = 1

-coordination number = 6

= 52% packing

Body centered cubic?

= two atoms

-8 corners x 1/8 (1 in the middle) = 2

-coordination number = 8

= 68% packing

Hexagonal?

= ABAB

-8 corners x 1/6 = 12

-coordination number = 12

= 74% packing

= close-packed

Face-centered?

= 4 atoms

-6 faces x 1/2 = 4

=cubic close packed

= ABCABC

-coordination number = 12

= 74% packing

Close-Packing

-hexagonal close packing

=layers in ABAB arrangement

=gives hexagonal unit cell

-cubic close packing

=layers in ABCABC arrangement

=give face centered cubic (FCC) unit cell

Metallic solids: properties of metals?

=shiny

=malleable

=ductile

=conduct heat and electricity

=often have BCC, HCP, or FCC structure; close packing allows heat/electrons to move between metal atoms

=high ability to conduct heat —> “cold” when you touch it at room temp; hot if it has been in the sun

=bendable

=“malleable/ductile” properties occur because atoms can slide past each other

*it does not have high heat capacity but can pass it to you fast (seat belt buckle example)

Alloy?

=mixture of metals

= “solvent” largest proportion of mixture

= “solute” metal present in smaller proportion

Types of Alloys?

1. substitutional: solute atoms substitute in the lattice of the solvent atom

*works best for atoms with similar sizes/properties

2. interstitial: solute atoms fit in holes between solvent atoms

*need to be small

*solute is often a non-metal (bonding has covalent character) - carbon steel

3. heterogeneous: not mixed uniformly (you can see clusters of each thing under the microscope)

*two phases

4. intermetallic: really a true compound instead of mixture

*composition cannot vary; atoms in definite order

*can write a chemical formula

Metals vs Ionic solids

1. molecular: compounds are molecules (ice) = relatively low melting points

Vs

2. ionic: can pack into lattices

-most common types:

1. cesium chloride structure

*Cl simple cubic (marbles) ; Cs in the holes (beads)

*the two ionic radii are similar

2. sodium chloride structure

*Cl FCC lattice ; Na in holes

3. zinc blende (ZnS) structure

*S2 FCC lattice ; Zn2+ in some small holes

*greatest difference between ion size

Solutions: general properties of solutions?

-most chemical reactions take place in solution

Solution?

=homogenous or uniform mixture of two or more substances

Solute vs Solvent

Solute= the substance in the mixture with less quantity

Solvent= the substance present with more quantity

Solutions can be?

=liquids, solids, and gases

Example of liquid solution?

Sugar water = sugar dissolved in water

*H20 = solvent (more amount)

*Sugar = solute (less amount)

Example of solid solution?

Metal alloys = brass

*brass = homogeneous mixture of copper and zinc

Example of gas solution?

Gas solutions = air

*air = series of gases (including O2, H2, CO2) dissolved in nitrogen gas, N2

-nitrogen=solvent (air solution)/(more amount)

-other gases=solutes (less amount)

General properties of liquid solutions

1. aqueous

2. true solution

Aqueous solution?

=clear/coloured but transparent (think stain glass)

-solution where the solvent (most amount) is water

-clear/transparent without visible particulates of solute

-colourless or coloured, depending in on solvent and solute

True solution?

=no solids

-homogeneous with uniform properties

-solute cannot be isolated by filtration (solids that are not water soluble leave stuff behind)

-no “settling-out” of particles

-particles are individual molecules in continuous random motions (separate)

Volumes of solvent and solute are not additive:

-1 litre H20 + 1 litre is <greater than 2 litres

-dependent on how molecules “fit” together

-attractive and repulsive forces (hydrogen bonds) (marble and beads in a beaker example)

*liquids are ALWAYS in motion

Electrolytes in aqueous solution?

=electrolyte —> conducts electricity *needs dissolves ions

Strong electrolytes?

=strong acid list

-solvable ionic compounds =alkali earth metals

-string acids and bases (break apart into ions)

*dissolve vs disassociate

Weak electrolytes?

=weak acids and bases (only a few of the molecules break apart into ions)

Non-electrolytes?

=insoluble ionic compounds

=molecular compounds (covalent)

-does not make ions

-no conductivity

Solubility?

=often relative (only so much is there to dissolve)

-how much solute can dissolve in a solvent (usually mol/L)

(L = molarity)

-not an “either or” situation; it is relative

-“miscible” = able to completely dissolve/mix

*imiscible = not able to completely dissolve/mix

Factors that affect solubility?

=polarity of solute and solvent

-the more spent and solute are different, the less soluble they are

-“like dissolves like” or polar dissolves polar

=temperature

-increase in temp usually increase’s solubility/dissolvable

*example = sugar dissolves better in HOT water vs cold which increases solubility

Factors that affect solubility?

=pressure

-high effect on gas, little effect on liquid or solids

-for gas solutes (less amount) increased pressure of gas increases solubility

*gases = solvable when COLD

-solubility of gas in liquid is directly proportional to pressure

-carbonates sodas contain CO2 = bubbles of gas

Factors that affect solubility?

=entropy: disorder (it is favourable to mix things/stir things)

=intermolecular forces: solute/solvent must “like” each other more than themselves

-solvent/solute interactions must be equal or greater than solvent/solvent interactions vice versa

Solubility and Saturation

Saturated solution = dissolved as much as you can get

-maximum amount of solute that can be dissolved at a PARTICULAR TEMPERATURE

Cooling solution = 1) decreases in solubility 2) precipitate of excess solute

Unsaturated solution?

=could still dissolve more solute (more to “mix”)

Supersaturated solution? (NOT VERY COMMON)

=unstable situation

=contains more solute than can be dissolved at that temperature

=excess solute will precipitate —> saturated solution + precipitate

(Solution with the maximum amount of solute that can dissolve at that temperature + precipitate)

*solutions involve DYNAMIC EQUILIBRIUM BETWEEN DISSOLVED AND UNDISSOLVED PARTICLES

Solubility and equilibrium?

=a saturated solution is an example of a dynamic equilibrium

-dissolution (dissolving of the solute) and precipitation (undissolving of the solute) proceed at the same rate

*as many molecules dissolve as they also precipitate

-saturated solution is a solution that is in equilibrium with the undissolved solute

-crystallisation (lower temperature, supersaturation) a “seed crystal” is introduced; more solute goes into the undissolved state, and does not return to the solution

Supersaturated solution?

= and example of supersaturated solution is sodium, acetate crystallising from a supersaturated solution

-Undissolved and dissolved solute is in equilibrium: slowly, cool, saturated solution, forms, supersaturated, solution, small disturbance (seed, Crystal, scratching the flask) results in crystallisation

Concentration of solutions?

= amount of solute dissolved in a given amount of solution

Math = concentration equals the amount of solute divided by the amount of solution

Molarity?

= molarity equals moles of solute divided by litres of solution

Concentration of a solution?

=all concentration units are fractions

-Numerator indicates the quantity of the solute

-Denominator indicates the quantity of either the solution or the solvent

-They differ in the units used to express these two quantities

Molarity?

Molarity = moles of Solute divided by volume of solution (mol/L)

* ratio of quantity of solute, to quantity of solution

Mole fraction?

mole fraction= Moles of solute divided by

(moles of solute + moles of solvent)

Concentration of a solution: two

1. Molarity = moles of solute divided by mass of solvent in kilograms (mol/Kg)

* ratio of quantity of solute to quantity of solvent

2. mass percent = (grams of solute divided by grams of solution) times 100

*end Unit is Percent