Chemistry 2 Exam 1 - Rabson

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Exam 1 CH 10/11

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-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

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-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

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-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

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-mostly energies if attraction

-more attraction = less kinetic energy

-stringer attractive forces between particles

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-mostly kinetic energy

-less attraction = more kinetic energy

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

-lower boiling point

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

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Types of Intermolecular Forces

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

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

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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

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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

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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

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Solid to liquid?

Melting or fusion

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Liquid to solid?


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Solid to gas?


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Gas to liquid?


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Liquid to gas?

Vaporizing or evaporation or boiling

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Gas to solid?


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Heat/ energy exiting

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Heat/energy entering

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*Boiling points increase?

Going down the periodic table

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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

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Linear (small) molecule

-fewer dispersion forces

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Linear (large) molecule

-more dispersion forces

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Spherical molecule

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

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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

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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

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Formal dipoles?

=permanent dipole

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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

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= the resistance of a liquid to flow

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Viscous liquids?

  1. motor oil

  2. syrup

*number on motor oil increase with viscosity

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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

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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

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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)

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Vapor pressure as a function of temp and intermolecular forces

-boiling point (external pressure = vapor pressure)

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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

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-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

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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)

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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

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=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

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Sublimation reversed?


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Opposite of frozen?

=melting (fusion)

-phase change from solid to liquid

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Opposite of fusion (melting)?


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Heat of Fusion?

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

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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

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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)

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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

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Phase diagram left?


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Phase diagram middle?


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Phase diagram right?


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Types of Solids?

  1. ionic

  2. non-bonding

  3. metallic

  4. network covalent

  5. molecular solids

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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

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Non-bonding solids?

=only dispersion forces

=this only applies to noble gases at extreme low temperatures

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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

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Network covalent solids?

= long extended molecule (indefinite)

=not soluble

=high melting/boiling point

=might conduct

=held by covalent bonds (carbon)


  1. diamond = whole network is covalent

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

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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)

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Allotropes of carbon?

  1. diamond

  2. graphite

  3. grapheme

  4. fullerene, C60

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Carbon: diamond?

=poor conductor

=not a free moving electron

=hard/rigid substance

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Carbon: graphite?

=good conductivity

=ONE free delocalised electron

=soft, layers slide across each other




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Carbon: graphene?

=very good conductor

= ONE free delocalised electron moving across the layers easily

=lightest/strongest material replacing silicon in photovoltaic cell

=drug delivery


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Carbon: fullerene?



=surface is sphere, not planar

=electrons CANNOT flow easily (low viscosity)

=lower electron mobility

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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

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Simple cubic?

= one atom

-8 corners x 1/8 = 1

-coordination number = 6

= 52% packing

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Body centered cubic?

= two atoms

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

-coordination number = 8

= 68% packing

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-8 corners x 1/6 = 12

-coordination number = 12

= 74% packing

= close-packed

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= 4 atoms

-6 faces x 1/2 = 4

=cubic close packed


-coordination number = 12

= 74% packing

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-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

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Metallic solids: properties of metals?




=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


=“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)

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=mixture of metals

= “solvent” largest proportion of mixture

= “solute” metal present in smaller proportion

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Types of Alloys?

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

*works best for atoms with similar sizes/properties

  1. 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

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

*two phases

  1. intermetallic: really a true compound instead of mixture

*composition cannot vary; atoms in definite order

*can write a chemical formula

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Metals vs Ionic solids

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


  1. 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

  1. sodium chloride structure

*Cl FCC lattice ; Na in holes

  1. zinc blende (ZnS) structure

*S2 FCC lattice ; Zn2+ in some small holes

*greatest difference between ion size

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Solutions: general properties of solutions?

-most chemical reactions take place in solution

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=homogenous or uniform mixture of two or more substances

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Solute vs Solvent

Solute= the substance in the mixture with less quantity

Solvent= the substance present with more quantity

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Solutions can be?

=liquids, solids, and gases

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Example of liquid solution?

Sugar water = sugar dissolved in water

*H20 = solvent (more amount)

*Sugar = solute (less amount)

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Example of solid solution?

Metal alloys = brass

*brass = homogeneous mixture of copper and zinc

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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)

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General properties of liquid solutions

  1. aqueous

  2. true solution

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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

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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)

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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

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Electrolytes in aqueous solution?

=electrolyte —> conducts electricity *needs dissolves ions

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Strong electrolytes?

=strong acid list

-solvable ionic compounds =alkali earth metals

-string acids and bases (break apart into ions)

*dissolve vs disassociate

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Weak electrolytes?

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

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=insoluble ionic compounds

=molecular compounds (covalent)

-does not make ions

-no conductivity

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=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

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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


-increase in temp usually increase’s solubility/dissolvable

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

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Factors that affect solubility?


-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

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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

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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

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Unsaturated solution?

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

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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)


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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

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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

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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

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= molarity equals moles of solute divided by litres of solution

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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

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

* ratio of quantity of solute, to quantity of solution

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Mole fraction?

mole fraction= Moles of solute divided by

(moles of solute + moles of solvent)

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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

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

*end Unit is Percent

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