Chemistry 20 Final Review

1
Chemistry 20 Final Review

• Bonding Unit
• Gases Unit
• Solutions, Acids and Bases Unit
• Stoichiometry Unit

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Bonding Unit Topics

• Lewis Structures
• These are your electron diagrams for individual
  elements (showing valence e-)
• Ex.

S
Mg
3
Intramolecular Forces

• Remember these are the forces WITHIN a molecule
• What forces hold a molecule together?
• ANS BONDS
• What type of bonds are there?
• ANS COVALENT AND IONIC

4
Ionic Bond Structures

• Remember how to draw electron dot diagrams for
  ionic compounds
• They DO NOT SHARE electrons the metal looses
  its outer shell electrons and the non-metal gains
  to a full 8 e-
• Ex

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Covalent Bond Structures

• These are the bonds holding MOLECULAR compounds
  together
• They DO SHARE the electrons
• These were also called Lewis Structures
• The element that goes in the middle is the one
  with the most BONDING e-
• Examples

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eg) PH3
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Remember double and triple bonds

• Each element except hydrogen needs 8 electrons
  around it and there should be NO LONE PAIRS
• This is when double and triple bonds form
• Ex.

O
O
N
N
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Structural Diagrams and Shape Diagrams

• When there are 2 shared electrons between two
  elements in a molecule draw a line to show this
  bond
• Ex.

• Remember the shapes and shape codes
• Ex. tetrahedral, trigonal planar, pyramidal,
  linear, bent

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• the code has two numbers

1. the number of attached to the
central atom
atoms
2. the number of on the
central atom
lone pairs
CH4
eg) NH3(g)
3 - 1
4 - 0
pyramidal
tetrahedral
10
tetrahedral
4 0
CH4
3 0
trigonal planar
CH2O
3 1
pyramidal
NH3
2 1
bent
HNO
2 2
bent
H2O
all other codes are
linear

• Molecules take on these shapes due to the VSEPR
  theory - valence shell electron pair repulsion
• molecules adjust their shapes so that valence e-
  are as far away from each other as possible

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Polar vs. Nonpolar

• Remember electronegativities
• The number in each element box above the element
• It shows how badly an elements wants e-
• The higher the number, the stronger it pulls
• When two elements are bonded together and there
  is a difference in electroneg. then you have a
  polar bond
• Ex. H F (see next slide)

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arrow points towards element with higher
electronegativity (?-)
Bond Dipole Arrows
at the end that is ?
H F
?
?-
13
 

• you can use the difference in electronegativity
  between two atoms to determine the bond

Difference in Electronegativity
3.3
1.7
0.5
0
slightly polar covalent
mostly ionic
polar covalent
non-polar covalent
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Polar vs. Nonpolar Molecules
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• tetrahedral if all atoms
  attached have the same pull (in or out),
  if different atoms attached

nonpolar
polar

• trigonal planar if all
  atoms attached have the same pull (in or out),
  if different atoms attached

nonpolar
polar

• pyramidal as long as there is a
  difference in electronegativity between the atoms

polar

• bent

polar

• linear
  look at electronegativity difference

polar or nonpolar
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Examples
1. H2O
2. HCl
polar
polar
4. C2HI
3. C2H2
np
np
polar
nonpolar
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Intermolecular Forces

• These are the forces that cause attraction
  BETWEEN molecules
• They are weaker then bonding within a molecule
• They are responsible for the bp and mp of
  compounds since when you boil/ melt a molecule
  you are ONLY breaking these forces BETWEEN
  molecules
• The three intermolecular forces we talked about
  the occur between MOLECULAR compounds
• HB, DD, LD

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DD - Dipole - Dipole

• These attractions occur in POLAR molecular
  compouds ONLY
• The slightly positive end of one molecule is
  attracted to the slightly negative end of another
  molecule

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LD London Dispersion Forces

• These attractive force occurs between ALL
  molecular compounds
• It is caused by electrons in atoms and molecules
  constantly being in motion
• So sometimes one side of a molecule can have more
  electron then the other side
• This creates a temporary polar molecule
• An attraction then forms between the ends of
  these polar molecules
• Remember you have stronger LD forces as the
  molecule becomes larger or has more electrons

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HB Hydrogen Bonding

• These attractive forces occur in molecular
  compounds that H bonded to either N, O or F
• Draw the structural diagram of the molecular
  compound to make sure the H is actually bonded to
  the N, O or F

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• the hydrogen has such a low electroneg. in
  comparison to N, O and F so it has its electrons
  pulled so far away from it. This makes it able
  to be attracted not only to the

?? pole
but also to the
lone pairs
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Other melting/ boiling points

• Remember intermolecular forces only occur between
  molecular compounds and are weaker forces then
  intramolecular forces (bonds)
• So when melting molecular compounds only the LD,
  DD, and HB need to be overcome
• Metals and ionic compounds are attracted to one
  another by the bonds holding them together
• Metallic structures and ionic compounds therefore
  have high bp/ mp due to having to overcome their
  intramolecular forces (bonding)

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MP of Metals?

• Metals are solid at room temp. because metal
  atoms have very strong forces between them
• I.e. metallic bonding
• So in order to melt them you need to add LOTS of
  energy (high temp) to overcome these strong forces

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Metallic Bond Model
metal cations
sea of delocalized electrons
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Ionic Compounds

• Ionic compounds are also attracted to one another
  by strong forces (not quite a strong as metallic
  though)
• I.e. ionic crystals
• So in order to melt them you need to add quite a
  bit of energy (high temp) to overcome these forces

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Ionic Crystals

• ionic compounds have

crystal structure
oppositely charged ions

• they form so that
  are as
  as possible

close together

• this
  
  is called a

3-D array of alternating positive and
negative ions
crystal lattice
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Scale of Forces
very high
very low
LD
DD
network covalent
HB
ionic
covalent
Intermolecular Forces (between)
Intramolecular Forces (within)
London Dispersion
metallic wide range
Dipole Dipole
ionic
Hydrogen Bonding
covalent
network covalent eg) diamond, SiC, SiO2
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Order of bps

• Using the scale of forces you can order compounds
  based on their relative bps
• Ex. From Highest to Lowest
• Network covalent compound (ex. SiO2)
• Ionic compound
• Molecular compound with HB, DD, LD
• Molecular compound with DD, LD
• Molecular compound with LD (if 2 molecular
  compounds have LD only then bigger molecule or
  molecule with more electrons has higher bp)

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Gases Unit

• Remember your formulas!!!!!

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When to use what?

• Use Boyles Law when temp. is constant
• P1V1 P2V2
• Use Charles Law when pressure is constant
• V1/T1 V2/T2
• Use Combined Gas Law when all three variables
  change
• P1V1/T1 P2V2/T2
• Use PVnRT if given a mass or number of moles

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Some others

• If you are only given info about pressure and
  tempurature and its in a sealed container then V1
  V2 so using Combined Gas Law cancel out the
  volumes to get left with
• P1V1/T1 P2V2/T2
• P1/T1 P2/V2
• Law of Combining Volumes
• You can use a balanced equation and multiply by
  coefficients wanted/given to get the volume of
  one gas if you know the volume of another

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What volume of oxygen is used up if 100 mL of
steam is formed in a composition reaction?
What you are given
What are ? solving for
O2(g) 2H2(g) ? 2H2O(g)
?
x mL
100 mL
100 mL x 1
2
x mL 50.0 mL
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Convert 650 mmHg to kPa.
Ratio of what you are trying to find
Ratio of known values
101.325 kPa x 760 mmHg 650
mmHg x 86.6 kPa  
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Solutions, Acids and Bases
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Remember your formulas here too

• cn/v
• nm/M
• V1C1VfCf
• pH-logH30
• pOH-logOH-
• H3010-pH
• OH-10-pOH
• pHpOH14

Remember this is the dilution formula
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Experiments

• Remember in experiments there are always three
  variables
• Manipulated what you are changing
• Responding the response to the change
• Controlled what you keep the same

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Ex What effect does eating carrots have on
eyesight?

• Manipulated amount of carrots eaten
• Responding how well you can see
• Controlled Same types of carrots, not eating any
  other food that could effect eyesight

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Electrolytes?

• Compounds that conduct electricity in water
  because they break apart into ions (ex. ionic
  compounds, acids)
• Ex. NaCl ? Na(aq) Cl-(aq)
• Molecular compounds DO NOT break down into ions
  so they are non-electrolytes

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Solubility

• The ability to dissolve
• If the solution is holding as many solutes as
  possible the solution is SATURATED and adding
  anymore solute will NOT be able to dissolve
• A saturated solution usually has a small amount
  of UNDISSOLVED solute at the bottom. This is in
  constant EQUILIBRIUM with the solute that is
  dissolved in the solution (they switch places
  with each other all the time)

Dissolved
Equilibrium
Undissolved
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Standard Solution

• Remember how to prepare a standard solution
• Use formulas nm/M and cn/v to get the mass you
  need for the certain volume and concentration
• Steps
• Weigh out solute
• Dissolve in half amount of water in a beaker
• Pour into final volume volumetric flask
• Fill flask, and invert to mix

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Dilution

• When you have a solution that has too high of a
  concentration you can add water to dilute it
  (water it down so its not as strong)
• Ex. You have 100mL of a 5.0 mol/L solution. You
  add 500mL of water. What is the new
  concentration?
• Ci 5.0 mol/L
• Vi 0.1L
• Vf 0.6L
• Cf ?

ViCiVfCf Solve for Cf
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Dissociation and dissociation equations

• This is the same as dissolving
• When you have a compound and put it in water 4
  situation to know
• It doesnt dissolve
• Ex. C25H52(s) ? C25H52(s)
• It does dissolve and its ionic
• Ex. Ca(OH)2(s) ? Ca2(aq) 2OH-(aq)
• It does dissolve and its molecular
• Ex. C12H22O11(s) ? C12H22O11(aq)
• It does dissolve and its an acid (this is a
  special case because it is a molecular compound
  but it acts as an ionic compound)
• Ex. H2SO4(aq) ? 2H(aq) SO42-(aq)

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Concentration of Ions

• If asked to calculate the concentration of an ion
  in a solution ?1st write the dissociation
  equation then treat it like a solution stoich.
  Question (no volumes are needed since they all
  have the same volume so you dont need to
  calculate n first)
• Ex. Calculate the ion concentrations when you
  have 0.500 mol/L H2SO4(aq) ?
• g w
  w
• H2SO4(aq) ? 2H(aq) SO42-(aq)
• c0.500 mol/L c ? c?
• c of His 0.500 mol/L x 2/1 1.00
  mol/L
• c of SO42- is 0.500 mol/L x 1/1
  0.500 mol/L

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Remember your properties of acids/ bases
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Neutral Substances
Acids
Bases

sour
taste
bitter
taste

electrolytes
electrolytes

electrolytes, non-electrolytes

acids
neutralize
bases
neutralize

indicators
do not
react with

indicators
react with
affect indicators the same way
litmus -
litmus -
red
blue
bromothymol blue -
bromothymol blue -
blue
yellow
phenolphthalein -
phenolphthalein -
pink
colourless

react with to produce
metals
H2(g)

pH
greater than 7
pH
of 7

pH
less than 7
eg)
HCl(aq), H2SO4(aq)
eg)
eg)
Ba(OH)2(aq) NH3(aq)
NaCl(aq), Pb(NO3)2(aq)
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What is an acid?

• The Arrhenius definition of an acid is that it
  has H at the beginning of the compound and is
  (aq)
• Ex. HF(aq)
• The modified Arrhenius definition of an acid is
  it reacts with water to form H3O ions
• Ex. HCl(aq) H2O(l) ? H3O(aq) Cl-(aq)

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What is a base?

• The Arrhenius definition of a base is that it has
  OH- ions at the end of an IONIC compound
• Ex. NaOH(aq)
• The modified Arrhenius definition of a base is
  that it reacts with water to form OH- ions.
• Ex. Na2CO3(aq) HOH(l) ? NaOH(aq) H2CO3(aq)

OH- ions
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Strong acids/ bases

• Weak acids and bases dont 100 break down to
  form H3O ions and OH- ions (strong ones DO)
• Strong acids are listed on the back of your
  periodic table
• If they are NOT on that list they are a weak acid
• Strong bases have OH- ions in them OR a metal
  with oxygen
• Ex. NaOH and MgO
• Every other base is a weak base
• Ex. NH3 and Na2CO3

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Monoprotic vs. Polyprotic

• Monoprotic acids only have 1 H ion to give away
  (to water)
• Ex. HCl, HF
• Similarly monoprotic bases can only accept one H
  ion (from water) or has 1 OH- ion
• Ex. NaOH, ions with only 1- charge (ex. F-)
• Polyprotic acids have more than 1 H to give away
• Ex. H2SO4, H3PO4
• Similarly polyprotic bases can accept more than 1
  H ion
• Ex. compound with ions that have more then 1-
  charge (ex. CO32-, PO43-)

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Remember the pH scale and the pOH scale

• pH scale is 0-14
• 0 strong acid, 14 strong base
• pOH is opposite
• 0 strong base, 14 strong acid

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and the calculations

• Ex. what is pH if H30 0.05 mol/L
• pH-logH30
• - log0.05
• 1.3 (remember SD only 1 SD 1 after decimal
  place on pH and pOH)
• What is pOH if H30 is 0.90 mol/L
• pOH-logOH-
• But we dont know OH-
• pOH also 14 pH so lets find pH

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• pH-logH30
• pH-log0.9
• 0.04575749
• pOH 14 - 0.045749
• pOH 13.95 (2 SD 2 after decimal)
• What is OH- if pOH 6.7
• OH-10-pOH
• OH-10-6.7
• 2 x 10-7 mol/L (1 after decimal 1 SD)

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Stoich!!!
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Now put it all together

• If dealing with MASS its gravimetric stoichuse
  nm/M
• If dealing with GASES its gas stoichuse nm/M or
  PVnRT (or if just volumes use volumes directly
  and x wanted/given)
• If dealing with SOLUTIONS and CONCENTRATIONS its
  solution stoichuse nm/M and cn/v

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LR vs. ER

• Remember that during a stoichiometric reaction
  there are always 2 reactants one is being used
  all up during the reaction (LR) and one will have
  some leftovers (ER)
• You need to figure these out so you know how much
  product will be produced
• You do this by calculating n of each reactant and
  then dividing them by their coefficient the
  SMALLER is the LR
• The LRs n (of moles before you divided by the
  coefficient) is then what you use to calculate
  the product

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Example 1   When 80.0 g copper and 25.0 g of
sulphur react, which reactant is limiting and
what is the maximum amount of copper(I) sulphide
that can be produced?
16 Cu(s) 1 S8(s) ? 8Cu2S(s)
x g M 159.17 g/mol
m 80.0 g M 63.55 g/mol
m 25.0 g M 256.56 g/mol
n 80.0 g 63.55g/mol 1.25 mol
n 25.0 g 256.56g/mol 0.0974 mol
?8/16
n
1.25mol
0.629 mol
n/16 0.0786mol
n/1 0.0974 mol
\ excess
\ limiting
m (0.629mol ) ? (159.17 g/mol) 100.17 g 
100 g
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Yield and Error

• Remember these formulas
• Predicted calculated amount from a stoich.
  calculation
• Actual amount amount you weigh after experiment
• error actual predicted x 100
• predicted
• yield actual x 100
• predicted

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Calculate the error and yield for the
following predicted mass of ppt 100
g actual mass of ppt 93.5 g
error 93.5 g - 100 g x 100 100 g
-6.50
yield 93.5 g x 100 100 g 93.5
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Titration's

• These where just a type of experiment in which
  you use solution stoichiometry and you find the
  volume of one of your reactants to calculate the
  concentration.
• In order to do a solution stoich. question, you
  need to know 3 variables in order to solve for
  the 4th
• Ex. The 4 the c and v of one reactant, and the c
  and v of the other reactant (in order to solve
  the c of one you must know the other 3)
• In the question however you are only given 2
  variables the c and v of only one reactant
  after the titration experiment you will have a
  volume to use giving you all 3

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Ex. Calculate the concentration of HCl when it
is titrated into 10.0 mL of 0.50 mol/L Ca(OH)2?

• w g
• 2HCl(aq) Ca(OH)2(aq) ? CaCl2(aq) 2HOH(l)
• c x c0.50 mol/L
• v find during v10.0 mL
• titration n cv
  (0.50)(0.01) 0.005 mol x 2/1 0.01 mol
• cn/v
• c 0.01 mol/v Lets do the titration and see
  what v is.

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Titration

• When doing this experiment you want to make sure
  you have the correct volume so we do the
  experiment several times to make sure the volumes
  we are getting are all the same (or very close
  within 0.2 mL of each other)
• This is why you have several trials.

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Results
Exclude these results
5.3 5.2 5.3 5.3 mL 3
These 3 are closest together so use them (they
are within 0.2 mL of each other)
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Ex. Calculate the concentration of HCl when it
is titrated into 10.0 mL of 0.50 mol/L NaOH?

• w g
• 2HCl(aq) Ca(OH)2(aq) ? CaCl2(aq) 2HOH(l)
• c x c0.50 mol/L
• v find during v10.0 mL
• titration n cv
  (0.50)(0.01) 0.005 mol x 2/1 0.01 mol
• cn/v
• c 0.01 mol/v Lets do the titration and see
  what v is
• c 0.01/0.0053
• 1.88679
• 1.9 mol/L

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Lastly.

• Remember titration curves and where the
  equivalence point is and what it means (when the
  acid and base fully react to create a neutral
  solution)

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Strong Acid Titrated with Strong Base
14
pH
7
Endpoint
?
0
volume of titrant added (mL)
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Strong Base Titrated with Strong Acid
14
?
pH
7
Endpoint
0
volume of titrant added (mL)