Transcript AP CH 1 2 3 ppt Brown LeMay

A.P. Chemistry
Unit 1
Chapters 1*, 2, 3
Why take Chemistry?
Purpose of AP Chem at CG?
Introductory terms
• Table groups
Definition Review: Try to fill in
1. Chemistry
2. Chemical
Property
3. Physical
Property
4. Intensive
property
5. Extensive
property
Terms to recognize
6. Element
7. ATOM
8. Compound
9. Molecule
10. Formula Unit
TERMS
11. Pure substances
12. Chemical reaction
13. Physical change
14. Mixture
15. Homogeneous
16. Heterogeneous
17. Chemical symbol
18. Chemical formula
More terms
19. Chemical Equation
20. Reactants
21. Products
22. Coefficients
23. Subscripts
24. Matter
25. Law of conservation of
energy
26. Law of conservation of
matter
27. Name Steps in the
“Scientific Method”
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
Asking a question
Forming hypotheses
Researching previously found information
Designing experiments
Conducting experiments / collecting data
Determining variables: dependent and independent
Organizing and analyzing data
Stating conclusions
Considering sources of error
Communicating results
Planning future experiments
Classification of Matter
ALL Matter
Mixture
Pure Substance
Easily Separated
Homogeneous
Solutions
Heterogeneous
Always Homogeneous
Element
Compound
Atoms
Periodic Table
Molecules or Crystals
Formulaic
Atoms Molecules Elements
Compounds?
1.
2.
3.
4.
5.
Aluminum foil
Carbon Dioxide
Zinc
Graphite
Helium
6. oxygen gas
7. sodium chloride
8. water
9. chlorophyll
10. nitrogen
1.
2.
3.
4.
5.
E, A
C, M
E, A
E, A
E, A
6. E, M
7. C, “?M”
8. C, M
9. C, M
10. E, M
Solid, Liquid, Gas
(room temperature)
1. Ammonia
2. Gasoline
3. Graphite
4. H2O
5. Shaving cream
6. Aluminum
7. Ice cream
8. Helium
9. Bromine
10.Sugar
Solid, Liquid, Gas
(room temperature)
1. G
2. L
3. S
4. L
5. L & G …?
6. S
7. L
8. G
9. L
10. S
Pure Substance or Mixture?
1.
2.
3.
4.
5.
6.
7.
8.
Water
Hydrogen
Salt
Tea
Sodium
Sugar
Iron oxide
concrete
9. Raisin cookie
10. Gatorade
Pure Substance or Mixture?
1.
2.
3.
4.
5.
6.
7.
8.
P
P
P
M
P
P
P
M
9. M
10. M
Chemical or Physical Change?
1. Dog is groomed
2. Child gets taller
3. Gas forms when
Baking soda is mixed
with vinegar
4. Pencil is sharpened
5. Paper burns
6. Leaves turn color
7. Ice melts
8. Sugar dissolves in water
9. Cookie bakes in oven
10. Cake mix is combined
with water.
Chemical or Physical Change?
1.
2.
3.
4.
5.
6.
7.
P
C
C
P
C
C
P
8. P
9. C
10. P (?)
Metric fundamental units
• Kilogram
• Meter
• Second
• mole
• Kelvin
• Coulomb
Metric units for?....
1.
2.
3.
4.
5.
6.
Length
mass
area
volume
Density
Weight and Force
7. Energy
8. particles
9. Pressure
10. Current
11. Potential
12. Power
fundamentals?....
1.
2.
3.
4.
5.
6.
7.
yes
yes
yes
yes
yes
Kg*m/s2
Kg*m2/s2
8. yes
9. Kg*m/s2
10. C/s
11. Kg*m2/Cs2
12. Kg*m2/s3
13. yes
Metric units?
• Fundamental?
• Can you convert?
•
•
•
•
•
•
•
•
Deci
Centi
Milli
Kilo
Nano
Giga
Micro
Mega
AP Chemistry
Chapter 2
Atoms,
Molecules
and Ions
John Dalton
• English school
teacher
• 1766-1844
• Author of the
Modern
Atomic Theory
• loved studying the weather
• saw the applications for chemistry
in his ideas about the atmosphere.
• Proposed Atomic Theory: 1803
• Dalton's theory was presented in
New System of Chemical
Philosophy (1808-1827).
John Dalton
• Was colorblind
• Daltonism
• His eyes were used to
prove it is a brain
disorder
Postulates of Dalton’s atomic theory
Matter is made of atoms which stay the same
during a chemical change.
An element is a substance made of one type
of atom, each of which has the same
properties.
A compound is matter made of
two or more elements combined
in fixed proportions.
A chemical reaction involves
rearrangement of atoms into new
substances, but no loss or gain of any
atoms.
Law of definite proportions
• Molecules of
the same
compound are
all the same.
The same elements can make
many different compounds.
Atomic Structure
Modern Theory Says…
• Atoms are made
of:
• Protons
• neutrons
• electrons
• a small dense
nucleus of protons
and neutrons and
surrounding
electrons.
Famous experiments leading to this
view of the atom:
• Joseph John
Thomson:
• The Cathode Ray
Tube Experiment
(Cambridge; 1897);
• discoverer of the
electron, Nobel
Prize Winner.
The Cathode Ray Tube and
Thomson’s Plum Pudding atom
Robert Millikan
• The Oil Drop
Experiment
(USA, 1909)
• measured
charge on an
electron
• calculated the
mass of
electrons.
Millikan’s Experiment
• Used calculation to
determine the
charge on each
suspended droplet
• All were multiples
of 1.6x10-19
Coulomb
Ernest Rutherford
• Studied Gold Foil
Experiments by
Geiger and Marsden
(1911)
• atoms are mainly
empty space with a
small, massive,
dense, positively
charged nucleus.
• Also discovered the
proton.
Gold Foil Experiment Video?
• Led to the idea
of a nuclear
atom
• Led to the idea
that atoms are
mainly made of
empty space.
Modern Atomic theory
• the atom consists
of 40 fundamental
particles.
• The electron is a
quark, but the
proton and the
neutron are not.
ISOTOPES
• Atoms of the
same element
that have
different masses
(different # of
neutrons in the
nucleus).
ISOTOPES TO KNOW
• The three isotopes
of hydrogen:
• Protium = H-1
• Deuterium = H-2
Water made with
this is called heavy
water.
• Tritium = H-3 and
is radioactive!
AZ symbols
• A: mass #, nuclear
particles, P+N
• Z: atomic #, nuclear
charge, P
• Neutrons =?
• X = element symbol
Write an “AZ” nuclide symbol for
• strontium 90
• silicon 30
• radon 226
Atomic number vs mass number
Protons : nuclear charge
vs
protons + neutrons : number of nuclear particles
• Atomic weight vs atomic number
~Weight of protons, neutrons, electrons
Vs
Proton number: nuclear charge
Atomic particles
Atomic Weights
• All relative to the Carbon-12 isotope
• Carbon-12 is the mass standard
• One mole of carbon 12 = 12 grams
Atomic Weights
• Represent: average
mass of isotopes
and their percent
composition in
nature.
Measuring atomic weight
• Units are “amu”
• Atomic mass units
• One mole of amu =
1 gram
Measuring the weight of atoms
By Mass spectrometry.
• An unknown is
compared to a known
sample ( the standard).
• Particles are
accelerated through a
gas and bent by a
magnetic field.
• The curvature of their
pathway is measured
and mass is
calculated.
• F=Bvq.
Mass Spec and examples
• See worksheet
Example
• Given data for Chromium, determine
its average atomic weight.
• Isotope Mass
Frac. Abundance
• Cr-50
49.9461
0.0435
• Cr-52
51.9405
0.8379
• Cr-53
52.9407
0.0950
• Cr-54
53.9389
0.0236
51.9959 u
Copper’s two
isotopes are
mass numbers
63 and 65: What
percent
abundance is
each if the
average atomic
mass is 63.5?
No calculator!
• What is the mass in
grams of a mole of
titanium atoms?
• What is the mass of
one atom of Ca-40?
• Answer in g and
amu
History of the periodic table
Dmitri Mendeleev (Russia)
• Wrote periodic
law
• Chart based on
atomic weights.
John Newlands, Great Britain
• law of octaves.
• ridiculed b/c of
inconsistencies
Julius Lothar Meyer, Germany
• periodic law around the
same time as Mendeleev.
• not credited as
Mendeleev b/c ??
Predictions
undiscovered…..
Henry Mosely
• 1887-1915
• Studied with
Rutherford,
• measured nuclear
charge / atomic
number of
elements.
Henry Mosely
• Reordered periodic
table by at.# and it
is better
Henry Mosely
• Volunteered for
service in WWI.
Was a signal officer
for the British
Army and killed in
action at Gallipoli
in 1915.
The Periodic Table
• Periods:
• 7 horizontal Rows
• Families / Groups:
18 vertical
columns
Sketch and label
Binary
• Made of two
elements.
• Examples?
Molecular
• Made of molecules
• Nonmetal atoms
making compounds
• Covalent: shared
electrons.
Ionic
• Made with a metal
ion or ammonium
and an anion…..
Ionic or Molecular?
•
•
•
•
•
Na2CO3
C4H10
MgSO4
Al2 (SO4)3
KF
• CuBr2
• H2 O
• Li2O
• NH4I
• RbClO
Making ions
• Cations vs anions
• Only changes in
electrons
Formula unit vs molecule
• Ionic vs not ionic
• Low ratio of ions vs
formula of particle
Ionic charges
• Metals vs
nonmetals
• Metals make
cations
• Nonmetals make
anions
Ion names
• Metals vs
nonmetals
• The –ide ending for
nonmetals
Nomenclature practice time
• Use the ion sheets
• Use the flowchart
• Follow the rules
• Practice!
Name these Compounds
1.
2.
3.
4.
5.
6.
HClO
HF
H2O2
PbCrO4
LiC2H3O2
CO
7. CdI2
8. N2O5
9. CuSO4
10. SrBr2
11. H3PO4
12. Ca(NO3)2
Name these Compounds
1.
2.
3.
4.
5.
6.
Hypochlorous acid
Hydrofluoric acid
Dihydrogen dioxide
Lead II chromate
Lithium acetate
Carbon monoxide
7.
8.
9.
10.
11.
12.
Cadmium iodide
Dinitrogen pentoxide
Copper II sulfate
Strontium bromide
Phosphoric acid
Calcium nitrate
Write Formulas for these Compounds
1.
2.
3.
4.
5.
6.
Cobalt III hydroxide
Barium phosphate
Magnesium chloride
Aluminum iodide
Sodium oxide
Perchloric acid
7. Nitrous acid
8. Oxalic acid
9. Hydrobromic acid
10.Diphosphorous
pentoxide
11.Dinitrogen
monosulfide
12.Silver carbonate
Write Formulas for these Compounds
1.
2.
3.
4.
5.
6.
Co(OH)3
Ba3(PO4)2
MgCl2
AlI3
Na2O
HClO4
7. HNO2
8. H2C2O4
9. HBr
10. P2O5
11. N2S
12. Ag2CO3
Organic Chemistry Intro
•
carbon based with hydrogen
and oxygen mainly.
•
Hydrocarbons: C and H only
•
Alkanes: Hydrocarbons with
only singly bonded carbon
atoms.
1st rule: C bonds 4 times
•
•
Alcohols: Have an -OH
functional group
•
Isomers: Molecules with the
same formula but different
structures
•
R: general symbol for a
carbon chain
• If R = __Carbon, it is
called:
• 1 Carbon = meth(yl)
• 2 Carbons = eth(yl)
• 3 Carbons =prop(yl)
• 4 Carbons =but(yl)
• 5 Carbons =pent(yl)
• 6 Carbons =hex(yl)
• 7 Carbons =hept(yl)
• 8 Carbons =oct(yl)
• 9 Carbons =non(yl)
• 10 Carbons =dec(yl)
ORGANIC
NOMENCLATURE
Draw… and decide if isomers
exist.
• Propane
• 1-butanol
• pentane
Name each. If they have any
isomers, name one of them too.
• CH3CH2CH2CHOHCH2CH2CH2CH2CH3
• CH3CH2CH2CH2CH2CH3
• CH3OH
STOICHIOMETRY
• Using a
balanced
equation to
make theoretical
predictions.
• Beqs show COM!
Simple reactions
• Combination
(synthesis)
• Smaller reactants
make a more
complex product.
Simple Reactions
• Decomposition
• A reactant forms
simpler products
Classify a-e
• Combination
• a c
• Decompostion
• b d e
Combustion
• Hydrocarbon and
oxygen gas
reactants
• CO2 and H2O
products
Combustion: write reactions
• Octane
• 2-Hexanol
• 3-Heptanol
“weight” Units
• Molecular or
formula weight:
u or amu
• Molar mass:
g/mol
Formula weight or
molecular weight of:
•
•
•
•
•
•
•
•
A. carbon dioxide
44.0 amu
B. water
18.0 amu
C. oxygen gas
32.0 amu
D. Table salt
58.5 amu
Molar Mass of:
•
•
•
•
•
•
•
•
A. carbon dioxide
44.0 g/mol
B. water
18.0 g/mol
C. oxygen gas
32.0 g/mol
D. Table salt
58.5 g/mol
Mole
• 6.02x1023
• Used to count
particles
• A mole of miniature
marshmallows would
cover the USA to a
depth of 600 miles.
Mole Relationships
• 1 mole = 6.02x1023 at or mc or fu
• 1 gram = 6.02x1023 amu
• 1 mole = S at wt of any formula (g)
• 1 mole gas (STP) = 22.4 L
Practice Mole
Calculations
1. 0.0365 g
2. 1.0x1024 atoms
3. 23.6 mol
4. 1.1x10-21 amu
5. 250. Liters
6. 42.1%C, 6.4%H,
51.5%O
7. 51.9 %N
• Determine the
empirical and
molecular formula of
a compound found by
combustion to
contain 39.9%
Carbon, 6.7%
hydrogen and 53.4%
oxygen. The
molecular weight of
the compound is 120
amu.
• Empirical: CH2O
• Molecular: C4H8O4
Example Problem
• Pretend to have 100
grams …. Or….
• If grams are given, use
them!
• Change grams to moles
for each element.
• Look at mole ratios to
work out lowest whole
number subscripts.
• Use known molar mass to
find molecular formula
with integer multiplier.
Problem Solving
• A sample of a
compound is found
to contain 17.5 g
Na, 39.7 g Cr and
42.8 g O. What is
its empirical
formula?
• Na2Cr2O7
• Sorbic acid is added
to food as a mold
inhibitor. Its
composition is 64.3%
C, 7.20% H , and the
rest oxygen. Its
molecular weight is
112 u. What is the
molecular formula for
sorbic acid?
• C6H8O2
•Challenge Problem, AP level
An organic acid contains only C, H and
O. A 12.72 mg sample of the acid is
completely burned in oxygen. It yields
18.63 mg of carbon dioxide and 7.62 mg
of water. What is the mass percentage of
each element in the organic acid? What’s
the empirical formula?
• Review: Write an
equation for the
formation of carbon
dioxide from its
elements.
• C + O2 ----> CO2
• How many grams
carbon are needed to
produce 150. grams
of carbon dioxide?
• 40.9 grams
• Stoichiometry uses the
balanced equation ratios.
• Balanced equation
coefficients are about
particle to particle ratios.
• Coefficients mean moles
or
atoms/molecules/formul
a units
• Or (by thinking of
Avogadro’s Hypothesis)
volumes(any units) for
gaseous substances.
Problem Solving
AMEDEO AVOGADRO
Stoichiometry
• Write an equation for the reaction of sodium in water.
• 2Na + 2H2O ---> 2Na+ + 2OH- + H2
• How many molecules of hydrogen gas are produced by the
reaction of 0.25 grams of sodium METAL?
• 3.3x1021 molecules
Stoichiometry
• Write the equation for the dehydration of ethyl alcohol and
butanoic acid into ethyl butyrate, an ester.
• C2H5OH + C3H7COOH ---> C3H7COOC2H5 + H2O
• How many grams of water can be made from 8.22x1023
molecules of ethyl alcohol (ethanol)?
• 24.6 grams
Ethyl butyrate is the odor of
pineapples.
Limiting Reagents and Percent Yield
• Write an equation for the
synthesis of aluminum
chloride.
• 2Al + 3Cl2 ---> 2AlCl3
• If 3.00 g Al react with
13.0 g Cl2, how much
AlCl3 can be produced?
Use an IRF box
2 Al
I Initial
moles
R reacted
moles
F
final moles
3g=
0.111 moles
+ 3Cl2 ------> 2AlCl3
13 g =
0.183 moles
0
moles
• Determine the
limiting reagent.
• Use initial moles
compared to how
many are required
for each reaction.
• Low number limits
the process.
Use an IRF box
2 Al
I Initial
moles
R reacted
moles
F
final moles
3g=
0.111 moles
* Limiting
reagent!!
+ 3Cl2 ------> 2AlCl3
13 g =
0.183 moles
0
moles
Use an IRF box
2 Al
I Initial
moles
R reacted
moles
F
final moles
3g=
0.111 moles
**
All 0.111
moles used
0 moles left
over
+ 3Cl2 ------> 2AlCl3
13 g =
0.183 moles
0
moles
Fill in “R” Row
• Mole ratios in “R”
row must match the
reaction coefficient
ratios.
• The next coefficient
divided by LR
coefficient, multiplied
by limiting moles --->
“R” moles
Use an IRF box
2 Al
I Initial
moles
R reacted
moles
F
final moles
3g=
0.111 moles
0.111 moles
+ 3Cl2 ------> 2AlCl3
13 g =
0.183 moles
0
moles
0.167 moles
0.111 moles
Fill in “F” row
• Subtract for
reactants, add for
products
• Once the box is
filled in with
moles, any
question can be
answered.
Use an IRF box
2 Al
I Initial
moles
R reacted
moles
F
final moles
+ 3Cl2 ------> 2AlCl3
13 g =
0.183 moles
0
moles
0.111 moles
0.167 moles
0.111 moles
0 moles
0.016 moles
0.111 moles
3g=
0.111 moles
Now solve the problem!
• If 3.00 g Al react with
13.0 g Cl2, how much
AlCl3 can be
produced?
• 14.8 grams
• If 12.0 grams is
recovered, what is
the percent yield?
• 81.0%
Problem Solving
• Thinking of the balanced
equation as a recipe
might help.
• Determine how many
times the reaction
“recipe” can be carried
out with each amount of
moles.
• Reactant that can make
the fewest “batches” is
the limiting reagent.
Limiting Reagent
and Percent Yield
Lithium metal
• Write an equation for the
synthesis of lithium
hydroxide from lithium
oxide and water.
• Li2O + H2O ---> 2 LiOH
• If 42.0 grams lithium
oxide react with 20.0
grams water, how much
LiOH can be produced?
• 53.1 grams
• If 45.0 grams are
obtained, what is the
percent yield of the
experiment?
• 84.7%
• Write an equation for the preparation of hydrocyanic acid and
water from ammonia, methane and oxygen.
• 2NH3 + 3O2 + 2CH4 --> 2HCN + 6H2O
• How many grams of HCN can be obtained from the reaction of
25.0 grams ammonia, 75.0 grams oxygen and 25.0 grams
methane?
• 39.7 grams
• If 11.0 grams HCN is obtained, what is the percent yield of the
reaction process?
• 27.7 %
Limiting Reagents
and %Yield
That’s the end of Ch. 3!
• Mole day is coming SOON!
• Make the moleata!
• Talk to all other classes about donations for mole
day?
Mole Day
• CG Chemistry
T-Shirt Ideas
T-shirt Ideas
Build the Moleata!
A.P. Chemistry
Chapter 4
Chemical Reactions: An
Introduction
I. Ionic Solutions
• Deionized water
vs. tap water,
bath water, lake
water, ocean
water?
• Ions!
Svante Arrhenius
• Svante August Arrhenius b. Uppsala,
Sweden, February 19, 1859
• son of Svante Gustaf and Carolina
• educated at the Cathedral school;
Showed an aptitude in mathematics and
physics.
• 1876: University of Uppsala:
mathematics, chemistry and physics.
• 1881: Stockholm’s Academy of Sciences.
• twice married - in
1894 to Sofia
Rudbeck, (one son)
and 1905 to Maria
Johansson (one son
and two daughters)
• died at Stockholm,
October 2, 1927, and
is buried at Uppsala.
Ionic solutions are .....
• electrolytic
• Capable of
conducting electricity
• Many ionic solids are
electrolytic in water.
• the ions =
electrolytes (particles
that conduct
electricity)
What ions are found in a solution of...
•
•
•
•
KOH?
CaCl2?
(NH4) 2SO4?
Write equations
for what these
compounds do
in water.
• KOH (s) --> K+ (aq) + OH- (aq)
• CaCl2 (s) --> Ca2+ (aq) + 2Cl- (aq)
• (NH4) 2SO4 (s) ---> 2NH4+ (aq) + SO42- (aq)
Nonelectrolytic • Make no ions (electrolytes)
in solution.
substances
•
•
•
•
•
•
Examples:
sucrose, C12H22O11
methanol, CH3OH
urea, NH2CONH2
antifreeze, HOC2H4OH
All are molecular
Dissolve but do not make ions
• C12H22O11 (s)---->C12H22O11 (aq)
• CH3OH(l) ----->CH3OH (aq)
• NH2CONH2 (s)----->NH2CONH2 (aq)
• HOC2H4OH (l)----->HOC2H4OH (aq)
Strong and Weak Electrolytes
• refers to degree (%)
of ionization of
solute.
• Acids and bases
are described as
strong or weak.
Strong vs. Weak Acids
• Strong acids
ionize 100%
• Weak acids
ionize only
partially
•
•
•
•
•
•
•
nitric
Students must Know
perchloric
the 6 strong acids
sulfuric
hydrochloric
hydrobromic
Hydroiodic
Write equations
showing what
each strong acid
does in water.
The Strong Acids in Water
•
•
•
•
•
•
Nitric
HNO3 ---> H+ + NO3Perchloric HClO4 ---> H+ + ClO4Sulfuric H2SO4 ---> 2H+ + SO42Hydrochloric HCl ---> H+ + ClHydrobromic HBr ---> H+ + BrHydroiodic HI ----> H+ + I-
Dilute vs. Concentrated
Solutions
• refers to the amount
dissolved per
volume of solution.
• dilute solutions:
small amount
dissolved.
• Concentrated: more
dissolved
Writing Chemical Equations
• can be done:
• A. molecularly: show the whole mixture.
• B. Ionically
– complete ionic: indicates any electrolytes in
mix
– Net ionic: only shows species that changed.
• AP CHEM: requires net ionic equation
writing.
• two aqueous solutions
are mixed and one of the
products is insoluble.
• Aqueous: Dissolved in
Water.
• Precipitate: insoluble
species.
• Spectator Ions: in the
mixture but do not take
part in a reaction.
I. Precipitation
Reactions
• SOLUBLE OR NOT?
• Know: Solubility rules of
ionic compounds..... page
136.
• Knowing the solubility
song / chart helps.
Examples: Write the net ionic reactions!
• 1. Potassium Chloride and Silver
nitrate react in aqueous solution.
• Net ionic: Cl- + Ag+ ----> AgCl
• 2. Ammonium sulfate and calcium
chloride react in aqueous solution.
• Net ionic: Ca2+ + SO42- ----> CaSO4
• 3. Sodium carbonate and copper II
sulfate react in aqueous solution.
• CO32- + Cu2+ -----> CuCO3
More Practice. What will happen
when the following mix?
• NiCl2 and
Na3PO4
• NaCl and
Fe(NO3)2
• Al2(SO4)3 and
KOH
• Pb(C2H3O2)2 and
NH4Cl
Combustion: burns in O2
• hydrocarbons
• Products always
CO2 and H2O(l)
• exothermic
• Heat makes H2O
vaporize.
• Recognize some
organic alkanes
and alkenes,
alkynes and
alcohols.
• Octane
• 2-Hexene
• 1-butyne
• 3-Heptanol
Equations to
Balance
Acid-Base Basics
• Taste, feel ?
• pH ranges ?
• Ions they make in
solution ?
• Household
examples of each
• Page 139
The Strong Acids in Water
•
•
•
•
•
•
Nitric
HNO3 ---> H+ + NO3Perchloric HClO4 ---> H+ + ClO4Sulfuric H2SO4 ---> 2H+ + SO42Hydrochloric HCl ---> H+ + ClHydrobromic HBr ---> H+ + BrHydroiodic HI ----> H+ + I-
The Strong Bases in water:
• LiOH
• NaOH
• KOH
LiOH ---> Li+ + OHNaOH ---> Na+ + OHKOH ---> K+ + OH-
• Ca(OH)2
• Sr(OH)2
Ba(OH)2
Ca(OH)2 ---> Ca2+ + 2OHSr(OH)2 ---> Sr2+ + 2OHBa(OH)2 ---> Ba2+ + 2OH-
Acid and Base
Definitions
Arrhenius
Acids contain
hydrogen and
make hydrogen
ions in water.
Bases contain
hydroxide and
make hydroxide
ions in water.
Bronsted and Lowry
acids and bases:
Proton donors and
proton acceptors.
Indicators to know
• Indicators are
molecules that
change colors at
different pH levels.
• phenolphthalein
• Litmus
• Methyl red
• Others?
Reactions of Acids/Bases
• Learn to Write Net
ionic equations for
many examples!
Neutralization reactions
• acid + base  water + salt.
• salt: metal cation combined
with an anion (often soluble in
water)
• General Equation
• HA + MOH -----> H2O + MA
• make examples…to net ionic
 Some Salts + acids  gases.
 Carbonates + Acids ----> carbon dioxide, water and a
salt.
 General Equation
 MCO3 + HA ----> H2O + CO2 + MA
 Write examples….. To net ionic.
 Demo and test for gas: Acetic acid and sodium
(bi)carbonate
Carbonates and
Acids
• Sulfites + Acids -----> sulfur dioxide, water
and a salt.
• General:
• MSO3 + HA -------> H2O + SO2 + MA
• examples: Write net reactions.
• SO2 gas…. Stinky, irritating
Sulfites and
Acids
Sulfides and Acids
• A sulfide reacts with an acid to produce
hydrogen sulfide gas and a salt.
• General:
• MS +HA --> H2S + MA
• examples
• test for gas?
One Base Reaction
• Decompostition of
Hydroxides
• MOH MO + H2O
• Examples…..
Oxidation and Reduction
• Aka single displacement, synthesis or
decomposition reactions.
• electrons exchanged / atoms change
oxidation states (charges).
Define: Oxidation and Reduction
• Oxidation: the
loss of electrons
• Reduction: The
gain of electrons
• LEO says “GER”
Rules for deciding Oxidation State:
• Elements alone have an ox. state of zero.
• H is always 1+, unless it is with a metal as a
hydride.
• O is always 2- unless it is in a (rare) peroxide:
H2O2, Li2O2, K2O2, or Na2O2.
• All common /main group metals keep their
periodic pattern charge.
• Other semi and nonmetal elements’ oxidation
states are determined last.
• Sum of ox. States = charge on species.
Tell each elements’ Oxidation State:
1. H3PO4
2. KNO3
3. Ca(NO2)2
4. BrO2 5. BrO36. BrO4 7. CH4
8. NH4Cl
9. Cl2O
10.N2O
11.NO
12.NO2
13.P2O5
14.KMnO4
15.Fe2(SO4)3
16.Na2C2O4
Types of “redox” reactions
• Synthesis / Combination
reactions
• Decomposition
Reactions
• “single replacement”
reactions
• Combustion reactions
• Other complex
reactions
What is oxidized,
what’s reduced?
• Copper nitrate
solution reacts with
zinc metal to make
aqueous zinc nitrate
and metallic copper.
• Lithium metal reacts
with a cobalt II
chloride solution to
make metallic cobalt
and aqueous lithium
chloride.
What is oxidized, what’s reduced?
• Ca + O2 ---> CaO
• HgCl2 ----> Hg + Cl2
Will redox occur?
Using the “Activity Series”
• See the AP pages
for the reduction
potential list.
What is reduced and what is
oxidized? What are Ox. And Red.
Agents? Write 1/2 reax.
• Examples:
• iron nail in copper sulfate.
• Aluminum foil in tin II chloride
solution.
• Copper wire is placed in silver
nitrate solution.
CH. 19, section 1
• Balancing complex
redox reactions in
acid/base
environments
Last Topic:
Solution
Chem/Stoich
“Volumetric
Analysis”
Measuring Concentration of
Solutions
• can be done in
several different
ways,
including....
• Molarity
• moles dissolved
per liter of
solution.
• Molality: moles
dissolved per
kilogram of solvent
• Mass percentage:
mass of solute
compared to mass of
solution.
• Mole fraction: moles
dissolved compared
to moles of total
solution particles.
Other
Concentration
Definitions
Molarity Equation
• M=n
V
M = Molarity
n = moles dissolved solute
V = volume of solution in
liters
Molality Equation
• m=n
•
kg
• m = molality
• n = moles dissolved
solute
• kg = mass of solvent
in kilograms.
Mass Percent Equation
• Mass of solute___
• Mass of solution
• Express concentration
as a percentage.
• Any units for mass
will do.
Mole Fraction Equation
• Moles of solute ________________
• Moles of solute + moles of solvent
• Express result as a decimal number
Practice: Molarity Problems
• What is the molarity of
a solution containing
34.2 grams of sucrose
in 2.00 liters of
solution?
• 0.0500 M
• What volume of 0.65 M
HCl holds 3.0 grams
of HCl?
• 0.126 L or 126 mL
Molarity and Dilution Problems
• What mass of silver
nitrate must be added
to a flask to make 500.
mL of 0.025 M
solution?
• 2.1 g
• How many mL of 4.00
M acetic acid are
needed to make 500.
mL of 1.00 M
solution?
• 125 mL
More Problems!
• How many mL of 8.4
M KNO3 are needed
to make 3.00 Liters
of 2.5 M solution?
• 890 mL
• Suppose 200.mL of
water are added to
400.mL of 1.20 M
HNO3. What is the
molarity of the
resulting mixture?
• 0.800 M
•
A 1.000 L sample of
polluted water was
analyzed for lead II ion by
adding excess sodium
sulfate to it. The mass of
lead II sulfate
precipitating was 220.0
mg What is the mass of
lead in the water?
• 150.3 mg
• What would be the
concentration of lead in
the water?
• 0.000725 M
Stoichiometry with Molarity
= Quantitative Analysis
Solution Stoich
• A flask contains
water mixed with
some HCl. The
solution is
titrated with 0.225
M KOH until a pH
of 7 is reached.
15.20 mL of the
KOH solution are
needed. What is
the mass of the
HCl in the flask?
• 0.125 grams
Solution concentrations
• If 35.0 grams of
potassium nitrate are
dissolved in 55.0
grams of water, the
solution has a density
of 1.108 g/mL.
Determine the
molarity, molality,
mass%, and mole
fraction concentration
values of the solution.
Learn to balance complex redox
reactions: separate note page
Include Ch. 19 section 1 problems
with the chapter 4 problem set.
Booknotes: not required for 19.1
Chemical or Physical Property? #1-7
1.
2.
3.
4.
5.
6.
7.
It’s a liquid
The pH is 12
It burns in air
It tastes sweet
It is green
It weighs 5 lbs.
It bubbles in acids
Chemical or Physical Property? #1-7
1.
2.
3.
4.
5.
6.
7.
P
P
C
P (?)
P
P
C
Intensive vs. Extensive Properties?
1.
2.
3.
4.
5.
6.
7.
Its mass is 50 g.
It dissolves in oil.
Its density is 1.5g/ml
It is 6 inches long
It conducts electricity
It is acidic
It is at room temperature.
Intensive vs. Extensive Properties?
1.
2.
3.
4.
5.
6.
7.
E
I
I
E
I
I
E
Homogeneous or Heterogeneous?
1.
2.
3.
4.
5.
6.
7.
8.
Concrete
Jello
Muddy water
Diamond
Hair
Children in a class
Tossed salad
milk
Homogeneous or Heterogeneous?
1.
2.
3.
4.
5.
6.
7.
8.
He
Ho
He
Ho
He
He
He
Ho
Element, Compound or Mixture?
1.
2.
3.
4.
5.
6.
7.
8.
C
E
C
M
E
C
C
M
9. M
10. M
Name these elements
1.
2.
3.
4.
5.
6.
7.
Sb
As
Ni
Fe
Zr
Ra
Au
8. Na
9. Sr
10. Ag
11. Ba
12. P
13. F
14. Mg
Name these elements
1.
2.
3.
4.
5.
6.
7.
antimony
arsenic
nickel
iron
zirconium
radium
gold
8. sodium
9. strontium
10. silver
11. barium
12. phosphorus
13. fluorine
14. magnesium
Write symbols for these elements
1.
2.
3.
4.
5.
6.
7.
Aluminum
Tin
Rubidium
Argon
Helium
Neon
uranium
8. lead
9. potassium
10. calcium
11. zinc
12. chlorine
13. copper
14. tungsten
Write symbols for these elements
1.Al
2. Sn
3. Rb
4. Ar
5. He
6. Ne
7. U
8. Pb
9. K
10. Ca
11. Zn
12. Cl
13. Cu
14. W
Solubility Song
• Sing
• Make a Chart
What is the chemistry
of a soluble ionic
compound?
Solubility Quiz: Soluble or not?
1. Iron II hydroxide
2. Potassium
phosphate
3. Barium nitrate
4. Strontium sulfate
5. Calcium chloride
6. Silver acetate
Solubility Quiz: Soluble or not?
1.
2.
3.
4.
5.
6.
not
sol
sol
not
sol
not
Name two solutions
with soluble salts
that would
combine to form
the precipitates in
#1, 4 and 6
Ions to Know
• Thoughts on
patterns
What is the common ionic charge
for each element?
1.
2.
3.
4.
5.
6.
7.
calcium
argon
potassium
nitrogen
chlorine
aluminum
oxygen
What is the common ionic charge
for each element?
1.
2.
3.
4.
5.
6.
7.
2+
0
1+
313+
2-
Name these ions
•
•
•
•
•
•
•
SO42NO3PO33C2H3O2NH4+
S2O32C2O42-
Name these ions
•
•
•
•
•
•
•
Sulfate
Nitrate
Phosphate
acetate
ammonium
thiosulfate
oxalate
Write formulas for these ions
•
•
•
•
•
•
•
Carbonate
Nitrite
sodium
iodite
sulfite
sulfide
bromate
Write formulas for these ions
•
•
•
•
•
•
•
CO32NO2Na+
IO2SO32S2BrO3-
What is the charge on these ions?
•
•
•
•
•
•
•
PO3
N
AsO4
ClO2
S2O8
NH4
IO4
Something New…in Chapter 1
• Antoine LaVoisier
1743-1794
French Chemist
father of modern
chemistry
At age 28 married 13-year-old Marie-Anne who
translated from English for him and illustrated his
books; she was well educated in chemistry herself.
•burned P and S in air, and proved the products
weighed more than the reactants but the weight
gained was lost from the air.
•Thus established the Law of Conservation of Mass.
•1778: demonstrated the "air" responsible for
combustion; named this portion of air oxygen and
the other part of air “azote” (Greek for no life).
•discovered that hydrogen combined with oxygen
to produce water.
Antoinne
LaVoisier
•1787: invented the system of nomenclature still used
today.
•1789: published the first modern chemical textbook,
with his theories:
•a clear Law of Conservation of Mass
•There is no such thing as phlogiston
•a list of elements, including oxygen, nitrogen,
hydrogen, phosphorus, mercury, zinc, and sulfur…..
but also included light and caloric
Lavoisier: "I have tried...to arrive at the truth by
linking up facts; to suppress as much as possible the
use of reasoning, which is often an unreliable
instrument which deceives us, in order to follow as
much as possible the torch of observation and of
experiment."
Phlogiston Theory
• Ancient Greeks
thought there were
four substances in the
world: E,A, F, W
• In the 1600’s Johann
Becher added to the
list: Phlogiston is a
5th “element”… it’s in
any substance that
burns!
•LaVoisier
•worked as a tax collector
•beheaded during the French
revolution for using public money to
fund his research.
LaVoisier beheaded in
France.
A Word about ENERGY
• What is it?
• Law of Conservation
of Energy
Energy:part of chemical reactions
• Energy is required
to break bonds.
• Energy is released
as bonds form.
• This is true in all
physical and
chemical changes.
Reactions and Energy
• Endothermic
• Exothermic
• More about NRG in later
chapters.
9. What is the % Ag in your alloy?
Compare the percent (what is your
error%) you got to the theoretical
value for the % Ag in dimes made
before 1950.
Suggest an error to account for the
difference. Choose one with the
correct direction.
Results of the atomic theory
• It yields
definitions of:
• Elements
• Compounds
• Chemical
Reactions
Dalton’s Postulates lead to two laws:
• Conservation of Mass
• Multiple Proportions:If two
elements form more than
one compound, the mass
ratios of one of the
elements in one compound
to the same element in the
other compound is always
in a small whole number
ratio.
• Think of benzene and
methane
Mendeleev
• Born in Siberia 1834
• Youngest of 14 children.
• Hated everything in
school except science.
• Father died when he was
2, mother favored him as
a student and child, she
died after he got admitted
to university at age 15.
Mendeleev
• In 1855 was told he had
two years to live, probably
had tuberculosis.
• Worked as a professor of
chemistry at St.
Petersburg, Russia
• Organized known
elements according to
their properties and thus
discovered the periodic
law.
Mendeleev
• Meyer also discovered a
periodic law, but
Mendeleev published
first.
• Was a talented and
popular public speaker
• Married Feozva, had two
children… they did not get
along. Divorced her and
married Anna, with whom
he had four children.
• The Czar looked the other
way on his “bigamy”
Mendeleev
• His periodic law was most
accepted after it was
shown his predictions of
the existence of other
elements were correct.
• Eka-silicon and two others
were discovered.
• Died in 1907 at the age of
73.
Molecular vs.
Empirical formula
• Molecular formula
is the real formula
for a compound
• Empirical formula is
the lowest ratio of
elements in the
compound.
• Example:
• ethylene glycol is
C2H6O2 (molecular)
• empirical formula is
CH3O (lowest ratio)
Structural Formula
• Arrangement of the
atoms in a formula to
show what shape,
function or type of
molecule it is.
•
H-O-H
CH3COOH
• Many organic
molecules are
frequently written
structurally.
Organic Carbon Chain Classes
• Alkanes
• Alkenes
• Alkynes
The End of Chapter 1&2 notes
• Time really matters.
• T-shirt? Submit ideas
Asap
• Mole Day volunteers
needed: make a
moleata, run a
contest.
A.P. Chemistry: Chapter 3
• Calculations
with
Chemical
Formulas
and
Equations
What is a Mole?