Transcript No Slide Title

CHEMISTRY
Nitrogen monoxide
can be prepared by the
oxidation of ammonia
by the following
The Central
Science equation:
9th Edition
NH3 (g) + O2 (g) NO
(g) + H2O (g)
If 45.7 g of NH3 and 52.5 g of O2 react together,
how many g of NO will be formed?
Limiting reactant prob
David P. White
If 18.59 mol of H2 is burned in air, what is the
theoretical yield (in mol) of H2O?
If 7.50 mol H2O are formed, what is the
percent yield?
If 18.59 mol of H2 is burned in air, what is the
theoretical yield (in mol) of H2O?
H2 + O2
H2O
If 7.50 mol H2O are formed, what is the percent
yield?
Chapter 4:
All about Aqueous Solutions
Electrolytic Properties
• Water is a poor conductor of electricity
• Aqueous solutions of ions can conduct electricity.
• Three types of solutes:
• Strong electrolytes: (solute is all ions)
• Weak electrolytes: (some ions, mostly molecules)
• Non-electrolytes:
(no ions, all molecules)
Models of dissolution (figure 4.3)
Ionic substance in water
Molecular
substance
in H2O
Ionization vs. dissociation
Ions form in water in two ways
Dissociation: ionic substance dissociates (separates)
Ionization: molecular substance (no ions) reacts
with water to form ions
Strong and Weak Electrolytes
• Strong electrolytes:
Exist as 100% ions, conducts electricity
Nearly 100%
•Weak electrolytes:
•make a small % of ions when dissolved.
•ions in equilibrium with the molecule.
•Can be very soluble, just not ionized
Compounds in Solution
Ionic compounds are strong electrolytes
Polyatomic ions remain intact as ions
when dissolving in water
Molecular compounds remain intact as molecules
when dissolving in water (non-electrolytes)
no ions in solution = nothing to transport electric charge.
Picture of strong electrolyte
Strong acid: HCl(g) + H2O(l)
H3O+(aq) + A-(aq)
Picture of weak electrolyte
Weak acid: HA(aq) + H2O(l)
H3O+(aq) + A-(aq)
Soluble vs. electrolyte
Don’t confuse solubility and strong vs. weak electrolytes.
Electrolyte means only that substance exists as ions in
water
Ionic?
Soluble?
HCl
KOH
BaSO4
No
HC2H3O2
YES
Electrolyte?
Precipitation reactions
Double replacement reaction:
ions switch
Special case of double replacement reactions
also called exchange or metathesis
A solid (precipitate) forms in these reactions
??How do we know what
the precipitate will be??
Precipitation Reactions
Solubility rules to remember
Soluble only means
greater than 0.01 moles dissolve in 1 L of solution
1. All nitrates, acetates, ammonium and Group 1 salts are
soluble
2. Solubility of chlorides, bromides and iodides
(all soluble except Ag+ Pb2+ and Hg22+)
3. Hydroxides (all insoluble except rule 1, Ca, Sr, Ba)
4. What to know
•
Sulfates mostly soluble
•
Phosphates and carbonates (insoluble except rule 1)
Net Ionic reaction Concept check
Predict the products when NaOH (aq) is combined
with HCl (aq), write a balanced chemical equation
(including states of matter)
Write the total ionic equation
Write the net ionic equation
Net ionic reactions
• Molecular equation (or “complete” equation):
all species listed like molecules with full formulas:
HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq)
total ionic equation: lists all ions:
H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) 
H2O(l) + Na+(aq) + Cl-(aq)
•Cross out “spectators” or ions on both sides of the
arrow (lazy bums that don’t react)
H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) 
H2O(l) + Na+(aq) + Cl-(aq)
•Net ionic equation:
lists unique ions, only those that react:
H+(aq) + OH-(aq)  H2O(l)
Steps to write net ionic equations:
1. Write a balanced complete equation.
2. Dissociate any and only strong electrolytes.
3. Cross out spectators.
4. What’s left is the net ionic eq.
These forms of compounds do not make ions
molecules
weak electrolytes
Water
gases
precipitates
Keep together, they don’t exist as ions in
solution
Concept check:
What is the complete equation?
NiCl2(aq) + 2AgNO3(aq)  2AgCl(s) + Ni(NO3)2(aq)
What is the total ionic equation?
What is the net ionic equation?
A solution of lead (II) nitrate is mixed with a solution
of sodium chloride.
Write the balanced chemical equation, the total ionic
equation, and the net ionic equation for the reaction
Acids
• Acid = substance that ionizes to form H+ in solution
(HCl, HNO3, CH3CO2H, citric, vitamin C).
H+ also written as H3O+ (hydronium)
• one acidic proton= _______________
(HC2H3O2).
• two acidic protons = _______________
(H2SO4).
Bases
Bases
• Bases = form OH-, or react with acids
(NH3, Drano™, Milk of Magnesia™).
• Metal hydroxides are strong bases:
Ba(OH)2
Ba2+ + 2OH -
Bases
Bases 2
• Some molecules, like amines, are weak bases.
• Weak bases ionize in water to make OH -
Memorize strong acids and bases
•Memorize 7 strong acids:
HCl, HBr, HI
H2SO4, HNO3, HClO3, HClO4
OR
Memorize the _________ common weak acids
•Soluble hydroxides are strong bases
Reactions of acids and bases
•Neutralization: acid + base are mixed:
HNO3(aq) + KOH(aq)  ???
•Salt = ionic compound
cation from base
anion from acid.
•Neutralization of acid with metal hydroxide
produces water and a salt.
•Acids + carbonates = CO2 and H2O
Concept check
Which substance has the most ions form when
dissolved in aqueous solution?
1. NaC2H3O2
2. K2CO3
3. Na3PO4
Concept check
Classify each of the following as a
strong (1), weak (2) or non-electrolyte (3)
When dissolved in water:
NaCl
(NH4)2CO3
C12H22O11
HF
CH3OH
(weak acid)
Concept check:
A solution of nickel (II) chloride is mixed with a
solution of silver nitrate, a ppt forms.
What is the complete equation?
What is the total ionic equation?
What is the net ionic equation?
Concentration for calculations
20.0 mL * 0.0183 g/mL = 0.366 g cobalt (II) nitrate
Concentration 2
Molarity, or number of moles per liter of solution
Molarity = moles of solute
=M
Volume of solution (liters)
Molarity can be used like a conversion factor
Volume x moles = moles
L
Moles x
L
moles
= volume (L)
A 1.0 M (molar) NaCl solution has 1.0 mol or
58.5 g of NaCl dissolved in water to make 1.0 L
To make a 250.0 mL solution of 1.00 molar NaCl
1. Measure mass
14.58 g
2. Dissolve in some water
3. Fill volumetric flask
to known volume
0.2500 L
4. M = moles/L
14.58 g x 1 mol/58.54 g x 1/0.2500 L
= 1.000 mol/L
0.20 L of 2.0 M solution:
0.20 L x 2.0 mol/L = 0.40 moles total
Pour ½ in each beaker
in each beaker:
0.10 L of 2.0 M solution:
0.10 L x 2.0 mol/L = 0.20 moles
Same conc, fewer particles
0.10 L x 2.00 mol/L = ______________
Pour 0.10 L H2O into beaker
0.10 L of solution + 0.10 L H2O = 0.20 L
0.20 moles
=
0.20 L new solution
_____
Same # of moles, but more dilute
What is molarity if 0.450 mol of NaCl is dissolved to
make 0.3500 L of solution?
What is molar concentration when 3.18 g of NaNO3
is dissolved to make 150.0 mL of solution?
Molarity, moles/L
What is molar concentration when 3.18 g of
NaNO3 is dissolved to make 150.0 mL of
solution?
3.18 g NaNO3 x 1 mol/84.99g = 0.0374 mol
0.0374 mol/0.1500L =
Molarity, moles/L
How many g of potassium sulfate are required to
make 235.0 mL of 0.152 M solution?
Molarity, moles/L
How many g of potassium sulfate are required to
make 235.0 mL of 0.152 M solution?
The mole highway
From mass
To mass
To Moles
With g/mole ratio or M
With g/mole ratio or M
From Moles
mole
mole
Use mole ratio from equation
How many g of lead (II) iodide can be made by
mixing
25.0 mL of 0.230 M potassium iodide with
25.0 mL of 0.140 M lead (II) nitrate?
Write molecular and net ionic equations.
Draw a mental model.
0.879 M H2SO4 is added to
45.0 mL of 0.100 M NaOH,
until the acid is just completely neutralized.
How many mL of H2SO4 were added?
pH is a measure of the acidity of a solution
pH is calculated as pH = - log10[H+]
or [H+] = 10-pH
What is pH of solution if [H+] = 1.35*10-4 M?
-log 0.000135 = ______ (use 2 decimals)
What is [H+] if pH = 4.25?
pH scale
Figure 16.5
• “auto-ionization” of water
H2O(l) + H2O(l)
H3O+(aq) + OH−(aq)
55.6 M
Water treated like a molecule (very few ions)
Oxidation and reduction
Some reactions are a transfer of eMg(s) +2HCl(aq)  MgCl2(aq) + H2(g)
Write net ionic equation:
Mg(s) +2H+(aq)  Mg2+(aq) + H2(g)
Oxidation and reduction
Mg(s) +2H+(aq)  Mg2+(aq) + H2(g)
• In the above rxn, Mg(s) loses e-, H+ gains e• Oxidized: atom, molecule, or ion becomes more
positively charged.
• Reduced: atom, molecule, or ion becomes less
positively charged.
Figure 4.13
Activity series
• Some metals are easily oxidized (lose e–),
others are not.
• Activity series: list of metals in decreasing ease of
oxidation.
• Metals higher on the activity series are more active
lose e– more easily.
• Any metal can be oxidized by the ions of elements
below it.
A copper strip in silver (I) nitrate solution
Time = 0
Time = 60 min
Which is more active, copper or silver?
What is the chemical reaction equation?
Zinc is higher than copper.
Is zinc more active?
Will Cu2+ ions oxidize Zn?
Will Zn2+ ions oxidize Cu?
Oxidation numbers
• Oxidation numbers:
• A tool to judge whether a substance has been
oxidized or reduced.
–
–
–
–
–
Elements
Monatomic ion
Oxygen
Hydrogen
halogens
Oxidation reations
• What elements are oxidized or reduced
CH4 + 2O2  CO2 + 2H2O
Fe2O3 + 3CO  2Fe + 3CO2
Br2 + 2NaI  2NaBr + I2
Cu(OH)2 + 2HNO3  Cu(NO3)2 + 2H2O
a castle in Westphalia, Germany, built in 1702
Lincoln Castle, Lincolnshire, England
Sulfate deposition LINK
2.4 million tonnes
14.8 million tons
SOx from coal, (some from oil) NOx from cars
Pseudoephedrine HCl
Hydrochloride
Citrate
Sulfate
Tartrate
sildenafil citrate
Oxycodone HCl
Glucosamine sulfate