Download Ionic Equations

2/17/2009
Lecture 11: Stoichiometry between reactants;
electrolytes; precipitation reactions
• Readings for next class
– 4.4 Acid-Base Reactions
Electrolytes
• Ionic compounds and acids and bases form free ions
(DISSOCIATE) when they dissolve
NaCl(s) → Na+(aq) + Cl-(aq)
Ionic solid
Electrolyte solution
Conducts electricity
HCl(g) → H+(aq) + Cl-(aq)
Molecular
Electrolyte solution
electrolyte
• Example: Write the chemical equation for the
dissociation of aluminum sulfate in water
Ions in solution.MOV
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Electrolytes
• Strong electrolytes dissociate completely (or almost)
– Most soluble ionic compounds are strong electrolytes
– HCl(g), H2SO4(l) and HNO3(aq) are strong electrolytes
• Weak electrolytes form ions to only a small extent
– E.g. NH3(g), acetic acid (l) (CH3CO2H)
• Solutions of non-electrolytes have very low
conductivities, since they produce no ions
– Most molecular substances are non-electrolytes
– Exceptions are acidic and basic molecules such as HCl(g)
and NH3(g)
Precipitation Reactions
• Occur when TWO ELECTROLYTE SOLUTIONS are
mixed AND an INSOLUBLE IONIC COMPOUND is
formed.
• If both possible products are soluble there is NO
REACTION
• Write a balanced chemical equation for the reaction
of a sodium chloride solution with a lead(II) nitrate
solution
• Complete the following equation:
NaCl(aq) + KNO3(aq) → No reaction because
both possible products are soluble
Precipitation.MOV
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Figure 4.6:
Reaction of
Magnesium
Chloride and
Silver
Nitrate
Copyright © Houghton Mifflin
Company. All rights reserved.
Presentation of Lecture Outlines, 4a–5
Figure 4.5: Limestone Formations
Source: Royalty-Free/Corbis.
Copyright © Houghton Mifflin
Company. All rights reserved.
Presentation of Lecture Outlines, 4a–6
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Lecture 12: Solubility Rules, Ionic Equations,
Acids and Bases
• Readings for next class
– 4.5 Oxidation-Reduction Reactions
Solubility of Ionic Compounds
• Solubility can be measured as the concentration of a SATURATED
SOLUTION (i.e. when no more solid will dissolve)
• SOLUBLE COMPOUNDS (> 5 g will dissolve in 1 L of water)
– All NITRATES
– All ACETATES
– All AMMONIUM SALTS
– All GROUP 1 METAL SALTS (Li, Na, K, Rb, Cs)
– Most CHLORIDES AND SUFATES
• INSOLUBLE COMPOUNDS (< 5 g will dissolve in 1 L of water)
– Most CARBONATES
– Most HYDROXIDES
– Most SULFIDES
– AgCl, Hg2Cl2, PbCl2, SrSO4, BaSO4, PbSO4
• See Appendix 1 of Lab manual for more details
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Ionic Equations
• A NET IONIC EQUATION shows only the ions that
participate in a reaction. Other ions are called
SPECTATOR IONS
• For example:
Pb2+(aq) + 2Cl-(aq) → PbCl2(s)
This is a general ionic equation for reaction of any Pb2+
containing solution with any Cl- containing solution to form
a precipitate of PbCl2
Procedure for Writing Ionic Equations
• Write the “molecular equation” as if all substances
were molecular.
E.g. for NaCl(aq) + Pb(NO3)2(aq) →
• Write DISSOLVED ELECTROLYTES as dissociated ions
– This gives the TOTAL IONIC EQUATION
• Cancel ions that appear in equal numbers on each
side
– These are the SPECTATOR IONS
– This gives the NET IONIC EQUATION
– See previous side for the example above
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Ionic Equations
• Ionic equations must be balanced for mass and
charge
– Checking this will help you spot mistakes
• Net ionic equations are more general than
“molecular equations”
– Each one describes many possible examples with different
spectator ions
Acids and Bases
• Arrhenius definition:
– Acid - forms H+ (protons) in water
• E.g. HNO3(aq) → H+(aq) + NO3-(aq)
– Base - forms OH- in water
• e.g. NaOH(s) → Na+(aq) + OH-(aq)
(strong base)
nitric acid (strong acid)
sodium hydroxide
• Bronsted and Lowry definition:
– Acid – donates a proton to another species
– Base – accepts a proton from another species
• Eg. NH3(aq) + H2O(l) → NH4+(aq)+ OH-(aq)
base
acid
– Ionization of an acid is best described as transfer of a proton to water,
which acts as a base
• HCl(aq) + H2O(l) → H3O+(aq) + Cl-(aq)
Hydronium ion
acid base
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Strong and Weak Acids and Bases
• Strong acids and bases ionize completely
– E.g. HCl(aq) solution contains NO HCl molecules
– Also HNO3(aq), H2SO4(aq), HClO4(aq), HBr(aq), HI(aq)
– Group IA and IIA hydroxides (except Be(OH)2)
• E.g. KOH(aq) contains no KOH units
• Weak acids and bases ionize only partly
– Ammonia – solution mainly contains NH3 molecules, with
some NH4-(aq)
– Also: organic acids, HF(aq), HNO2(aq), H3PO4(aq)
– Write the chemical equation for ionization of aqueous
acetic acid. Which form predominates in the solution?
Lecture 13: Neutralization Reactions; OxidationReduction (Redox) reactions
• Readings for next class
– 4.6 Balancing Simple Oxidation-Reduction
Equations
– 4.7 Molar Concentration
– 4.8 Diluting Solutions
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Neutralization Reactions
• Reaction of an acid with a base to produce an ionic compound
(called a SALT)
– Water is usually produced as well
• For a strong acid and strong base, the NET IONIC EQUATION is:
H+(aq) + OH-(aq) → H2O(l)
– The salt is formed from the spectator ions
– Write “molecular”, total, and net ionic equations for the neutralization
of aqueous sulphuric acid with potassium hydroxide
• Neutralization causes weak acids and bases to ionize and
react completely, but they are left in molecular form in ionic
equations:
– Write “molecular”, total, and net ionic equations for the neutralization
of aqueous acetic acid with sodium hydroxide
Polyprotic acids
• Have more than on ionizable proton
• React in multiple steps
– Write “molecular equations” for the reaction of aqueous sulfuric acid
with NaOH(aq)
(a) in two steps
(b) for the overall reaction
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Formation of a gas
• Carbonates, sulfites, and sulfides react with acids to produce
gases
– Carbonates produce CO2(g)
– Sulfites produce SO2(g)
– Sulfides produce H2S(g)
• Write “molecular”, total, and net ionic equations for the
reaction of aqueous sodium carbonate with HCl(aq)
– Note that gases produced are molecular (i.e. not ionized)
• Write “molecular”, total, and net ionic equations for the
reaction of solid CuS with HCl(aq)
– Note the CuS is insoluble, and therefore not ionized
Oxidation and Reduction
• OXIDATION – loss of electrons from a substance
• REDUCTION – Gain of electrons by a substance
• A reaction that involves oxidation and reduction is called a
REDOX REACTION
• Write the molecular, total ionic and net ionic equations for the
reaction of zinc with hydrochloric acid to produce hydrogen
and a solution of zinc chloride
– Identify the OXIDIZING AGENT (OXIDANT) and REDUCING
AGENT (REDUCTANT)
– Write HALF EQUATIONS for the oxidation and reduction
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Lecture 14: Oxidation Numbers; Reaction Types;
Balancing Redox Equations
• Readings for next class
– None
– Study material from chapters 1, 2, and 3 for the
test on Thursday
– Attempt the following problem:
• Combustion of a 0.173 g sample of an organic solid
containing only C, H and O produced 0.437 g of CO2 and
0.077 g of H2O. Calculate the empirical formula of the
compound.
OXIDATION NUMBER
OXIDATION NUMBER, ON - charge on atom if all bonding is
considered to be ionic
Used to identify redox reactions and the oxidation and reduction
processes
Rules
1. ON = 0 for all elements. e.g. for H in H2
2. ON = charge of a monotomic ion. e.g. +1 for Na+, -2 for S23. ON = +1 for hydrogen (H) in all compounds except ionic hydrides (e.g. NaH)
4. ON = -2 for oxygen (O) in all compounds except H2O2 (ONO = -1) and other
peroxides
5. Sum of oxidation numbers must be
(a) zero for molecules
(b) ionic charge for ions
Determine the oxidation number of each element in AlCl3, NH3 and ClO4- .
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Use of Oxidation Numbers
• OXIDATION – ON increases by the number of electrons
involved (n)
• REDUCTION – ON decreases by the number of electrons
involved (n)
• Identify the oxidation and reduction processes in the
following reaction:
6KI(aq) + KClO3(aq) + 6HCl(aq) → 3I2(s) + 7KCl(aq) +3H2O(l)
How many electrons are involved in the oxidation?
How many electrons are involved in the reduction?
Types of Reactions
• COMBINATION REACTIONS
– Two substances combine to make a new substance
Al and Br2.MOV
– When elements combine, a redox process occurs
Write a chemical equation for the combination of hydrogen and
carbon to form methane (CH4). Identify the oxidant and reductant.
• DECOMPOSITION REACTIONS
Oxidation of Zn.MOV
– One substance reacts to give two or more substances
– Usually requires heat or a catalyst
– Carbonates produce CO2, perchlorates produce O2
Write a chemical equation for the thermal decomposition of sodium
perchorate to produce NaCl and oxygen gas. Identify any changes in
oxidation states.
Write a chemical equation for the thermal decomposition of sodium
carbonate to produce sodium oxide and carbon dioxide. Identify any
changes in oxidation states.
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2/17/2009
Lecture 15: Combustion Reactions;
Displacement Reactions; Balancing Redox
Equations
• Readings for next class
– 5.1 Gas Pressure
– 5.2 Empirical Gas Laws
Important Notice
• Assignments are due at 1:00pm in your A tutorial, and NOT
in the dropboxes outside of C-4009. These will only be used
in the event of a storm
• There are 10 different versions of each assignment. You are
only required to submit 1.
• Assignments will only be marked if they are handed in on
the original assignment sheet. Photocopies or separate
sheets are NOT acceptable. If you lose or misplace an
assignment sheet, please come to C-4009 to get a new
version.
• On the top of your assignment, please indicate when your
lab is, and what bench number you sit at in the lab. This is
for our sorting and recording purposes.
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Types of redox reactions
• COMBUSTION REACTIONS
– A substance reacts rapidly with oxygen (often produces a flame)
– Always redox reactions
– One or more oxides are produced
– Organic compounds produce CO2 and H2O
Write a balanced chemical equation for the combustion of C6H12O6(s).
Write a balanced chemical equation for the combustion of aluminum.
Identify changes in oxidation state in the above reactions.
• DISPLACEMENT REACTIONS
– An element reacts with a compound to give a new element and new
compound
– Always redox reactions
Write a chemical equation for the displacement of hydrogen from
water by calcium metal.
Thermite.mov
An activity series for predicting the occurrence of displacement reactions
Table 4.6
• Elements only
displace other
element that are
below them in the
series
• Complete the
following:
Zn(s) + CuCl2(aq) →
Cu(s) + ZnCl2(aq) →
Ag+ and Cu.MOV
Zn in acid and Cu2+.MOV
You need to learn this order
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Balancing Redox Equations
• Use the Half-Reaction Method
– Write the reduction and oxidation half-equations with the
correct number of electrons to account for the change in
oxidation state
– Multiply the complete half-equations, as needed to have
equal numbers of electrons for the reduction and
oxidation
– Add the two half-equations and cancel the electrons
Balance the following equation by the half-reaction method
Al(s) + ZnCl2(aq) → AlCl3(aq) + Zn(s)
Lecture 16: Solution Concentrations; Dilution;
Quantitative Analysis
• Readings for next class
– 5.3 The Ideal gas law
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SOLUTION CONCENTRATION
• In a SOLUTION, a SOLUTE is dissolved in a SOLVENT
• MOLARITY (Molar Concentration) - moles of solute per litre of
solution
MOLARITY = moles of solute/litres of solution
• UNITS are mol L-1 or M
• A VOLUMETRIC FLASK is used to make solutions of known
concentration
• Example:
What concentration of CuCl2is
obtained by dissolving
17.87 g of CuCl2 in 500.00 mL
of water?
Dilution
Concentrated solution + solvent ⇒ dilute solution
• MOLES OF SOLUTE DOESN’T CHANGE
• Initial moles of solute (MiVi) = final moles of solute
(MfVf)
• What volume of 2.63 M NaOH(aq) is required to
prepare 500.0 mL of 0.100 M NaOH(aq)?
Answer: 1.90 x 10-2 L
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Quantitative Analysis
Determination of the amount of a
species in a sample
• Gravimetric Analysis – the amount of a species is
determined from the mass of a product
– Precipitation is often used
• Volumetric Analysis (TITRATION) - the amount of a
species is determined from the volume of solution
needed for complete reaction
– Most commonly and acid-base reaction
(neutralization) is used
Problems
The silver in a 10.0 g sample of ore was oxidized and
dissolved to give 25.0 mL of Ag+(aq) solution. The Ag+
was then quantitatively precipitated as AgCl(s). If the
mass of AgCl obtained was 0.549 g, what was the
concentration of Ag+ in the solution?
Answer: 0.153 M
21.20 mL of 0.1000 M NaOH were required to
neutralize 10.00 mL of H2SO4(aq). What was the
initial H2SO4 concentration?
Answer: 0.1060 M
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Stoichiometry for Solution Reactions
Given volume and molarity, or mass
Mol of substance
Chemical equation
Mol required
Mass, molarity or volume required
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