Chapter 3
Chemical Reactions
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Chemical and Physical Properties• Chemical Changes
– rusting or oxidation– chemical reactions
• Physical Changes– changes of state– density, color, solubility, melting, boiling
– Extensive Properties: depend on quantity– Intensive Properties: do not depend on quantity
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States of Matter• Changes from one state to
another: Physical Change•heating•cooling
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Physical Change vs. Chemical Change
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Physical Change vs. Chemical Change
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Chemical Equations
Symbolic representation of a chemical reaction (chemical change) that shows:
1. -reactants on left side of reaction2. -products on right side of equation3. -relative amounts of each using coefficients
H2 + O2 H2O
for a reaction to occur molecules, atoms, ions must interact with one another in the appropriate orientation under the right conditions
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Chemical Equations• Are an attempt to show on paper
what is happening at the molecular level
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Chemical Equations
• Look at the information an equation provides:
• • reactants products
1 formula unit 3 molecules 2 atoms 3 moles (molecule/mole) (moles/f.u.) (moles/f.u.)
(molecules.f.u.)
the states of matter also listed
)(CO 3 + Fe(s)2 CO(g) 3 +(s) OFe 232 g
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Chemical Equations
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Chemical Equations• Law of Conservation of Matter
– Matter is neither created nor destroyed in a chemical reaction• -There is no detectable change in quantity of
matter in an ordinary chemical reaction• -Balanced chemical equations must always include
the same number of each kind of atom on both sides of the equation
OH 4 CO 3 O 5 HC 22283
Balancing equations is a skill acquired only with a lot of practice!!!
By working many problems
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Balancing Composition Reactions
Na(s) + Cl2(g) NaCl(s)
Mg(s) + O2(g) MgO(s)
Al(s) + Br2(l) AlBr3(s)
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Balancing Reactions On Your Own
P4(s) + O2(g) P4O10(s)
CO(g) + O2(g) CO2(g)
P4(s) + Cl2(g) PCl3(l)
SO2(g) + O2(g) SO3(g)
P4O6(g) + O2(g) P4O10(s)
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Balancing Decomposition Reactions
N2O(g) N2(g) + O2(g)
H2O2(aq) H2O(l) + O2(g)
AgBr(s) Ag(s) + Br2(l)
NH4HCO3(s) NH3(g) + H2O(g) + CO2(g)
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Balancing Displacement Reactions on Your Own
AgNO3(aq) + Cu(s) CuNO3(aq) + Ag(s)
Al(s) + H2SO4(aq) Al2(SO4)3(aq) + H2(g)
Cl2(g) + NaI(aq) I2(s) + NaCl(aq)
CaCl2(aq) + Na3PO4(aq) NaCl(aq) + Ca3(PO4)2(s)
Ca(OH)2(aq) + HNO3(aq) Ca(NO3)2(aq) + H2O(l)
Ca(NO3)2(aq) + K2CO3(aq) KNO3(aq) + CaCO3(s)
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Law of Conservation of Matter
Combustion reaction: the burning of a fuel in oxygen producing oxides or oxygen containing compounds– -NH3 burns in oxygen to form nitrogen
monoxide and water
OH 6 + NO 4 O 5 + NH 4
or
OH 3 + NO 2 O + NH 2
223
2225
3
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Law of Conservation of Matter
• C7H16 burns in oxygen to form carbon dioxide and water.
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Solutionsa mixture of two or more substances dissolved in
anotherSolute: substance present in the smaller amount that is dissolved
by the solventSolvent: substance present in the larger amount that dissolves the
solute
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Properties of Aqueous Solutions
• Electrolytes – produce ions in solution and conduct
electricity
– Strong electrolytes • ionize or dissociate 100% in water
– NaCl(s)Na+(aq) + Cl-(aq)
– Weak electrolytes • ionize or dissociate much less than 100% in
water– HF(l) H+(aq) + F-(aq)
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Strong Electrolytesconduct electricity extremely well
in dilute aqueous solutions– -ionize in water 100%
Examples:1. HCl, HNO3, etc
• strong soluble acids2. NaOH, KOH, etc
• strong soluble bases3. NaCl, KBr, etc
• soluble ionic salts
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Strong Ionic Salts
3(aq)2(aq)
100% OHs23
-(aq)(aq)
100% OH(s)
NO 2Ca )Ca(NO
ClNaNaCl
2
2
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Weak Electrolytes
conduct electricity poorly in aqueous solutions-ionize much less than 100% in water
Examples:1. CH3COOH, (COOH)2
• weak acids
2. NH3, Fe(OH)3 • weak bases
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Properties of Aqueous Solutions
Nonelectrolytes solutes that do not conduct electricity in water – do not “ionize”
• Examples:• C2H5OH – ethanol• Sugars – glucose, sucrose, etc.
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Aqueous Solution Conductivity
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Solubility
• maximum amount of solute that can dissolve in a given amount of solvent– -defined as the amount of solute that dissolves in
100 g solvent
• Unsaturated Solution: • contains less than the maximum amount that
dissolves• Saturated solution: • contains the maximum amount that dissolves• Supersaturated solution: • contains more than the maximum amount that
normally dissolves
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Solubility
Rules for determining solubility:• soluble (dissolves) vs. insoluble (does not
dissolve)
Figure 5.3 on page 179
OH- and O2-, except Ba2+
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Soluble Insoluble Exceptions1. Group IA and ammonium salts (Li+, Na+, K+, NH4
+)___________
2. Acetates, nitrates, chlorates, perchlorates (CH3COO-, NO3
-, ClO3-,
ClO4-)
___________
3. most chlorides, bromides, and iodides (Cl-, Br -, I-)
Salts formed with Ag+,
Hg2+, Pb2+
4. most fluorides (F-) Salts formed with Group IIA
5. most sulfates (SO42-) Salts formed with Group IIA
(Ca2+, Sr2+, Ba2+), Ag+, Hg2+, Pb2+
6. most carbonates, phosphates, sulfides (CO3
2-, PO4
3-, S2-)
Salts formed with Group IA and NH4
+ (rule #1)
7. most oxides (O2-) _______________________
8. most hydroxides (OH-) Salts formed with Group IA and Ca2+, Sr2+
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Solubility
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Metathesis Reactions
two ionic aqueous solutions are mixed and the ions switch partners
AX + BY AY + BX
Metathesis reactions remove ions from solution in 3 ways:
1. form H2O – neutralization (acid-base reactions)2. form an insoluble solid (precipitation reactions)3. form a gas
• -Ion removal is the driving force of metathesis reactions
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Precipitation ReactionsThree representation:
1. 1. Molecular equation
2. 2. Total ionic equation
aq3saqaq3 NaNOAgClNaCl AgNO
Ag+(aq) + NO3-(aq) + Na+ (aq) + Cl-(aq)
AgCl(s) + Na+ (aq) + NO3-(aq)
3. Net ionic equation
Ag+(aq) + Cl-(aq) AgCl(s)
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Precipitation Reactions
• 1. Molecular equation
(s)3)aq(3aq)(32(aq)23 CaCO +KNO 2 COK + )Ca(NO
2.Total ionic reaction
s3-
aq3aq
-2aq3aq
-aq3
2aq
CaCO NO 2K 2
COK 2 NO 2 Ca
3. Net ionic reaction
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Arrhenius Acids
substances that generate H3O+ (H+) in aqueous solutions
-Strong acids ionize 100% in water
-aqaq
%100g
-aqaq3
%100g
Cl H HCl
Cl OH HCl
orwater
-
aq3aqOH
3
-aq3aq3
100%2 3
NO + H HNO
or
NO + OH OH HNO
2
(l)
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Substances that donate protons (H+)
• Strong Acids
• Formula Name1. HCl hydrochloric acid2. HBr hydrobromic acid3. HI hydroiodic acid4. HNO3 nitric acid5. H2SO4 sulfuric acid6. HClO3 chloric acid7. HClO4 perchloric acid
Bronsted-Lowry Acids
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aq3-aq3
7% 23 OH + COOCH OH COOHCH
aq-aq3
7%
3 H + COOCH COOHCH
Acids
•-Weak acids ionize <100% in water
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•Common Weak Acids
•Formula Name
1.HF hydrofluoric acid
2.CH3COOH acetic acid (vinegar)
3.HCN hydrocyanic acid
4.HNO2 nitrous acid
5.H2CO3 carbonic acid (soda water)
6.H3PO4 phosphoric acid
Acids
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• Substance that produce OH- ions in aqueous solution (water)
– Strong bases ionize 100% in water
(aq)OH 2 + (aq)Ba Ba(OH)
(aq)OH + (aq)K KOH-+2
2
-+
Arrhenius Bases
• Weak bases are covalent compounds that ionize <100% in water
-(aq)aq4
100%
2g3 OH + NH OH + NH
(l)
C C
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Substances that accept protons (H+)
• Strong bases:1. LiOH, NaOH, KOH, RbOH, CsOH,
Ca(OH)2, Sr(OH)2
2. Notice that they are all hydroxides of IA and IIA metals
Bronsted-Lowry Bases
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Acid-Base (neutralization) Reactions
form water and salt (ionic compound)– acid + base salt + water
• 1. Molecular equation
)(2 (aq)(aq)(aq) OH + KBr KOH + HBr
2. Total ionic equation
)(2-aqaq
-aqaq
-aqaq OH + Br+KOH+K+Br+H
3. Net ionic equation
l
(l)
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Acid-Base (neutralization) Reactions
1. Molecular equation
)(2aq)(23(aq)3(aq)2 OH 2 + )Ca(NOHNO 2 + Ca(OH)
2. Total ionic equation
)(2-
aq32aq
-aq3aq
-aq
2aq OH 2 +NO 2+ CaNO 2+ H 2+OH 2+Ca
3. Net ionic equation
(l)
(l)
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There are four acid-base reaction combinations
that are possible:
1. strong acids – strong bases2. weak acids – strong bases3. strong acids – weak bases4. weak acids – weak bases
Acids and Bases
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• Polyprotic acids: •Have more than 1 hydrogen ion that it can
donate to a base
1 mol sulfuric acid reacts with 1 mol sodium hydroxideH2SO4(aq) + NaOH(aq) NaHSO4(aq) + H2O(l)
1 mol sulfuric acid reacts with 2 mols sodium hydroxide
H2SO4(aq) + 2NaOH(aq) Na2SO4(aq) + 2H2O(l)
Acids and Bases
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Gas Forming Reactions
H2CO3 H2O(l) + CO2 (g)
H2SO3 H2O(l) + SO2 (g)
NH4OH NH3(g) + H2O(l)