chemical reactions · evidence for chemical reactions. the chemical equation ... atoms on left...
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Chemical reactions
ClassificationsReactions in solution
Ionic equations
Learning objectives
Distinguish between chemical and physical change
Write and balance chemical equations
Describe concepts of oxidation and reduction
Classify reaction according to types of reactants and products
Distinguish among strong, weak and non-electrolytes
Identify common acids and bases by from chemical formula
Predict formation of precipitates by application of solubility rules
Write total and net ionic equations from balanced molecular equations
Chemical vs physical redux
Physical: No new substance!
Chemical: New substance formed!
Evidence for chemical reactions
The chemical equation
aA + bB = cC + dD
Law of Conservation of Matter states that matter is
neither created nor destroyed
Means: atoms on left equals atoms on right
Reactant
side
Product
sidecoefficientELEMENT or
COMPOUND
Chemical book-keeping
Keys to balancing equations:
“Have I gained or lost any atoms?”
Put down the correct formula for each
reactant or product
Formulas cannot be changed in order to
balance the equation
Reaction of hydrogen with oxygen to produce
water: reactants are H2 and O2, product is H2O
Count the atoms: 4 H and 2 O 4 H and 2 O
The big number
multiplies every
atom after it
The subscript
only multiplies
the atom before it
2 H2 + O2 → 2 H2O
Molecular representation of the
reaction
Balance the equations
CH4 + O2 = CO2 + H2O– CH4 + 2O2 = CO2 + 2H2O
C3H8 + O2 = CO2 + H2O– C3H8 + 5O2 = 3CO2 + 4H2O
N2 + H2 = NH3
– N2 + 3H2 = 2NH3
Do balancing equation exercises
One approach to classification
Oxidation – reduction: focusing on
electrons
Oxidation is loss of electrons
Reduction is gain of electrons
Oxidation is always accompanied by
reductionThe total number of electrons is kept constant
Oxidizing agents oxidize and are
themselves reduced
Reducing agents reduce and are
themselves oxidized
Redox in chemistry
All reactions involve rearrangement of
atoms and molecules
Some reactions involve rearrangement of
atoms and molecules and electrons
– Photosynthesis, respiration, combustion...
These are called redox reactions
Any reaction involving elements must be
redox
Combination (synthesis)
reactionsElement + element compound (redox)– S + O2 → SO2
– Metal + nonmetal binary ionic compound
– Nonmetal + nonmetal binary covalent compound
Compound + element compound (redox)– CO + O2 → CO2
Compound + compound compound– SO2 + H2O →H2SO3
Decomposition reactions
Compound element +
element (redox)
– HgO → Hg + O2
Compound element +
compound (redox)
– PCl5 → PCl3 + Cl2
Compound compound
+ compound
– CaCO3 → CaO + CO2
Single replacement (displacement)
Element displaces another
element from compound
(redox)
Cu + 2 AgNO3 → Cu(NO3)2 + 2 Ag
Predicting single replacement
reactions: the activity series
Element higher in the
will displace one
lower in the series
The element higher is
a stronger reducing
agent
The element lower is
a stronger oxidizing
agent
Three types of double
displacement reaction
Compounds
exchanging partners
– Usually ionic
compounds in solution
Precipitation
Acid-base
neutralization
Gas formation
Precipitation
Identify ions and swap them
BaCl2 + Na2SO4 → BaSO4 + 2 NaCl
Acid – base neutralization:
special case of double replacement
KOH(aq) + HNO3(aq) = KNO3(aq) + H2O(l)
Product is liquid water not a solid
BASE ACID SALT WATER
Gas formationProduct is either a gas or is unstable and
decomposes to a gas
CaCO3(s) + 2 HCl(aq) = CaCl2(aq) + H2O(l) + CO2(g)
Writing balanced molecular equations
for double replacement reactions
Use correct formulae
– Metal ion charge predicted
from group number
– Use table for correct
formula and charge for
polyatomic ions
Identify as solid (s), gas
(g), liquid (l) or dissolved
(aq)
Balance: atoms (groups)
on left = atoms (groups)
on right
Balancing double replacement equations
It’s very much a matter of states – show
them!
Pb(NO3)2(aq) + 2KI(aq) = 2KNO3(aq) + PbI2(s)
Balance polyatomic ions as units:
– Pb2+, K+, I-, NO3-
Left hand side Right hand side1 Pb2+ 1 Pb2+
2 NO3- 2 NO3
-
2 K+ 2 K+
2 I- 2 I-
Molecular equation for reaction of
Na2SO4 + Ba(NO3)2
Combustion
Element or compound
reacting with oxygen
(redox)
– CH4 + O2 → CO2 + H2O
Associated with
production of heat and
light
Involves hydrocarbons
(CxHy), nonmetals (S) and
metals (Mg)
Sorting solution reactions:
dissolved species
Electrolytes:
– Ionic compounds produce ions in solution
(NaCl, NH4NO3 etc.)
Non-electrolytes:
– Covalent compounds do not produce ions in
solution (CH3OH, C6H12O6 etc.)
Electrolytes: distinguishing by
strength
All soluble substances that produce ions are
electrolytes
Strong electrolytes are characterized by
complete dissociation in water
Weak electrolytes dissociate to a much smaller
extent.
Strong, weak or non electrolyte?
All soluble salts are strong electrolytes
Strong acids and bases are strong
electrolytes
Weak acids and bases are weak
electrolytes
Insoluble compounds are non-electrolytes
Molecular compounds are non-electrolytes
Classification of electrolytes
Strong
electrolytes
Weak
electrolytes
Non-
electrolytes
ACIDS:
HCl, HBr, HI
HClO4, HNO3, H2SO4
ACIDS:
HF, H3PO4,
CH3CO2H
Molecular
covalent
compounds:
H2O,
CH3OH,
C12H22O11
(sucrose)
Most organic
compounds
and
INSOLUBLE
salts
SALTS:
KBr, Na3PO4
SALTS:
None
BASES:
NaOH, Ba(OH)2
BASES:
NH3
Flow chart for determining type of
electrolyte
3. Ionic or covalent?
1. Soluble in H2O?
2. Acid or base?
Nonelectrolyte
3. Weak or strong?
Strong
electrolyteWeak
electrolyte
Yes No
NoYes
Weak CovStrong Ionic
Recognizing acids and basesAcids usually have H at the beginning of the
formula – HCl
Bases usually have OH in the formula – NaOH
– Not in organic compounds though - CH3OH
Acid formula Name Base formula Name
HCl Hydrochloric acid NaOH Sodium hydroxide
H2SO4 Sulfuric acid KOH Potassium
hydroxide
H3PO4 Phosphoric acid Ba(OH)2 Barium hydroxide
HNO3 Nitric acid NH3 Ammonia
HClO4 Perchloric acid (CH3)3N Trimethylamine
CH3CO2H Acetic acid
HCO2H Formic acid
Citric acid
The strong acids and bases
Strong acids
(Only six)
Strong bases
(g1A and g2A)
HCl Hydrochloric acid LiOH Lithium hydroxide
HBr Hydrobromic acid NaOH Sodium hydroxide
HI Hydroiodic acid KOH Potassium
hydroxide
HNO3 Nitric acid Ca(OH)2 Calcium
hydroxide
H2SO4 Sulfuric acid Sr(OH)2 Strontium
hydroxide
HClO4 Perchloric acid Ba(OH)2 Barium hydroxide
Solubility roolsGroup I and ammonium (NH4
+) compounds soluble
Nitrates (NO3-), acetates (CH3CO2
-) soluble
Chlorides, bromides and iodides generally soluble
{except Pb(II), Ag(I) and Hg(I)}
Sulphates (SO42-) generally soluble (except g2A and
Pb2+)
Carbonates (CO32-), phosphates (PO4
3-) generally
insoluble (except gIA)
Hydroxides (OH-), sulphides (S2-) generally insoluble
(except gIA and gIIA)
Total ionic equations
Pb(NO3)2(aq) + K2CrO4(aq) = 2KNO3(aq) + PbCrO4(s)
Total ionic equation
Dissolved substances:
– Strong electrolytes show as ions
– Weak or non- electrolytes show as molecular
formula
Pb2+(aq) + 2NO3-(aq) + 2K+(aq) + CrO4
2-(aq) =
2K+(aq) + 2NO3-(aq) + PbCrO4(s)
Net ionic equations
Spectator ions are those ions that do not
undergo a change; they do not participate
in the chemical change and are the same
on both sides of the equation
Remove all spectator ions from the
equationPb2+(aq) + 2NO3
-(aq) + 2K+(aq) + CrO42-(aq) =
2K+(aq) + 2NO3-(aq) + PbCrO4(s)
Net ionic equations
Pb2+(aq) + CrO42-(aq) = PbCrO4(s)
Mass and charge must still balance, although overall charge may not be neutral in a net ionic equation
Net ionic equation for reaction of
Na2SO4 + Pb(NO3)2