Chemical Reactions

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Chemical Reactions

• Chemistry is about reactions with molecules
  colliding and forming new molecules.
• A number of reactions can be classified as
  "types, allowing some general predictions on
  outcomes of reactions.

Understanding the mechanism of reactions is
important to our understanding of processes such
as acid rain, corrosion, stain removers.
Dissolution Reactions Precipitation
Reactions Acid-Base Reactions Oxidation-Reduction
Reactions
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• Dissolution Reactions
• In dissolution reactions two or more compounds
  disperse into each other to form a homogenous
  phase.
• The starting compounds could be different phases
  (e.g. a solid and a liquid), but the outcome of
  dissolution is a homogenous phase a SOLUTION
• In dissolution reactions, the compound of lower
  concentration is called the SOLUTE and the higher
  concentration component is the SOLVENT.

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• During dissolution the solvent interacts with the
  solute such that for the solute, the interactions
  between the solute and solvent dominate over the
  solute-solute interactions and solvent-solvent
  interactions.
• Dissolution reactions are considered to be
  intermediate between a chemical and physical
  process.
• In terms of it being considered to be a chemical
  process, solute-solute interactions are broken up
  and replaced by solute-solvent interactions.
• On the other hand the solution that results
  cannot be expressed as a chemical formula and
  hence the outcome of dissolution cannot be
  represented as a typical chemical equation.

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• To write an equation for a dissolution reaction
  the solvent is left out and the change in state
  of the solute denoted.
• For example dissolving sucrose in water
• C12H22O11(s) ? C12H22O11 (aq)
• s - solid
• aq - aqueous solution.

Note for dissolution reactions, the solvent need
not be water, nor necessarily a liquid Other
examples of common liquid solvents are , but
benzene (C6H6), acetone (CH3COCH3), carbon
tetrachloride (CCl4), methanol (CH3OH)
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• Dissolution of Ionic Compounds
• Most ionic compounds dissolve easily in water.
• As we have seen ionic compounds, like NaCl, have
  rigid lattices defined by the oppositely charged
  ions.
• For the ionic compound to dissolve in water, the
  water molecules must overcome the strong
  interactions that exist between the oppositely
  charged species so that ion-water
  (solute-solvent) interactions dominate over
  ion-ion interactions (solute-solute)

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• The negative end of the water molecule (O)
  interacts with the positive ions in the crystal
  and the positive end of water (H) interacts with
  the negative end
• When a water molecule encounters an ion it
  orients itself so that the appropriate "side" of
  the water molecule interacts with the ion (for
  negative ions, the H points toward the ion and
  for positive ions, the O).

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• Having oriented itself in this way, the water
  molecule essentially pulls this ion out of the
  crystal lattice.
• Other water molecule surround this one ion and
  screens the ion from the oppositely charged ions
  in the crystal.
• Hence the water molecules, solvate the ion, and
  the solvated ion then moves through the solution.
  
• By this process the ionic compound dissolves in
  water and is said to DISSOCIATE INTO ITS IONS.
• Dissolution reactions for ionic compounds are
  written as
• NaCl(s) ? Na(aq) Cl-(aq)

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• Dissolution of Covalent compounds
• Neutral molecules do not have charges by which
  they can interact with the solvent.
• However, they can interact with the solvent
  through their polarity.
• Since the polarity of a molecule is due to a
  charge separation within the molecule, solute
  molecules that are polar can dissolve in polar
  solvents in much like the way ions dissolve in
  water (a polar solvent).

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• For polar solutes, the polar solvent molecules
  orient themselves around the solute molecules so
  that the more positive end of the solvent is
  oriented towards the more negative solute
  molecule and the more negative end of the solvent
  orients itself towards the more positive end of
  the molecule.
• In this way the solute molecules are solvated
  replacing the solute-solute interactions by
  solute-solvent interactions.
• The solute molecules remain intact, but each
  solute molecule is solvated by solvent molecules.
  
• Dissolution of molecular compounds can be written
  as
• C12H22O11(s) ? C12H22O11 (aq)

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• For a solute molecule to go into solution, the
  solvent molecules must solvate the solute
  molecule so that the solute-solvent interaction
  dominate over the solute-solute interactions.
• Hence, solute molecules dissolve in solutions of
  the same polarity - like dissolves like.
• So polar solutes dissolve in polar solvents,
  non-polar solutes dissolve in non-polar solvents
• Soaps or surfactants are designed so that the
  soap molecule dissolves in water, yet can
  interact with a non-polar oil molecule in a
  grease stain

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a) Dissolution of an ionic compound
b) Dissolution of a covalent compound
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• Solubilities
• Ethanol and water dissolve in each other and are
  said to be miscible.
• As more ethanol is added, at some point the
  ethanol concentration is larger than the water
  concentration and the solute and solvent switch.
• Water and ethanol have infinite solubilties in
  the other.
• Solubility is defined as the amount of a solute
  that can dissolve in a fixed amount of solvent,
  at a given temperature
• Solubility varies with temperature - generally
  higher the temperature larger is the solubility.

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• As NaCl is added to water, a point is reached
  when the NaCl does dissolve, but remains as a
  solid in the salt solution.
• The point at which the NaCl stops dissolving in
  the salt solution, defines the solubility of NaCl
  in water, at that temperature.
• The solution is said to be saturated with NaCl.
• For a liquid solute dissolved in a liquid
  solvent, at the saturation point, a new layer is
  formed, with the new layer contains the solute
  with some of the original solvent dissolved.

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• Electrolytes and Non-Electrolytes

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NaCl(s) --gt Na(aq) Cl-(aq) When the battery
is turned on the Na ions flow toward the
negative plate (anode) and the Cl- ions to the
positive plate (cathode).
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• The flow of ions constitutes a current. The
  circuit is now complete, current flows through
  the circuit, and the bulb turns on.
• NaCl is called an electrolyte.
• Electrolyte a compound which when dissolved in
  a solvent dissociates to form ions in solution.
• Typically electrolytes are ionic compounds since
  they dissolve in solution to form ions.
• Example K2SO4
• Some covalent compounds (like acids and bases)
  can dissociate in solution to form ions.

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• Electrolytes are characterized as being strong or
  weak.
• The strength of an electrolyte depends on the
  degree to which the compound dissociates in water
  to form ions.
• Hence ionic compounds like NaCl and K2SO4 which
  dissociate completely in water are strong
  electrolytes.
• Weak electrolytes do not dissociate extensively
  in water- consequently the conductance of a
  solution of a weak electrolyte in low.
• Non-electrolytes do not dissociate in solution to
  form ions and hence their solutions do not
  conduct electricity.

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• Precipitation Reactions

2 KI(aq) Pb(NO3)2(aq) --gt PbI2 (s) 2 KNO3 (aq)
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• The reaction between the KI and Pb(NO3)2 results
  in the formation of PbI2 which has a very low
  solubility in water and forms a solid
  precipitate.
• 2 KI(aq) Pb(NO3)2(aq) --gt PbI2 (s) 2 KNO3
  (aq)
• For compounds insoluble in water, the attraction
  between the oppositely charged ions in the solid
  crystal are too strong to be overcome by solvent
  water molecules.
• KNO3, being soluble in water, exists in solution
  as K and NO3- ions.

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• The reaction
• 2 KI(aq) Pb(NO3)2(aq) --gt PbI2 (s) 2 KNO3
  (aq)
• is also called a METATHESIS reaction.
• In a metathesis reaction atoms or groups of atoms
  are switched.
• In this example K and Pb2 switch anions.
• NOTE Metathesis reactions do not have to result
  in precipitation.

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• Precipitation reactions result when
• 1) an amount exceeding the compounds solubility
  in a particular solvent is added to the solvent
• 2) Removal of solvent - example by evaporation.
• 3) Changing the solvent - since solubilities vary
  from solvent to solvent, changing the solvent can
  result in precipitation
• 4) Changing the temperature - solubilities vary
  with temperature. Cooling a solution of a
  compound can result in the compound precipitating
  out of solution.

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• Predicting Precipitation Reactions
• To determine if a product of a reaction between
  two ionic compounds will result in a precipitate
  being formed, check the solubilities of the
  compounds formed by the reaction.

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• Example If aqueous solutions of AgNO3 and NaCl
  are mixed will a precipitate form? If yes,
  identify the precipitate.
• AgNO3 (aq) NaCl (aq) --gt AgCl (?) NaNO3 (?)

AgCl is insoluble, NaNO3 is soluble AgNO3 (aq)
NaCl (aq) --gt AgCl (s) NaNO3 (aq)
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• Net Ionic Equations
• AgNO3(aq) NaCl(aq) --gt AgCl(s) NaNO3(aq)
• If we re-write the above equation in terms of the
  species that actually exist in solution, the
  equation is

Ag(aq) NO3-(aq) Na(aq) Cl-(aq) --gt
AgCl(s) Na NO3-(aq)
This is the complete IONIC equation In writing
the complete ionic equation we see that Na and
NO3- exist in the same form on both sides of the
equation, whereas the Ag and Cl- have reacted to
form solid AgCl.
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• Both Na and NO3- are called SPECTATOR IONS since
  they are present in solution but do not
  participate directly in the chemical reaction.
• They provide the Ag and Cl- ions which
  reacted, but they themselves do not directly
  participate in the chemical reaction.
• Hence, the complete ionic equation can be
  re-written to show only those species which are
  directly involved in the chemical reaction

Ag(aq) Cl-(aq) --gt AgCl(s)
This is the NET IONIC equation.
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• Note Net ionic equations, as with any chemical
  equation must be balanced, in terms of mass and
  charge.
• So for a reaction between Pb2(aq) and I-(aq) to
  form PbI2(s) must be written as
• Pb2(aq) 2 I-(aq) --gt PbI2(s)

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• Example Write the net ionic equation for the
  precipitation reaction that occurs when aqueous
  solutions of calcium chloride and sodium
  carbonate are mixed.
• First write the chemical formulas of the reactants

Next, determine what the products of the reaction
will be and which product is the precipitate.
The products of this reaction are NaCl and
CaCO3. From the solubility table, determine if
NaCl and CaCO3 are soluble in water or not. Write
the net ionic equation - make sure it is balanced
in terms of mass and charge.
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• Hence, the equation for the reaction is
• CaCl2(aq) Na2CO3(aq) --gt NaCl(aq) CaCO3 (s)
• CaCl2(aq) Na2CO3(aq) --gt 2 NaCl(aq) CaCO3 (s)
• Ca2(aq) 2Cl-(aq) 2 Na CO32- (aq) --gt
  2 Na(aq)
  2Cl-(aq) CaCO3 (s)
• Canceling the spectator ions, gives the net ionic
  equation
• Ca2(aq) CO32- (aq) --gt CaCO3 (s)
• Check that the ionic equation is balanced in
  terms of mass and charge

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• Acid-Base Reactions
• Acids and bases are probably one of the more
  commonly encountered compounds
• Acids are found in fruit (citric acid), vinegar
  (acetic acid), in our stomachs (HCl), and in acid
  rain (H2SO4, HNO3).
• Examples of bases - aqueous solutions of ammonia
  (used in household cleanser) antacids like
  Tums and Rolaids
• Acids and bases are also electrolytes they
  dissociate in water to form ions.

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• Arrhenius Acids and Bases
• Arrhenius acids are compounds which in aqueous
  solution dissociate to form H ions.
• (The H ion is also referred to as a proton since
  a H ion does not have an electron and the charge
  is due to the single proton in the nucleus)
• Hence acids are H donors or proton DONORS.

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• Examples
• HCl(aq) --gt H(aq) Cl-(aq)
• HNO3(aq) --gt H(aq) NO3-(aq)
• Acids like HCl and HNO3 are called MONOPROTIC
  acids since every molecule of HCl or HNO3
  produces one H
• Acids like HCl and HNO3 completely dissociate in
  water
• These acids are are STRONG ACIDS, and hence also
  strong electrolytes

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• H2SO4(aq) --gt H(aq) HSO4-(aq)
• The HSO4- formed can dissociate further producing
  a H. However, HSO4- is not as strong an
  electrolyte as H2SO4 and not all the HSO4- ions
  dissociate.

Dissociation of HSO4- is incomplete
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H2SO4 is a strong electrolyte since it
completely dissociates in water and hence is a
strong acid. HSO4- is a weak electrolyte since
it does not completely dissociate in water and
hence a weak acid. An aqueous solution of H2SO4
contains H, HSO4- and SO42- ions.
H2SO4 is called a DIPROTIC acid, since each
molecule of H2SO4 can produce up to 2 H
ions. Polyprotic acids - produce more that 2 H
/molecule of acid
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• Arrhenius Bases
• An Arrhenius base is a compound that when
  dissolved in water dissociates to produce an OH-
  (hydroxide) ion.

NaOH(s) --gt Na(aq) OH-(aq)
Compounds that do not contain OH- can still be
bases as long as when the dissolve in water the
chemical reaction that results produces OH-
ions. For example, if NH3 is dissolved in water,
it can react with the water to produce OH-
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• In this example, note that the reaction is not
  complete as indicated by the double arrows.
• NH3 is a weak electrolyte and hence a weak base.
  
• Also, note that NH3 received a H from water, and
  so in this reaction the water acts as an acid
  giving up an H to NH3.

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producing both H and OH- ions.

• However the extent to which pure water
  dissociates is very,very small, and a very small
  number of water molecules dissociate.
• In fact, pure water is not considered to be an
  electrolyte and does not conduct electricity.
• Compounds, like water, that can exhibit both
  acidic and basic properties are called
  AMPHOTERIC.
• Other examples of amphoteric compounds are amino
  acids which exhibit both acidic and basic
  properties.

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• Reactions between acids and bases
• When solutions of an acid and a base are mixed, a
  NEUTRALIZATION reaction occurs.
• The products of the neutralization reaction have
  neither acidic nor basic properties.
• Example
• HCl(aq) NaOH(aq) --gt H2O(l) NaCl(aq)
• In general
• acid base --gt salt water

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• The complete ionic equation for the reaction
  between HCl and NaOH is
• H(aq) Cl-(aq) Na(aq) OH-(aq) --gt
  H2O(l) Na(aq)
  Cl-(aq)
• Therefore, the net ionic equation is
• H(aq) OH-(aq) --gt H2O(l)

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• Problem
• Milk of magnesia is essentially Mg(OH)2 and is
  insoluble in water. Adding HCl to a suspension
  of milk of magnesia dissolves it leaving behind a
  clear solution.
• Write the overall equation and the net ionic
  equation.
• Overall equation
• Mg(OH)2(s) 2 HCl(aq) --gt MgCl2(aq) 2 H2O(l)
• Complete ionic equation
• Mg(OH)2(s) 2H(aq) 2Cl-(aq) --gt
  
  Mg2(aq) 2Cl- (aq) 2H2O(l)
• Net ionic equation
• Mg(OH)2(s) 2H(aq) -gt Mg2(aq) 2H2O(l)