Double Replacement Reactions – In a double replacement reaction, two soluble compounds react and form an insoluble precipitate as demonstrated in class during M-L #8.
Written using generic symbols, it can be shown as:
AX + BY → AY + BX
(A and B represent cations – usually metals)
(X and Y represent anions – nonmetal elements or polyatomic ions)
Specific steps for writing and balancing double replacement
reactions:
1. Determine the solubility of the two reactant salts (salt is a
chemist’s term for any ionic compound).
a. If one or both of the reactant salts are insoluble − STOP − no
reaction can occur because the insoluble salt can not dissociate
(break apart) and therefore cannot make the double replacement
“switcheroo”, therefore, if either salt is insoluble write; "No reaction
can occur because a reactant is insoluble."
b. Both salts must be soluble to proceed.
2. Determine the identity of the product salts by making the
“switcheroo”.
a. If given the names of the reactant salts, make the double
replacement switch to predict the names of the products.
b. If given the formulas for the reactant salts, make the double
replacement switch to predict the products, being careful to carry
just the ions not any subscripts that are there to equalize the
charge in the compound.
3. Look up the solubility of the products.
a. If both products are soluble − STOP - no reaction occurs because
there is no chemical change to the products, all ions are still just
dissolved in the water. Write; "No reaction occurs because both
products are soluble."
b. If one of the products is insoluble, it would appear as the formation
of an insoluble salt, and should be labeled (ppt) precipitate.
4. Make a skeleton equation by writing out the correct chemical
formulas of both the reactants and the products, setting the
subscripts as a result of the charges of the individual ions using
the criss-cross method. (Do not set the subscripts of the products the
same as the subscripts in the reactants – you must consider the
charges of the ions that are attaching to each other.)
5. Balance the skeleton equation changing only the coefficients out
in front of the formulas.
6. Label the appropriate product (s) or (ppt). And label the soluble
salts as (aq).
7. Convert the overall equation (sometimes called a molecular
equation) into a “net” ionic equation by dropping out the
spectator ions.
a. Any ion that is aqueous on both sides “drops out” of the equation.
There is no need for (s) or (ppt) in the net ionic equation because
representation of ions with their respective charges tell us that they
must be in solution, and any compounds written together will
assumed to be solids (or possibly liquids or gases, as appropriate.)
b. Only the ions that are part of the precipitate are included in the
equation. The spectator ions drop out.
Single Replacement Reactions – In a single replacement reaction one lone
element replaces another element in a compound.
AX + B → BX + A
(for replacement of hydrogen: HX + B → BX + H2 )
perhaps it would be helpful to show the charges of ions that carry a
charge
A+X− + B → B+X− + A
(for replacement of hydrogen: H+X− + B → B+X− + H2 )
Metal B (which has no charge in element form) replaces metal A (a
positive ion in the compound AX) to form a new compound BX and atoms
of element A (which has no charge). Remember that in this type of single
replacement reaction, A and B are both metals that form cations,
positively-charged ions when combined with X which is a nonmetal and
has formed an anion, a negatively-charged ion. When A and B are alone
or uncombined they are atoms as an element and have no charge.
Examples: Consider using the carefully outlined steps above to avoid making mistakes.
1. sodium chloride is combined with silver sulfide.
step 1 sodium chloride - soluble, silver sulfide – insoluble thus no reaction can occur.(step 1a)
2. zinc hydroxide is combined with potassium chromate.
step 1 zinc hydroxide – insoluble thus no reaction can occur. (step 1a)
3. sodium sulfate is combined with tin(II) nitrate.
step 1 sodium sulfate - soluble, tin(II) nitrate - soluble
step 2 sodium nitrate + tin(II) sulfate
step 3 both products are soluble thus no reaction occurs since no precipitate is formed.
4. ammonium sulfide is combined with sodium chloride.
step 1 ammonium sulfide - soluble, sodium chloride - soluble
step 2 ammonium chloride - soluble, sodium sulfide
step 3 both products are soluble thus no reaction occurs since no precipitate is formed.
5. calcium nitrate is combined with potassium sulfite.
step 1 calcium nitrate - soluble, potassium sulfite - soluble
step 2 calcium sulfite + potassium nitrate
step 3 calcium sulfite – insoluble (ppt) thus a reaction occurs
step 4
Ca(NO3)2 + K2SO3 → CaSO3 + KNO3
Ca2+(aq) + NO3−(aq) + K+(aq) SO32−(aq) → Ca2+ SO32−(ppt) + K+(aq) + NO3−(aq)
steps 5+6
Ca(NO3)2 + K2SO3 → CaSO3 (ppt) + 2 KNO3(aq)
step 7
Ca2+(aq) + SO32−(aq) → CaSO3(s)
6. sodium carbonate is combined with aluminum nitrate.
step 1 sodium carbonate - soluble, aluminum nitrate - soluble
step 2 sodium nitrate + aluminum carbonate
step 3 aluminum carbonate - insoluble (ppt) thus a reaction occurs
step 4
Na2CO3 + Al(NO3)3 → Al2(CO3)3 + NaNO3
Na+ + CO32− + Al3+ + NO3− → Al3+ CO32− (ppt) + Na+ + NO3−
steps 5+6
3 Na2CO3 (aq) + 2 Al(NO3)3 (aq) → Al2(CO3)3 (ppt) + 6 NaNO3(aq)
step 7
3 CO32−(aq) + 2 Al3+(aq) → Al2(CO3)3 (s)
Please note the aluminum ions are written as 2 Al3+ not Al23+. This is to represent the ions
as separated, not stuck together.
Steps to Follow to Successfully Write SR Reactions that Occur in Solution:
Typically, you will be given the reactants and asked to predict the products.
A For the replacement to occur, the reactant compound (AX) must be soluble which allows the
compound to separate into parts and therefore allows the “switcheroo” to take place.
B For the replacement to occur the reacting lone element must be "more active" than the
cation that is being replaced out of the compound. (If it is hydrogen that is being replaced,
pay attention to whether the hydrogen is in water or an acid.)
Determining if the reacting element is more active than the cation in the compound can be
accomplished by looking at the Activity Series. If the elemental reactant is less active than
"who" it is attempting to replace, the replacement cannot occur - write NR for no reaction.
C If steps A + B allow a reaction to occur, write the names for all the substances that are
formed.
• Write formulas and set the subscripts in the reactant and product compounds.
o Remember the subscripts of ionic compounds are a result of the criss-cross of the
charges on each ion which is a result of electrons lost and gained.
o Remember that subscripts cannot be added to make an equation balance.
o Remember that any hydrogen gas formed is diatomic.
• Also note that when H is replaced from water, the remaining negative ion is OH− (not
O2−).
D After the subscripts are set forming the skeleton equation, balance the equation.
E Turn the balanced equation into a net ionic equation. Aqueous ions that are in the same
form on both sides of the equation are called spectator ions and should be removed.
It is not always metals (or hydrogen) that are doing the replacing:
The Halogens can get in on the action too:
There is a second type of single replacement reaction in which the anion
(negative ion) is replaced from the compound. This type of reaction will
only be done for compounds involving the halogens (The second last
column of elements: F, Cl, Br, I, At). One anion replaces another. Written
using generic symbols, it is:
Y2 + A+Z− → A+Y− + Z2
• Element Y (a halogen) will replace Z (a less active halogen in the
compound AZ) to form a new compound AY and the free element
Z. Remember that Y and Z have no charge when they are free
elements, but they are both anions when combined with A (a
cation).
• There is no need for a separate activity series because the order in
the periodic chart will serve as the activity series – the higher on
the chart, then the more active the element.
• The solubility of the reactant compound does not matter since the
halogen will be used in either gas or liquid form and therefore can
react directly with the solid, liquid, or gaseous compound.
Remember that you will
have the activity series
right as a reference.
Examples: Consider if a reaction can occur, and if it does, write a balanced chemical
equation and then a net ionic equation.
1. nickel (II) sulfide is combined with magnesium
a. Remember to check solubility of the compound first. No reaction occurs because nickel
sulfide is insoluble and if the nickel and sulfide can’t separate then no replacement can
occur.
2. magnesium chloride is combined with cobalt
a. magnesium chloride is soluble
b. Using the activity series cobalt is below magnesium, it is less active and therefore can not
replace magnesium.
3. nickel (III) chloride is combined with magnesium
a. Nickel chloride is soluble. No need to consider the solubility of the metal.
b. Check the activity series to see if the magnesium is more active than the nickel. It is, so a
reaction can occur.
c. NiCl3(aq) + Mg(s) → Ni(s) + MgCl2(aq) this is the unbalanced skeleton equation.
d. Balance: 2 NiCl3(aq) + 3 Mg(s) → 2 Ni(s) + 3 MgCl2(aq)
e. Net ionic: 2 Ni3+(aq) + + 3 Mg(s) → 2 Ni(s) + 3 Mg2+(aq)
4. magnesium is combined with aqueous hydrochloric acid
a. the word “aqueous” tells you that the hydrochloric acid is soluble.
b. magnesium can replace hydrogen from an acid
c. Mg(s) + HCl(aq) → MgCl2(aq) + H2(g) Be sure and note that the hydrogen must form diatomic
molecules when it is replaced. Free hydrogen
atoms cannot exist
d. Balance: Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)
e. Net Ionic: Mg(s) + 2H+(aq) → Mg2+(aq) + H2(g)
5. calcium is dropped into water
a. it is silly to even think about if water as aqueous – it’s water.
b. calcium is more active than hydrogen (even hydrogen in water)
c. Ca(s) + H2O(L) → Ca(OH)2(s) + H2(g)
d. Balance: Ca(s) + 2 H2O(L) → Ca(OH)2(s) + H2(g) Calcium hydroxide is only slightly soluble so it
can be written as a solid.
e. Net ionic: Ca(s) + 2 H2O(L) → Ca(OH)2(s) + H2(g) The net ionic equation is the same as the
complete equation because no particle is in the
same form on both sides of the reaction.
6. chlorine gas is bubbled through aqueous hydrogen bromide
a. In this reaction it is the halogens that will exchange places. Because the reacting diatomic
halogen will usually be liquid or gas, it can react directly with the compound in whatever form
the compound comes in: solid, liquid, aqueous or gas. Thus it is not necessary to concern
yourself with the solubility of the reacting compound.
b. chlorine is more reactive than bromine. The halogens are more reactive as you proceed up
the periodic chart.
c. Cl2(g) + HBr(aq) → HCl(aq) + Br2(L)
d. Balance: Cl2(g) + 2 HBr(aq) → 2 HCl(aq) + Br2(L)
e. Net Ionic: Cl2(g) + 2 Br−(aq) → 2 Cl−(aq) + Br2(L)
Synthesis Reactions – Synthesis reactions occur when two substances join together to produce one, more complex compound. The starting substances can be elements or
compounds. The product compound has more atoms than the reactant substances. Written using generic symbols, it can be shown as:
A + B → AB
Examples: Predict the products, write the skeleton equation, then balance
Important Points to Remember (i.e. memorize)
• Notice that in ALL synthesis reactions there is only one substance on A. magnesium metal reacts with oxygen gas (produces magnesium oxide)
skeleton: Mg(s) + O2(g) → MgO(s)
the product (right-hand) side. This is our definition of a synthesis
balanced: 2 Mg(s) + O2(g) → 2 MgO(s)
reaction.
B. hydrogen gas reacts with oxygen gas (produces water)
• You should notice some patterns right:
skeleton: H2(g) + O2(g) → H2O(L)
o metal chlorides react with oxygen to produce the metal chlorate.
balanced: 2 H2(g) + O2(g) → 2 H2O(L)
o metal oxides react with carbon dioxide to produce the metal
C. potassium reacts with chlorine gas (produces potassium chloride)
carbonate. (not redox)
skeleton: K(s) + Cl2(g) → KCl(s)
o metal oxides react with water to produce a metal hydroxide solution.
balanced: 2 K(s) + Cl2(g) → 2 KCl(s)
(not redox)
o (nonmetal oxides react with water to produce acids, you will not be D. iron reacts with oxygen gas (the iron will be III) (produces iron(III) oxide)
skeleton: Fe(s) + O2(s) → Fe2O3(s)
asked to predict these products.)
balanced: 4 Fe + 3 O2(g) → 2 Fe2O3 (s)
• Notice that three of these reactions (a, b, and d) would also fit into the
combustion category as well as this synthesis category as long as E. calcium oxide reacts with carbon dioxide to produce calcium carbonate
skeleton: CaO(s) + CO2(g) → CaCO3(s)
the reaction occurred fast.
already balanced
• Notice that two elements are combining in each example to produce
F. sodium oxide reacts with carbon dioxide to produce sodium carbonate
one product.
Na2O(s) + CO2(s) → Na2CO3(s)
• Don't forget about the “9 + H” (H2 N2 O2 F2 Cl2 Br2 I2 P4 S8)
already balanced
• Notice that unreacted metals are written as single atoms (no
G. potassium chloride reacts with oxygen gas to produce potassium chlorate
subscripts).
skeleton: KCl(s) + O2(g) → KClO3(s)
Synthesis can also be
balanced: 2 KCl + 3 O2(g) → 2 KClO3(s)
• two compounds making a single more complex compound
• a compound and an element joining together to make a more complex H. barium chloride reacts with oxygen gas to produce barium chlorate
skeleton: BaCl2(s) + O2(g) → Ba(ClO3)2(s)
compound
balanced: BaCl2(s) + 3 O2(g) → Ba(ClO3)2(s)
Decomposition Reactions – During decomposition, one reactant compound splits apart into two (or more) substances. (Essentially this is the reverse of the synthesis reaction.)
The product substances can be elements or compounds. Generally heat is used to induce the decomposition. Written using generic symbols, it is usually shown as:
AB → A + B
Decomposition can split one compound into its elements. Decomposition can also split one compound into two simpler compounds (or a compound and an element).These are
some examples of compounds decomposing into their elements. Practice writing the skeleton equations, then take the time to practice balancing.
a. solid mercury(I) oxide decomposes into its elements (a gas and a liquid)
Important Points to Remember (i.e. memorize)
skeleton: Hg2O(s) → Hg(L) + O2(g)
• Always remember the “9 + H” (H2 N2 O2 F2 Cl2 Br2 I2 P4 S8). All
balanced: 2 Hg2O(s) → 4Hg(L) + O2(g
other elements will be written as Y (no subscript) unless you
b. liquid water decomposes into its gaseous elements
are told otherwise.
skeleton: H2O(L) → H2(g) + O2(g)
• Notice how, in every case left, there is only one substance on
balanced: 2 H2O(L) → 2 H2(g) + O2(g)
the reactant (left-hand) side. This is always the case in a
decomposition reaction.
c. magnesium chloride decomposes into its elements (a gas and a solid)
• The reaction is always redox if at least one element is formed.
skeleton: MgCl2(s) → Mg(s) + Cl2(g)
• You may notice some patterns above - the first three are the
already balanced
reverse of the three synthesis reactions that you should
memorize
d. iron(II) sulfide decomposes into its solid elements
The following decompositions occur upon heating.
skeleton: FeS(s) → Fe(s) + S8(s)
o Chlorate salts decompose into chloride salts and oxygen
balanced: 8 FeS(s) → 8 Fe(s) + S8(s)
gas.
e. liquid hydrogen peroxide (H2O2) decomposes into liquid water and oxygen gas.
o Carbonate salts decompose into oxide salts and carbon
skeleton: H2O2(L) → H2O(L) + O2(g)
dioxide.
balanced: 2 H2O2(L) → 2 H2O(L) + O2(g)
o Hydroxide salts decompose into oxide salts and water.
o Hydrogen peroxide decomposes into water and oxygen gas.
f. solid calcium carbonate decomposes into solid calcium oxide and gaseous carbon dioxide.
• As required by the decomposition of hydrogen peroxide in
skeleton: CaCO3(s) → CaO(s) + CO2(g)
class, some reactions require a catalyst, which is a
already balanced
substance that changes the speed of a chemical reaction
g. solid sodium carbonate decomposes into solid sodium oxide and gaseous carbon dioxide.
without itself undergoing a permanent chemical change in the
skeleton: Na2CO3(s) → Na2O(s) + CO2(g)
process.
already balanced
h. solid potassium chlorate decomposes into solid potassium chloride and oxygen gas
skeleton: KClO3(s) → KCl(s) + O2 (g)
balanced: 2 KClO3(s) → 2 KCl(s) + 3 O2(g)
i. solid barium chlorate decomposes into solid barium chloride and oxygen gas
skeleton: Ba(ClO3)2(s) → BaCl2(s) + O2(g)
balanced: Ba(ClO3)2(s) → BaCl2(g) + 3 O2(g)
j. magnesium hydroxide decomposes into magnesium oxide and water
skeleton: Mg(OH)2(s) → MgO(s) + H2O(L)
already balanced
Combustion Reactions – Combustion, at its most general, can mean the fast reaction of oxygen gas (O2) with any substance (that serves as the fuel) usually producing heat and
light (flame) and oxide products. Generic reaction:
X + O2(g) → X?O? (g) + Energy
Burning Organic Compounds - Our society depends on energy, which we get primarily from combustion of very particular types of compounds called organic compounds (any
compound containing carbon). So most commonly we will use the term combustion to mean the reaction of oxygen with a compound containing carbon and hydrogen, and maybe
oxygen as well. These compounds fall into two general categories:
• hydrocarbons (C,H compounds)
• carbohydrates (C,H,O compounds in which the H and O are present in a 2:1 ratio)
• various other organics
A common synonym for combustion is "burn". Uncontrolled combustion happening VERY rapidly would be an explosion. Written using generic symbols, combustion of organic
compounds could be shown as:
CxHy + O2(g) → CO2(g) + H2O(g) + Energy
An important hint when balancing combustion equations is to always balance the oxygen last.
Burning Compounds that contain more than just C H O - Notice that when burning hydrocarbons and carbohydrates the products are all the same, in every reaction. Some
carbon hydrogen, oxygen compounds may also contain other elements. When you burn these types of compounds, the other elements end up as oxides also. Burning compounds
that contain nitrogen (burns to form NO2) or sulfur (burns to form SO2).
Burning Compounds that don’t contain any C H or O - Other substances that don’t contain carbon, hydrogen or oxygen can be burned. In class we have burned magnesium.
Since there is no carbon or hydrogen present, water and carbon dioxide cannot be products. Instead the product is magnesium oxide (note that the product is MgO, because
magnesium is Mg2+ and oxygen is O2−).
Incomplete Combustion - There are still more complexities with the concept of combustion as you probe even deeper. Sometimes all of the carbon in the fuel can’t “get” all of the
oxygen that it would like to bond with. For example, an inefficient combustion caused by not having enough oxygen or occurring low burn temperatures, yields the usual products,
CO2 and H2O as well as carbon monoxide, CO and elemental carbon (soot), C. It is impossible to balance an incomplete combustion because there is no way of knowing just how
much of each of the carbon monoxide, carbon, and carbon dioxide forms. In all of the reactions that you write in this first-year chemistry course, unless told otherwise, you can
assume “complete” combustion occurs.
So why is carbon monoxide SO dangerous? - Actually, carbon monoxide is no more toxic to your cells than carbon dioxide, however, if you are exposed to a mix of air that
contains even very small quantities of carbon monoxide, the hemoglobin on your red blood cells will preferentially choose carbon monoxide NOT oxygen gas. This will mean that
you would be asphyxiated much faster, than when exposed to a mix of gases that contains air and carbon dioxide.