Chapter 11: Chemical Reactions (Phelps)
Chemical Formulas
chemical formula: a formula listing the number and type of each element in a molecule or a “formula unit.” E.g., C6H12O6. Sometimes called the molecular formula.
A formula of C6H12O6 means 6atoms of carbon, 12atoms of hydrogen, and 6atoms of oxygen.
empirical formula: a chemical formula that has been reduced to lowest terms. E.g., if the chemical formula is C6H12O6 then the empirical formula would be C1H2O1 or just CH2O.
molecule: a group of elements that are chemically bonded together and behave as a single substance.
formula unit: if a substance makes large crystals and not individual molecules, a formula unit is the “molecule” that would be represented by the empirical formula.
In a chemical formula, the subscript after a symbol indicates the number of atoms of that symbol. If the subscript is after parentheses, it means that many of everything in parentheses. For example, the formula Ca3(PO4)2 means 3atoms of calcium, and 2groups of (1phosphorus atom + 4oxygen atoms). This means we have a total of 2phosphorus atoms and 8 (=4×2) oxygen atoms in the molecule.
Chemical Equations
chemical equation: a set of symbols that describe a chemical reaction. For example:
This equation states that 2molecules of H2 (gas) and 1molecule of O2 (gas) react to produce 2molecules of H2O (liquid) and heat (energy).
reactants: the starting materials; chemicals (and things like energy) that react. In a chemical equation, the reactants are before the arrow (on the left).
products: chemicals (and other things like energy) that are produced. In a chemical equation, the products are after the arrow (on the right).
reaction conditions: anything that doesn’t take part in the reaction, but is needed to make the reaction happen. In a chemical equation, reaction condition information is placed above and/or below the arrow.
Symbols Used in Chemical Equations
The following symbols are commonly used in chemical equations:
Symbol / Meaning(s) / solid
(ℓ) / liquid (A script “L” is used to avoid confusion with the number 1.)
(g) / gas
(↑) / gas (formed by the reaction and not collected)
(aq) / aqueous (dissolved in water)
(ppt) or (↓) / solid (formed by the reaction)
(cr) / crystalline (solid is in the form of crystals, not powder or lump)
Δ / heat
In the reaction:
the Δ means heat was needed to get the reaction started, even though it produced more heat once it got going.
Types of Chemical Reactions
There are many types of chemical reactions. Five of the most common are:
Synthesis (combination): two or more reactants combine to form a single product.
A + B à C
decomposition: one reactant disintegrates (decomposes) to form two or more products:
A à B + C
single replacement (sometimes called single displacement): atoms of one element replace atoms of another in a compound:
A + BC à AC + B
Most often, AC and BC are ionic compounds, which means A and B are metals, and C is a non-metal or negative polyatomic ion.
As an analogy, imagine that BC are a couple, and C breaks up with B to go out with A.
Predicting the Products
of Chemical Reactions
Single and double replacement reactions usually involve ionic compounds (and sometimes water, which we treat as the ionic compound H+OH−).
In a single replacement reaction, atoms of an element react with a compound, replacing the atom of the same type. Metals replace metals; non-metals replace non-metals. For example:
Na + CaCl2 à NaCl + Ca
(Na replaces Ca in the compound.)
KBr + Cl2 à KCl + Br2
(Cl replaces Br in the compound.)If an element reacts with a compound, you can predict what the products are going to be, because the element simply replaces the other element of the same type. For example, if you were given the problem:
Ca + NaCl à ?
Calcium is a metal, so it will replace sodium. This means calcium will end up with chloride (CaCl2), and sodium will end up by itself (Na). The reaction is therefore:
Ca + NaCl à CaCl2 + Na
Notice that we have to balance the charges every time we put two ions together. Na needed only one Cl atom to be balanced, but Cl needs two.
In a double replacement reaction, the two ions of the same type switch places, as in:
KCl + MgO à MgCl2 + K2O
(K and Mg are trading places;
K is now with O and Mg is now with Cl.)
Notice again that we had to balance the charges. We needed only one K+ ion with Cl−, but we need 2K+ ions with O2−. Similarly, Mg2+ needs only one O2− ion, but it needs two Cl− ions.
If we had the problem:
NH4OH + Ca3(PO4)2 à ?
we would swap NH4+ with Ca2+. Balancing the charges, NH4+ would now go with PO43− to form (NH4)3PO4, and Ca2+ would now go with OH− to form Ca(OH)2. This gives the equation:
NH4OH + Ca3(PO4)2 à (NH4)3PO4 + Ca(OH)2
Watch out for H+ and OH− getting together to form HOH (which we write as H2O), e.g.:
HCl + Ca(OH)2 à CaCl2 + H2O
double replacement (sometimes called double displacement): when two positive ions (or two negative ions) switch with each other to form two new compounds:
As an analogy, imagine that AX and CY are two couples. A and B switch boyfriends, so B is now going out with X and A is now going out with Y.
combustion: a special kind of reaction in which a hydrocarbon (a compound containing carbon and hydrogen) reacts with O2 (burns, or “combusts”) to form CO2 and H2O. For example:
C3H8(ℓ) + 5O2 (g) à 3CO2 (g) + 4H2O (g)+ heat
The internal combustion engine in your car is a special reactor in which octane (C8H18) combusts in a cylinder, producing heat. The heat makes the gases inside the cylinder expand. (Remember the gas laws!) This pushes the piston down, which makes the car go.
When Can A Reaction Happen?
In class, you saw some demonstrations. The first one was the following reaction between aluminum metal and copper(II) chloride:
Al(s) + CuCl2(aq) → AlCl3(aq) + Cu(s) (1)
The reverse reaction (the same reaction in the opposite direction) is:
Cu(s) + AlCl3(aq) → CuCl2(aq) + Al(s) (2)
However, this didn’t happen in the demo. Reaction(1) proceeded until all of the aluminum metal and CuCl2 were used up, and there was only AlCl3 and copper metal in the beaker. Once that happened, the beaker was set for reaction(2), but reaction(2) didn’t happen.
The second demo was the reaction between sodium carbonate (Na2CO3) and calcium chloride (CaCl2):
Na2CO3(aq) + CaCl2(aq) → NaCl(aq) + CaCO3(ppt) (3)
However, once the calcium carbonate is formed, it doesn’t redissolve. I.e., reaction(3) happens, but the reverse reaction(4), doesn’t:
CaCO3(s) + NaCl(aq) → CaCl2(aq) + Na2CO3(aq) (4)
Why might reactions (1) and (3) happen, but not reactions (2) and (4)?
Activity Series
In the reaction between aluminum metal and copper(II) chloride:
Al(s) + CuCl2(aq) → AlCl3(aq) + Cu(s) + heat
the beaker got hot. This means the reaction gave off heat, which was lost to the surroundings (the water that the chemicals were dissolved in, the beaker, the air, your hand). Once the energy was given off, the chemicals didn’t have enough energy to go the other direction. In other words:
Cu(s) + AlCl3(aq) → N.R.
(N.R. stands for “No Reaction”.)
For single replacement reactions, there is a list, called the activity series, which lists metals in order, based on how much energy they give off when they combine with a negative ion to form an ionic compound. A metal that’s higher on the list can replace anything lower on the list (more energy is given off), but a metal that’s lower on the list doesn’t have enough energy to replace one that’s higher up.
Aluminum is higher than copper on the activity series, which means aluminum can replace copper, but copper can’t replace aluminum.
Solubility
In class, you saw a demonstration of the reaction between sodium carbonate (Na2CO3) and calcium chloride (CaCl2):
Na2CO3(aq) + CaCl2(aq) → NaCl(aq) + CaCO3(ppt) (1)
However, once the calcium carbonate is formed, it doesn’t redissolve. I.e., reaction(1) happens, but the reverse reaction(2), doesn’t:
CaCO3(s) + NaCl(aq) → CaCl2(aq) + Na2CO3(aq) (2)
This is because of the way ionic compounds behave when they are dissolved in water.
If an ionic compound dissolves in water, it dissociates (splits) into its ions. This means that the “(aq)” after an ionic compound really means that the compound is dissociated, and is floating around in the solution as separate positive and negative ions.
For example, CaCl 2 splits into one Ca2+ ion and two Cl− ions. The Ca2+ ions are attracted to the negative part of the H2O molecule (the oxygen atoms), and Cl− ions are attracted to the positive parts (the hydrogen atoms).
This means that:
If: / Then:attraction to H2O > attraction to each other / compound dissociates & dissolves
attraction to each other > attraction to H2O / compound stays together and does not dissolve
The higher the charge, the stronger the attraction. This comes from Coulomb’s Law:
where F is the force of attraction between the two ions, k is a constant, d is the distance between the two ions, and q1 and q2 are the two charges.
In general, compounds with ions that have low charges (+1 or −1) tend to be soluble, because the attraction between the ions is weaker.
Ion Description / Some Examples+1 (any) / NH4+, all alkali metals (Na+, K+, etc.)
−1 (weak-to-moderate—anything except F− or OH−) / most −1 polyatomic ions (NO3−, ClO3−, CN−, C2H3O2−), most halogens (Cl−, Br−, I−)
−2 (weak only) / most SO42− compounds
You only need one of the ions to have a weak charge, because once the weak ion is associated with water molecules, the stronger ion has a much harder time finding another ion than it has finding an H2O molecule.
In terms of chemical reactions, this means that double displacement reactions only happen if the two ions stick together hard enough to precipitate. If all of the ions stay dissolved, you just have water with a bunch of ions in it, but not a chemical reaction.
Balancing Chemical Equations
Remember from Dalton’s theory of the atom:
“Atoms are neither created nor destroyed in any chemical reaction.”
In the equation:
S + O2 à SO3
There are 2 O atoms on the left, but 3 on the right. We can’t change the actual molecules that take part in the reaction, so we need a different number of each molecule.
Easiest:
S + 1½ O2 à SO3
But we can’t have ½ of a molecule of O2, so we need twice as many of each molecule, to get rid of the fractions:
2 S + 3 O2 à 2 SO3
This works because there are 2 atoms of S and 6 atoms of O on each side.
How to balance equations:
1. Figure out which elements to balance first, middle, last:
· Save for last: any element by itself
· Do first: elements that appear in one molecule on each side (if you haven’t already saved them for last).
· In the middle: everything else
2. Start with one element on the “First” list. Add coefficients to make it balance.
3. Pick another element. (Work your way through the “First,” then “Middle,” then “Last” lists.) Start with elements that already have at least one coefficient, but still need at least one.
4. Repeat step #3 until everything is balanced.
Notes:
· Polyatomic ions usually stay together.
· If you need a fractional number of a molecule, put in the fraction temporarily, then immediately multiply all of your coefficients by the denominator to get rid of the fractions.
Example:
H2SO4 + HI à H2S + I2 + H2O
1. Make lists:
a. Save for last: I
b. Do first: S,O
c. In the middle: H
2. Balance S & O (order doesn’t matter):
a. Let’s start with S:
1 H2SO4 + HI à 1 H2S + I2 + H2O
b. Then O:
1 H2SO4 + HI à 1 H2S + I2 + 4 H2O
3. Balance H:
1 H2SO4 + 8 HI à 1 H2S + I2 + 4 H2O
4. Balance I:
1 H2SO4 + 8 HI à 1 H2S + 4 I2 + 4 H2O
For the final answer, leave out any coefficient of 1:
H2SO4 + 8 HI à H2S + 4 I2 + 4 H2O