Today is Friday (!),April 20th, 2018

Pre-Class:

What’s that?

In This Lesson:Chemical Reactions

(Lesson 3 of 4)

Today’s Agenda

• Where we are and where we’ve been.

• Chemical Reactions

• Balancing Chemical Equations

• Types of Chemical Reactions

• Where is this in my book?

– P. 346 and following…

By the end of this lesson…

• You should be able to write, balance, classify, and predict chemical reactions.

A Wide-Range Review

• Way back, in a time called the beginning of the semester and a place called here, we talked about matter.

• We looked at its forms and properties.

• We learned how to measure it and describe it.

• We talked about what it’s made of (atoms) and what atoms are made of too.

A Wide-Range Review

• Then we talked about electrons.

• We talked about where we might find them at any given time and how they’re arranged in elements.

• We talked about how chemists arrange elements in the periodic table.

• We learned how elements bond with one another.

A Wide-Range Review

• We also learned how to name the combinations they form.

• We learned another scale for measurement: the mole.

• We learned how to measure proportions and write formulas.

• Now, we’re going to learn how to write formulas for entire chemical reactions.

Chemical Reactions

• A chemical reaction is when a set of chemicals is changed into another set of chemicals.

• Reactions can be endergonic or exergonic.

– Endergonic: Energy absorbed.

– Exergonic: Energy released.

• Exergonic reactions can happen spontaneously.

• More on this to come in the next unit.

Chemical Reactions

• Previously, we discussed how reactions are shown in basic form:

• Reactants (starting stuff) are shown on the left of the equation.

• Products (ending stuff) are shown on the right of the equation.

• The arrow means “yields.”

• Example:– Reactant(s) Product(s)

– 4Fe + 3O2 2Fe2O3

Just checking…

• Identify the reactant(s):– Na + Cl NaCl

• Na, Cl

• Identify the reactant(s):– 2H2O 2H2 + O2

• H2O

• Identify the product(s):– Na + Cl NaCl

• NaCl

• Identify the product(s):– 2H2O 2H2 + O2

• H2 + O2

One other thing…

• Don’t forget the symbols we covered back in the beginning of the semester:

– (s) means a chemical is in solid form.

– (l) means a chemical is in liquid form.

• Most liquids will be H2O for us.

– (g) means a chemical is in gaseous form.

• Watch for BrINClHOF elements!

– (aq) means a chemical is in aqueous form –dissolved in water.

• Acids are always (aq).

Lastly…

• Don’t forget the BrINClHOF (diatomic) elements:– Bromine (Br2)– Iodine (I2)– Nitrogen (N2)– Chlorine (Cl2)– Hydrogen (H2)– Oxygen (O2)– Fluorine (F2)

• When these elements are on their own, they bond to themselves. YOU MUST REMEMBER THIS!!!!!!!!!!!!!!!!!!!!!1

Types of Reactions

• Now let’s talk about the types of chemical reactions.

• For this class, you’ll need to know five of them:– Combination Reactions (also known as Synthesis)

– Decomposition Reactions

– Single Replacement Reactions

– Double Replacement Reactions

– Combustion Reactions

• Let’s take a look at them in chemistry terms as well as…prom terms.

1: Synthesis

+

“I told you they were together!”

1. Combination (Synthesis) Reactions

• In combination/synthesis reactions, two or more chemicals combine to make a new compound.

– A + X AX

• Examples include:

– Reactions with oxygen and sulfur.

– Reactions of metals with halogens.

– Reactions with oxides.

2: Decomposition

“Well we all saw that coming.”

+

2. Decomposition Reactions

• In decomposition reactions, a single compound breaks down into two or more simpler substances.– AX A + X

• Examples include (KNOW THESE):– [metal]CO3 [metal]oxide + CO2

– [metal]OH [metal]oxide + H2O

– [metal]ClO3 [metal]chloride + O2

– Acids [nonmetal]oxide + H2O

3: Single Replacement

“Scandalous!”

++

3. Single Replacement Reactions

• In single replacement reactions, a lone element takes the place of an element in a compound.

– A + BX AX + B

– BX + Y BY + X

• Examples include:

– Metals replacing metals.

– Hydrogen in water being replaced by a metal.

– Hydrogen in acid being replaced by a metal.

– Halogens being replaced by more reactive halogens.

Activity Series

• When we talked about Single Replacement Reactions, I mentioned “more reactive halogens.”

• As it turns out, elements (not just halogens) have varying degrees of reactivity.

• Chemists have created lists called Activity Series that detail how (in our class’s case) metals and halogens react with one another, assuming they even react at all.– Also known as Reactivity Series. Makes sense…

Activity Series of Metals

• Metals can replace other metals if they’re above the metal they’re replacing.

• Metals can replace Hydrogen in acids if they’re above Hydrogen.

• Metals can replace Hydrogen in water if they’re above Magnesium.

Activity Series of Halogens

• A halogen can replace another halogen in a compound if it is above the one it’s replacing.

• Example:– 2NaCl (s) + F2 (g) 2NaF (s) + Cl2 (g)

– MgCl2 (s) + Br2 (g) NO REACTION

• Note that the halogen activity series is the same as the group order of halogens on the table.

Activity Series: Reaction or No?

• Cu + MgSO4 Mg + CuSO4

– No reaction (Magnesium is above Copper).

• Pb + ZnSO4 Zn + PbSO4

– No reaction (Zinc is above Lead).

• Fe + 2AgNO3 2Ag + Fe(NO3)2

– Reaction (Iron is above Silver).

• 2Al + 3H2O Al2O3 + 3H2

– No reaction (Aluminum is not above Magnesium).

Single Replacement Technicalities**Technically, this is important.

• When a Group I or Group II metal (alkali/alkaline earth) reacts with water, they only replace one of the hydrogens.

• Examples:

– K + H2O KOH + H2

– Mg + H2O Mg(OH)2 + H2

• In other words, they form hydroxides, not oxides.

Single Replacement Exceptions**Not covered on any quiz/test

• Please know there are cases where the activity series isn’t so “cut and dry.”

– For example, magnesium can sometimes react with water if it’s water vapor, not liquid water.

– Alternatively, reactions of metals with acids can sometimes be influenced by whether the acid is a strong or weak acid.

• As in, copper can react with strong acids like nitric acid (HNO3) but won’t react so readily with something like acetic acid (HC2H3O2).

• As with everything in chemistry, there are exceptions.

Aside: Coins

• There’s an interesting phenomenon with pocket change relating to the Activity Series:– Pennies tend to become very dull relatively quickly, yet

quarters and other “silvery” coins tend not to. What’s the deal?

• As it turns out, pennies are plated in copper, while other coins are plated in nickel. – Because nickel is higher on the list, it’s less likely to be

replaced and thus tarnish.

– Copper, on the other hand, is a bit of a chemical “weakling.”

http://www.harpercollege.edu/tm-ps/chm/100/dgodambe/thedisk/series/3postlab.htm

http://www.usmint.gov/about_the_mint/index.cfm?action=coin_specifications

Single Replacement Reaction Example

• Thermite!

– Used often for welding train rails together (and other such heat-intensive reactions), thermite is a relatively simple single replacement reaction.

• It can reach 2480 °C (or 4500 °F).

– Two videos:

• Al + Fe2O3 Al2O3 + Fe

• Al + MnO2 Al2O3 + Mn

4: Double Replacement

“Gross!”

+ +

4. Double Replacement Reactions

• In double replacement reactions, ions of two compounds flip places in aqueous solutions, forming two new compounds.

– AX + BY AY + BX

• Typically, one of the compounds formed is:

– A precipitate (a solid or a bubblin’ gas).

– A molecular compound (usually water).

• One of the products must be insoluble!

Reminder: Dissociation

PbNO3

NO3

Pb2+NO3

-

NO3-

Bound ions in… …component ions out.

Predicting States of Matter

• Remember that double replacement reactions occur in solutions.

• To predict the states of matter resulting from a double replacement reaction, first write the equation.

• Then, use your solubility rules.– FYI, when they say “salts involving,” just think of it

as saying “ionic compounds involving…”

– FYI, when they say “halides,” just think of it as saying “halogens…”

Solubility Example

• NiNO3 (aq) + KBr (aq) → KNO3 (?) + NiBr (?)

• In what states are potassium nitrate and nickel (I) bromide?

• According to your solubility rules:– 1./2. All salts of Group IA and nitrates are soluble, so

potassium nitrate is.

– 3. All salts of halides (halogens) are soluble, except… so nickel (I) bromide is.

• NiNO3 (aq) + KBr (aq) → KNO3 (aq) + NiBr (aq)– So no reaction, since both of them are soluble.

• Remember, one must be insoluble!

Solubility Example

• Ba(NO3)2 (aq) + (NH4)3PO4 (aq) → NH4NO3 (?) + Ba3(PO4)2 (?)

• In what states are ammonium nitrate and barium phosphate?

• According to your solubility rules:– 1./2. All salts of ammonium and nitrates are soluble so

ammonium nitrate is.– 5. All salts of …phosphate… are insoluble except… so

barium phosphate is not.

• Ba(NO3)2 (aq) + (NH4)3PO4 (aq) → NH4NO3 (aq) + Ba3(PO4)2 (s)– One compound is insoluble, so there will be a reaction.

Precipitates

• In the previous example, Ba3(PO4)2 “fell out” of solution.– In other words, it took on a solid form and was no

longer dissolved.

– We would expect it to collect at the bottom of the container.

• This is an example of precipitation, or the formation of a precipitate.– A precipitate is a solid or gas substance that “falls

out” of an aqueous solution.• A precipitate could be water, but this is less common.

Precipitate Video/Demo

• KI (aq) + Pb(NO3)2 (aq) KNO3 (?) + PbI2 (?)

• Check your solubility rules for the phase of the two products.– 1./2. All salts of Group IA and nitrates are soluble,

so potassium nitrate is.

– 3. All salts of halides are soluble except those of lead (II), so lead (II) iodide is insoluble.

• KI (aq) + Pb(NO3)2 (aq) KNO3 (aq) + PbI2 (s)

• [Video]

Remembering Solubility Rules?

• The Solubility Song!

– Link available in Chemistry Links.

– Lyrics available in Worksheets and Keys.

5: Combustion

+

This one’s hard to picture. Basically, oxygen reacts with something, usually releasing a lot of light and/or heat.

5. Combustion Reactions

• In combustion reactions, a substance reacts with oxygen, typically releasing a large amount of energy in the form of heat and light.

• When the reactants are only oxygen and a hydrocarbon, carbon dioxide and water are the products.– Hydrocarbons are compounds made of only hydrogen,

carbon, and/or oxygen.

• Examples include:– C3H8 (g) + 5O2 (g) 3CO2 (g) + 4H2O (g)– P4 (s) + 5O2 (g) P4O10 (s)

• This is a combustion and synthesis reaction!

Combustion Reaction Demo

• C2H5OH + 3O2 2CO2 + 3H2O

Aside: Great Moments in Science

• Meet Pilatre de Rozier:

• Mr. Rozier wanted to test the flammability of hydrogen, so he inhaled some, then exhaled over an open flame.

• Result?– Singed eyebrows.

http://www.sciencephoto.com/image/228189/350wm/H4180269-Pilatre_de_Rozier,_French_balloonist-SPL.jpg

Identifying Chemical Reactions

• Let’s practice identifying chemical reactions:

– Chemical Reactions Packet, Page 2, Upper Section

• Don’t worry about balancing them yet.

Balancing Chemical Equations

• In addition to identifying chemical reactions, they also need to be balanced.

• According to the Law of Conservation of Mass/Matter, the mass of the reactants must equal the mass of the products.

– Atoms are conserved.

• So, all chemical formulas must show the same AMOUNTS OF ATOMS on both sides of the arrow.

– No elements can appear or disappear, either.

• Unless it’s a nuclear reaction.

PhET

• Balancing Chemical Equations

Skeleton Equations

• Up till this point in the semester, sometimes we’ve been writing equations that are not balanced, just to describe which elements are reacting.

• These unbalanced equations are called skeleton equations.– Think “bare bones” equations.

• From now on, we’ll need to balance our equations, so here are some directions.

How to Balance Chemical Equations

• Under the arrow, vertically list each element.

– Don’t use any additional subscripts.

• Put a box around each term in the equation.

• Use coefficients to balance each side.

– NOT subscripts.

• Balance hydrogen second-to-last and oxygen last.

– How to remember this?

Balancing Chemical Equations Example

• ___Al + ___O2 ___Al2O3

Al

O

1 2

2 3

2

2

3 2

6 6

4

4

4

☺

Important Note

• Keep in mind that, like in empirical formulas, the coefficients in a balanced equation should not be able to be reduced.

• In other words:

– 4Na + 2Cl2 4NaCl should really be

– 2Na + Cl2 2NaCl

• Even if it’s balanced, it has to be reduced to the lowest whole number ratio.

Balancing Synthesis Reactions1 of 3

1. ___CaO + ___H2O ___Ca(OH)2

– CaO + H2O Ca(OH)2

2. ___P4 + ___O2 ___P2O5

– P4 + 5O2 2P2O5

3. ___Ca + ___O2 ___CaO

– 2Ca + O2 2CaO

4. ___Cu + ___S8 ___ CuS

– 8Cu + S8 8CuS

Balancing Synthesis Reactions2 of 3

5. ___S8 + ___O2 ___SO2

– S8 + 8O2 8SO2

6. ___H2 + ___N2 ___NH3

– 3H2 + N2 2NH3

7. ___H2 + ___Cl2 ___HCl

– H2 + Cl2 2HCl

8. ___Ag + ___S8 ___Ag2S

– 16Ag + S8 8Ag2S

Balancing Synthesis Reactions3 of 3

9. ___Cr + ___O2 ___Cr2O3

– 4Cr + 3O2 2Cr2O3

10.___Al + ___Br2 ___AlBr3

– 2Al + 3Br2 2AlBr3

11.___Na + ___I2 ___NaI

– 2Na + I2 2NaI

12.___H2 + ___O2 ___H2O

– 2H2 + O2 2H2O

Balancing Decomposition Reactions1 of 3

1. ___BaCO3 ___BaO + ___CO2

– BaCO3 BaO + CO2

2. ___MgCO3 ___MgO + ___CO2

– MgCO3 MgO + CO2

3. ___K2CO3 ___K2O + ___CO2

– K2CO3 K2O + CO2

4. ___Zn(OH)2 ___ZnO + ___H2O

– Zn(OH)2 ZnO + H2O

Balancing Decomposition Reactions2 of 3

5. ___Fe(OH)2 ___FeO + ___H2O

– Fe(OH)2 FeO + H2O

6. ___Ni(ClO3)2 ___NiCl2 + __O2

– Ni(ClO3)2 NiCl2 + 3O2

7. ___NaClO3 ___NaCl + ___O2

– 2NaClO3 2NaCl + 3O2

8. ___KClO3 ___KCl + ___O2

– 2KClO3 2KCl + 3O2

Balancing Decomposition Reactions3 of 3

9. ___H2SO4 ___H2O + ___SO3

– H2SO4 H2O + SO3

10.___H2CO3 ___H2O + ___CO2

– H2CO3 H2O + CO2

11.___Al2O3 ___Al + ___O2

– 2Al2O3 4Al + 3O2

12.___Ag2O ___Ag + ___O2

– 2Ag2O 4Ag + O2

Balancing Single Replacement Reactions1 of 2

1. ___AgNO3 + ___Ni ___Ni(NO3)2 + ___Ag

– 2AgNO3 + Ni Ni(NO3)2 + 2Ag

2. ___AlBr3 + ___Cl2 ___AlCl3 + ___Br2

– 2AlBr3 + 3Cl2 2AlCl3 + 3Br2

3. ___NaI + ___Br2 ___NaBr + ___I2

– 2NaI + Br2 2NaBr + I2

4. ___Ca + ___HCl ___CaCl2 + ___H2

– Ca + 2HCl CaCl2 + H2

Balancing Single Replacement Reactions2 of 2

5. ___Mg + ___HNO3 ___Mg(NO3)2 + ___H2

– Mg + 2HNO3 Mg(NO3)2 + H2

6. ___ Zn + ___H2SO4 ___ZnSO4 + ___H2

– Zn + H2SO4 ZnSO4 + H2

7. ___K + ___H2O ___KOH + ___H2

– 2K + 2H2O 2KOH + H2

8. ___Na + ___H2O ___NaOH + ___H2

– 2Na + 2H2O 2NaOH + H2

1. ___AlI3 + ___HgCl2 ____AlCl3 + ____HgI2

– 2AlI3 + 3HgCl2 2AlCl3 + 3HgI2(s)

2. ___HCl + ___NaOH ___NaCl + ___H2O

– HCl + NaOH NaCl + H2O

3. ___BaCl2 + ___H2SO4 ___BaSO4 + ___HCl

– BaCl2 + H2SO4 BaSO4 + 2HCl

4. ___Al2(SO4)3 + ___Ca(OH)2 ___Al(OH)3 + ___CaSO4

– Al2(SO4)3 + 3Ca(OH)2 2Al(OH)3 + 3CaSO4

Balancing Double Replacement Reactions1 of 3

Balancing Double Replacement Reactions2 of 3

5. ___AgNO3 + ___K3PO4 ___Ag3PO4 + ___KNO3

– 3AgNO3 + K3PO4 Ag3PO4 + 3KNO3

6. ___CuBr2 + ___AlCl3 ___CuCl2 + ___AlBr3

– 3CuBr2 + 2AlCl3 3CuCl2 + 2AlBr3

7. ___Ca(C2H3O2)2 + ___Na2CO3 ___CaCO3 + __NaC2H3O2

– Ca(C2H3O2)2 + Na2CO3 CaCO3 + 2NaC2H3O2

8. ___NH4Cl + ___Hg2(C2H3O2)2 __NH4C2H3O2 + ___Hg2Cl2– 2NH4Cl + Hg2(C2H3O2)2 2NH4 C2H3O2 + Hg2Cl2

Balancing Double Replacement Reactions3 of 3

9. ___Ca(NO3)2 + ___HCl ___CaCl2 + ___HNO3

– Ca(NO3)2 + 2HCl CaCl2 + 2HNO3

10. ___FeS + ___HCl ___FeCl2 + ___H2S

– FeS + 2HCl FeCl2 + H2S

11. ___Cu(OH)2 + ___HC2H3O2 ___Cu(C2H3O2)2 + ___H2O

– Cu(OH)2 + 2HC2H3O2 Cu(C2H3O2)2 + 2H2O

12. ___Ca(OH)2 + ___H3PO4 ___Ca3(PO4)2 + ___H2

– 3Ca(OH)2 + 2H3PO4 Ca3(PO4)2 + 6H2

Balancing Combustion Reactions1 of 3

1. ___CH4 + ___O2 ___CO2 + ___H2O

– CH4 + 2O2 CO2 + 2H2O

2. ___C2H6 + ___O2 ___CO2 + ___H2O

– 2C2H6 + 7O2 4CO2 + 6H2O

3. ___C3H8 + ___O2 ___CO2 + ___H2O

– C3H8 + 5O2 3CO2 + 4H2O

Balancing Combustion Reactions2 of 3

4. ___C4H10 + ___O2 ___CO2 + ___H2O

– 2C4H10 + 13O2 8CO2 + 10H2O

5. ___C5H12 + ___O2 ___CO2 + ___H2O

– C5H12 + 8O2 5CO2 + 6H2O

6. ___C6H14 + ___O2 ___CO2 + ___H2O

– 2C6H14 + 19O2 12CO2 + 14H2O

Balancing Combustion Reactions3 of 3

7. ___C2H4 + ___O2 ___CO2 + ___H2O

– C2H4 + 3O2 2CO2 + 2H2O

8. ___C2H2 + ___O2 ___CO2 + ___H2O

– 2C2H2 + 5O2 4CO2 + 2H2O

9. ___C6H6 + ___O2 ___CO2 + ___H2O

– 2C6H6 + 15O2 12CO2 + 6H2O

Balancing Equations Practice

• Chemical Reactions Packet (key online)

• You must complete at least 3 problems in every section on this packet (except page 4).

• You must score 30 points or higher:– Page 1, Upper Section: 1 point each

– Page 1, Lower Section: 2 points each

– Page 2, Upper Section: 1 point each

– Page 2, Lower Section: 2 points each• Includes Page 3 – must have states of matter.

Combustion Reaction Details

• Chemical Reactions Packet, Page 4

Predicting Products

• Here’s a little conceptual question:– If methane (CH4) combusts, what are the reactants

and what are the products?

• Since combustion reactions always use oxygen gas as a reactant and form water and carbon dioxide as products (if the other reactant is a hydrocarbon), the skeleton equation would look like this:– Skeleton: CH4 + O2 CO2 + H2O

– Balanced: CH4 + 2O2 CO2 + 2H2O

Predicting Products [reminder]

• Typical reaction processes:

– Metal carbonates break down to metal oxides and CO2.

– Metal hydroxides break down to metal oxides and H2O.

– Metal chlorates break down to metal chlorides and O2.

– Acids break down to [nonmetal] oxides and H2O.

– Group IA and IIA metals reacting with water replace only one of the hydrogen atoms.

Keep an eye out…

• Some reactions are multi-stage; the product(s) may further decompose.

• The following products easily decompose in solution:– Carbonic Acid (H2CO3) makes H2O + CO2

– Ammonium Hydroxide (NH4OH) makes H2O + NH3

– Sulfurous Acid (H2SO3) makes H2O + SO2

• Notice how you can see the H2O come out of the initial product:– H2CO3 H2O + CO2

Keep an eye out…

• Also, know these:

– H3PO4 H2O + P2O5

– H3PO3 H2O + P2O3

– HNO3 H2O + N2O5

– HNO2 H2O + N2O3

– H2CO3 H2O + CO2

– NH4OH H2O + NH3

– H2SO3 H2O + SO2

• Sorry guys, it’s how it is.

• Notice how you can see the H2O come out of the initial product:

– H2CO3 H2O + CO2

Predicting Products Special Practicespecial practice is special

• H2CO3

– H2CO3 H2O + CO2

• Ca + H2O

– 2Ca + 2H2O 2CaOH + H2

• KClO3

– 2KClO3 2KCl + 3O2

• HNO3

– 2HNO3 N2O5 + H2O

Predicting Products

• Once you can spot the type of the reaction, you can easily predict the products.

• Here’s a website to help:

– http://www.files.chem.vt.edu/RVGS/ACT/notes/Practice_Predicting.html

– Also linked on my Chemistry Links page.

Predicting Products BIG HINTS

• Don’t worry about subscripts when predicting products.– Figure out who’s together and/or who’s alone.

– THEN write subscripts.

• Also, don’t forget the diatomic (BrINClHOF) elements.

• Finally, make sure you’re replacing cationswith cations (written first) and anions with anions (written second).

Reaction Type Summary

+ Double Replacement

+ Single Replacement

Decomposition

+ Synthesis

Combustion+ O2

Quick Interlude

• What kind of reaction is this?• C6H12O6 + 6O2 6CO2 + 6H2O

• That’s right, it’s combustion, since oxygen is a reactant.– Notice also that since we are combusting a hydrocarbon,

the products are water and carbon dioxide.

• Most importantly, what is this reaction?– Yep, it’s cellular respiration. You know, the thing all your

cells, and nearly all cells everywhere, are doing right now?– Your body runs off combustion! That’s why they call it

“burning Calories.

More Predicting Products Practice

• Now for reactions other than combustion:

– Equations Worksheet, Lower Section

• First, just figure out the type of reaction.

• Then predict the products.

• Then balance.

– Predict the Products worksheet

• Hey Honors kiddies – try those difficult ones at the bottom.

NOTE

• The next section will NOT be covered on Unit 4 Quiz 4 (but will be on the Unit 4 Test).

Reminder: Dissociation

PbNO3

NO3

Pb2+NO3

-

NO3-

Bound ions in… …component ions out.

Net Ionic Equations

• There’s one thing I haven’t mentioned yet, and it concerns single or double replacement reactionsand ionic compounds only.– It’s what ionic compounds actually do when they’re in

solution.

• Recall that in solution, dissolved ionic compounds dissociate into their component ions.

• Example:– CuSO4 in solution becomes Cu2+ and SO4

2-.– NaCl in solution becomes Na+ and Cl-.

• This allows them to conduct electricity.

Dissociation Practice

• Into what does H2SO4 dissociate?

• What’s the compound made of, and how many of each?

• 2H and 1SO4

• Write their charges in:

• 2H+ and 1SO42-

• Ta-da!

Net Ionic Equations

• As a result of these compounds dissociating, however, they really don’t actually interact with one another as they would outside of solution.

• We write ionic equations to express this new system of interaction for compounds in solution (aq).

• Let’s try an example…

Net Ionic Equations

• Imagine we react sodium sulfate and barium chloride.

• Na2SO4 + BaCl2

• What are the products (double rep. reaction)?

• Na2SO4 + BaCl2 NaCl + BaSO4

• That’s the skeleton equation. Balanced?

• Na2SO4 + BaCl2 2NaCl + BaSO4

• States of matter (use solubility table)?

• Na2SO4 (aq) + BaCl2 (aq) 2NaCl (aq) + BaSO4 (s)

Na2SO4 (aq) + BaCl2 (aq) 2NaCl (aq) + BaSO4 (s)

• That up there is the molecular equation.

• Now for the procedure:

• Separate (aq) compounds into their ions (no more subscripts except for polyatomic ions).

• Use coefficients to indicate how many of each on each side:• 2Na+ + SO4

2- + Ba2+ + 2Cl- 2Na+ + 2Cl- + BaSO4 (s)

– That right there is the complete ionic equation.• Note: there should be (aq) for all the other terms – I just

couldn’t fit it all on one line.

2Na+ + SO42- + Ba2+ + 2Cl- 2Na+ + 2Cl- + BaSO4 (s)Remember, (aq) around all the other terms but BaSO4

• Notice something about the equation?– “2Na+” and “2Cl-” appear on both sides of the equation,

unchanged.

• These are spectator ions; ions in solution that do not take part in the reaction.– Like fans at a sports game, they’re around, but they really

don’t have that much to do with the score.

– In fact, they don’t really even bond with each other.

• The ions that are not spectating are the driving force.– In this case, SO4

2- and Ba2+ are the driving forces behind the reaction. More on this in a little bit…

2Na+ + SO42- + Ba2+ + 2Cl- 2Na+ + 2Cl- + BaSO4 (s)Remember, (aq) around all the other terms but BaSO4

• So let’s try something. Let’s eliminate the spectator ions from the equation, since they really didn’t actively participate anyway.

• Ba2+ (aq) + SO42- (aq) BaSO4 (s)

• That right there is our net ionic equation, an equation describing a single- or double replacement reaction (aq) in which one compound becomes insoluble and precipitates.

Double Replacement Reactions

• There’s one more thing to take note of here.

• Double replacement reactions happen because of one or more driving forces:– Formation of a solid precipitate.

– Formation of a gas precipitate.

– Formation of a molecular compound (mostly water).

• If none of these happen, the reaction won’t happen either.

It starts like this…

Ag NO3

Ag+NO3

-

NaCl

Na+Cl-

Silver Nitrate Solution Sodium Chloride Solution

When mixed…

Ag+

NO3-

Na+

Cl-

BUT! The silver ions and chloride

ions come together to form

an insoluble product (AgCl).

…the sodium and nitrate ions can’t form an insoluble product (NaNO3

can dissolve), so they stay in

solution. The silver chloride “falls

out” of solution and precipitates

as a solid.

Ag Cl

A Particularly Interesting Example

• (NH4)2SO4 (aq) + Ca(OH)2 (aq) ?

• You would probably say:

– (NH4)2SO4 (aq) + Ca(OH)2 (aq) CaSO4 (aq) + NH4OH (?)

• According to your solubility rules chart, that last phase of matter should be (aq), and therefore that reaction shouldn’t happen.

• BUT! Remember that NH4OH decomposes to NH3

and H2O, both of which become the driving forces:

– (NH4)2SO4 + Ca(OH)2 CaSO4 (aq) + NH3 (aq) + H2O (aq)

Closure: Net Ionic Equations

• Now let’s practice:

– Molecular, Complete, Net Ionic Equations Worksheet

• Any 3. Two must be from #5 on.

• #7/#11 feature decomposition on the product side.

• #9 is weird and worth a try.

– Driving Forces worksheet

• Any 3.

• #2 features decomposition on the product side.

– Lab – Double Replacements

• All yo’ reactions.