Today is Friday (!), April 17 th, 2015 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 Reactions Types of Chemical Reactions Where is this in my book? – P. 321 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 + 3O 2  2Fe 2 O 3

Just checking… Identify the reactant(s): – Na + Cl  NaCl Na, Cl Identify the reactant(s): – 2H 2 O  2H 2 + O 2 H 2 O Identify the product(s): – Na + Cl  NaCl NaCl Identify the product(s): – 2H 2 O  2H 2 + O 2 H 2 + O 2

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 H 2 O 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 (Br 2 ) – Iodine (I 2 ) – Nitrogen (N 2 ) – Chlorine (Cl 2 ) – Hydrogen (H 2 ) – Oxygen (O 2 ) – Fluorine (F 2 ) 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]CO 3  [metal]oxide + CO 2 – [metal]OH  [metal]oxide + H 2 O – [metal]ClO 3  [metal]chloride + O 2 – Acids  [nonmetal]oxide + H 2 O

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) + F 2 (g)  2NaF (s) + Cl 2 (g) – MgCl 2 (s) + Br 2 (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 + MgSO 4  Mg + CuSO 4 – No reaction (Magnesium is above Copper). Pb + ZnSO 4  Zn + PbSO 4 – No reaction (Zinc is above Lead). Fe + 2AgNO 3  2Ag + Fe(NO 3 ) 2 – Reaction (Iron is above Silver). 2Al + 3H 2 O  Al 2 O 3 + 3H 2 – 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 + H 2 O  KOH + H 2 – Mg + H 2 O  Mg(OH) 2 + H 2 In other words, they form hydroxides, not oxides.

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.”

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 compounds must be insoluble!

Reminder: Dissociation Ca Cl Ca 2+ Cl - 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 table. – 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 NiNO 3 (aq) + KBr (aq) → KNO 3 (?) + NiBr (?) In what states are potassium nitrate and nickel (I) bromide? According to your solubility table: – 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. NiNO 3 (aq) + KBr (aq) → KNO 3 (aq) + NiBr (aq) – So no reaction, since both of them are soluble. Remember, one must be insoluble!

Solubility Example Ba(NO 3 ) 2 (aq) + (NH 4 ) 3 PO 4 (aq) → NH 4 NO 3 (?) + Ba 3 (PO 4 ) 2 (?) In what states are ammonium nitrate and barium phosphate? According to the solubility table, – 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(NO 3 ) 2 (aq) + (NH 4 ) 3 PO 4 (aq) → NH 4 NO 3 (aq) + Ba 3 (PO 4 ) 2 (s) – One compound is insoluble, so there will be a reaction.

Precipitates In the previous example, Ba 3 (PO 4 ) 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(NO 3 ) 2 (aq)  KNO 3 (?) + PbI 2 (?) Check your solubility tables 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(NO 3 ) 2 (aq)  KNO 3 (aq) + PbI 2 (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, 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: – C 3 H 8 (g) + 5O 2 (g)  3CO 2 (g) + 4H 2 O (g) – P 4 (s) + 5O 2 (g)  P 4 O 10 (s) This is a combustion and synthesis reaction!

Combustion Reaction Demo C 2 H 5 OH + 3O 2  2CO 2 + 3H 2 O

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.

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.

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 + ___O 2  ___Al 2 O 3 Al O

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: – 2Na + 2Cl  2NaClshould really be – Na + Cl  NaCl Even if it’s balanced, it has to be reduced to the lowest ratio.

Balancing Synthesis Reactions 1 of 3 1.___CaO + ___H 2 O  ___Ca(OH) 2 – CaO + H 2 O  Ca(OH) 2 2.___P 4 + ___O 2  ___P 2 O 5 – P 4 + 5O 2  2P 2 O 5 3.___Ca + ___O 2  ___CaO – 2Ca + O 2  2CaO 4.___Cu + ___S 8  ___ CuS – 8Cu + S 8  8CuS

Balancing Synthesis Reactions 2 of 3 5.___S 8 + ___O 2  ___SO 2 – S 8 + 8O 2  8SO 2 6.___H 2 + ___N 2  ___NH 3 – 3H 2 + N 2  2NH 3 7.___H 2 + ___Cl 2  ___HCl – H 2 + Cl 2  2HCl 8.___Ag + ___S 8  ___Ag 2 S – 16Ag + S 8  8Ag 2 S

Balancing Synthesis Reactions 3 of 3 9.___Cr + ___O 2  ___Cr 2 O 3 – 4Cr + 3O 2  2Cr 2 O 3 10.___Al + ___Br 2  ___AlBr 3 – 2Al + 3Br 2  2AlBr 3 11.___Na + ___I 2  ___NaI – 2Na + I 2  2NaI 12.___H 2 + ___O 2  ___H 2 O – 2H 2 + O 2  2H 2 O

Balancing Decomposition Reactions 1 of 3 1.___BaCO 3  ___BaO + ___CO 2 – BaCO 3  BaO + CO 2 2.___MgCO 3  ___MgO + ___CO 2 – MgCO 3  MgO + CO 2 3.___K 2 CO 3  ___K 2 O + ___CO 2 – K 2 CO 3  K 2 O + CO 2 4.___Zn(OH) 2  ___ZnO + ___H 2 O – Zn(OH) 2  ZnO + H 2 O

Balancing Decomposition Reactions 2 of 3 5.___Fe(OH) 2  ___FeO + ___H 2 O – Fe(OH) 2  FeO + H 2 O 6.___Ni(ClO 3 ) 2  ___NiCl 2 + __O 2 – Ni(ClO 3 ) 2  NiCl 2 + 3O 2 7.___NaClO 3  ___NaCl + ___O 2 – 2NaClO 3  2NaCl + 3O 2 8.___KClO 3  ___KCl + ___O 2 – 2KClO 3  2KCl + 3O 2

Balancing Decomposition Reactions 3 of 3 9.___H 2 SO 4  ___H 2 O + ___SO 3 – H 2 SO 4  H 2 O + SO 3 10.___H 2 CO 3  ___H 2 O + ___CO 2 – H 2 CO 3  H 2 O + CO 2 11.___Al 2 O 3  ___Al + ___O 2 – 2Al 2 O 3  4Al + 3O 2 12.___Ag 2 O  ___Ag + ___O 2 – 2Ag 2 O  4Ag + O 2

Balancing Single Replacement Reactions 1 of 2 1.___AgNO 3 + ___Ni  ___Ni(NO 3 ) 2 + ___Ag – 2AgNO 3 + Ni  Ni(NO 3 ) 2 + 2Ag 2.___AlBr 3 + ___Cl 2  ___AlCl 3 + ___Br 2 – 2AlBr 3 + 3Cl 2  2AlCl 3 + 3Br 2 3.___NaI + ___Br 2  ___NaBr + ___I 2 – 2NaI + Br 2  2NaBr + I 2 4.___Ca + ___HCl  ___CaCl 2 + ___H 2 – Ca + 2HCl  CaCl 2 + H 2

Balancing Single Replacement Reactions 2 of 2 5.___Mg + ___HNO 3  ___Mg(NO 3 ) 2 + ___H 2 – Mg + 2HNO 3  Mg(NO 3 ) 2 + H 2 6.___ Zn + ___H 2 SO 4  ___ZnSO 4 + ___H 2 – Zn + H 2 SO 4  ZnSO 4 + H 2 7.___K + ___H 2 O  ___KOH + ___H 2 – 2K + 2H 2 O  2KOH + H 2 8.___Na + ___H 2 O  ___NaOH + ___H 2 – 2Na + 2H 2 O  2NaOH + H 2

Balancing Double Replacement Reactions 1 of 3 1.___AlI 3 + ___HgCl 2  ____AlCl 3 + ____HgI 2 – 2AlI 3 + 3HgCl 2  2AlCl 3 + 3HgI 2 (s) 2.___HCl + ___NaOH  ___NaCl + ___H 2 O – HCl + NaOH  NaCl + H 2 O 3.___BaCl 2 + ___H 2 SO 4  ___BaSO 4 + ___HCl – BaCl 2 + H 2 SO 4  BaSO 4 + 2HCl 4.___Al 2 (SO 4 ) 3 + ___Ca(OH) 2  ___Al(OH) 3 + ___CaSO 4 – Al 2 (SO 4 ) 3 + 3Ca(OH) 2  2Al(OH) 3 + 3CaSO 4

Balancing Double Replacement Reactions 2 of 3 5.___AgNO 3 + ___K 3 PO 4  ___Ag 3 PO 4 + ___KNO 3 – 3AgNO 3 + K 3 PO 4  Ag 3 PO 4 + 3KNO 3 6.___CuBr 2 + ___AlCl 3  ___CuCl 2 + ___AlBr 3 – 3CuBr 2 + 2AlCl 3  3CuCl 2 + 2AlBr 3 7.___Ca(C 2 H 3 O 2 ) 2 + ___Na 2 CO 3  ___CaCO 3 + __NaC 2 H 3 O 2 – Ca(C 2 H 3 O 2 ) 2 + Na 2 CO 3  CaCO 3 + 2NaC 2 H 3 O 2 8.___NH 4 Cl + ___Hg 2 (C 2 H 3 O 2 ) 2  __NH 4 C 2 H 3 O 2 + ___Hg 2 Cl 2 – 2NH 4 Cl + Hg 2 (C 2 H 3 O 2 ) 2  2NH 4 C 2 H 3 O 2 + Hg 2 Cl 2

Balancing Double Replacement Reactions 3 of 3 9.___Ca(NO 3 ) 2 + ___HCl  ___CaCl 2 + ___HNO 3 – Ca(NO 3 ) 2 + 2HCl  CaCl 2 + 2HNO 3 10.___FeS + ___HCl  ___FeCl 2 + ___H 2 S – FeS + 2HCl  FeCl 2 + H 2 S 11.___Cu(OH) 2 + ___HC 2 H 3 O 2  ___Cu(C 2 H 3 O 2 ) 2 + ___H 2 O – Cu(OH) 2 + 2HC 2 H 3 O 2  Cu(C 2 H 3 O 2 ) 2 + 2H 2 O 12.___Ca(OH) 2 + ___H 3 PO 4  ___Ca 3 (PO 4 ) 2 + ___H 2 – 3Ca(OH) 2 + 2H 3 PO 4  Ca 3 (PO 4 ) 2 + 6H 2

Balancing Combustion Reactions 1 of 3 1.___CH 4 + ___O 2  ___CO 2 + ___H 2 O – CH 4 + 2O 2  CO 2 + 2H 2 O 2.___C 2 H 6 + ___O 2  ___CO 2 + ___H 2 O – 2C 2 H 6 + 7O 2  4CO 2 + 6H 2 O 3.___C 3 H 8 + ___O 2  ___CO 2 + ___H 2 O – C 3 H 8 + 5O 2  3CO 2 + 4H 2 O

Balancing Combustion Reactions 2 of 3 4.___C 4 H 10 + ___O 2  ___CO 2 + ___H 2 O – 2C 4 H O 2  8CO H 2 O 5.___C 5 H 12 + ___O 2  ___CO 2 + ___H 2 O – C 5 H O 2  5CO 2 + 6H 2 O 6.___C 6 H 14 + ___O 2  ___CO 2 + ___H 2 O – 2C 6 H O 2  12CO H 2 O

Balancing Combustion Reactions 3 of 3 7.___C 2 H 4 + ___O 2  ___CO 2 + ___H 2 O – C 2 H 4 + 3O 2  2CO 2 + 2H 2 O 8.___C 2 H 2 + ___O 2  ___CO 2 + ___H 2 O – 2C 2 H 2 + 5O 2  4CO 2 + 2H 2 O 9.___C 6 H 6 + ___O 2  ___CO 2 + ___H 2 O – 2C 6 H O 2  12CO 2 + 6H 2 O

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 (CH 4 ) 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: CH 4 + O 2  CO 2 + H 2 O – Balanced: CH 4 + 2O 2  CO 2 + 2H 2 O

Predicting Products [reminder] Typical reaction processes: – Metal carbonates break down to metal oxides and CO 2. – Metal hydroxides break down to metal oxides and H 2 O. – Metal chlorates break down to metal chlorides and O 2. – Acids break down to [nonmetal] oxides and H 2 O. – 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 (H 2 CO 3 ) makes H 2 O + CO 2 – Ammonium Hydroxide (NH 4 OH) makes H 2 O + NH 3 – Sulfurous Acid (H 2 SO 3 ) makes H 2 O + SO 2 Notice how you can see the H 2 O come out of the initial product: – H 2 CO 3  H 2 O + CO 2

Keep an eye out… Also, know this: – 2 H 3 PO 4  3 H 2 O + P 2 O 5 – 2 H 3 PO 3  3 H 2 O + P 2 O 3 – 2 HNO 3  H 2 O + N 2 O 5 – 2 HNO 2  H 2 O + N 2 O 3 Sorry guys, it’s just how it is.

Predicting Products Special Practice special practice is special H 2 CO 3  – H 2 CO 3  H 2 O + CO 2 Ca + H 2 O  – 2Ca + 2H 2 O  2CaOH + H 2 KClO 3  – 2KClO 3  2KCl + 3O 2 HNO 3  – 2HNO 3  N 2 O 5 + H 2 O

Predicting Products Once you can spot the type of the reaction, you can easily predict the products. Here’s a website to help: – actice_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 cations with cations (written first) and anions with anions (written second).

Reaction Type Summary  + Double Replacement  + Single Replacement  Decomposition  + Synthesis Combustion  + O2O2

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.

Reminder: Dissociation Ca Cl Ca 2+ Cl - 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 reactions and 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: – CuSO 4 in solution becomes Cu 2+ and SO – NaCl in solution becomes Na + and Cl -. This allows them to conduct electricity.

Dissociation Practice Into what does H 2 SO 4 dissociate? What’s the compound made of, and how many of each? – 2H and 1SO 4 Write their charges in: – 2H + and 1SO 4 2- 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. – Na 2 SO 4 + BaCl 2  What are the products (double rep. reaction)? – Na 2 SO 4 + BaCl 2  NaCl + BaSO 4 That’s the skeleton equation. Balanced? – Na 2 SO 4 + BaCl 2  2NaCl + BaSO 4 States of matter (use solubility table)? – Na 2 SO 4 (aq) + BaCl 2 (aq)  2NaCl (aq) + BaSO 4 (s)

Na 2 SO 4 (aq) + BaCl 2 (aq)  2NaCl (aq) + BaSO 4 (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 + + SO Ba Cl -  2Na + + 2Cl - + BaSO 4 (s) – That right there is the complete ionic equation.

2Na + + SO Ba Cl -  2Na + + 2Cl - + BaSO 4 (s) 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, SO 4 2- and Ba 2+ are the driving forces behind the reaction. More on this in a little bit…

2Na + + SO Ba Cl -  2Na + + 2Cl - + BaSO 4 (s) So let’s try something. Let’s eliminate the spectator ions from the equation, since they really didn’t actively participate anyway. – Ba 2+ + SO 4 2-  BaSO 4 (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.

Net Ionic Equations Now let’s practice: – Molecular, Complete, Net Ionic Equations Worksheet Any 3. Two must be from #5 on. – Driving Forces worksheet Any 3. – Lab – Double Replacements All yo’ reactions.

Closure ml – Also linked on my Chemistry Links page. You may try 5, 10, or 15 questions (let’s go with 15 for today), and you may set the difficulty at easy, intermediate, or hard. Your goal is to do 15 easy questions, 15 intermediate questions, and 15 hard questions. – Let me know when you’re done.