Chapter 11 Chemical Reactions

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Transcript Chapter 11 Chemical Reactions

1 Chapter 11

“Chemical Reactions”

Section 11.1

p. 321 Describing Chemical Reactions

 All chemical reactions… have two parts:

1. Reactants

start with = substances you  

2. Products

= end with reactants turn into products Reactants  Products 3

Products Reactants - Page 321

In a chem rxn  Atoms not created or destroyed (Law of Conservation of Mass)  rxn described in a: #1.

sentence

every item is a word Copper reacts with chlorine to form copper (II) chloride.

#2.

word equation

Copper + chlorine  some symbols used copper (II) chloride 5

Symbols in equations? – Text page 323  arrow ( →) separates reactants from products (points

products)  –Read as: “reacts to form” or yields plus sign = “and”  (s) after formula = solid:  (g) = gas:  (l) = liquid: Fe

(s)

2(g)

(l)

Symbols used in equations  (aq) after formula = dissolved in water, aqueous solution: NaCl

(aq)

is salt water solution   used after product - indicates gas

produced

: H

2 ↑

 used after product - indicates solid

produced

: PbI

2 ↓

Symbols used in equations ■

double arrow indicates a reversible reaction (more later)

■   

, heat supplied to rxn shows that

■

indicates catalyst supplied

(here, platinum is catalyst) 8

What is a catalyst?

 substance that speeds up rxn, w/o being changed or used up in rxn  Enzymes - biological or protein catalysts in your body 9

#3. The Skeleton Equation  Uses formulas and symbols to describe rxn –but doesn’t indicate how many; means they’re NOT balanced  All chem equations are

description

of rxn 10

Write a skeleton equation for: Solid iron (III) sulfide reacts with gaseous hydrogen chloride to form iron (III) chloride and hydrogen sulfide gas.

Nitric acid dissolved in water reacts with solid sodium carbonate to form liquid water and carbon dioxide gas and sodium nitrate dissolved in water.

Now, read these equations:

Fe (s) + O 2(g)



Fe 2 O 3(s) Cu (s) + AgNO 3(aq)



Ag (s) + Cu(NO 3 ) 2(aq) NO 2(g) N 2(g) + O 2(g)

#4. Balanced Chemical Equations  Atoms can’t be created or destroyed in an ordinary reaction: –All atoms we

start with

we must

end up with (balanced!)

 balanced equation has same # of each element

on both sides

of equation 13

1) 2) 3) 4) Rules for balancing: Assemble

correct formulas

for all reactants and products, using “+” and “→” Count # of atoms of each type on both sides Balance elements

one at a time coefficients

by adding (the numbers in front) where you need more until LAST! save balancing the H and O (hint: I prefer to save O until the

very

last) Double-Check make sure balanced 14



Never

change coefficients ) subscript (only change – changing subscript (formula) describes different chemical – H 2 O different than H 2 O 2 

Never

put coefficient in only in front

middle

of formula; 2 NaCl is okay, but Na 2 Cl is not.

Practice Balancing Examples 

_AgNO 3 + _Cu



_Cu(NO 3 ) 2



_Mg + _N 2



_Mg 3 N 2



_P + _O 2



_P 4 O 10



_Na + _H 2 O



_H 2



_CH 4 + _O 2



_CO 2 + _H 2 O

Balancing Equations 17 Balancing Chemical Reactions Mark Rosengarten – 8:21

Section 11.2

p. 330

Types of Chemical Reactions

Types of Reactions 

5 major types



predict the products

 predict whether or not they will

happen at all



How?

We recognize them by their

reactants

#1 - Combination Reactions  Combine = put together  2 substances combine to make one cmpd (also called “synthesis”)   Ca + O 2 SO 3  CaO + H 2 O  H 2 SO 4  predict products, especially if

reactants are 2 elements

 Mg + N 2 20 (symbols, charges, cross)

   Complete and balance: Ca + Cl 2 Fe + O 2 Al + O 2    (assume iron (II) oxide is the product)  Remember

first step

…write

correct formulas –

you can still change

subscripts

balancing!

at this point, but not while 

Then balance

21 coefficients only by changing just

#1 – Combination Reactions  Additional Notes: a) Some nonmetal oxides

with H 2 O

- produces acid: SO 2 + H 2 O  (how “ acid rain ” forms) H 2 SO 3

react

b) Some metallic oxides

react with H 2 O

- produces base: CaO + H 2 O  Ca(OH) 2 22

#2 - Decomposition Reactions  decompose = fall apart 

one reactant

breaks apart into 2 or more elements or cmpds  NaCl Na + Cl 2  CaCO 3 CaO + CO 2  Note:

energy

(heat, sunlight, electricity, etc.) usually required 23

#2 - Decomposition Reactions  predict products if binary cmpd (made of 2 elements) –It breaks apart into the elements:  H 2 O  HgO 24

#2 - Decomposition Reactions  If cmpd has > 2 elements you must be given one of products –other product from the missing pieces   NiCO 3 CO 2 + ___ H 2 CO 3 (aq)  CO 2 + ___

heat

#3 - Single Replacement Reactions  One element replaces another (new dance partner)  Reactants must be

an element & cmpd

 Products will be a different element  26 and different cmpd  Na + KCl  F 2 + LiCl  K + NaCl LiF + Cl 2 (Cations switched) (Anions switched)

#3 Single Replacement Reactions  Metals replace other metals (they can also replace H)  K + AlN   Zn + HCl   Think of water as: HOH –Metals replace first H, then combines w/ hydroxide (OH).

 Na + HOH  27

#3 Single Replacement Reactions  can even tell

whether or not

replacement rxn will happen: single –b/c some chemicals more “active” others –More active

replaces

less active than  list – p. 333 Metals Activity Series of  Higher on list replaces lower 28

Higher activity Lower activity

The “Activity Series” of Metals

Lithium Potassium Calcium Sodium Magnesium Aluminum Zinc Chromium Iron Nickel Lead Hydrogen Bismuth Copper Mercury Silver Platinum Gold 1) Metals can replace other metals, if they are above metal trying to replace (i.e. Zn will replace Pb) 2) Metals above H can replace H in acids.

3) Metals from Na upward can replace hydrogen in H 2 O

The “Activity Series” of Halogens Higher Activity

Fluorine Chlorine Bromine Iodine

Lower Activity Halogens can replace other halogens in compounds, if they are above halogen they are replacing

2NaCl (s) + F 2(g)



MgCl 2(s) + Br 2(g)



2NaF (s) + Cl 2(g)

#3 Single Replacement Reactions  Practice:

Fe + CuSO

 

Pb + KCl

 

Al + HCl

 31

#4 - Double Replacement Reactions 

Two things replace each other.

–

Reactants must be two ionic

compounds

, in aqueous solution

 

NaOH + FeCl 3

 –

positive ions change place (dance partners) NaOH + FeCl 3



Fe +3 OH + Na +1 Cl -1 = NaOH + FeCl 3



Fe( OH) 3 + Na Cl

#4 - Double Replacement Reactions  Have certain “

driving forces

” , or reasons – only happens if one product: a) doesn’t dissolve in water & forms solid (a “ precipitate ”), or b) is gas that bubbles out, or c) is molecular compound (usually water) 33

Complete and balance:  assume all of the following reactions actually take place:

CaCl 2 CuCl 2 + NaOH + K 2 S

 

KOH + Fe(NO 3 ) 3



(NH 4 ) 2 SO 4 + BaF 2

 34

How to recognize which type?



Look at the reactants : E + E = Combination C = Decomposition E + C = Single replacement C + C = Double replacement

36 Practice Examples:  

H 2 + O 2 H 2 O

   

Zn + H 2 SO 4 HgO

    

KBr + Cl 2 AgNO 3



+ NaCl



Mg(OH) 2 + H 2 SO 3



 #5 – Combustion Reactions Combustion means “

add oxygen

”  Normally, a cmpd composed of only C, H, (and maybe O) is reacted with oxygen – called “ burning ”  Complete combustion, products are CO 2 and H 2 O  If incomplete, products are CO (or possibly just C) and H 2 O 37

Combustion Reaction Examples: 

C 4 H 10 + O 2

 

C 4 H 10 + O 2



(assume complete) (incomplete)



C 6 H 12 O 6 + O 2



(complete)



C 8 H 8 + O 2



(incomplete)

SUMMARY: An equation...

 Describes a rxn  Must be balanced (follows the Law of Conservation of Mass)  only balance by changing coefficients 

special symbols

to indicate physical state, catalyst or energy required, etc. 39

 5 major types Reactions  We can tell what type they are by looking at reactants  Single Replacement happens based on the

Activity Series

 Double Replacement happens if one product is: 1) a precipitate (an insoluble solid), 2) water (a molecular compound) , or 3) a gas 40

Section 11.3

p. 342 Reactions in Aqueous Solution Co(NO 3 ) 2 NiCl 2 CuSO 4 KMnO 4 K 2 Cr 2 O 7 K 2 CrO 4

Net Ionic Equations  Many reactions occur in water- that is, in

aqueous solution

 When dissolved in water, many ionic cmpds “dissociate”, or separate, into cations & anions  Now write ionic equation 42

Net Ionic Equations  Example (needs to be a double replacement reaction) AgNO 3 + NaCl  AgCl + NaNO 3 1. this is the full balanced equation 2. next, write it as ionic equation by splitting the cmpds into their ions: Ag

+ NO 3

+ Na

+ Cl



AgCl

+ Na

+ NO 3

Note that the AgCl did not ionize, because it is a

“precipitate” (Table 11.3 p. 344)

Net Ionic Equations 3.

simplify

by crossing out ions not directly involved (called spectator ions)

Ag 1+ + Cl 1-



AgCl

This is called the

net ionic equation

Let’s talk about precipitates before we do some other examples 44

  Predicting the Precipitate Insoluble salt is a precipitate [note Figure 11.11, p.342 (AgCl)] General solubility rules are found: a) Table 11.3, p. 344 in textbook b) Reference section - page R54 (Table B.9) 45

Let’s do some examples together of net ionic equations, starting with these reactants:

BaCl

+ AgNO

→ 47

NaCl + Ba(NO

)

→ 48

Pb(NO 3 ) 2(aq) + H 2 SO 4(aq)  49