Types of Reactions Chemistry

Learning Objective



TLW identify and quantify changes that occur during chemical reactions (TEKS 8)

Agenda



Review of Basics



Define Types of Chemical Reactions and Real-World Examples



Examples of the Types of Chemical Reactions



Group Practice



Individual Practice



Labs

Chemical Reactions – General Stuff 

Reactants products are added together to make



Reactants are found on left side of chemical equation and products are on the right side



Compounds are formed – substances made up of two or more elements that chemically combined (not easily separated or impossible to return to original element)

Chemical Reactions – General Stuff  

Valence electrons and periodic trends determine how various atoms of elements react (or don’t react) Catalysts are “outside agents” that may be added to speed up reactions or cause reactions to perform differently under different conditions (such as varying temperatures).

–

They are not used up in the reaction

Chemical Reactions – General Stuff 

Don’t forget about the laws of conservation of mass and energy … neither can be created nor destroyed

–

Reactants are transformed

–

Products are formed

–

Chemical energy converts to heat, light, electricity, sound

Chemical Reactions – General Stuff 

Also – Law of Definite Proportions , which states in any sample of a chemical compound the elements are always combined in the same

proportion by mass

H 2 O H = 1 g x 2 = 2 g O = 16 g Total 18 g H = 2 g / 18 g = 11% O = 16 g / 18 g = 89%

Chemical Reactions – General Stuff 

From Dalton’s Laws (remember him?)



Law of Multiple Proportions states whenever two elements form more than one compound different masses of one element that combine with the same mass of the other element are in the ratio of small whole number H 2 O H 2 O 2

Chemical Reactions – General Stuff  

Law of Definite Proportions

– describes composition of one compound

Law of Multiple Proportions

– compares composition of two different compounds containing same elements

Types of Reactions

        

Combination (a.k.a. Synthesis, Addition) Decomposition Single Replacement (a.k.a. Single Displacement) Double Replacement (a.k.a. Double

Displacement)

Combustion Neutralization Precipitation Reduction/Oxidation Reactions (Redox) Energy Producing – Exothermic, Endothermic, Light

Combination (aka Synthesis, Addition)

 

Two or more elements or substances combine to form a new compound. A + B



AB



Examples –

–

Formation of rust

–

Air pollutant sulfur dioxide

– –

Polymerization (plastics) Photosynthesis (plants)

Combination



A + B



AB



Where A and B are elements and AB is a compound



Note that only one compound exists on the RIGHT SIDE…

Combination Examples

  

4Fe(s) + 3O 2 (g)



S(s) + O 2 (g)



SO x 2Fe (g) 2 O 3 (s) 2Na + Cl 2



2NaCl



Teacher Demo – need a better one…..

Decomposition

 

A single compound is broken down to produce two or more smaller compounds and/or elements.

AB



A + B



Example –

–

Water with electricity into hydrogen and oxygen (electrolysis)

–

Baking soda with heat

Decomposition



AB



A + B



Where AB is a compound and A & B are elements or other compounds



Note that only one compound exists on the LEFT SIDE…

Decomposition



Examples:



2H 2 O (l)



2H 2(g) + O 2(g)



NaHCO 3(s) heat



2H 2(g) + NaCO 3



Teacher demo – baking soda and heat

Single Replacement (Single Displacement)

   

One element replaces (displaces) a similar element in a compound Produces heat (is exothermic) A + BC



AC + B Example –

–

If you place an iron nail into a beaker of copper (II) chloride you will begin to see reddish copper forming on the iron.

–

Iron replaces (displaces) copper in the solution and the copper falls out of solution as a metal

Single Replacement

  

A + BC



AC + B Where A and B are elements and BC and AC are compounds Can have more than 2 reactants and/or products



Example: 2HCl (l) Fe (s) + Zn (s) + CuCl 2(l)

 

ZnCl 2(l) Cu (s) + H + FeCl 2(g) 2(l)

Single Replacement



All single replacement reactions are

exothermic



They give off heat and occur rapidly



Teacher demo – Zinc plus hydrochloric acid or iron nail in copper(II)chloride

Double Replacement (Double Displacement)

    

Ions from two compounds in solution exchange to produce two new compounds AB + CD



AD + CB One compound usually forms a precipitate that settles out of the solution, a gas that bubbles out, or a molecular compound like water The other compound formed often remains dissolved in the solution Examples –

–

Baking soda and vinegar

–

Dried fruit

Double Replacement

    

AB + CD



AD + CB Where AB, CD, AD, & CB are all compounds Can have more than 2 reactants and/or products Examples: 2HCl + 2NaOH



2NaCl + 2H 2 O



Na 2 SO 3(aq) + 2HCl (aq)



2NaCl (aq) + H 2 O (l) + SO 2(g)



Teacher demo – the ever popular baking soda and vinegar or cleaning pennies with vinegar using salt as a catalyst

Combustion

Carbon substances combine with oxygen, releasing large amounts of energy, in the form of heat, light, etc. Carbon dioxide and water are also typical products C x H y + O 2



CO 2 + H 2 O Examples –

–

Natural gas to heat a house

–

Hydrogen powered cars

General Formula

C x H y + O 2



CO 2 + H 2 O

hydrocarbon oxygen carbon water dioxide

Examples CH 4 + 2O 2



CO 2 + 2H 2 O 2H 2 (g) + O 2 (g)



2H 2 O(l)



Teacher demo – lighting a match

Neutralization

 

Double Replacement Reaction where wn acid and a base react to form water and a salt General formula HA + BOH



H 2 O + BA

acid base water salt

Examples – HCl(aq) + NaOH(s)



H 2 O(l) + NaCl(s) H 2 SO 4 (aq) + Ca(OH) 2 (aq)



H 2 O(l) + CaSO 4 (s)

Precipitation

 

When solutions are saturated, adding additional ions will cause a precipitate to form Solid – usually sinks to bottom, but can float



Solubility product (K sp ) can be used to predict formation of precipitates

–

If ion-product concentration > K sp then a precipitate will form

–

If ion-product concentration < K sp then a precipitate will not form



Teacher demo – baking soda and vinegar… again (how boring)



Practice Worksheet – Identify 5 basic reactions

Reduction/Oxidation

Discovery Video - Electrochemistry Reduction/Oxidation Reactions A. Reaction in which

electrons

transferred are B. Commonly called the

redox

reaction C. One element is electrons

reduced

– it

gains

One element is electrons

oxidized

– it

loses

Redox reactions have

radicals

fragments of molecules with at least one electron for

bonding

Ex. of radicals = Styrofoam

Example of redox reaction

Rust

= iron reacts with oxygen

Fe 2 O 3

Fe loses 3 electrons O gains 2 electrons

Rules for Assigning Oxidation Numbers

Rule 1 The oxidation number of any uncombined element is 0

Example: The oxidation number of Na (s) is 0.

Rule 2 The oxidation number of a monatomic ion equals the charge on the ion.

Example: The oxidation number of Cl

-

is -1.

Rule 3

The more electronegative element in a binary compound is assigned the number equal to the charge it would have if it were an ion.

Example: The oxidation number of O in NO is -2.

Rule 4

The oxidation number of fluorine in a compound is always -1.

Example: The oxidation number of F in LiF is -1.

Rule 5

Oxygen has an oxidation number of -2 unless it is combined with F, when it is +2, or it is in a peroxide, such as H 2 O 2 , when it is -1.

Example: The oxidation number of O in NO

2

is -2.

Rule 6

The oxidation state of hydrogen in most of its compounds is +1 unless it is combined with a metal, in which case it is -1.

Example: The oxidation number of H in LiH is -1.

Rule 7

In compounds, Group 1 and 2 elements and aluminum have oxidation numbers of +1, +2, and +3, respectively.

Example: The oxidation number of Ca in CaCO 3 is +2.

Rule 8 The sum of the oxidation numbers of all atoms in a neutral compound is 0.

Example: The oxidation number of C in CCO 3 is +4.

Rule 9

The sum of the oxidation numbers of all atoms in a polyatomic ion equals the charge of the ion.

Example: The oxidation number of P in H 2 PO 4 is +5.

Exothermic Reactions

 

In many reactions, less energy is required to break the bonds in the reactants than is released when bonds form to make new products In these reactions some type of heat or light is released and they are called exothermic



Exothermic reactions can be detected by a rise in temperature

Exothermic Reactions



What are examples of exothermic reactions you are familiar with?



Demonstration

Endothermic Reactions



Sometimes more energy is required to break bonds in the reactants than is released to form new products



The are called endothermic reactions



You can detect these reactions by a decrease in temperature

Endothermic Reactions



What are examples of endothermic reactions you are familiar with?



Demonstration

Group Practice ~ Name Those Reactions

A Group Activity



Types of Chemical Reactions – Sorting Mats link

Individual Practice



Identifying Types of Reactions Worksheet link



Crossword Puzzle

Looking Ahead   Labs – – Conservation of Mass – Types of Chemical Reactions – Empirical Formula Determination – Predicting the Amount of Product in a Reaction – Identifying Relationships between Reactants and Products in a Reaction – Predicting the Products of a Reaction – Precipitation Reactions – Energetic Reactions – exothermic and endothermic experiments – Qualitative and Quantitative Analysis More on calculating definite proportions and multiple proportions