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Ionic Chemical Reactions. Lecture 12 (Ch. 10) 10/1/12 HW: 1, 9, 35, 39, 41, 45, 51, 61. Introduction. To date, we have learned about: Chemical elements and their classifications Valence electron configurations and their effect on chemical properties and reactivity Chemical bonding

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Ionic chemical reactions

Ionic Chemical Reactions

Lecture 12 (Ch. 10)

10/1/12

HW: 1, 9, 35, 39, 41, 45, 51, 61


Introduction
Introduction

  • To date, we have learned about:

    • Chemical elements and their classifications

    • Valence electron configurations and their effect on chemical properties and reactivity

    • Chemical bonding

    • Molecular geometries

  • Now, we will begin to learn about chemical reactions involving ionic compounds


Ionic compounds recap
Ionic Compounds (recap)

Na+

Electrostatic interactions between NaCl molecules holds them together in a lattice, which is why all ionic compounds are solid at room temperature.

Cl-


Dissociation of ionic compounds in water
Dissociation of Ionic Compounds in Water

NaCl (s) NaCl(aq)

H2O (L)

  • An ionic compound fully dissolves in water to form an aqueous solution

  • The compound will split into cations and anions.

water

NaCl (s)

NaCl (aq) Na+(aq) + Cl-(aq)

=


Polyatomic ions
Polyatomic Ions

  • Polyatomic ions are covalent molecules with an overall charge. These molecules behave as normal ions.

  • Polyatomic ions do NOT break apart in water

  • Ex. Al(PO4) (s) Al3+ (aq) + PO43-(aq)

H2O (L)

Aluminum Phosphate

Aluminum

cation

Phosphate anion


You DEFINITELY want to know these (Polyatomic Ions)

2-

1-

3-

1+


Types of reactions
Types of Reactions

  • Chemical reactions involving ionic compounds can be classified as one of the following:

  • combination reactions

  • decomposition reactions

  • single replacement reactions

  • double replacement reactions


Combination Reactions

  • In a combination reaction, multiple reactants combine to form a single product

    • The reaction may occur between two elements

    • Or between an element and a compound

    • Or between two compounds


Combination reaction
Combination Reaction

(E+E) 3Li(s) + P(g) Li3P(s)

(E+C) 2Na(s) + Cl2(g) 2NaCl(s)

(C+C) SO3(g) + H2O(l) H2SO4(aq)


Examples
Examples

  • Predict the products of the following combination reactions. Also, predict the phase of each reactant and product.

  • Hints: You will form ionic compounds. Also, pay attention to possible polyatomic ions.

  • Li (s) + ½O2 (g)

    MgO (s) + CO2 (g)

2

Li2O (s) Lithium oxide

MgCO3 (s) Magnesium Carbonate


Decomposition reaction
Decomposition Reaction

In a decomposition reaction,

  • one substance splits into two or more simpler substances

    2HgO(s) 2Hg(l) + O2(g)

    2KClO3(s) 2KCl(s) + 3O2(g)


Single replacement reaction
Single Replacement Reaction

In a single replacement reaction,

  • An element reacts with a salt, and two elements switch places. When one metal replaces another in an ionic compound, this is also called a transmetallation reaction.

    Zn(s) + 2AgCl (aq) ZnCl2(aq)+ 2Ag(s)

    Fe(s) + CuSO4(aq) FeSO4(aq) + Cu(s)


Transmetallations
Transmetallations

  • Transmetallations occur because one metal is more active (less stable) than the other.

  • In the reaction below, Zn displaces Ag because Zn is more active :

    Zn(s) + 2AgCl (aq) ZnCl2(aq)+ 2Ag(s)

  • A metal of greater activity will displace a less active metal. The opposite will NOT occur. An activity series is provided on pg. 325 of the text.


Group examples
Group Examples

Activity Series

Predict the products. Include phase. Balance if necessary.

Li (s) + Ca(ClO4)2 (aq)

Na (s) + ZnSO4(aq)

K (s) + LiCl (aq)


Single replacement reactions involving metals and strong acids
Single Replacement Reactions involving Metals and Strong Acids

  • The acids above are known as the strong acids. They are referred to as “strong” because they fully dissociation in water. KNOW THESE.

  • When a metal reacts with a strong acid, the metal replaces the hydrogen atom to yield an ionic compound and hydrogen gas.

STRONG ACIDS

Zn(s) + 2HCl (aq) ZnCl2(aq) + H2(g)


Zn and hcl combine in a single replacement reaction
Zn and AcidsHCl Combine in a Single Replacement Reaction


Single replacement reactions don t just involve metals
Single Replacement Reactions Don’t Just Involve Metals Acids

  • A more reactive nonmetal can also replace a less reactive one, as shown below.

    F2(g) + 2KCl --> 2KF (s) + Cl2 (g)

  • The general rule of thumb with nonmentals is: reactivity increases up a group.

  • F > Cl > Br > I


Double replacement reaction
Double Replacement Reaction Acids

In a double replacementresult,

  • two salts react, and the anions exchange places

    AgNO3(aq) + NaCl(aq)AgCl(s)+ NaNO3(aq)

    ZnS(s) + 2HCl(aq) ZnCl2(aq) + H2S(g)


Example
Example Acids

  • Balance the following double replacement reactions

A. CaBr2 (aq) + K2CO3(aq)

B. NH4Cl (aq) + MgSO4 (aq)


Review
Review Acids

Classify each of the following reactions as:

Combination, decomposition, single replacement, or double replacement

A. 2Al(s) + 3H2SO4(aq) Al2(SO4)3(aq) + 3H2(g)

B. Na2SO4(aq) + 2AgNO3(aq) Ag2SO4(s) + 2NaNO3(aq)

C. 2NaClO3(s) 2NaCl(s) + 3O2(g)

D. 3Mg(s) + N2(g) Mg3N2 (s)


Most double replacement reactions are precipitation reactions
Most Double Replacement Reactions are Precipitation Reactions

  • An easy way to identify a chemical reaction is if there is a change in phase.

  • In a Precipitation Reaction, an insoluble ionic product is formed.

  • In the figure to the left, Na2S (aq) and Cd(NO3)2 (aq) undergo double replacement to form CdSand NaNO3 .

  • CdS is insoluble (does not dissociate). The result is the formation of a solid product.


Solubility rules
Solubility Rules Reactions

  • All group 1 and ammonium salts are soluble!

  • All nitrates, acetates, and perchlorates are soluble

  • Ag, Pb, and Hg(I) salts are all insoluble (except for those mentioned in 2)

  • Carbonates, sulfides, oxides, and phosphates are insoluble (except group 1)

  • MOST hydroxides are insoluble, EXCEPT for hydroxides of Ba, Ca, and Sr (and group 1)

  • All sulfates are soluble EXCEPT for Ca and Ba


Examples of precipitates
Examples of Precipitates Reactions

  • Use solubility rules to predict the products of the following double replacement reactions. If there is no change of phase, say ‘no reaction’:

    • BaCl2 (aq) + Na2SO4(aq)

    • MgBr2 (aq) + K2CO3 (aq)

    • NaCH3COO (aq) + CaBr2 (aq)


Net ionic equations
Net Ionic Equations Reactions

  • It is proper practice to use NET IONIC EQUATIONSwhen describing a double replacement reaction, especially one involving the formation of a precipitate

  • Ex. Na2S(aq) + Cd(NO3)2(aq) 2NaNO3(aq) + CdS(s)

  • Since we know that ionic solutions dissociate in water, we can rewrite the equation above in ionic form:

2Na+(aq) + S2-(aq) + Cd2+(aq) + 2NO3-(aq) CdS(s) + 2Na+(aq) + 2NO3-(aq)

The ions in red undergo a chemical reaction, as indicated by the change in phase. The remaining ions are called SPECTATOR IONSbecause they are not involved in the reaction in any way.


Net ionic equations1
Net Ionic Equations Reactions

  • The spectators ions cancel out. The remaining reactants and products comprise the net ionic equation.

Na+(aq) + S2-(aq) + Cd2+(aq) + NO3-(aq) CdS(s) + Na+(aq) + NO3-(aq)

Cd2+(aq) + S2-(aq) CdS(s)

NET IONIC EQUATION


Example1
Example Reactions

  • Identify the spectator ions, then write the net ionic equation corresponding to the following reactions:

  • Na2S(aq) + 2HCl(aq) 2NaCl (aq) + H2S(g)

  • 2AgClO4(aq) + (NH4)2SO4 ?


Part ii introduction to red ox reactions
Part II. Introduction to Red-Ox Reactions Reactions

  • Single replacement reactions are examples of red-ox (reduction-oxidation) reactions

  • A reduction process corresponds to a process in which the oxidation state (charge) of an element/ion becomes more negative during the course of a reaction

  • In an oxidationprocess, the oxidation state of an element/ion becomes more positiveduring a reaction


Introduction to red ox reactions
Introduction to Red-Ox Reactions Reactions

  • Consider the following single replacement reaction:

    Zn(s) + Cu SO4 (aq) Zn SO4 (aq) + Cu (s)

On the reactant side, we have elemental Zn. The charge on any pure element is 0

On the product side, we have a Zn2+ ion. Since the charge of Zn has gone from 0 to 2+, Zn has undergone an oxidation. Zn loses 2 electrons. Where did they go???

On the product side, we have elemental Cu, so Cu has undergone a reduction from 2+ to 0 by taking electrons from Zn.

On the reactant side, we have a Cu2+ ion.


Oxidizing and reducing agents
Oxidizing and Reducing Agents Reactions

Zn(s) + Cu SO4 (aq) Zn SO4 (aq) + Cu (s)

  • We have identified the reduction and oxidation processes in the reaction above

    Zn0 Zn2+ + 2e-

    Cu2+ + 2e-  Cu0

RED-OX REACTIONS

  • Because Zn gets oxidized, it is the reducing agent. In other words, the oxidation of Zn causes the reduction of Cu2+

  • Because Cu2+ gets reduced, it is the oxidizing agent. Zn is oxidized because Cu2+ takes electrons away from more active Zn.


Zn cu transmetallation
Zn-Cu ReactionsTransmetallation

Zn(s) + Cu SO4 (aq) Zn SO4 (aq) + Cu (s)


Example of red ox reactions in everyday life rust
Example of Red-Ox Reactions in Everyday Life: Rust Reactions

Reduced

4Fe(s) + 3O2(g) 2Fe2O3(s)

Oxidized


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