1. LECTURE 3: IONIC COMPOUNDS (Ch. 6)

2. Ionic Compounds
The nucleus of an atom is unchanged by chemical reactions (number of protons never changes)
However, electrons are readily added and lost and ions are formed
When a metal reacts with a nonmetal, ions form and attract. The result is an ionic compound.
Let’s consider the formation of a very common ionic compound, NaCl (s)

3. Ionic Compounds
We know that Na(s) and Cl2(g) react together to form NaCl (s), but how?
The most important thing to know about chemical reactions is that atoms undergoing a reaction will always seek to reach a noble gas configuration
Let’s look at the electron configurations of Na and Cl

4. Mechanism of an Ionic Reaction
Na: [Ne] 3s1 Cl: [Ne] 3s2 3p5
For Na, the nearest noble gas is Ne. To reach the Ne configuration, it needs to lose a single electron.
Na ( [Ne] 3s1 ) ---> Na+ ([Ne]) + e-
11 e-
10 e-
1st Ionization Energy
11 p+
11 p+
Neutral Na atom
Na+ cation

5. Mechanism of an Ionic Reaction
Na: [Ne] 3s1 Cl: [Ne] 3s2 3p5
For Cl, the nearest noble gas is Ar. To reach the Ar configuration, it needs to gain a single electron.
Cl ([Ne] 3s2 3p5) + e- ---> Cl- ([Ar])
Electron affinity describes the energy change associated with the addition of an electron.
18 e-
17 e-
1st Electron Affinity
17 p+
17 p+
Cl- anion
Neutral Cl atom

6. Mechanism of an Ionic Reaction
Na and Cl can simultaneously achieve a noble gas configuration if an electron is transferred from the metal (Na) to the nonmetal (Cl)
[Ne] 3s1 + [Ne] 3s2 3p5 ---> Na+ Cl-
[Ne]
[Ar]
IONIC COMPOUND
Lewis dot structure of the product
Na+
Cl-

7. Predicting and Balancing Charge
So now, we understand that ionic compounds form when metal and nonmetal ions interact
We also see why sodium chloride is NaCl, not NaCl2 or Na2Cl, etc.
The overall charge of any complete molecule must be zero.
Since the Na loses an electron to become Na+, and Cl gains an electron to become Cl-, only one of each ion is needed to balance the charge.
In ionic compounds, the metal is always positively charged (cation) and the nonmetal is always negatively charged (anion)

8. Predicting and Balancing Charge
GP1
metals
nonmetals 1+ 2+ 3+ 3- 2- 1-

9. Dissolving Ionic Compounds in Water
Ionic compounds completely dissociate in water, forming individual ions. Ions become completely ‘hydrated’.
Na+ Cl Na+(aq) + Cl- (aq)
H2O (L)
Here, NaCl is the solute, water is the solvent. The symbol (aq) refers to an aqueous solution (i.e. species dissolved in water).

10. Dissolving Ionic Compounds in Water
Na+
Cl-
Water molecules “solvate” ionic compounds, ripping the ions apart.
The negative oxygen atoms (red) attract to the positive Na+, and the positive hydrogens attract to the negative Cl-

11. Electrolytes
Ions in solution are capable of conducting electric current (hence, the term electrolyte).
Non-ionic compounds (e.g. covalent) do not exhibit this property because they do not dissociate

12. Ionic Radii
Cations are smaller than their neutral atom counterparts, and anions are larger
Anions have large electron clouds because the excess of negative charge causes repulsion between electrons. The excess positive charge in cations draws the electron cloud closer to the nucleus.
+ e-
- e-
Cation, X+
Neutral X
Anion, X-

13. Ionic Radii

14. Energy Changes In Reactions
The electrostatic attraction, or the electrical attraction between positive and negative ions, is what holds an ionic compound together
When two ions form an ionic compound, there is an overall change in energy.
We can calculate this energy by considering:
the ionization energy of the metal
the electron affinity of the nonmetal
the coulombic energy of attraction between the cation and anion

15. Energy Changes In Reactions
Lets revisit the reaction: Na(g) + Cl(g)  NaCl(s)
Ignore the monatomic chlorine
To form NaCl, there are 3 steps
Form Na+ (ionization energy)
Form Cl- (electron affinity)
Join them together (coulombic energy)

16. Energy Changes In Reactions
1. (Ionization of Na) Na(g)  Na+(g) + e- ΔEI = aJ *Positive sign means energy is added. 2. (Ionization of Cl) Cl(g) + e-  Cl-(g) ΔEEA = aJ *Negative sign means energy is released. 3. (Coulombic energy) Na+(g) + Cl-(g)  Na+Cl-(s) ?

17. Coulombic Energy
The third step is to join the two ions, as shown below.
rNa = 102 pm
rCl = 181 pm
The equation shown above is Coulomb’s Law, which gives the energy change (Ec) that results when two ions come together.
Q1 and Q2 are the charges of the metal and nonmetal
d is the distance between the nuclei. This is the sum of the ionic radii.
k is a constant. (231 aJ•pm)

18. Solve…. Negative energy change indicates a favorable process
Negative energy change indicates a favorable process
GP2: Refer to ionic radii table

19. Electron Configurations of Transition Metals
When a transition metal forms an ion, electrons are first removed from the preceding s-orbital.
Fe: [Ar] 4s2 3d6
Fe2+: [Ar] 3d6
Fe3+: [Ar] 3d5
If the ionization of a transition metal results in an unpaired s-electron, that electron will move into the valence d orbital
Ni: [Ar] 4s2 3d8
Ni+: [Ar] 4s1 3d8 ---> [Ar] 3d9

20. Electron Configurations of Transition Metals
Whenever possible, a transition metal will transfer its own valence s-electrons into the preceding d-orbital to obtain a completely filled or half-filled d-orbital. This makes d4 and d9 configurations unlikely when s-electrons are present.
Cu: [Ar] 4s2 3d9  [Ar] 4s1 3d10
Cr: [Ar] 4s2 3d4  [Ar] 4s1 3d5
Cr+: [Ar] 4s1 3d4  [Ar] 3d5

21. Electron Configurations of Transition Metals
Transition metals can have multiple positive ionic charges. To distinguish, a roman numeral is placed in front of a transition metal in a compound to identify its charge.
Ex. FeCl2 ---> Here, Fe is 2+. So, we name this compound:
Iron (II) chloride
FeCl3 ---> Here, Fe is 3+.
Iron (III) chloride
GP 3

22. Electronegativity
An important concept related to ionic and covalent bonding is : electronegativity
Electronegativity is the ability of an atom to attract electrons to itself. A more electronegative atom is in greater need of electrons, and will attract them more strongly than a less electronegative atom (think “tug-of-war”)
Thus, when there is a difference in electronegativities between atoms in a molecule, the electrons are NOT equally shared.

23. Electronegativity Increases Up and To the Right

24. Polar Covalent (Partially polar) Difference In Electronegativity
Although we have discussed ionic and covalent bonds, most chemical bonds are neither purely ionic or purely covalent
Most compounds are an intermediate between the two.
Bond Type
Covalent
Polar Covalent (Partially polar)
Ionic (Totally Polar)
Difference In Electronegativity
< 0.4
0.4 – 2.0
> 2.0

25. Electronegativity Cl- Na+
Let’s consider NaCl. The difference in electronegativity between Na and Cl is:
Because the difference in electronegativity is so big, the Na electron is completely pulled away by the Cl atom. So, the molecule is totally polar (ionic)
This is why there is no actual bond in ionic compounds, only coulombic attraction.
3.16 – = ionic
Cl-
Na+

26. - + Electronegativity H Cl Let’s look at another molecule, like HCl:
The HCl molecule is not purely ionic or covalent, but BOTH.
The electron density is unevenly distributed, such that more of the electron cloud is on the Cl than on the H.
Therefore, the H and Cl have partial charges. The arrow depicts the direction of electron “pull”, or dipole. Any molecule with a net dipole has polarity.
3.16 – 2.1 = polar covalent - + δ δ H Cl
Partial positive character
Partial negative character

27. Lecture 3: Part 2 Polyatomic Ions, Ionic Reactions and Redox Reactions
Polyatomic ions are covalent molecules that possess charge and behave as normal ions in solution.
When a salt containing a polyatomic ion is dissolved in water, the polyatomic ions themselves DO NOT break apart. They are simply separated from the counter-ion.
Example: Phosphate (PO43-)
Na3PO4 (s) Na+ (aq) + PO43-(aq)
H2O (L)
Sodium Phosphate
Sodium
cations
Phosphate anion

28. KNOW YOUR POLYATOMIC IONS !!!!!
Charge
Name
Structure -1
Hydroxide
OH-
Cyanide
CN-
Bicarbonate
HCO3-
Acetate
CH3COO-
Nitrate
NO3-
Nitrite
NO2-
Perchlorate
ClO4-
Charge
Name
Structure -2
Carbonate
CO32-
Oxalate
C2O42-
Sulfate
SO42-
Sulfite
SO32-
Charge
Name
Structure +1
Ammonium
NH4+
Charge
Name
Structure -3
Phosphate
PO43-
GP4

29. Ionic Reactions
Chemical reactions involving ionic compounds can be classified as one of the following:
combination reactions
decomposition reactions
single replacement reactions
double replacement reactions

30. 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
3Li(s) + P(g)  Li3P(s)
Ca(s) + Cl2(g)  CaCl2(s)
SO3(g) + H2O(l)  H2SO4(aq)
MgO(s) + CO2(g)  MgCO3 (s)
GP5

31. Decomposition Reactions
2HgO(s) 2Hg(l) + O2(g) 2KClO3(s) 2KCl(s) + 3O2(g)

32. Single Replacement Reactions
In a single replacement reaction,
When one metal replaces another, this is also called a transmetallation reaction.
Zn(s) + 2AgCl (aq)  ZnCl2(aq) + 2Ag(s)
Fe(s) + CuSO4(aq)  FeSO4(aq) + Cu(s)

33. Single Replacement
Transmetallation reactions occur because one metal is less stable (more active) in its elemental state than the other.
In the reaction below, Zn displaces Ag+ because Zn is more reactive:
Zn(s) + 2AgCl (aq) ZnCl2(aq) + 2Ag(s)
More active metals prefer to exist as aqueous ions instead of uncharged elements.
An activity series is provided on pg. 325 of the text listing metals in order of reactivity.

34. Group Work Activity Series
As shown in the table, Li, which has the highest reactivity, is the most likely to react. It will swap with any metal ion below it in the table.
Ex. Predict the products.
Li (s) + Ca(ClO4)2 (aq)
Na (s) + ZnSO4(aq)
K (s) + LiCl (aq)

35. Single Replacement Reactions Involving Acids
When a metal reacts with a binary acid (HX), the metal replaces the hydrogen atom to yield an ionic compound and hydrogen gas.
Zn(s) + 2HCl (aq) ZnCl2(aq) + H2(g)

36. Redox Reactions
Single replacement reactions are examples of red-ox (reduction-oxidation) reactions
A reduction process corresponds to a process in which a species receives electrons. The charge of the element/ion becomes more negative.
In an oxidation process, a species will lose electrons, causing its charge to become more positive.

37. Redox 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.

38. Oxidizing and Reducing Agents
Zn(s) + Cu SO4 (aq) Zn SO4 (aq) + Cu (s)
We have identified the reduction and oxidation processes in the reaction above
O: Zn0  Zn2+ + 2e-
R: Cu2+ + 2e-  Cu0
RED-OX HALF 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+ accepts electrons from the more reactive Zn.

39. Oxidizing and Reducing Agents
Zn(s) + Cu SO4 (aq) Zn SO4 (aq) + Cu (s)

40. Ex. Rust Formation
Reduced
4Fe(s) + 3O2(g) 2Fe2O3(s)
Oxidized
GP6

41. Double Replacement Reactions
In a double replacement result,
two salts react, and the anions exchange places
AgNO3(aq) + NaCl(aq) AgCl(s) + NaNO3(aq)
ZnS(s) + 2HCl(aq) ZnCl2(aq) + H2S(g)

42. Examples Balance the following double replacement reactions
A. CaBr2 (aq) + K2CO3(aq)
B. NH4Cl (aq) + MgSO4 (aq)

43. Most Double Replacement Reactions Yield Precipitates
An easy way to identify a chemical reaction is if there is a change in phase.
In a Precipitation Reaction, an insoluble (solid, does not dissolve) ionic product is formed.
In the figure to the left, Na2S (aq) and Cd(NO3)2 (aq) undergo double replacement to form CdS and NaNO3 .
CdS is insoluble

44. Net Ionic Equations
It is proper practice to use NET IONIC EQUATIONS when there is a change in phase
Ex. Na2S(aq) + Cd(NO3)2(aq) NaNO3(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 IONS because they are not involved in the reaction in any way.

45. Net Ionic Equations
Na+(aq) + S2-(aq) + Cd2+(aq) + NO3-(aq) CdS(s) + Na+(aq) + NO3-(aq)
The spectators ions cancel out. The remaining reactants and products comprise the net ionic equation.
In order to write a net ionic equation, you must know which ionic compounds are insoluble. The solubility rules enable this.
Cd2+(aq) + S2-(aq) CdS(s)
NET IONIC EQUATION

46. Solubility Rules All group 1 and ammonium salts are soluble!
All nitrates, acetates, and perchlorates are soluble
With the exception of all anions mentioned in #2, Ag, Pb, and Hg(I) salts are all insoluble
With the exception of those cations mentioned in #1, carbonates, sulfides, oxides, and phosphates are insoluble
With the exception of Ca, Sr, and Ba, and those cations mentioned in #1, all hydroxides are insoluble
All sulfates are soluble EXCEPT for Ca, Sr, Ba, and those cations mentioned in #3.

47. Examples
Use solubility rules to predict the products of the following double replacement reactions. Write the net ionic reaction. If there is no reaction, write ‘no reaction’:
MgBr2 (aq) + K2CO3 (aq)
NaCH3COO (aq) + CaBr2 (aq)