1. Solutions Unit
Honors Chemistry

2. Naming Acids Review: Ex) HCl hydrochloric acid
A. Binary – H +one anion Prefix “hydro”+ anion name +“ic”acid
Ex) HCl hydrochloric acid
Ex) H3P hydrophosphoric acid
B. Tertiary – H + polyatomic anion no Prefix “hydro”
(oxo) end “ate” = “ic” acid end “ite” = “ous” acid
Ex) H2SO4 sulfuric acid
Ex) H2SO3 sulfurous acid

3. Properties of Acids and Bases:
Taste
Touch
Reactions with Metals
Electrical Conductivity
Acid
sour
looks like water, burns, stings
Yes-produces H2 gas
electrolyte in solution
Base
(alkali)
bitter
looks like water, feels slippery
No Reaction

4. Indicators: Turn 1 color in an acid and another color in a base.
Litmus Paper: Blue and Red
An aciD turns blue litmus paper reD
A Base turns red litmus paper Blue.
Phenolphthalein: colorless in an acid and pink in a base
pH paper: range of colors from acidic to basic
pH meter: measures the concentration of H+ in solution

6. Reactions
Neutralization: A reaction between an acid and base. When an acid and base neutralize, water and a salt (ionic solid) form.
Acid Base → Salt Water
Ex) HCl NaOH → NaCl HOH

7. Arrhenius Definition (1884):
A. An acid dissociates in water to produce more
hydrogen ions, H+.
HCl  H Cl-1
B. A base dissociates in water to produce more hydroxide ions, OH-.
NaOH  Na OH-1
C. Problems with Definition:
• Restricts acids and bases to water solutions.
• Oversimplifies what happens when acids
dissolve in water.
• Does not include certain compounds that have
characteristic properties of acids & bases.
Ex) NH3 (ammonia) doesn’t fit

8. Bronsted-Lowry Definition (1923):
A. An acid is a substance that can donate hydrogen ions. Ex) HCl → H+ + Cl-
Hydrogen ion is the equivalent of a proton.
Acids are often called proton donors.
Monoprotic (HCl), diprotic (H2SO4) , triprotic (H3PO4)
B. A base is a substance that can accept hydrogen ions. Ex) NH3 + H+ → NH4+
Bases are often called proton acceptors.
C. Advantages of Bronsted-Lowry Definition
•Acids and bases are defined independently of how
they behave in water.
•Focuses solely on hydrogen ions.

9. Hydronium Ion:
Hydronium Ion – H3O+ This is a complex ion that forms in water.
H H2O  H3O+1
To more accurately portray the Bronsted-Lowry, the hydronium ion is used instead of the hydrogen ion.

10. STRONG Acid/Base versus WEAK Acid/Base
Strength refers to the % of molecules that form IONS.
A strong acid or base will completely ionize (>95% as ions). This is represented by a single () arrow.
HNO3 + H2O  H3O+ + NO3-
A weak acid or base will partially ionize (<5% as ions). This is represented by a double (↔) arrow.
HOCl + H2O ↔ H3O+ + ClO-

11. HF < HCl < HBr < HI increasing strength
7 Strong Acids
HNO3 H2SO4 HClO3
HClO4 HCl HBr HI
8 Strong Bases
LiOH NaOH KOH
RbOH CsOH Ca(OH)2
Sr(OH)2 Ba(OH)2

12. Strength vs. Concentration
Strength refers to the percent of molecules that form ions
Concentration refers to the amount of solute dissolved in a solvent. Usually expressed in molarity.

13. Ionization of Acids & Bases
H2SO4  2 H+ + SO4-2
Sulfuric acid
H3PO3 
Phosphorous acid
Ca(OH)2 
Calcium hydroxide
3 H+ + PO3-3
Ca OH-1

14. Conjugate Acid-Base Pairs: A pair of compounds that differ by only one hydrogen ion
Acid donates a proton to become a conjugate base.
Base accepts proton to become a conjugate acid.
A strong acid will have a weak conjugate base.
A strong base will have a weak conjugate acid.

15. Acid (A), Base (B), Conjugate Acid (CA), Conjugate Base (CB)
NH3 + H2O ↔ NH OH-
HCl + H2O ↔ Cl H3O+
Base and Conjugate Acid are a Conjugate Pair.
Acid and Conjugate Base are a Conjugate Pair. B A CA CB A B CB CA

16. AciDonates & Bases accept
H2O H2O ↔ H3O OH−
B A CA CB
H2SO OH− ↔ HSO4− H2O
A B CB CA
HSO4− H2O ↔ SO4− H3O+
A B CB CA
OH− H3O+ ↔ H2O H2O
B A CA CB

17. The Self-ionization of Water & pH
1. Water is amphoteric, it acts as both an acid and a base in the same reaction.
Ex) H2O(l) + H2O(l) ↔ H3O+(aq) + OH-(aq)
Keq = equilibrium constant = [H3O+] [OH-]
Because reactants and products are at equilibrium, liquid water is not included in the equilibrium expression
@ 25C, [H3O+] = 1 x 10-7 M and [OH-] = 1 x 10-7 M
Kw = ion product constant or equilibrium constant for water
Kw = [H3O+] [OH-] = 1 x M2
1.0 x M2 = [1.0 x 10-7 M] [1.0x10-7 M]
1.0 x = [H3O+] [OH-]

18. Acids: [H3O+] > 1 x 10-7 M Bases: [OH-] > 1 x 10-7 M
Using Kw in calculations: If the concentration of H3O+ in the blood is 4.0 x 10-8 M, what is the concentration of OH­ ions in the blood? Is blood acidic, basic or neutral?
Kw = [H3O+] [OH-]
1.0 x M2 = [4.0 x 10-8 M] [OH-]
2.5 x 10-7 M = [OH-]
slightly basic

19. The pH scale (1909): the power of Hydrogen
Measure of H3O+ in solution.
pH = -log[H3O+]
Range of pH: 0-14
pH < 7: acid
pH > 7: base
pH = 7: neutral
pOH = -log[OH-]
pH + pOH = 14

20. OH- H+ 14 1 pH [H3O+] [OH-] 14 1x10-14 1x100 13 1x10-13 1x10-1 1211
1x10-11
1x10-3 10
1x10-10
1x10-4 9
1x10-9
1x10-5 8
1x10-8
1x10-6 7
1x10-7 6 5 4 3 2 1 14 H+
OH- 1

21. Significant Digits Rule
The number of digits AFTER THE DECIMAL POINT in your answer should be equal to the number of significant digits in your original number
Ex -log[8.7x10-4M]
Calc Answer =
Sig Fig pH = 3.06

22. Concentration Percent concentration by mass (mass % Molarity (M)
(solute/solution) x 100% = % Concentration
Molarity (M)
Moles of solute/Liters of solution = mol/L
Molality (m)
Moles of solute/mass of solvent = mol/kg
ppm and ppb
Used for very dilute solutions
Dilution – a process in which more solvent is added to a solution
How is this solution different?
Volume, color, molarity
How is it the same?
Same mass of solute, same moles of solute
In Dilution ONLY – M1V1 = M2V2

23. Dissolution Process Ionic Compounds NaCl(s)  Na+1(aq) + Cl-1(aq)
For dissolution to occur, must overcome solute attractions and solvent attractions.
Dissociation Reaction: the separation of IONS when an ionic compound dissolves (ions already present)
Try calcium chloride
electrolyte
nonelectrolyte
Dissolving NaCl in water
hexahydrated for Na+1;
most cations have 4-9 H2O molecules
6 is most common
Solvation: process of solvent molecules
surrounding solute
Hydration: solvation with water

24. V. Solution Stoichiometry
A. Many reactants are introduced to a reaction chamber as a solution.
B. The most common solution concentration is molarity.
molarity = mol/liter
C. Examples
Excess lead(II) carbonate reacts with 27.5 mL of 3.00M nitric acid. Calculate the mass of lead(II) nitrate formed
PbCO3 + 2HNO3  Pb(NO3)2 + H2CO3

25. 2. Calculate the volume, in mL, of a 0
2. Calculate the volume, in mL, of a molar solution of sulfuric acid required to react completely with g of sodium carbonate according to the equation below.
H2SO Na2CO3  Na2SO4 + CO2 + H2O

26. Energy Changes Heat of solution = Hsoln Endothermic Exothermic
Solute particles separating in solid
Solvent particles moving apart to allow solute to enter liquid
Energy absorbed
Exothermic
Solvent particles attracted to solvating solute particles
Energy released