Test 3 Review: Acids & Bases

Acid and Base models

Vocabulary

• Amphoteric - can donate or accept a proton. It can act as either acids or bases but it’s not limited to the Bronsted-Lowry theory.

• Amphiprotic - substance can act as both Bronsted-lowry acids or Bronsted-Lowry bases

  • H₂O

All amphiprotic things are amphoteric but not all amphoteric are amphiprotic

• Polyprotic - can donate more than one proton

• Conjugate base - an acid that has lost its hydrogen ion

• Conjugate acid - a base that has gained a hydrogen ion

• Conjugate pair - A conjugate acid-base pair differ by a proton (H⁺)

Water Ions

• Water is H⁺ and OH^-

• The usual solvent for acids and bases is water

• Water breaks down to form a hydrogen ion (H⁺) and a hydroxide ion (OH^-)

• H2O(l) → H⁺(aq) + OH^-(aq)

• Self-ionizes: 2 water molecules react to form a hydronium ion (H₃O⁺) and a hydroxide ion (OH^-)

• H2O(I) + H2O(I) → H₃O⁺(aq) + OH‍^-(aq)

Hydronium ions

• H₃O⁺ is a hydrated hydrogen ion (H⁺)

  • water is attached to an H⁺ by a covalent bond

• H₃O⁺ and H⁺ can be used interchangeably to represent a hydrogen ion in an aqueous solution

• H2O(l) → H⁺(aq) + OH^-(aq)

Ions in a solution

• All aqueous solutions contain hydrogen ions (H) and hydroxide ions (OH)

  • An acidic solution contains more H than OH

  • A basic solution contains more OH than H

  • A neutral solution contains equal concentrations of both H and OH

Bases

• The net result of adding a base to water is that the solution will have a higher concentration of hydroxide ions (OH) than pure, neutral water

• Pure neutral water has a concentration of hydroxide ions that is 1×10^-7 M

Adding acid to water

• The net result of adding an acid to water is that the solution will have a higher concentration of hydrogen ions (H) than pure, neutral water.

• Pure neutral water has a concentration of hydrogen ions that is 1×10^-7 M

Concentrations

• The typical range for the concentration of hydrogen ions found in nature is from 1×10⁰ M to 1×10^-14 M

pH

• Range from 0-14

  • 0.0 is strongly acidic

  • 14.0 is strongly basic

  • 7.0 is neutral

• Acids pH between 0 - 7

• bases pH between 7 - 14

Two definitions of acids and bases

• Arrhenius

• Brownsted-Lowry

Arrhenius's definition of acids and bases

• In the early 1800s, Arrhenius developed the idea that the conductivity of solutions was due to the presence of ions in the solution

• This was considered ridiculous until charged particles were discovered in the late 1890s

• Acids: Produce hydrogen ions in aqueous solutions

• Bases: produce hydroxide ions

Bronsted-Lowry’s definition of acids and bases

• Chemicals that give up H⁺ are acids (proton donors)

• Chemicals that accept H⁺ are bases (proton acceptors)

• All Bronsted-Lowry acids contain hydrogen

  • They must be able to donate a hydrogen ion to another species

• All Bronsted-Lowry bases must have a lone pair of electrons to form a bond with the hydrogen ion donated to them by the Bronsted-Lowry acid

• Amines act as Bronsted-Lowry bases

Combining theories

• Acids still always have hydrogen

• Arrhenius bases have to have an OH that dissociates

• Bronsted-lowry bases need an unshared pair of electrons (to accept the proton)

• Water can act as an acid or a base (it is amphoteric)

Reminders about acids and bases

1. Most acid-base reactions occur in aqueous solutions

2. The strength of an acid depends on the extent to which hydrogen is ionized when the acid dissolves in water. This is called dissociation

3. Acids are considered strong if they essentially dissociate 100%. There are six strong acids - assume all others are weak acids and do not dissociate extensively

4. Strong acids:

   • HCl

   • HBr

   • HI

   • HNO₃

   • H₂SO₄

   • HClO₄

5. Weak acids: everything else

6. Bases attract the H ions from acids. Therefore, they often contain hydroxide. Soluble hydroxides are strong bases.

7. Soluble hydroxides contain 1A cations plus barium hydroxide and strontium hydroxide

Common Acids

• Hydrochloric acid - HCl

• Nitric acid - HNO₃

• Sulfuric acid - H₂SO₄

• Ethanoic acid - CH₃COOH

• Carbonic acid - H₂CO₃

• Phosphoric acid - H₃PO₄

• Benzoic acid - C₆H₅COOH

Acids can be classified as monoprotic, diprotic, or triprotic depending on the number of hydrogen ions, or protons, that they can donate in acid-base reactions

monoprotic

diprotic

triprotic

Although Ethanoic acid has 4 hydrogens, it is monoprotic because it can only donate 1 hydrogen.

Common bases

• Calcium oxide - CaO

• Copper (II) oxide - CuO

• Sodium hydroxide - NaOH

• Lithium hydroxide - LiOH

• Calcium hydroxide - Ca(OH)₂

• Barium hydroxide - Ba(OH)₂

• Ammonia - NH₃

• Ethanamine - CH₃CH₂NH₂

Bases that are soluble in water are known as alkalis.

Summary

• In the Arrhenius model, acids produce hydrogen ions and bases produce hydroxide ions in solution

• In the Bronsted-Lowry model, acids are proton donors and bases are proton acceptors

• Strong acids and bases dissociate 100% in solution. Weak acids and bases do not.

• Stronger acid = weaker conjugate base

• Stronger base = weaker conjugate acid

pH and pOH

pH

• The pH of a substance is a measure of the concentration of hydrogen ions, [H⁺], in solution.

• The negative logarithm of the hydrogen ion activity

• Each number represents a tenfold change in the acidity or alkalinity of a substance

• The logarithmic scale makes it easier to record and graph very small numbers

• The Danish biochemist Soren Sorenson proposed the term pH to refer to the “power of hydrogen ions”. He defined the “p” as the negative of the logarithm, -log, of [H+]

• Used to measure the acidity of water and soil, which is important for the health of animals and crops

• The higher the concentration of H ions in solution, [H+], the lower the pH value

• The lower the [H+], the higher the pH value

• A higher concentration of OH^- ions, [OH⁻], corresponds to a higher pH value

• A lower concentration of OH^- ions, [OH⁻], corresponds to a lower pH value

• pH values below 0 and above 14 are possible

• pH = −log[H⁺(aq)]

• If the [H+] is equal to [OH^-], the solution is neutral

  • [H+] = [OH-]

    • Neutral

    • pH = 7

• If the [H+] is greater than the [OH-], the solution is acidic

  • [H⁺] > [OH⁻]

    • Acidic

    • pH <7

• If the [OH-] is greater than the [H+], the solution is basic, or alkaline

  • [OH⁻] > [H⁺]

    • Basic

    • pH >7

Square brackets are used to show the concentration of a substance

Another logarithmic scale is the Richter scale which is used to measure earthquakes

Logarithmic scale

• This scale is logarithmic to the base 10

• A change of one pH unit represents ten times the change in hydrogen ion concentration

  • Example:

    • An increase of two pH units represents a 100 times decrease in [H⁺]

• The pH of a solution is directly related to the concentration of hydrogen ions present

  • Values for pH are stated with many decimal places as the # of significant figures in the H⁺ concentration

  • Examples:

    • 0 - 1×10⁰ or 1

    • 1 - 1×10^-1 or 0.1

    • 2 - 1×10^-2 or 0.01

    • 5 - 1×10 or 0.00001

    • 6 - 1×10^-6 or 0.000001

    • 7 - 1×10^-7 or 0.0000001

    • 8 - 1×10^-8 or 0.00000001

    • 13 - 1×10^-13 or 0.0000000000001

    • 14 - 1×10^-14 or 0.00000000000001

• A strong acid's pH is the exponent's negative if the coefficient is 1.

• Every time that we want to state the strength of an acid we need to express the [H⁺] using scientific notation or write out a string of zeros.

pH calculations

• for strong acids or bases, the [H⁺] or [OH^-] in solution is equal to the initial concentration of the strong acid or base

• Strong acids and bases completely dissociate in solution

• Rearranging the equation for pH allows us to calculate the concentration of H⁺ ions in a solution from its pH value:

  • [H⁺(aq)] = 10^−pH

pOH calculations

• The pOH of a solution can be calculated by taking the negative log of the hydroxide ion concentration, [OH^-], using the following equation:

  • pOH = −log[OH⁻(aq)]

Summary

• pH is related to [H+] by the formula: pH = -log [H+]

• [H+] is related to pH by the formula: [H+] = 10^-pH

• pH is related to pOH by the formula: pH + pOH = 14

Acid-base Indicators

• an acid-base indicator is a weak acid or a weak base in which the dissociated and undissociated forms have different colors

• Examples of common indicators:

  • Litmus paper

  • Universal indicator

Litmus paper

• In an acidic solution, blue litmus changes to red and in an alkaline solution, red litmus changes to blue

Universal indicator

• Universal indicator is actually a mixture of indicators that produces different colors in solutions of different pH

  

  Indicators may be used to measure the pH of a solution, by making use of the fact that their color changes with pH.

pH meters

• A pH probe and meter produces a more accurate method of measuring pH

• The electrode of the pH meter is placed in the solution to be tested and a voltage is generated that is converted into a pH meter reading displayed on the screen

The ionic product of water

• Water molecules dissociate however, only to a very small extent; this is known as the auto-ionisation of water

• H₂O (l) ⇌ H⁺ (aq) + OH^– (aq)

  • The equilibrium constant expression (K_c) for the above reaction is:

    • K_c=[H+][OH−]/[H2O]

• The position of equilibrium for the dissociation of HO lies very far to the left, so the concentration of the water is effectively constant. This leads to a new expression known as the ionic product of water (K_w):

  • K_w=[H+][OH−]

• like any equilibrium expression, K_w is temperature-dependent, therefore, at a certain temperature the product of [H⁺] and [OH^-] is a constant.

  • as [H⁺] of a solution increases, [OH^-] decreases (and vice versa)

  • A neutral solution is any which has [H⁺] equal to [OH^-]

  • Acidic solution - [H⁺] > [OH^–]

  • Basic solution - [OH^–] > [H⁺]

Vocabulary

• Ionic product of water - the product of the [H⁺] and [OH^-] in water at a particular temperature

• Auto-ionisation of water - The ability of water to react with itself and form two ions

Oxide trends

Bonding and structure

• The period 3 oxides show a gradual trend of decreasing metallic character across the period

• Solids under standard conditions because of strong electrostatic attraction between the ions, a feature of ionic bonding

• The bonding in period 3 oxides changes from ionic to covalent across the period

Acid-base properties of the period 3 oxides

• Aluminium oxide is amphoteric (it can act as both an acid and a base) although it does not react with water

• Across the period (left to right) the acidity increases

• Basic oxides dissolve in water to produce basic solutions

  • Na₂O (s) + H₂O (l) → 2NaOH (aq)

  • MgO (s) + H₂O (l) → Mg(OH)₂ (aq)

• Acidic oxides dissolve in water to produce acidic solutions

  • P₄O₁₀ (s) + 6H₂O (l) → 4H₃PO₄ (aq)

  • SO₃ (g) + H₂O (l) → H₂SO₄ (aq)

  • 3NO₂ (g) + H₂O (l) → 2HNO₃ (aq) + NO (g)

  • Cl₂O (g) + H₂O (l) → 2HClO (aq)

• Silicone dioxide does not react with water because of its giant covalent structure, but does react with strong bases such as sodium hydroxide

Test 4 Review: Redox & Electrochemistry

Rules for determining Oxidation states

• The oxidation state of an uncombined element is zero. this applies regardless of the structure of the element

  • Example: Cl₂, Xe, S₈

• The sum of the oxidation states of all the atoms or ions in a neutral compound is zero

• The sum of the oxidation states of all the atoms in an ion is equal to the charge of the ion

• The more electronegative element in a substance is assigned a negative oxidation state. The less electronegative element is assigned a positive oxidation state

  • Electronegativity is greatest at the top-right of the periodic table and decreases toward the bottom-left

• Group 1 metals always have a +1 oxidation state

• Group 2 metals always have a +2 oxidation state

• Oxygen always has a -2 oxidation state UNLESS it is in a peroxide or F₂O

  • In a peroxide, the oxidation state of oxygen is -1

  • In F₂O, the oxidation state of oxygen is +2

• Hydrogen always has an oxidation state of +1 UNLESS it is in a metal hydride

  • When hydrogen is in a metal hydride, it has an oxidation state of -1

• Fluorine always has an oxidation state of -1

• Chlorine always has an oxidation state of -1 UNLESS it is in a compound with O or F

  • Chlorine adopts a wide variety of oxidation states so it is safer to assume the oxidation state is not -1

Acid and Base half-reactions

Step 1. Determine whether or not it is an oxidation-reduction reaction

Step 2. Separate the reaction into half-reactions

• The substance being reduced will have electrons as reactants

• The substance being oxidized will have electrons as products

In some cases, it is important to note which half-reaction is being oxidized and which is being reduced

Acidic conditions

• Acidic conditions usually imply a solution with an excess of H⁺ concentration, making the solution acidic

• the balancing starts by separating the reaction into half-reactions

• However, instead of immediately balancing the electrons, balance all the elements in the half-reactions that are not hydrogen and oxygen

• Then, add H₂O molecules to balance any oxygen atoms

• Next, balance the hydrogen atoms by adding protons

• Now, balance the charge by adding electrons and scale the electrons (multiply by the lowest common multiple) so that they will cancel out when added together

• Finally, add the two half-reactions and cancel out common terms

Half-reaction method: acidic

Each reaction is balanced by adjusting coefficients and adding H₂O, H⁺, and e^- in this order:

1. Determine, using oxidation states, what is being oxidized and reduced

2. Balance elements in the equation other than O and H

3. Balance the oxygen atoms by adding the appropriate number of water (H₂O) molecules to the side that has fewer oxygen atoms

4. Balance the hydrogen atoms (including those added in step 2 to balance the oxygen atoms) by adding H⁺ ions to the opposite side of the equation

5. Add up the charges on each side. Make them equal by adding enough electrons (e^-) to the more positive side.

   (rule of thumb: e- and H are almost always on the same side)

6. The e^- in each half-reaction must be made equal; if they are not equal, they must be multiplied by the appropriate integers (the lowest common multiple) to be made the same

7. The half-equations are added together, canceling out the electrons to form one balanced equation. Common terms should also be canceled out

8. The equation can now be checked to make sure that it is balanced

Basic conditions

• Bases dissolve into OH^- ions in solution; hence, balancing redox reactions in basic conditions requires OH^-

• Follow the same steps as for acidic conditions. The only difference is adding hydroxide ions to each side of the net reactions to balance any H⁺

• OH^- and H⁺ ions on the same side of a reaction should be added together to form water

• Again, any common terms can be canceled out

Half reaction Method: Basic

The same steps with the addition of 6b. Each reaction is balanced by adjusting coefficients & adding H₂O, H⁺, e^- and OH^- in this order:

1. Determine, using oxidation states, what is being oxidized and reduced

2. Balance elements in the equation other than O and H

3. Balance the oxygen atoms by adding the appropriate number of water (H₂O) molecules to the side that has fewer oxygen atoms

4. Balance the hydrogen atoms (including those added in step 2 to balance the oxygen atom) by adding H⁺ ions to the opposite side of the equation

5. Add up the charges on each side. Make them equal by adding enough electrons (e^-) to the more positive side (e^- and H⁺ are almost always on the same side)

6. The e^- in each half-reaction must be made equal; if they are not equal, they must be multiplied by appropriate integers (the lowest common multiple) to be made the same

7. The half-equations are added together, canceling out the electrons to form one balanced equation

   • Add the appropriate number of OH^- to neutralize all H⁺ and convert into water molecules

8. Common terms should also be canceled out

9. The equation can now be checked to make sure that it is balanced

Summary

• Acidic half-reactions are balanced by adjusting coefficients & adding H₂O, H⁺, and e^- in that order

• Alkaline half-reactions are balanced by adjusting coefficients & adding H₂O, H⁺, e^-, and OH^- in that order

Electrolytic cells

The electrical conductivity of ionic compounds when molten or dissolved in aqueous solutions is explained by the fact that mobile ions move a particular direction in an electric field.

Electrolysis - The process by which a compound is broken down into its constituent elements using electricity.

• An electrolytic cell is composed of a molten or aqueous electrolyte, a battery, and two electrodes, the anode and the cathode.

• An electric current enters and leaves via the two electrodes, which are usually made of graphite (carbon) or an unreactive metal such as platinum

  • Such metals are know as being inert as they do not take part in the reaction

• The anode is connected to the positive terminal of the battery

  • Negative ions, or anions, are attracted to the anode

• The cathode is connected to the negative terminal of the battery

  • Positive ions, or cations, are attracted to the cathode

• When the ions reach the surface of the electrodes they undergo either oxidation or reduction and are discharged

In an electrolytic cell, the anode is positively charged and the cathode is negatively charged.

Voltaic vs. Electrolytic cells

Voltaic cells

• Oxidation occurs at the negative anode

• Reduction occurs at the positive cathode

• Involves an exothermic spontaneous redox reaction

• The cell converts chemical energy into electrical energy

• The cathode is positive and the anode is negative during discharge

• The cell uses two separate aqueous solutions connected by a salt bridge and an external circuit

• Electric current is conducted by the electrons in the external circuit and the movement of ions in the salt bridge

Electrolytic cells

• Oxidation occurs at the positive anode

• Reduction occurs at the negative cathode

• Involves an endothermic non-spontaneous redox reaction

• The cell converts electrical energy into chemical energy

• The cathode is negative and the anode is positive during electrolysis

• The electrolyte is a molten liquid (or an aqueous solution)

• Electric current is conducted by the electrons in the external circuit and the movement of ions in the electrolyte

New: Kinetics & Equilibrium