Acids and Bases

Acids and Bases

• Acids
  • Always have hydrogen as the first element in their chemical formula
  • Examples: HCl, HNO3, H2SO4
  • H+ (also written H3O+, the hydronium ion) is the ion responsible for their properties
  • Electrolytes that conduct electricity in solution
  • React with metals to produce H2(g) and an ionic compound
  • React with carbonates to produce CO2(g), H2O (I) and an ionic compound
  • Taste/feel is sour/watery
  • pH is less than 7.0
  • Colour with phenolphthalein is colourless
  • Red/Blue litmus paper turns Red
  • Colour with Bromothymol blue is Yellow
• Bases
  • Usually have a metal as the first element and the negative ion is OH- in their chemical formula
  • Examples: NaOH, Mg(OH)2, Zn(OH)2
  • OH- (the hydroxide ion) is the ion responsible for their properties
  • Electrolytes that conduct electricity in solution
  • No reaction with metals
  • No reaction with carbonates
  • Taste/Feel is bitter/slippery
  • pH is more than 7.0
  • Colour with phenolphthalein is Pink
  • Red/Blue litmus paper turns Blue
  • Colour with Bromothymol blue is Blue

Common acid examples

• Sulfuric acid is used in the production of fertilisers and is a chemical found in acid rain.
• Hydrofluoric acid is a highly corrosive substance which, as a gas, is a severe poison and acts as a catalyst in oil refining.
• Cola contains phosphoric acid, which has been linked to the lowering of bone density in various studies.
• Lemon juice is about 5% citric acid, a weak organic acid that gives lemons their sour taste.
• Various acids are formed during the fermentation process in beer production and the addition of CO2 also causes the pH to lower slightly.
• Levels of acids found in coffee can be affected by the altitude where the coffee was grown and the minerals present in the soil.
• Milk goes sour over time due to bacteria producing lactic acid as part of a fermentation process.

Naming Acids

• Acids are written with hydrogen as their first element, although there are some exceptions.
• The name of the acid depends on the name of its anion (negative ion).
• Acids must be in aqueous solutions, dissolved in H₂O.
• HCl (s) is not an acid, but HCl (aq) is an acid.

Binary Acids

• Binary acids are made with two types of elements: hydrogen + non-metal
  • Write the root of the non-metal name.
  • Add the prefix hydro- to the root name.
  • Add the ending –ic acid to the root name.
  • Examples:
    • HCl → Hydrochloric acid
    • HF → Hydrofluoric acid
    • HBr → Hydrobromic acid

Ternary Acids

• These acids are made of hydrogen and a polyatomic ion
  • Look at the ending of the polyatomic anion.
  • Replace the ending -ite with -ous if the anion ends in -ite.
  • Replace the ending -ate with -ic if the anion ends in -ate.
  • Add the word acid.
• Examples:
  • H2SO4 contains the sulfate ion, named Sulfuric acid.
  • HNO2 contains the nitrite ion, named Nitrous acid.
  • H2CO3 contains the carbonate ion, named Carbonic acid.
• It's not called sulfic acid, because that sounds weird

Acid Formulae

• Determine if there is a polyatomic ion. If yes, it is a ternary acid, otherwise it is a binary acid with hydro_ic acid.
• Then identify the ions involved.
• Determine how many H+ ions are required so that the net charge of the compound is zero by using the crossover rule.
• The same way ionic compounds are named.

Acid Formulae Examples

• Boric acid: not "hydro___ic acid", so must contain a polyatomic ion.
  • Boric acid → borate polyatomic ion → BO33- and H+ → H3BO3
• Hydroiodic acid, contains "hydro____ic acid”, so must be a binary acid.
  • Hydroiodic acid → H+ and I- → HI

Acid Examples

• HCl: Chloride (Cl-), hydrochloric acid
• HClO3: Chlorate (ClO3-), chloric acid
• H2SO3: Sulfite (SO32-), sulfurous acid
• H2CrO4: Chromate (CrO42-), chromic acid
• CH3COOH*: Acetate (CH3COO), acetic acid (*Acetic acid's formula is an exception to the general rule)

Acidic Oxides

• When a nonmetal combines with oxygen, it produces a nonmetal oxide.
• Most nonmetal oxides, such as sulfur dioxide (SO₂) or nitrogen monoxide (NO) are gases at room temperature and produce acidic solutions when dissolved in water.
• Acids produce H+ (or H3O+) ions in solution
  • But nonmetals oxides do not have hydrogen in their formulae, so how do they produce an acid?
  • The production of an acid from a nonmetal oxide occurs in two steps:
    • Synthesis reaction resulting in nonmetal oxide
    • Ionization of the ionic compound

Carbon Dioxide as Acidic Oxide

• Synthesis: CO2 (g) + H2O (l) → H2CO3 (aq)
• Ionization: the hydrogen carbonate ionizes in water:
  • H2CO3 (aq) → H+ (aq) + CO32- (aq)

Nitrogen Oxides as Acidic Oxides

• Nitrogen makes up about 78% of our air, and most of it is harmless N2, but some of it is found in nitrogen oxides
• Naturally occurring nitrogen oxides result from lightening strikes, but most it is a by-product of the combustion of fuels in cars: -N2 (g) + O2 (g) → 2 NO (g)
• NO is released into the atmosphere through the car's exhaust pipe, mixes with oxygen in the air to form nitrogen dioxide: 2 NO (g) + O2 (g) → 2 NO2 (g)
• Nitrogen dioxide is a large component of air pollution and smog
• Nitrogen dioxide can produce two types of acid when it reacts with water:
  • 2NO2 (g) + H2O (l) → HNO3 (aq) + HNO2 (aq)
• Both nitric acid (HNO3) and nitrous acid (HNO₂) will ionize in water and produce hydrogen ions.

Sulphur Oxides as Acidic Oxides

• Sulphur is an impurity in many fossil fuels like coal and crude oil.
• When these fuels are burned, sulphur dioxide is produced: S (s) + O2 (g) → SO2 (g).
• The sulphur dioxide further reacts with oxygen in the atmosphere to make sulphur trioxide: SO2 (g) + O2 (g) → SO3 (g)
• Sulphur trioxide reacts with the water in the atmosphere, creating sulphuric acid (a big part of acid rain), which ionizes to release hydrogen ions.
  • SO3 (g) + H2O → H2SO4 (aq)

Naming Bases

• Bases are written with the name of the metal cation first, even if the cation is a polyatomic ion.
• Then the word hydroxide is added at the end, as it is the anion for all bases.
• NaOH: sodium (Na+), Sodium hydroxide
• NH4OH: ammonium (NH4+), ammonium hydroxide
• Mg(OH)2: magnesium (Mg2+), Magnesium hydroxide

Base Formulae

• Steps for writing chemical formulae:
  • Identify the ions involved, one must be hydroxide.
  • Determine how many OH- ions are required so that the net charge of the compound is zero by using the crossover rule.
  • The same way we name ionic compounds!
• Example 3: Sodium Hydroxide: Na+ and OH- → NaOH

Basic Oxides

• When a metal combines with oxygen, it produces a metal oxide.
• Metallic oxides generally react with water to form metal hydroxides.
• When the metal hydroxides are put into solution, the ions dissociate (break into) their ions.
• Since one of the ions is an OH-, it form a base.

Arrhenius Theory

• This first theory came to be in the late 1800's, by Swedish chemist Svante Arrhenius.
• Arrhenius wondered why some solutions conduct an electric current while others do not.
  • Salt dissolved in water results in a solution that can conduct electricity.
  • Sugar dissolved in water results in a solution that cannot conduct electricity.
• Arrhenius suggested that certain kinds of substances break apart in water to form ions, which results in solutions that are able to conduct an electric current.
• Arrhenius observed that acidic and basic solutions always conduct an electric current.
• Arrhenius acids must contain a hydrogen atom that can become a hydrogen ion.
  • HCl(g) → H+(aq) + Cl-(aq)
• When acids combine with water, they ionize.
• Arrhenius bases must release hydroxide ions (OH(aq)) in a water solution:
  • NaOH(s) → Na+(aq) + OH (aq)
• When bases dissolve in water, they form ions (ionize), but bases dissociate (a special word for bases that means to break apart the ions).

Brønsted-Lowery Theory

• Arrhenius' theory doesn't work for all acids and bases because ammonia, NH3(g), does not contain hydroxide ions, but it forms a basic solution when dissolved in water.
• The Brønsted-Lowery theory works for molecules like ammonia which chemically reacts with water to produce a hydroxide ion.
• Chemists Johannes Nicolaus Brønsted and Thomas Martin Lowry independently developed definitions of acids and bases based on the compounds' abilities to either donate or accept protons (H+ ions) (1923).
• Acids are defined as proton donors and bases are defined as proton acceptors.
• A compound that acts as both a Brønsted-Lowry acid and base is called amphoteric (e.g., Water)
• Water is not the only amphoteric substance out there, but it is the one we will commonly use
• The positively charged hydrogen ion is strongly attracted to the oxygen atoms of the water molecules H+(aq) + H2O(l) → H3O+(aq)
• H3O+(aq) is known as the hydronium ion.

Acids and Bases in General

• Acids are a substance that reacts with water to produce H3O+(aq) (hydronium) ions in aqueous solution. Bases are a substance that dissolves in water or reacts with water to produce OH-(aq) (hydroxide) ions in aqueous solution.

Strong vs. Weak Acids

• Strong acids are acids that ionize nearly 100% in water and are very good electrolytes in solution.
• Their pH is very low (close to 0 or 1).
• Examples: perchloric acid, hydrochloric acid, hydrobromic acid, hydroiodic acid, nitric acid, sulfuric acid
• Acid strength increases down a group on the periodic table
  • HF (weak) < HCl < HBr < HI.
• Acid strength increases across a period (left → right)
  • H3P < H2S < HCl.
• Acid strength increases as the number of oxygen atoms increases
  • HClO < HClO2 < HClO3 < HClO4.
• Weak acid is an acid that ionizes only slightly in solution (less than 1%): Almost all of the acid is found in solution as intact molecules with the H+ ion still attached.
• They are very poor electrolytes in solution, have a pH between 4.5 and about 6.9 (depending on acid concentration)
• Organic (or carbon based) acids are examples of weak acids such as acetic acid, ascorbic acid (vitamin C), malic acid (grapes), citric acid (citrus fruits)

Strong vs. Weak Bases

• Strong bases are bases that dissociate nearly completely in solution and are very good electrolytes in solution with a very high pH (close to 13 or 14).
• Arrhenius bases are examples of strong bases such as lithium hydroxide, sodium hydroxide, potassium hydroxide, calcium hydroxide, strontium hydroxide, barium hydroxide.
• A weak base is a base that dissociates only slightly in solution (less than 1%) and almost all of the base is found in solution as intact 'molecules' with the OH still attached.
• Many weak bases are insoluble in water and very poor electrolytes in solution.
• They have a pH between 7.1 and about 9.5 (depending on concentration).
• A Brønsted-Lowery base (like ammonia) is an example of a weak base.

Strength vs. Concentration

• Strong and weak describe how much an acid ionizes, or a base dissociates in solution.
• Concentrated and dilute describe how much solute is in solvent, or how much acid or base is in water.
• Common units for concentration include mol/L, mol/kg, ppm, ppb
• If there were more acid or base in water, then the solution would be concentrated; a dilute solution would have less acid or base in water.
• Weak acids and bases can also be made into concentrated and dilute solutions.
• The word “weak” can be misleading because concentrated solutions of weak acids or bases can still be very dangerous!
• Hydrogen peroxide is a weak acid and can be dilute form so it's safe to clean skin cuts. In concentrated form, it is so explosive that it can be used as rocket fuel!

pH

• pH stands for power of hydronium.
• The pH scale is used to measure how acidic or basic a solution is; runs from 0 to 14.
• From 0 to 6.9 are acids, from 7.1 to 14 are bases.
• Substances with a pH of 7 are neutral.

Interpreting pH Values

• Acids produce hydronium ions in solution; therefore, an acidic solution has a hydronium ion concentration that is greater than its hydroxide ion concentration.
• Bases produce hydroxide ions in solution; therefore, a basic solution has a hydroxide ion concentration that is greater than its hydronium ion concentration.
• Each step on the pH scale is actually 10X
  • A substance with a pH of 4 is 10x more acidic than a substance with a pH of 5
  • A substance with a pH of 4 is 100x more acidic than a substance with a pH of 6!
• Grapes, with a pH of 3 are 10000 more acidic than pure water at a pH of 7!

Measuring pH Values

• There are 2 common ways of measuring pH:
  • Universal indicator turns different colours at specific pH values (it is dried on paper to produce a tool called pH paper): A drop is placed of the substance you are testing on the paper, and then compare the colour produced on the paper with a colour chart to determine the pH value
  • A pH meter, which is the most accurate way to measure pH uses a pH probe (electronic) that is placed in solution, that will generate a pH value reading.
• pH can be mathematically determined if the concentration of H3O+ or OH- is known, using three simple formulae:
  • pH = -log[H3O+]
  • pOH = -log[OH-]
  • pH + pOH = 14
• Square brackets [ ] means “concentration of"
• Water will spontaneously form the ions H3O+ and OH- which is referred to as the self-ionization of water:
  • 2H2O ⇄ H3O+ + OH-
• Since acids and bases need to be in aqueous solutions (i.e., dissolved in water) there will be H3O+ and OH- ions in every solution.
  • [H3O+] = [OH-] indicates a neutral solution
  • [H3O+] > [OH-] indicates an acidic solution
  • [H3O+] < [OH-] indicates a basic solution
• Example 1: a solution has the H3O+ concentration of 5.4 x 10-12 mol/L
  • pH = -log[H3O+] = -log[5.4×10-12] = 11.26
  • This solution is a base because the pH is more than 7.
• Example 2: a solution has the HO- concentration of 1.4 x 10-8 mol/L:
  • pOH = -log[OH-] = -log[1.4×10−8] = 7.85
  • pH = 14 - pOH = 6.15: this solution is an acid because the pH is less than 7
• If we only know the concentration of either H3O+ or OH- we can mathematically determine the missing concentration with the simple formula:
  • Kw = [H3O+][OH-]
  • Where Kw = 1.0 x 10-14
• Example 3: A solution has an OH- concentration of 8.9 x 10-6 mol/L:
  • Kw = [H3O+][OH-]
  • [H3O+] = Kw/[OH-] = 1.0×10-14/8.9×10-6 = 1.0×10-9
  • Since [H3O+] < [OH-] this is a basic solution

Neutralization

• Many people have baking soda in refrigerators, also know as sodium bicarbonate to removes the odors caused by spoiling foods/ the smelly breakdown products of many foods (acids!).

• Baking soda neutralizes the acids of these breakdowns!

• Neutralization reactions occur when an acid reacts with a base, in aqueous solution, to produce an ionic compound and water

• An ionic compound is sometimes known as a "salt", and it is made of the anion from an acid, from a base, and are a pH neutral

• Example 1: Nitric acid reacts with sodium hydroxide: HNO3(aq) + NaOH(aq) → HOH(l) + NaNO3(aq)

• Example 2: Sulphuric acid reacts with barium hydroxide: H2SO4(aq) + Ba(OH)2(aq) → HOH(ł) + BaSO4(aq)

• Neutralization reactions can create more than just water and an ionic compound - when you neutralize an acid with a compound that contains either the carbonate or the bicarbonate anion, you also produce carbon dioxide.

• Example 1. HCl(aq) + NaHCO3(s) → NaCl(aq) + H2O(l) + CO2(g)

• Example 2. 2HCl(aq) + CaCO3(s) → CaCl2(aq) + H2O(l) + CO2(g)

• It is possible to have a neutralization reaction when a strong acid reacts with a metal oxide (i.e. a basic oxide) to also produce an ionic compound and water: HCl (aq) + Na2O (aq) → H2O (l) + 2NaCl (aq) strong acid + metal oxide → water + ionic compound

• It is also possible to have a neutralization reaction when a strong base reacts with a nonmetal oxide (i.e. an acidic oxide) to also produce an ionic compound and water: 2NaOH (aq) + CO2 (aq) → H2O (l) + Na2CO3 (aq) strong base + nonmetal oxide → water + ionic compound