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Acids Vs. Bases:

• Generally Acids are H+ ions attached to a nonmetal anion, or a polyatomic ion.

  • Will usually have an H in front

• Generally Bases are Ionic Compounds with metal cations attached to a OH- ion.

  • OH- are hydroxide ions.

  • Will usually end in OH

→ Acids, Bases and Salts are known as electrolytes. Which means that when they are dissolved in water they split into their individual ions allowing for electrical conductivity.

The Theories:

→ There are 2 ones that we use to define Acids and Bases. (They build off of each other)

Arrhenius:

• Acids → When dissolved in water it will release H+ ions into the solution, increasing the concentration of H+ ions.

  • H+ ions are also called protons, because that’s all it is. (a Hydrogen ion has no electrons).

• Bases → When dissolved in water it will release OH- ions (Hydroxide Ions) increasing their concentration in a solution.

  • Base can also be referred to as Alkalinity.

• The Regents prefers this theory because it is easier to understand.

Bronsted-Lowry:

• Acids → Any substance that will donate a Hydrogen Ion. (H+)

  • Has an extra H+

• Bases → Any substance that will accept a Hydrogen Ion. (H+)

  • Has a spot for the H+ ion to attach

• Conjugate Pairs →A base in the solution accepts the H+ Ion, becoming a conjugate Acid.

  • Works the same way for conjugate Bases.

Naming:

Acids:

• For Acids with no oxygen you use the prefix “hydro”

  • The add the second element's name but change the ending to “ic”

    • They finally add “Acid” at the end.

  • Should look like hydro(element)ic Acid

  • Example: HCl (aq)

    • Hydrochloric Acid

  • Example: HI (aq)

    • HydroIodic Acid

• For Acids with Oxygen you don’t use a prefix

  • The first part is just the polyatomic Ion, then it will either end in “ic” or “ous”

    • If the polyatomic ends in “ate” then the acid ends in “ic”

    • If the polyatomic ends in “ite” then the acid ends in “ous”

      • Finally we add “Acid to the very end”

  • The pneumonic to remember this is:

    • I ATE organIC apples, despITE being poisonOUS.

  • Example: H2SO4 (aq)

    • The polyatomic is Sulfate so the acid will end in “ic”

      • Sulfuric Acid

Neutralization and Acid Base Equations:

Neutralization/ Titration:

→ The names can be used interchangeably.

• A neutralization reaction will usually look like

  • Acid + Base → Salt + Water

    • remember salt is an Ionic compound.

  • The acronym to remember it is SWAB.

• When the reaction is done the acid and base neutralize each other, so the pH returns to 7.

• Titration is when you add an acid or a base with an unknown volume or Molarity, to another acid or base, with a known volume or Molarity, until it neutralizes so you can find the unknown variable

  • When they are neutralized the moles of both the acid and base should be equal so from that we get the equation:

    • (MA)(VA) = (MB)(VB)

    • M= Molarity

    • V= Volume

    • A= Acid

    • B= Base

      • Equation is on the back of the reference table.

Other Equations:

• When you add an Acid with a more reactive metal, it will produce H2 and a salt.

  • The more reactive metals can be found on Table J, in the reference table.

    • Any metals shove H2 will react with Acids, anything below it won;t

    • Cu, Ag, and Au won’t react with Acids.

  • The acronym to remember these equations is: MASH

  • Example:

  • 2K + 2HCl → H2 + 2KCl

Strength (pH/pOH), and Indicators:

Strength:

→ The more Ions which can dissociate the stronger the acid or the base.

• Monoprotic- Has one H+ or OH- ion.

  • HCl

• Diprotic- Has two H+ or OH- ions.

  • H2SO4

    • This pattern for names continues on

• Organic Acids like Ethanoic Acid are generally weaker.

pH and pOH scale:

pH Scale:

→ Expression of how many H+ ions are in a solution. (The Acidity)

• 0-6 is very acidic (0 is the most acidic)

• 7 is neutral (really any number between 6-8)

• 8-14 is very basic. (14 is the most Basic)

  • It is calculated using: -log[H+]

  • [H+] means concentration of H+ ions in the solution.

  • [ ] always mean concentration

  • The higher the concentration of H+ ions, the lower the pH.

• The change between each number is x10.

  • so going from a pH= 1 to pH= 2

  • the concentration decreases by 10

  • Going from pH=2 to pH= 4

    • the concentration decreases by 100

pOH Scale:

→ Expression of how many OH- ions are in a solution. (The Alkalinity)

• opposite of pH

• found using a similar method: -log[OH-]

  • 0-6 is basic

  • 7 is neutral (really anything 6-8)

  • 8-14 is Acidic

• You could also subtract your pH from 14 and it will give you the pOH.

Indicators:

• Change colors when Hydrogen is lost or gained. (Used to show the general pH)

  • There are many different ones which have unique behaviors

    • You can find the most common indicators and their colors for certain pH’s on the reference table (Table M)

Next Unit: Unit 9- Redox/Electrochemistry