knowt logo

Holt Chemistry Chapter 15: Acids and Bases

Holt Chemistry Chapter 15: Acids and Bases 


15.1 What Are Acids and Bases? 

  • Acids are electrolytes and can conduct electric currents 
  • Acids react with metals. They usually produce hydrogen gas 
  • Strong acids will ionize completely in a solution 
    • This is a completion reaction 
    • Strong acids are better conductors of electricity 
    • HCl is a strong acid
    • HBr is a strong acid 
    • HI is a strong acid 
    • HNOis a strong acid 
    • H2SOis a strong acid 
    • HClO4 is a strong acid 
    • HIO4 is a strong acid 
  • Weak acids will release few hydrogen ions in aqueous solutions 
    • They have smaller a Keq than the strong acids 
    • CH3COOH is a weak acid 
    • HCN is a weak acid 
    • HF is a weak acid
    • HNO2 is a weak acid 
    • H2SO3 is a weak acid 
    • HOCl is a weak acid 
    • H3PO4 is a weak acid 
  • The Arrhenius definition of an acid is limited to aqueous solutions 
  • Arrhenius states that an acid is any substance that, when added to water, increases the hydronium ion concentration 
  • Acids are usually gases or liquids 
  • The hydronium ion is H3O+
  • Bases are also electrolytes
  • Bases are usually solids 
  • Bases feel slippery because they react with the oil in your hands, converting the bases into soaps 
  • Some bases are insoluble while others are soluble
  • Very soluble bases are called alkalis 
  • Strong bases ionize completely in a solvent 
    • NaOH is a strong base 
    • KOH is a strong base 
    • Ca(OH)2 is a strong base 
    • Ba(OH)2 is a strong base 
    • Na3POis a strong base 
  • Weak bases release few hydroxide ions in aqueous solutions 
    • NH3 is a weak base 
    • Na2CO3 is a weak base 
    • K2CO3 is a weak base 
    • C6H5NHis a weak base 
    • (CH3)3N is a weak base 
  • The hydroxide ion is OH-
  • The Arrhenius definition of a base is a substance, when dissolved in water, increases the hydroxide ion concentration 
  • The Arrhenius definitions are limited to aqueous solutions
  • The Arrhenius definitions cannot classify substances that sometimes act as bases and sometimes act as acids 
  • The Bronsted-Lowry definition of an acid is any substance that can donate a proton 
  • The Bronsted-Lowry definition of a base is any substance that can accept a proton 
  • The Bronsted-Lowry proton is most commonly referred to as H+
  • The Bronsted-Lowry definitions are not limited to aqueous solutions
  • A conjugate acid in an acid that forms when a base gains a proton 
  • A conjugate acid is a base that forms when an acid loses a proton 
  • Conjugate acids pair up with bases 
  • Conjugate bases pair up with acids 
  • Acids and bases are always reactants 
  • Conjugate acids and conjugate bases are always products 
  • Amphoteric describes a substance, such as water, that has properties of both acids and bases 

15.2 Acidity, Basicity, and pH

  • When a water molecule donates a proton to another water molecule, both a hydronium ion and hydroxide ion form. This is known as the self-ionization of water 
    • The self-ionization of water is a special equilibrium constant 
    • It is represented with the symbol Kw
    • The equation to find the self-ionization of water is K= [H3O+] x [OH-]
  • The concentrations of hydronium [H3O+] are equal to the concentrations of hydroxide [OH-]
  • [H3O+] = [OH-] = 1.00 x 10-7
  • Kw = 1.00 x 10-14
  • Using this formula, you can calculate an unknown concentration 
  • The concentration of hydronium ions is expressed as acidity 
  • The concentration of hydroxide ions is expressed as basicity
  • If the concentrations are equal then the solution is neutral 
  • pH stands for the power of hydronium 
  • The pH is on a negative logarithmic scale, meaning a lower pH has a higher hydronium concentration 
  • pOH stands for power of hydroxide 
  • A pH of 7 is neutral 
  • Below are different equations to solve for pH, pOH, and concentrations 
    • pH = -log[H3O+
    • [H3O+] = 10-pH
    • pOH = -log[OH-]
    • [OH-] = 10-pOH
    • pH + pOH = 14
  • Indicators change color, the color will then tell you what pH the substance is 
  • pH meters can also be used 

15.3 Neutralization and Titrations 

  • A neutralization reaction is when and acid and a base form water and a salt 
    • This reaction will form equal concentrations of hydronium and hydroxide 
  • When the neutralization reaction is complete, this is known as the equivalence point 
    • This is when the concentrations are equal 
  • This is different than the endpoint 
  • The endpoint is when a color change occurs 
  • However, if a correct indicator is used, the endpoint can be the same as the equivalence point 
  • Titration is the method used to determine the concentration of an unknown substance. It is done by adding concentration of a known solution until a color change occurs in the mixture 
  • The solution added is called the titrant 
  • The equipment used in a titration are two burets, a titration flask, and an indicator 
  • The solution of a known concentration is called a standard solution 
  • Titration curves show pH plotted against volume 
    • The curve will be steep at the equivalence point 
  • Indicators have transition range. They do not cover the whole pH scale, but rather small sections 
    • Thymol blue ranges from 1.2-2.8
    • Methyl orange ranges from 3.1-4.4
    • Litmus ranges from 5.0-8.0
    • Bromthymol blue ranges from 6.0-7.6 
    • Phenolphthalein ranges from 8.0-9.6
    • Alizarin yellow ranges from 10.1-12.0 
  • n = cV 
    • n is the amount of solute in moles 
    • c is the concentration in moles per liter 
    • V is the volume in liters 
    • (c H3O+)(V H3O+) = (c OH-)(V OH-)
    • Therefore cV = cV

15.4 Equilibria of Weak Acids and Bases 

  • Some acids are able to donate protons better, therefore they are stronger acids 
  • Some bases accept protons better, therefore they are stronger bases 
  • The general principle states In an acid-base reaction, the conjugate base of the stronger acid is the weaker base, and the conjugate acid of the stronger base is the weaker acid 
  • The acid-ionization constant describes the ionization of acid in water 
    • Ka
    • The formula can be expressed has [H3O+][other product] / [reactant] = Ka
  • A buffer solution is made from a weak acid and its conjugate base
  • Buffers neutralize small amounts of acids or bases 
  • Buffers can be found in blood, foods, milk, and more 

Holt Chemistry Chapter 15: Acids and Bases 


15.1 What Are Acids and Bases? 

  • Acids are electrolytes and can conduct electric currents 
  • Acids react with metals. They usually produce hydrogen gas 
  • Strong acids will ionize completely in a solution 
    • This is a completion reaction 
    • Strong acids are better conductors of electricity 
    • HCl is a strong acid
    • HBr is a strong acid 
    • HI is a strong acid 
    • HNOis a strong acid 
    • H2SOis a strong acid 
    • HClO4 is a strong acid 
    • HIO4 is a strong acid 
  • Weak acids will release few hydrogen ions in aqueous solutions 
    • They have smaller a Keq than the strong acids 
    • CH3COOH is a weak acid 
    • HCN is a weak acid 
    • HF is a weak acid
    • HNO2 is a weak acid 
    • H2SO3 is a weak acid 
    • HOCl is a weak acid 
    • H3PO4 is a weak acid 
  • The Arrhenius definition of an acid is limited to aqueous solutions 
  • Arrhenius states that an acid is any substance that, when added to water, increases the hydronium ion concentration 
  • Acids are usually gases or liquids 
  • The hydronium ion is H3O+
  • Bases are also electrolytes
  • Bases are usually solids 
  • Bases feel slippery because they react with the oil in your hands, converting the bases into soaps 
  • Some bases are insoluble while others are soluble
  • Very soluble bases are called alkalis 
  • Strong bases ionize completely in a solvent 
    • NaOH is a strong base 
    • KOH is a strong base 
    • Ca(OH)2 is a strong base 
    • Ba(OH)2 is a strong base 
    • Na3POis a strong base 
  • Weak bases release few hydroxide ions in aqueous solutions 
    • NH3 is a weak base 
    • Na2CO3 is a weak base 
    • K2CO3 is a weak base 
    • C6H5NHis a weak base 
    • (CH3)3N is a weak base 
  • The hydroxide ion is OH-
  • The Arrhenius definition of a base is a substance, when dissolved in water, increases the hydroxide ion concentration 
  • The Arrhenius definitions are limited to aqueous solutions
  • The Arrhenius definitions cannot classify substances that sometimes act as bases and sometimes act as acids 
  • The Bronsted-Lowry definition of an acid is any substance that can donate a proton 
  • The Bronsted-Lowry definition of a base is any substance that can accept a proton 
  • The Bronsted-Lowry proton is most commonly referred to as H+
  • The Bronsted-Lowry definitions are not limited to aqueous solutions
  • A conjugate acid in an acid that forms when a base gains a proton 
  • A conjugate acid is a base that forms when an acid loses a proton 
  • Conjugate acids pair up with bases 
  • Conjugate bases pair up with acids 
  • Acids and bases are always reactants 
  • Conjugate acids and conjugate bases are always products 
  • Amphoteric describes a substance, such as water, that has properties of both acids and bases 

15.2 Acidity, Basicity, and pH

  • When a water molecule donates a proton to another water molecule, both a hydronium ion and hydroxide ion form. This is known as the self-ionization of water 
    • The self-ionization of water is a special equilibrium constant 
    • It is represented with the symbol Kw
    • The equation to find the self-ionization of water is K= [H3O+] x [OH-]
  • The concentrations of hydronium [H3O+] are equal to the concentrations of hydroxide [OH-]
  • [H3O+] = [OH-] = 1.00 x 10-7
  • Kw = 1.00 x 10-14
  • Using this formula, you can calculate an unknown concentration 
  • The concentration of hydronium ions is expressed as acidity 
  • The concentration of hydroxide ions is expressed as basicity
  • If the concentrations are equal then the solution is neutral 
  • pH stands for the power of hydronium 
  • The pH is on a negative logarithmic scale, meaning a lower pH has a higher hydronium concentration 
  • pOH stands for power of hydroxide 
  • A pH of 7 is neutral 
  • Below are different equations to solve for pH, pOH, and concentrations 
    • pH = -log[H3O+
    • [H3O+] = 10-pH
    • pOH = -log[OH-]
    • [OH-] = 10-pOH
    • pH + pOH = 14
  • Indicators change color, the color will then tell you what pH the substance is 
  • pH meters can also be used 

15.3 Neutralization and Titrations 

  • A neutralization reaction is when and acid and a base form water and a salt 
    • This reaction will form equal concentrations of hydronium and hydroxide 
  • When the neutralization reaction is complete, this is known as the equivalence point 
    • This is when the concentrations are equal 
  • This is different than the endpoint 
  • The endpoint is when a color change occurs 
  • However, if a correct indicator is used, the endpoint can be the same as the equivalence point 
  • Titration is the method used to determine the concentration of an unknown substance. It is done by adding concentration of a known solution until a color change occurs in the mixture 
  • The solution added is called the titrant 
  • The equipment used in a titration are two burets, a titration flask, and an indicator 
  • The solution of a known concentration is called a standard solution 
  • Titration curves show pH plotted against volume 
    • The curve will be steep at the equivalence point 
  • Indicators have transition range. They do not cover the whole pH scale, but rather small sections 
    • Thymol blue ranges from 1.2-2.8
    • Methyl orange ranges from 3.1-4.4
    • Litmus ranges from 5.0-8.0
    • Bromthymol blue ranges from 6.0-7.6 
    • Phenolphthalein ranges from 8.0-9.6
    • Alizarin yellow ranges from 10.1-12.0 
  • n = cV 
    • n is the amount of solute in moles 
    • c is the concentration in moles per liter 
    • V is the volume in liters 
    • (c H3O+)(V H3O+) = (c OH-)(V OH-)
    • Therefore cV = cV

15.4 Equilibria of Weak Acids and Bases 

  • Some acids are able to donate protons better, therefore they are stronger acids 
  • Some bases accept protons better, therefore they are stronger bases 
  • The general principle states In an acid-base reaction, the conjugate base of the stronger acid is the weaker base, and the conjugate acid of the stronger base is the weaker acid 
  • The acid-ionization constant describes the ionization of acid in water 
    • Ka
    • The formula can be expressed has [H3O+][other product] / [reactant] = Ka
  • A buffer solution is made from a weak acid and its conjugate base
  • Buffers neutralize small amounts of acids or bases 
  • Buffers can be found in blood, foods, milk, and more