🏠 Home➗ Math🧪 Science🏛️ History📺 Arts & Humanities🤝 Social Studies💻 Engineering & Technology💰 Business📚 Other📓 Study Guides🏆 Leaderboard💯 All Tags❓ Unanswered🔀 Random🎒 Acids and BasesAcids and bases will neutralize one another to form liquid water and a salt. Describe the general properties of acids and bases, comparing the three ways to define them
Each of these acids ionize essentially 100% in solution. By definition, a strong acid is one that completely dissociates in water; in other words, one mole of the generic strong acid, HA, will yield one mole of H+, one mole of the conjugate base, A−, with none of the unprotonated acid HA remaining in solution. By contrast, however, a weak acid, being less willing to donate its proton, will only partially dissociate in solution. At equilibrium, both the acid and the conjugate base will be present, along with a significant amount of the undissociated species, HA. Two key factors contribute to overall strength of an acid:
Acid strengths are also often discussed in terms of the stability of the conjugate base. Stronger acids have a larger Ka and a more negative pKa than weaker acids.
A−(aq)+H2O(aq)→AH(aq)+OH−(aq)\text{A}^-(\text{aq})+\text{H}_2\text{O}(\text{aq})\rightarrow \text{AH}(\text{aq})+\text{OH}^-(\text{aq})A−(aq)+H2O(aq)→AH(aq)+OH−(aq) Thus, deprotonated water yields hydroxide ions, which is no surprise. The concentration of hydroxide ions increases as pH increases. Most alkali metal and some alkaline earth metal hydroxides are strong bases in solution. These include:
HCl(aq)+NaOH(aq)→H2O(l)+NaCl(aq)\text{HCl}(\text{aq})+\text{NaOH}(\text{aq})\rightarrow \text{H}_2\text{O}(\text{l})+\text{NaCl}(\text{aq})HCl(aq)+NaOH(aq)→H2O(l)+NaCl(aq) This reaction is called a neutralization reaction.
Acids + Bases Made Easy! Part 1 - What the Heck is an Acid or Base? - Organic Chemistry - YouTube: Ever wondered what the heck an Acid or Base actually is? Were you ever super confused in high school or college chemistry? In this video I introduce to you guys what the heck an Acid and Base really is forgetting the Lewis or Bronstead/Lowry definitions and then we'll go more in depth in parts 2,3, and 4. An Arrhenius acid dissociates in water to form hydrogen ions, while an Arrhenius base dissociates in water to form hydroxide ions. Recall the Arrhenius acid definition and its limitations.
NH2−\text{NH}_2^-NH2− ) will readily deprotonate ammonia. Thus, the Arrhenius definition can only describe acids and bases in an aqueous environment. An Arrhenius acid-base reaction is defined as the reaction of a proton and an hydroxide ion to form water:H++OH−→H2O\text{H}^++\text{OH}^-\rightarrow \text{H}_2\text{O}H++OH−→H2O Thus, an Arrhenius acid base reaction is simply a neutralization reaction.Chemistry 12.1 What are Acids and Bases? (Part 1 of 2) - YouTube: This introduction to acids and bases discusses their general properties and explains the Arrhenius definitions for acids and bases. A Brønsted-Lowry acid is any species capable of donating a proton; a Brønsted-Lowry base is any species capable of accepting a proton. Differentiate Brønsted-Lowry and Arrhenius acids.
acid + base ⇌\rightleftharpoons⇌ Here, a conjugate base is the species that is left over after the Brønsted acid donates its proton. The conjugate acid is the species that is formed when the Brønsted base accepts a proton from the Brønsted acid. Therefore, according to the Brønsted-Lowry definition, an acid-base reaction is one in which a conjugate base and a conjugate acid are formed (note how this is different from the Arrhenius definition of an acid-base reaction, which is limited to the reaction of H+ with OH- to produce water). Lastly, note that the reaction can proceed in either the forward or the backward direction; in each case, the acid donates a proton to the base. Consider the reaction between acetic acid and water:H3CCOOH(aq)+H2O(l)⇌H3CCOO−(aq)+H3O+(aq)\text{H}_3\text{CCOOH}(\text{aq})+\text{H}_2\text{O}(\text{l})\rightleftharpoons \text{H}_3\text{CCOO}^-(\text{aq})+\text{H}_3\text{O}^+(\text{aq})H3CCOOH(aq)+H2O(l)⇌H3CCOO−(aq)+H3O+(aq) Here, acetic acid acts as a Brønsted-Lowry acid, donating a proton to water, which acts as the Brønsted-Lowry base. The products include the acetate ion, which is the conjugate base formed in the reaction, as well as hydronium ion, which is the conjugate acid formed. Note that water is amphoteric; depending on the circumstances, it can act as either an acid or a base, either donating or accepting a proton. For instance, in the presence of ammonia, water will donate a proton and act as a Brønsted-Lowry acid:NH3(aq)+H2O(l)⇌NH4+(aq)+OH−(aq)\text{NH}_3(\text{aq})+\text{H}_2\text{O}(\text{l})\rightleftharpoons \text{NH}_4^+(\text{aq})+\text{OH}^-(\text{aq})NH3(aq)+H2O(l)⇌NH4+(aq)+OH−(aq) Here, ammonia is the Brønsted-Lowry base. The conjugate acid formed in the reaction is the ammonium ion, and the conjugate base formed is hydroxide.Chemistry 12.1 What are Acids and Bases? (Part 2 of 2) - YouTube: This lesson continues to describe acids and bases according to their definition. We first look at the Brønsted-Lowry theory, and then describe Lewis acids and bases according to the Lewis Theory Water is capable of acting as either an acid or a base and can undergo self-ionization. Explain the amphoteric properties of water.
H2O+H2O⇌H3O++OH−{\text{H}}_{2}\text{O} + {\text{H}}_{2}\text{O} \rightleftharpoons {\text{H}}_{3}{\text{O}}^{+} + \text{O}{\text{H}}^{-}H2O+H2O⇌H3O++OH− This is an example of autoprotolysis (meaning "self-protonating") and it exemplifies the amphoteric nature of water (ability to act as both an acid and a base ).Note that the self-ionization of water is an equilibrium reaction: H2O+H2O⇌H3O++OH−KW=1.0×10−14{\text{H}}_{2}\text{O} + {\text{H}}_{2}\text{O} \rightleftharpoons {\text{H}}_{3}{\text{O}}^{+} + \text{O}{\text{H}}^{-}\quad\quad\quad \text{K}_\text{W}=1.0\times10^{-14}H2O+H2O⇌H3O++OH−KW=1.0×10−14 Like all equilibrium reactions, this reaction has an equilibrium constant. Because this is a special equilibrium constant, specific to the self-ionization of water, it is denoted KW; it has a value of 1.0 x 10−14. If we write out the actual equilibrium expression for KW, we get the following: KW=[H+][OH−]=1.0×10−14\text{K}_\text{W}=[\text{H}^+][\text{OH}^-]=1.0\times 10^{-14}KW=[H+][OH−]=1.0×10−14 However, because H+ and OH- are formed in a 1:1 molar ratio, we have: [H+]=[OH−]=1.0×10−14=1.0×10−7 M[\text{H}^+]=[\text{OH}^-]=\sqrt{1.0\times 10^{-14}}=1.0\times 10^{-7}\;\text{M}[H+]=[OH−]=1.0×10−14=1.0×10−7M Now, note the definition of pH and pOH:pH=−log[H+]\text{pH}=-\text{log}[\text{H}^+]pH=−log[H+] pOH=−log[OH−]\text{pOH}=-\text{log}[\text{OH}^-]pOH=−log[OH−] If we plug in the above value into our equation for pH, we find that:pH=−log(1.0×10−7)=7.0\text{pH}=-\text{log}(1.0\times 10^{-7})=7.0pH=−log(1.0×10−7)=7.0 pOH=−log(1.0×10−7)=7.0\text{pOH}=-\text{log}(1.0\times 10^{-7})=7.0pOH=−log(1.0×10−7)=7.0 Here we have the reason why neutral water has a pH of 7.0; it represents the condition at which the concentrations of H+ and OH- are exactly equal in solution. We have already established that the equilibrium constant KW can be expressed as:KW=[H+][OH−]\text{K}_\text{W}=[\text{H}^+][\text{OH}^-]KW=[H+][OH−] If we take the negative logarithm of both sides of this equation, we get the following:−log(KW)=−log([H+][OH−])-\text{log}(\text{K}_\text{W})=-\text{log}([\text{H}^+][\text{OH}^-])−log(KW)=−log([H+][OH−]) −log(KW)=−log[H+]+−log[OH−]-\text{log}(\text{K}_\text{W})=-\text{log}[\text{H}^+]+-\text{log}[\text{OH}^-]−log(KW)=−log[H+]+−log[OH−] pKW=pH+pOH\text{pK}_\text{W}=\text{pH}+\text{pOH}pKW=pH+pOH However, because we know that pKW = 14, we can establish the following relationship: pH+pOH=14\text{pH}+\text{pOH}=14pH+pOH=14 This relationship always holds true for any aqueous solution, regardless of its level of acidity or alkalinity. Utilizing this equation is a convenient way to quickly determine pOH from pH and vice versa, as well as to determine hydroxide concentration given hydrogen concentration, or vice versa.Self-ionization of Water: Explanation of self-ionization of water and the formation of hydronium and hydroxide ions. The acid dissociation constant (Ka) is the measure of the strength of an acid in solution. Compare and contrast acid strengths using Ka and pKa values.
HA(aq)⇌H+(aq)+A−(aq)\text{HA}(\text{aq}) \rightleftharpoons \text{H}^+(\text{aq}) + \text{A}^-(\text{aq})HA(aq)⇌H+(aq)+A−(aq) In the above reaction, HA (the generic acid), A- (the conjugate base of the acid), and H+ (the hydrogen ion or proton) are said to be in equilibrium when their concentrations do not change over time. As with all equilibrium constants, the value of Ka is determined by the concentrations (in mol/L) of each aqueous species at equilibrium. The Ka expression is as follows: Ka=[H+][A−][HA]\text{K}_\text{a}=\frac{[\text{H}^+][\text{A}^-]}{[\text{HA}]}Ka=[HA][H+][A−] Acid dissociation constants are most often associated with weak acids, or acids that do not completely dissociate in solution. This is because strong acids are presumed to ionize completely in solution and therefore their Ka values are exceedingly large. Due to the many orders of magnitude spanned by Ka values, a logarithmic measure of the acid dissociation constant is more commonly used in practice. The logarithmic constant (pKa) is equal to -log10(Ka).The larger the value of pKa, the smaller the extent of dissociation. A weak acid has a pKa value in the approximate range of -2 to 12 in water. Acids with a pKa value of less than about -2 are said to be strong acids. A strong acid is almost completely dissociated in aqueous solution; it is dissociated to the extent that the concentration of the undissociated acid becomes undetectable. pKa values for strong acids can be estimated by theoretical means or by extrapolating from measurements in non-aqueous solvents with a smaller dissociation constant, such as acetonitrile and dimethylsulfoxide.
Acetic acid is a weak acid with an acid dissociation constant Ka=1.8×10−5\text{K}_\text{a}=1.8\times 10^{-5}Ka=1.8×10−5 pKa=−log(1.8×10−5)=4.74\text{pK}_\text{a}=-\text{log}(1.8\times 10^{-5})=4.74pKa=−log(1.8×10−5)=4.74 A p-scale is a negative logarithmic scale. Convert between pH and pOH scales to solve acid-base equilibrium problems.
H2O⇌H+(aq)+OH−(aq)\text{H}_2\text{O}\rightleftharpoons \text{H}^+(\text{aq})+\text{OH}^-(\text{aq})H2O⇌H+(aq)+OH−(aq) This reaction has a special equilibrium constant denoted KW, and it can be written as follows: KW=[H+][OH−]=1.0×10−14\text{K}_\text{W}=[\text{H}^+][\text{OH}^-]=1.0\times 10^{-14}KW=[H+][OH−]=1.0×10−14 Because H+ and OH- dissociate in a one-to-one molar ratio,[H+]=[OH−]=1.0×10−14=1.0×10−7[\text{H}^+]=[\text{OH}^-]=\sqrt{1.0\times 10^{-14}}=1.0\times 10^{-7}[H+]=[OH−]=1.0×10−14=1.0×10−7 If we take the negative logarithm of each concentration, we get:pH=−log[H+]=−log(1.0×10−7)=7.0\text{pH}=-\text{log}[\text{H}^+]=-\text{log}(1.0\times 10^{-7})=7.0pH=−log[H+]=−log(1.0×10−7)=7.0 pOH=−log[OH−]=−log(1.0×10−7)=7.0\text{pOH}=-\text{log}[\text{OH}^-]=-\text{log}(1.0\times 10^{-7})=7.0pOH=−log[OH−]=−log(1.0×10−7)=7.0 Here we have the reason that neutral water has a pH of 7.0 -; this is the pH at which the concentrations of H+ and OH- are exactly equal. Lastly, we should take note of the following relationship:pH+pOH=14\text{pH}+\text{pOH}=14pH+pOH=14 This relationship will always apply to aqueous solutions. It is a quick and convenient way to find pH from pOH, hydrogen ion concentration from hydroxide ion concentration, and more.Generically, this p-notation can be used for other scales. In acid -base chemistry, the amount by which an acid or base dissociates to form H+ or OH- ions in solution is often given in terms of their dissociation constants (Ka or Kb). However, because these values are often very small for weak acids and weak bases, the p-scale is used to simplify these numbers and make them more convenient to work with. Quite often we will see the notation pKa or pKb, which refers to the negative logarithms of Ka or Kb, respectively. Interactive: pH: Test the pH of things like coffee, spit, and soap to determine whether each is acidic, basic, or neutral. Visualize the relative number of hydroxide ions and hydronium ions in solution. Switch between logarithmic and linear scales. Investigate whether changing the volume or diluting with water affects the pH. Or you can design your own liquid! pH and pOH: This lesson introduces the pH scale and discusses the relationship between pH, [H+], [OH-] and pOH. CC licensed content, Shared previouslyCC licensed content, Specific attribution |