akohol

ALKOHOL, FENOL & ETER

Beberapa senyawa Alkohol & Eter yang sering dijumpai dalam kehidupan sehari-hari.

Dietileter: sebagai pemati rasa

Etanol: sebagai campuran minuman keras.

Metanol: bahan dasar spiritus

2-propanol: sebagai bakteriosida

Kedudukan atom O pada eter & alkohol, mirip kedudukan O pada air.

Gugus R pada alkohol dan eter dapat berupa aril atau alkil.

Oleh karena itu, kedua senyawa ini banyak dijumpai di alam, baik hasil sintesis maupun terjadi secara alami.

Alkohol dan eter merupakan isomer Yaitu: mempunyai rumus molekul yang sama, tetapi mempunyai struktur yang berbeda.

Rumus molekul umum: CnH2n+2O

Berdasarkan struktur, alkohol digolongkan  atas dasar atom C  yang mengikat gugus hidroksil.

1.                 1.  Alkohol Primer [1^o], gugus –OH terikat pada atom C primer, yaitu atom karbon yang mengikat satu atom C lain.  

2. Alkohol Sekunder [2^o], gugus –OH terikat pada atom C sekunder, yaitu atom karbon yang mengikat dua atom C lain.
3. Alkohol Tersier [3^o], gugus –OH terikat pada atom C tersier, yaitu atom karbon yang mengikat tiga atom C lain.

Alkohol dengan gugus R berupa aril, pada umumnya disebut fenol.

A.      Kenaikan Titik didih.

Titik didih senyawa organik biasanya berbanding lurus dengan berat molekul (Mr).

Karena adanya ik.hidrogen, maka dibutuhkan energi tambahan untuk memutus ik.hidrogen tsb. 

  air (Mr=18) mempunyai titik didih 100 ^oC, sementara aseton (Mr=58) mempunyai titik didih 58 ^oC

2.      Ethanol (Mr=46)& Methoxymethane (Mr=46), namun mempunyai TD yg sangat berbeda. 

     .Methanol (Mr=32) mempunyai TD yg lebih tinggi dari prophane (Mr=44).

B.      Kelarutan dalam air

Kelarutan alkohol & eter dalam air juga disebabkan oleh adanya ikatan hidrogen.

Semakin tinggi rantai alkil, alkohol & eter semakin sukar larut. Hal ini disebabkan oleh adanya steric hindrance. Dalam senyawa organik tidaklah mudah mendefinisikan suatu senyawa melepaskan elektron atau menangkap elektron untuk menyatakan proses oksidasi atau reduksi.

Aturan sederhana definisi oksidasi atau reduksi dalam senyawa organik adalah :

Teroksidasi

Jika sebuah molekul memperoleh oksigen [O] atau kehilangan hidrogen [H].

Tereduksi

Jika sebuah molekul kehilangan oksigen [O] atau memperoleh hidrogen [H].

problem in the acid and base

1. There are no vacant orbital on the HCl that can accept electron pairs. Why, then, HCl is a Lewis acid?
my opinion:
Chlorine is more electronegative than hydrogen, and this means that the hydrogen chloride molecule will be polar. Electrons in the hydrogen-chlorine bond will be attracted to the chlorine, producing hydrogen which is slightly positive and slightly negative chlorine. 
Lone pairs on the nitrogen contained in the ammonia molecule are drawn to the slightly positive hydrogen atom in HCl. After the lone pair of nitrogen close to the hydrogen atom, electrons in the hydrogen-chlorine bond still would refuse to the chlorine.
Finally, the coordination bond is formed between the nitrogen and hydrogen, and chlorine is lost out as a chloride ion.

Simak

Baca secara fonetik

strenght of acid and base

Acid-Base Pairs, Strength of Acids and Bases, and pH

Conjugate Acid-Base Pairs	Strong and Weak Acids and Bases	The Acid Dissociation Equilibrium Constant, Ka
The Relative Strengths of Conjugate Acid-base Pairs	Comparing Relative Strengths of Pairs of Acids and Bases	The Leveling Effect of Water
The Advantages of the Brnsted Definition	pH As A Measure of the Concentration of the H3O+ Ion	Table: pH of Common Acids and Bases

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Conjugate Acid-Base Pairs

Acids and bases exist as conjugate acid-base pairs. The term conjugate comes from the Latin stems meaning "joined together" and refers to things that are joined, particularly in pairs, such as Brnsted acids and bases.

Every time a Brnsted acid acts as an H⁺-ion donor, it forms a conjugate base. Imagine a generic acid, HA. When this acid donates an H⁺ ion to water, one product of the reaction is the A^- ion, which is a hydrogen-ion acceptor, or Brnsted base.

HA	+	H2O		H3O+	+	A-
acid						base

Conversely, every time a base gains an H⁺ ion, the product is a Brnsted acid, HA.

A-	+	H2O		HA	+	OH-
base				acid

Acids and bases in the Brnsted model therefore exist as conjugate pairs whose formulas are related by the gain or loss of a hydrogen ion.

Our use of the symbols HA and A^- for a conjugate acid-base pair does not mean that all acids are neutral molecules or that all bases are negative ions. It signifies only that the acid contains an H⁺ ion that isn't present in the conjugate base. Brnsted acids or bases can be neutral molecules, positive ions, or negative ions. Various Brnsted acids and their conjugate bases are given in the table below.

Typical Brnsted Acids and Their Conjugate Bases

Acid			Base
H3O+			H2O
H2O			OH-
OH-			O2-
HCl			Cl-
H2SO4			HSO4-
HSO4-			SO42-
NH4+			NH3
NH3			NH2-

A compound can be both a Brnsted acid and a Brnsted base. H₂O, OH^-, HSO₄^-, and NH₃, for example, can be found in both columns in the table above. Water is the perfect example of this behavior because it simultaneously acts as an acid and a base when it forms the H₃O⁺ and OH^- ions.

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Strong and Weak Acids and Bases

Many hardware stores sell "muriatic acid" a 6 M solution of hydrochloric acid HCl(aq) to clean bricks and concrete. Grocery stores sell vinegar, which is a 1 M solution of acetic acid: CH₃CO₂H. Although both substances are acids, you wouldn't use muriatic acid in salad dressing, and vinegar is ineffective in cleaning bricks or concrete.

The difference between the two is that muriatic acid is a strong acid and vinegar is a weak acid. Muriatic acid is strong because it is very good at transferring an H⁺ ion to a water molecule. In a 6 M solution of hydrochloric acid, 99.996% of the HCl molecules react with water to form H₃O⁺ and Cl^- ions.

HCl(aq) + H₂O(l) H₃O⁺(aq) + Cl^-(aq)

Vinegar is a weak acid because it is not very good at transferring H⁺ ions to water. In a 1 M solution, less than 0.4% of the CH₃CO₂H molecules react with water to form H₃O⁺ and CH₃CO₂^- ions.

CH₃CO₂H(aq) + H₂O(l) H₃O⁺(aq) + CH₃CO₂^-(aq)

More than 99.6% of the acetic acid molecules remain intact.

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The Acid Dissociation Equilibrium Constant, K_a

The relative strengths of acids is often described in terms of an acid-dissociation equilibrium constant, K_a. To understand the nature of this equilibrium constant, let's assume that the reaction between an acid and water can be represented by the following generic equation.

HA(aq)

+

H2O(l)

H3O+(aq)

+ A-(aq)

In other words,some of the HA molecules react to form H₃O⁺ and A^- ions, as shown in the figure below.

By convention, the concentrations of these ions in units of moles per liter are represented by the symbols [H₃O⁺] and [A^-]. The concentration of the HA molecules that remain in solution is represented by the symbol [HA].

The value of K_a for acid is calculated from the following equation.

When a strong acid dissolves in water, the acid reacts extensively with water to form H₃O⁺ and A^- ions. (Only a small residual concentration of the HA molecules remains in solution.) The product of the concentrations of the H₃O⁺ and A^- ions is therefore much larger than the concentration of the HA molecules, so K_a for a strong acid is greater than 1.

Example: Hydrochloric acid has a K_a of roughly 1 x 10⁶.

Weak acids, on the other hand, react only slightly with water. The product of the concentrations of the H₃O⁺ and A^- ions is therefore smaller than the concentration of the residual HA molecules. As a result, K_a for a weak acid is less than 1.

Example: Acetic acid has a K_a of only 1.8 x 10^-5.

K_a can therefore be used to distinguish between strong acids and weak acids.

Strong acids:		Ka > 1
Weak acids:		Ka < 1

The Relative Strengths of Conjugate Acid-base Pairs

• Strong acids have a weak conjugate base.

Example: HCl is a strong acid. If HCl is a strong acid, it must be a good proton donor. HCl can only be a good proton donor, however, if the Cl^- ion is a poor proton acceptor. Thus, the Cl^- ion must be a weak base.

HCl(g)	+	H2O(l)		H3O+(aq)	+	Cl-(aq)
Strong acid						Weak base

• Strong bases have a weak conjugate acid.

Example: Let's consider the relationship between the strength of the ammonium (NH₄⁺) and its conjugate base, ammonia (NH₃). The NH₄⁺ ion is a weak acid because ammonia is a reasonably good base.

NH4+(aq)	+	H2O(l)		H3O+(aq)	+	NH3(aq)
Weak acid						Good base

                                    Comparing Relative Strengths of Pairs of Acids and Bases

The value of K_a for an acid can be used to decide whether it is a strong acid or a weak acid, in an absolute sense. It can also be used l to compare the relative strengths of a pair of acids.

Example: Consider HCl and the H₃O⁺ ion.

HCl		Ka = 1 x 106
H3O+		Ka = 55

These K_a values suggest that both are strong acids, but HCl is a stronger acid than the H₃O⁺ ion.

A high proportion of the HCl molecules in an aqueous solution reacts with water to form H₃O⁺ and Cl^- ions. The Brnsted theory suggests that every acid-base reaction converts an acid into its conjugate base and a base into its conjugate acid.

There are two acids and two bases in this reaction. The stronger acid, however, is on the left side of the equation.

HCl(g)	+	H2O(l)		H3O+(aq)	+ Cl-(aq)
stronger acid				weaker acid

The general rules suggest that the stronger of a pair of acids must form the weaker of a pair of conjugate bases. The fact that HCl is a stronger acid than the H₃O⁺ ion implies that the Cl^- ion is a weaker base than water.

Acid strength:		HCl > H3O+
Base strength:		Cl- < H2O

Thus, the equation for the reaction between HCl and water can be written as follows.

HCl(g)	+	H2O(l)		H3O+(aq)	+	Cl-(aq)
stronger acid		stronger base		weaker acid		weaker base

It isn't surprising that 99.996% of the HCl molecules in a 6 M solution react with water to give H₃O⁺ ions and Cl^- ions. The stronger of a pair of acids should react with the stronger of a pair of bases to form a weaker acid and a weaker base.

Let's look at the relative strengths of acetic acid and the H₃O⁺ ion.

CH3CO2H		Ka = 1.8 x 10-5
H3O+		Ka = 55

The values of K_a for these acids suggest that acetic acid is a much weaker acid than the H₃O⁺ ion, which explains why acetic acid is a weak acid in water. Once again, the reaction between the acid and water must convert the acid into its conjugate base and the base into its conjugate acid.

But this time, the stronger acid and the stronger base are on the right side of the equation.

CH3CO2H(aq)	+	H2O(l)		H3O+(aq)	+	CH3CO2-(aq)
weaker acid		weaker base		stronger acid		stronger base

As a result, only a few of the CH₃CO₂H molecules actually donate an H⁺ ion to a water molecule to form the H₃O⁺ and CH₃CO₂^- ions.

The magnitude of K_a can also be used to explain why some compounds that qualify as Brnsted acids or bases don't act like acids or bases when they dissolve in water. When the value of K_a for an acid is relatively large, the acid reacts with water until essentially all of the acid molecules have been consumed. Sulfuric acid (K_a = 1 x 10³), for example, reacts with water until 99.9% of the H₂SO₄ molecules in a 1 M solution have lost a proton to form HSO₄^- ions.

H₂SO₄(aq) + H₂O(l) H₃O⁺(aq) + HSO₄^-(aq)

As K_a becomes smaller, the extent to which the acid reacts with water decreases.

As long as K_a for the acid is significantly larger than the value of K_a for water, the acid will ionize to some extent. Acetic acid, for example, reacts to some extent with water to form H₃O⁺ and CH₃CO₂^-, or acetate, ions.

CH₃CO₂H(aq) + H₂O(l) H₃O⁺(aq) + CH₃CO₂^-(aq)

As the K_a value for the acid approaches the K_a for water, the compound becomes more like water in its acidity. Although it is still a Brnsted acid, it is so weak that we may be unable to detect this acidity in aqueous solution.

Some potential Brnsted acids are so weak that their K_a values are smaller than water's. Ammonia, for example, has a K_a of only 1 x 10^-33. Although NH₃ can be a Brnsted acid, because it has the potential to act as a hydrogen-ion donor, there is no evidence of this acidity when it dissolves in water.

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The Leveling Effect of Water

All strong acids and bases seem to have the same strength when dissolved in water, regardless of the value of K_a. This phenomenon is known as the leveling effect of water the tendency of water to limit the strength of strong acids and bases. We can explain this by noting that strong acids react extensively with water to form the H₃O⁺ ion. More than 99% of the HCl molecules in hydrochloric acid react with water to form H₃O⁺ and Cl^- ions, for example,

HCl(g) + H₂O(l) H₃O⁺(aq) + Cl^-(aq)

and more than 99% of the H₂SO₄ molecules in a 1 M solution react with water to form H₃O⁺ ions and HSO₄^- ions.

H₂SO₄(aq) + H₂O(l) H₃O⁺(aq) + HSO₄^-(aq)

Thus, the strength of strong acids is limited by the strength of the acid (H₃O⁺) formed when water molecules pick up an H⁺ ion.

A similar phenomenon occurs in solutions of strong bases. Strong bases react quantitatively with water to form the OH^- ion. Once this happens, the solution cannot become any more basic. The strength of strong bases is limited by the strength of the base (OH^-) formed when water molecules lose an H⁺ ion.

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The Advantages of the Brnsted Definition

The Brnsted definition of acids and bases offers many advantages over the Arrhenius and operational definitions.

• It expands the list of potential acids to include positive and negative ions, as well as neutral molecules.
• It expands the list of bases to include any molecule or ion with at least one pair of nonbonding valence electrons.
• It explains the role of water in acid-base reactions: Water accepts H⁺ ions from acids to form the H₃O⁺ ion.
• It can be expanded to include solvents other than water and reactions that occur in the gas or solid phases.
• It links acids and bases into conjugate acid-base pairs.
• It can explain the relationship between the strengths of an acid and its conjugate base.
• It can explain differences in the relative strengths of a pair of acids or a pair of bases.
• It can explain the leveling effect of water the fact that strong acids and bases all have the same strength when dissolved in water.

Because of these advantages, whenever chemists use the words acid or base without any further description, they are referring to a Brnsted acid or a Brnsted base.

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pH As A Measure of the Concentration of the H₃O⁺ Ion

Pure water is both a weak acid and a weak base. By itself, water forms only a very small number of the H₃O⁺ and OH^- ions that characterize aqueous solutions of stronger acids and bases.

H2O(l)	+	H2O(l)		H3O+(aq)	+	OH-(aq)
base		acid		acid		base

The concentrations of the H₃O⁺ and OH^- ions in water can be determined by carefully measuring the ability of water to conduct an electric current. At 25^oC, the concentrations of these ions in pure water is 1.0 x 10^-7 moles per liter.

[H₃O⁺] = [OH^-] = 1.0 x 10^-7 M (at 25C)

When we add a strong acid to water, the concentration of the H₃O⁺ ion increases.

HCl(aq) + H₂O(l) H₃O⁺(aq) + Cl^-(aq)

At the same time, the OH^- ion concentration decreases because the H₃O⁺ ions produced in this reaction neutralize some of the OH^- ions in water.

H₃O⁺(aq) + OH^-(aq) 2 H₂O(l)

The product of the concentrations of the H₃O⁺ and OH^- ions is constant, no matter how much acid or base is added to water. In pure water at 25^oC, the product of the concentration of these ions is 1.0 x 10^-14.

[H₃O⁺][OH^-] = 1.0 x 10^-14

The range of concentrations of the H₃O⁺ and OH^- ions in aqueous solution is so large that it is difficult to work with. In 1909 the Danish biochemist S. P. L. Sorenson suggested reporting the concentration of the H₃O⁺ ion on a logarithmic scale, which he named the pH scale. Because the H₃O⁺ ion concentration in water is almost always smaller than 1, the log of these concentrations is a negative number. To avoid having to constantly work with negative numbers, Sorenson defined pH as the negative of the log of the H₃O⁺ ion concentration.

pH = -log [H₃O⁺]

the concept of pH compresses the range of H₃O⁺ ion concentrations into a scale that is much easier to handle. As the H₃O⁺ ion concentration decreases from roughly 10⁰ to 10^-14, the pH of the solution increases from 0 to 14.

If the concentration of the H₃O⁺ ion in pure water at 25^oC is 1.0 x 10^-7 M, the pH of pure water is 7.

pH = -log [H₃O⁺] = -log (1.0 x 10^-7) = 7

When the pH of a solution is less than 7, the solution is acidic. When the pH is more than 7, the solution is basic.

Acidic:		pH < 7
Basic:		pH > 7