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Reactions in Solution

Reactions in Solution. The most important substance on earth is water. In chemistry, water is necessary for many reactions to take place. Table salt (NaCl) when put into water dissolves into its ions, Na + and Cl - .   Water is the solvent and NaCl, is the solute .

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Reactions in Solution

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  1. Reactions in Solution The most important substance on earth is water. In chemistry, water is necessary for many reactions to take place. Table salt (NaCl) when put into water dissolves into its ions, Na+ and Cl-.   Water is the solvent and NaCl, is the solute. The mixture is an aqueous solution. The Water Molecule allows many substances to be dissolved in them. One side of the water molecule is negatively charged, and the other side is positively charged. Water is a polar molecule.

  2. Electricity and Solutions • A useful characteristic of solutions is the ability to conduct electricity. To determine if a solution has the ability to conduct electricity, an electrical conductivity apparatus is used. An electrical conductivity apparatus is basically a battery and light bulb setup which lights up when electricity is conducted through the solution. • Electrolytes are substances that produce ions upon dissolving. There are two ways to provide these mobile ions for conducting purposes. • Dissociation of Ionic Compounds • Ionisation of Polar Covalent Molecular Substances

  3. Electrolytes 1. Dissociation of Ionic Compounds: Ionic compounds are made of cations and anions. These ions are locked into position in their crystal structure and are not able to move around. In water, the water molecules, are attracted to the ions. The ions are said to be dissociated, and able to carry electrical particles to conduct current.   K3PO4 + H2O ----> 3K+ (aq) + PO4-3 (aq) Such substances are said to be electrolytes. Salts that are completely soluble in water are strong electrolytes. Salts that are slightly soluble are weak electrolytes at best. The strength of an electrolyte is measured by its ability to conduct electrical current.

  4. Electrolytes 2. Ionisation of Polar Covalent Molecular Substances Polar molecular substances are substances whose atoms are covalently bonded. Each molecule has a net molecular dipole moment and thus a positive and a negative end. Polar water molecules can line up around polar molecule. If this dipole-dipole interaction can overcome the dissociation energy of a bond the molecule will fragment with bonding electrons going with the most electronegative atom in the broken bond, creating ions. (Electronegativity is the electron attracting ability of an atom) Such polar molecular compounds are called electrolytes. An example of a strong electrolyte is any of the strong acids, such as HBr. H-Br + H2O ----> H3O+ (aq) + Br- (aq)

  5. Electrolytes Some polar molecular substances have such strong covalent bonding that water is only able to overcome these stronger dissociation energies in a portion of the molecules. For example,a weak acid such as ethanoic acid, CH3-COO-H, dissolves in water with only a small percentage of the molecules being ionized.  CH3COOH + H2O H3O++CH3COO- Non-electrolytes are substances that do not produce ions when they dissolve. This results when polar molecular substances are large enough and their covalent bonding is strong enough so that water is not able to break any of the covalent bonds during the solvation process. As a result, the neutral molecules are solvated (separated by solvent water molecules) without any ionization occurring.

  6. Acids and Bases • The properties of acids include the following: • Taste sour (but don't taste them!!) • Their water solutions conduct electrical current (electrolytes) • They react with bases to form salts and water • Turns Blue Litmus Paper to Red • The properties of bases include the following: • Have a slippery feel between the fingers • Have a bitter taste (but don't taste them!!) • React with acids to form salts and water • Turns Red Litmus Blue • Their water solutions conduct electrical current (electrolytes)

  7. Acids and Bases Arrhenius in 1884 discovered that acids give off H+ ions and allow for a good flow of electricity through a solution. Arrhenius also discovered that bases give off OH- ions and OH- ions also allow for a good flow of electricity through the solution. Traditionally Professor Arrhenius defined: Acid released Hydrogen ion (as Hydronium ions, H3O+) in water solution. Base produced Hydroxide ion in water solution. The limitations on these definitions were: 1. The need for water 2. The need for a protic acid 3. The need for Hydroxide bases

  8. Bronsted/Lowry acids and bases Bronsted and Lowry defined these two terms the following: Acid-Proton donor Base-Proton acceptor These definitions are not as restrictive as Arrhenius’ definitions. • No need for water although it can be present, it need not be. • Bases do not have to be Hydroxide compounds. However, one restriction still remaining is the need for a protic acid. Each Bronsted acid is coupled to a conjugate base to constitute a CONJUGATE ACID-BASE PAIR CH3COOH + H2O H3O++CH3COO-

  9. Lewis Acids and Bases G.N. Lewis defined these in an even less restrictive manner: Acid- Electron pair acceptor Base- Electron pair donor In this set of definitions there is no longer a need for a protic acid. In other words only electron exchange must occur. These definition sets are NOT contradictory. A Proton donor is the same as an electron acceptor. A Proton acceptor is the same as an electron donor. Also the first set of definitions are less inclusive so that all of the Arrenhius acids are found under the Bronsted definition but not all Bronsted acids will be Arrenhius acids. All Arrenhius and Bronsted acids will be under the Lewis definition but not all Lewis acids will be Bronsted or Arrenhius acids.

  10. Acid and Base Strength Strong acids (memorise) dissociate completely in water HClO4, HCl, HBr, HI, HNO3 and H2SO4 Strong bases are the metal hydroxides of Group 1 and 2 E.g. LiOH, NaOH, KOH, Ba(OH)2, Mg(OH)2 etc Weak acids and bases are not completely ionised in solution CH3COOH + H2O H3O++CH3COO- Ka is an equilibrium constant called the acid dissociation constant

  11. Acid and Base Strength :NH3 + H2O NH4++OH- (a molecular base) The magnitude of the Ka or Kb, using water as a common proton donor/acceptor, determines the strength of the acid or base In general (for acids) HA + H2O H3O++A- Water is AMPHOTERIC. It can act as an acid or a base

  12. Acid and Base Strength Ka ~1010 1x10-2 6.8x10-4 1.75x10-5 9.5x10-8 5.7x10-10 4.7x10-11 1.8x10-16 Stronger Acid Levelling Effect Each acid will transfer a proton to a base below it in a mixed solution HClO4 ClO4- H2SO4 HSO4- HCl Cl- H3O+ H2O HSO4- SO42- HF F- CH3COOH CH3COO- H2S HS- NH4+ NH3 HCO3- CO32- H2O OH- Stronger Base

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