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Chapter 13
 

Chapter 13

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    Chapter 13 Chapter 13 Presentation Transcript

    • Chapter 13 Solutions
    • What is the solution?
      • A solution is just a Homogeneous Mixture!
      • One material (solute) is completely dissolved in another (solvent)
    • How do Solutions Work?
      • Compounds are either polar or non-polar.
      • A polar solvent only will dissolve a polar solute.
      • A non-polar solvent will only dissolve a non-polar solute
      • “ Like dissolves like”
    • Why “Like dissolves like”?
      • Polar compounds will hang tightly to their own kind (like magnets).
      • They will not surround (as a solvent) or allow themselves to be surrounded (as a solute) by non-polar compounds
    • Concentrate on the Solution
      • Concentration is a measure of how much solute is present in the solution.
        • Qualitative: Described without numbers
        • Quantitative: Described WITH numbers.
    • Qualitative Concentrations
      • Unsaturated: The solvent can still dissolve more solute.
      • Saturated: The solution contains the maximum amount of solute.
      • Supersaturated : The solution has dissolved more than the normal maximum amount of solute
        • VERY unstable!
    • Chapter 13 Separating Mixtures
    • Separating Mixtures
      • Mixtures are separated based on their physical properties.
      • Mixtures may be separated by decanting, pouring off the liquid, with a centrifuge, filtering, or evaporation
    • Mixtures
      • If the boiling points of the components are different , distillation can separate them based on their boiling points.
      • As one component reaches its boiling point, it evaporates from the mixture and is allowed to cool and condense. This is called a distillate.
        • This process continues until all the desired components have been separated from the mixture.
    • Chapter 13 Concentration and Molarity
    • Concentration
      • In a solution, the solute is distributed evenly throughout the solvent. This means that any part of a solution has the same ratio of solute to solvent as any other part of the solution.
        • This ratio is the concentration of the solution.
        • The concentration is the amount of a particular substance in a given quantity of a solution
    •  
    • Chapter 13 Physical Properties of Solutions
    • Electrical Conductivity
      • Some substances conduct electricity and some cannot.
      • The conductivity of a substance is described as its ability to conduct an electric current.
      • The conductivity of a substance depends on whether it contains charged particles, and these particles must be able to move.
    • Electrical Conductivity
      • Electrons move freely within a metal, thus allowing it to conduct electricity.
      • An aqueous solution of ionic compounds such as NaCl contains charged ions, which can move about. Solutions of ionic compounds conduct electricity.
      • Pure water does not conduct electricity.
    • Electrical Conductivity
      • An electrolyte is a substance that dissolves in a liquid solvent and provides ions that conduct electricity.
      • Strong electrolytes completely dissociate into ions and conduct electricity well.
      • Weak electrolytes provide few ions in solution.
      • Covalent compounds may be strong electrolytes, weak electrolytes, or nonconductors.
    •  
    • Electrical Conductivity
      • The extent to which electrolytes dissociate into ions is indicated by the conductivity of their solutions.
      • The sugar sucrose does not ionize at all in solution.
        • It is a nonelectrolyte and does not conduct electricity.
      • A nonelectrolyte is a liquid or solid substance that does not allow the flow of an electric current, either in solution or in its pure state, such as water or sucrose.
    •  
    • Colligative Properties
      • The physical properties of water are changed when substances dissolve in it.
      • Salt can be added to icy sidewalks to melt the ice.
        • The salt actually lowers the freezing point of water.
        • Ice is able to melt at a lower temperature than it normally would.
      • This change is called freezing-point depression.
    •  
    • Colligative Properties
      • Nonvolatile solutes such as salt also increase the boiling point of a solvent.
      • This change is called boiling-point elevation.
        • For example, glycol in a car’s radiator increases the boiling point of water in the radiator, which prevents overheating.
        • It also lowers the freezing point, preventing freezing in cold weather.
    •  
    • Colligative Properties
      • Any physical effect of the solute on the solvent is a colligative property.
      • A colligative property is a property of a substance or system that is determined by the number of particles present in the system but independent of the properties of the particles themselves.
    • Colligative Properties
      • Any solute, whether an electrolyte or a nonelectrolyte, contributes to the colligative properties of the solvent.
      • The degree of the effect depends on the concentration of solute particles (either molecules or ions) in a certain mass of solvent.
        • The greater the particle concentration is, the greater the boiling-point elevation or the freezing-point depression is.
    •  
    • Chapter 14 Reversible Reactions and Equilibriums
    • Complete Reactions
      • If enough oxygen gas is provided for the following reaction, almost all of the sulfur will react:
      • S 8 + 8O 2 -> 8SO 2
      • Reactions such as this one, in which almost all of the reactants react, are called completion reactions.
      • In other reactions, called reversible reactions , the products can re-form reactants.
    • Reversible Reactions Solid calcium sulfate, the product, can break down to make calcium ions and sulfate ions in a reaction that is the reverse of the previous one.
    • Reversible Reactions
      • The reactions occur at the same rate after the initial mixing of CaCl 2 and Na 2 SO 4 .
      • The amounts of the products and reactants do not change.
      • Chemical equilibrium is a state of balance. the rate of a forward reaction equals the rate of the reverse reaction.
    •  
    •  
    • Equilibrium
      • In equilibrium, an atom may change from being part of the products to part of the reactants many times.
      • But the overall concentrations of products and reactants stay the same.
      • For chemical equilibrium to be maintained, the rates of the forward and reverse reactions must be equal.
    • Chapter 14 Systems and Stress
    • STRESS
    • Le Chatelier’s Principle
      • Stress is another word for something that causes a change in a system at equilibrium.
      • Chemical equilibrium can be disturbed by a stress, but the system soon reaches a new equilibrium.
      • Le Châtelier’s principle states that when a system at equilibrium is disturbed, the system adjusts in a way to reduce the change.
    • Le Chatelier’s Principle
      • Chemical equilibria respond to three kinds of stress:
        • changes in the concentrations of reactants or products
        • changes in temperature
        • changes in pressure
      • When a stress is first applied to a system, equilibrium is disturbed and the rates of the forward and backward reactions are no longer equal.
    • Le Chatelier’s Principle
      • The system responds to the stress by forming more products or by forming more reactants.
      • A new chemical equilibrium is reached when enough reactants or products form.
      • At this point, the rates of the forward and backward reactions are equal again.
    •  
    • Concentrate
      • Increase the amount of reactants, the equilibrium will shift towards products.
      • Increase the amount of products, the equilibrium will shift towards the reactants.
    • Getting Hot!
      • If the forward reaction is exothermic, cooling the system will force the reaction forward.
      • If the forward reaction is endothermic, heating the system will force the reaction forward.
    • Feel the Pressure
      • Increasing pressure will shift the equilibrium to the side with less atoms.
    •  
    • Practical Le Chatelier
      • The chemical industry makes use of Le Châtelier’s principle in the synthesis of ammonia by the Haber Process.
      • High pressure is used to drive the following equilibrium to the right .
    • Chapter 15 Acids and Bases
    • Acids
      • These substances can recognized as acidic by their tart, sour, or sharp taste.
      • These substances contain dissolved compounds that chemists describe as acids .
      • Many other acids, such as sulfuric acid or hydrochloric acid, are highly caustic and should not be put to the taste test.
    • Acids
      • Acids are electrolytes, so their solutions in water are conductors of electric current.
      • Like other electrolytes, hydrogen chloride dissociates to produce ions.
      • HCl( g ) + H2O( l )  H3O+( aq ) + Cl−( aq )
      • The hydronium ion, H3O+, is able to transfer charge through aqueous solutions much faster than other ions do.
    •  
    • Acids
      • Another property shared by aqueous solutions of acids is that they react with many metals.
      • All metals that are above hydrogen in the activity series react with acids to produce hydrogen gas.
      • 2H3O+( aq ) + Zn( s )  2H2O( l ) + H2( g ) + Zn2+( aq )
    • Acids
      • Some electrolytes are strong and others are weak, depending on whether they dissociate completely or partially.
      • When a weak acid is dissolved in water, only a small fraction of its molecules are ionized at any given time.
    • Bases
      • Bases are another class of electrolytes. Unlike acids, which are usually liquids or gases, many common bases are solids.
      • Solutions of bases are slippery to the touch, but touching bases is an unsafe way to identify them.
        • The slippery feel comes about because bases react with oils in your skin, converting them into soaps.
    • Bases
      • Some bases, such as magnesium hydroxide, Mg(OH)2, are almost insoluble in water.
      • Other bases, such as potassium hydroxide, are so soluble that they will absorb water vapor from the air and dissolve in the water.
      • A base that is very soluble in water is called an alkali, a term that describes the Group 1 metals of the periodic table.
    • Bases
      • The alkali metals react with water to form hydroxides that are water-soluble alkalis. These are called basic or Alkaline.
      • Just as acids may be strong or weak depending on whether they ionize completely or reach an equilibrium between ionized and un-ionized forms, bases are also classified as strong or weak.
    • Bases
      • Both strong and weak bases generate hydroxide ions when they dissolve in water.
      • Many oxides, carbonates, and phosphates are bases, too.
    • Classifications
      • Arrhenius acid – Produces Hydronium ions
      • Arrhenius Base – Produces Hydroxide ions
    • Classification
      • Brønsted-Lowry acids – Donate Protons
      • Brønsted-Lowry acids – Accept Protons
    • Conjugate Acids and Bases Look again at the equation for the reversible reaction of ammonia, NH 3 , with water:
          • base acid c. acid c. base
      • Water donates a proton to ammonia, so it is an acid.
      • Ammonia accepts the proton, so it is a base.
    • Amphoteric
      • Some species are both an acid and a base and can both donate and accept protons.
      • Such species are described as amphoteric.
      • Amphoteric describes a substance, such as water, that has the properties of an acid and the properties of a base.
    •  
    • Self-Ionization
      • Water is both an acid and a base. It both gives and receives protons.
      [H3O+] = [OH−] = 1.00 × 10 −7 M
    • Self-Ionization
      • An equilibrium-constant expression relates the concentrations of species involved in an equilibrium.
      • The relationship for the water equilibrium is simply
      • [H3O+][OH−] = Kw
      • The value of Kw can be found from the known concentrations of the hydronium and hydroxide ions in pure water.
      • Kw = (1.00 × 10 −7 )(1.00 × 10 −7 ) = 1.00 × 10 −14
      • The product of these two ion concentrations is always a constant.
    •  
    • Chapter 15 pH, or more math!
    • Meaning of pH
      • When acidity and basicity are exactly balanced such that he numbers of H 3 O + and OH − ions are equal, we say that the solution is neutral.
      • Pure water is neutral because it contains equal amounts of the two ions.
    • Meaning of pH
      • A solution made by dissolving 0.100 mol of NaOH in 1.00 L of water has a hydroxide ion concentration of 0.100 M.
      • The hydronium ion concentration can be calculated using Kw .
    • Meaning of pH
      • In 1909, Danish chemist Søren Sørensen proposed using the negative logarithm of [H 3 O + ] as the index of basicity and acidity.
      • He called this measure the pH or p ower of H ydrogen.
      • pH can be calculated by the following mathematical equation:
      • pH = −log [H 3 O + ] OR [H 3 O + ] = 10 −pH
    • The Meaning of pH, continued Calculating pH from [H 3 O + ], continued
      • The pH equation may be rearranged to calculate the hydronium ion concentration from the pH.
      • [H 3 O + ] = 10 −pH
      • Because pH is related to powers of 10, a change in one pH unit corresponds to a tenfold change in the concentrations of the hydroxide and hydronium ions.
      Section 2 Acidity, Basicity, and pH Chapter 15
    •