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Net Ionic Equation

Net Ionic Equation. Net ionic equations are used to show only the chemicals and ions involved in a chemical reaction in order to simplify information about a reaction. The ions that are not involved in the reaction are called spectator ions and are removed from the reaction.

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Net Ionic Equation

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  1. Net Ionic Equation • Net ionic equations are used to show only the chemicals and ions involved in a chemical reaction in order to simplify information about a reaction. • The ions that are not involved in the reaction are called spectator ions and are removed from the reaction. NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(l) Na+ + OH- + H + +Cl-  Na + +Cl- + H2O(l) OH- + H +  H2O(l)

  2. Acids • Arrhenius definiton of acids • A substance that produces H+ ions • Brønsted-Lowry definition of acids • An acid is a substance that can transfer a hydrogen ion (H+), also called a proton, or Hydronium (H3O+) to another substance. It is often called a “proton donor.” Notice the slight difference in wording, but essentially they mean the same thing… Ex: HCl  H+ + Cl- HNO3  H+ + NO3- HCH3CO2  H+ + CH3CO2-

  3. The Hydronium Ion • H3O+ is called the hydronium ion. • It is another way to understand how hydrogen ions, H+, move around in a solution. • Hydronium ions are created when an acid is dissolved into water. HCl(g) + H2O(l)  H3O+(aq) + Cl-(aq) HNO3 (s) + H2O(l)  H3O+(aq) + NO3-(aq)

  4. Naming Acids • Binary acids – usually acids that are made of hydrogen and one other substance. • Name these by using hydro- for the hydrogen and the name of the second element with ending of –ic, then following it up with the word, acid. Ex: HCl – hydrochloric acid Usually binary means 2, but exceptions are made when it is an acid of a polyatomic ion that does not contain oxygen Ex: HCN – hydrocyanic acid

  5. Naming Acids • Oxyacids – acids of polyatomic ions that contain oxygen. • Name these by noting the name of the oxyanion and following it with the word acid. If it ends in –ate, then the acid is named by the oxyanion and the suffix of –ic. If it ends in –ite, then the acid is named by the oxyanion and the suffix of –ous. Ex: HBrO3 – bromic acid HBrO2 – bromous acid

  6. What we know about acids • Acids are well known as substances that are corrosive, sour tasting, turn litmus paper a red color, and react with bases and metals to form salts. • Common acids that you deal with every day include: Acetic acid, carbonic acid, citric acid, maleic acid, fumaric acid, ascorbic acid, acetylsalicylic acid, amino acids, hydrochloric acid

  7. Bases • Arrhenius definition of bases • A substance that produces OH- ions • Brønsted-Lowry definition of bases: • a substance that can accept a hydrogen ion (H+) from a proton donor. It is often called a “proton acceptor.” Notice that these two definitions are very different. Ex: NaOH + H+ Na+ + H2O NH3 + H+  NH4+ Mg(OH)2 + 2H+  Mg+ + 2H2O NaHCO3 + H+  Na+ + H2CO3 There are no naming rules for bases

  8. What we know about bases • Bases are commonly known as substances that are corrosive, have a slippery feel, taste bitter, turn litmus paper blue, and react with acids to form salts. • Basic solutions are often called “alkaline solutions.” • Common bases that you may have dealt with include: Ammonia, sodium bicarbonate, sodium hydroxide, magnesium hydroxide, calcium hydroxide, calcium carbonate (coffee is acidic… you’ll notice this when you drink bad coffee)

  9. Strong vs Weak • Acid and base strengths are determined by their abilities to dissociate (ions breaking apart when dissolved) • Strong acids and bases completely dissociate. Not only are these powerful chemicals, but are also extremely good electrolytes (ie: they conduct electricity extremely well) • Weak acids and bases do not completely dissociate meaning that fewer ions enter into solution. A noticeable feature of weak acids and bases are that they are poor electrolytes (ie: they do not conduct electricity well)

  10. pH • pH is defined as the measurement of the hydrogen ion concentration present in a solution. • Mathematically, it stands for: pH = -log [H+] • The typical pH scale ranges from 0 to 14 • 0 - 6.9 Acidic range. (0 = strongest) • 7 Neutral (not acidic or basic) • 7.1 – 14 Basic range. (14 = strongest base) • The pH of some common materials: ACIDSBASES HCl -1 to 0.1 Human blood 7.35 to 7.45 Vinegar 2.2 Sea water 7.36 to 8.21 Lemon Juice 2.2 to 2.4 Eggs 7.6 to 8 Apple juice 2.9 to 3.3 Baking soda 8.0 Acid rain 5.2 Borax 9.2 Human saliva 6.3 to 6.6 Soda lye 13.0 to 15.0 Distilled water 7.0

  11. Water has neutral pH • Water has a natural ability to dissociate into H3O+ and OH- ions. H2O + H2O  H3O+ + OH- • Kw, the self-ionization constant of water, says there will always be a natural balance of H3O+ to OH- ions in water. Kw= [OH-] x [H3O+] = 1.00 x 10-14 • Pure water has 1.00 x 10-7M OH- and 1.00 x 10-7M H3O+ (do the math… why is the pH scale 0-14? Why is 7 neutral?)

  12. Acid Dissociation Constant • Strength of acids are determined by numerical values called the acid dissociation constant (Ka) • This basically tells you how well an acid will dissociate in water • The bigger the number, the stronger the acid • The smaller the number, the weaker the acid

  13. Acid and Conjugate Base • We can show the dissociation of an acid and write it as an ionic reaction H2S  H+ + HS- Ka = 8.9x10-8 • Notice that in this equation, H2S donates a H+ ion, so is thus defined as an acid. • The reverse reaction can also occur H+ + HS- H2S Kb = • Now notice in this next reaction, HS- accepts a H+ ion so it fits the definition of a base. We can say that HS- is the conjugate base of H2S

  14. Base Dissociation Constant • Similarly, base strength is characterized by how well bases will gain H+ ions. • This is characterized numerically by the base dissociation constant Kb • The larger the value of Kb, the stronger the base • The smaller the value of Kb, the weaker the base.

  15. Base and Conjugate Acid • We can show the dissociation of a base and write it as an ionic reaction NH3 + H2O  NH4+ + OH- Kb = 2.5x10-5 • Notice that in this equation, NH3accepts a H+ ion, so is thus defined as a base. • The reverse reaction can also occur NH4+ + OH- NH3 + H2O Ka = Now notice in this next reaction, NH4+ donates a H+ ion so it fits the definition of an acid. We can say that NH4+ is the conjugate acid of NH3

  16. Practice • Balance the following equation: NaOH(aq) + H2SO4(aq)  Na2SO4 (aq) + H2O(l) • What are the acid, base, conjugate acid, conjugate base of the above reaction? • Write the net ionic equation for the above reaction • How does this net ionic equation exemplify the reaction of acids with bases?

  17. NaOH + H2SO4Na2SO4 + H2O How do you identify the acid in a chemical equation? • Look at the reactants • Which reactant has Hydrogen? • Which reactant has Hydrogens to be given away? NaOH or H2SO4 How do you identify the base in a chemical equation? • Look at the reactants • Which reactant has hydroxide? • Which reactant may want to absorb hydrogen? NaOH or H2SO4

  18. How do you identify the conjugate acid in a chemical equation? • Look at the products. • Which product has Hydrogen? • Which product has Hydrogens that can be removed to recreate the original base? Na2SO4 or H2O • Look at the products. • Which product has looks like it’s missing Hydrogens? • Which product will recreate the original acid if we give hydrogens back to it? Na2SO4 or H2O

  19. Neutralization • Neutralization occurs when an acid reacts with a base to produce a salt and water. • A complete neutralization occurs when all the available H+ ions from the acid are reacted, and no excess base is present. In other words, the amount of acid = the amount of base in the reaction solution. This is also known as the equivalence point • An acid-base indicator is a substance whose color is affected by acidic and basic solutions. We can use indicators to monitor pH changes in an acid-base reaction.

  20. Titration • Titration is a technique that uses acid-base neutralization to determine the amount of a substance in solution as it reacts with a known amount of acid or base. • Ex: I have a 10mL solution of NaOH of unknown concentration. It reacts completely with 35mL of 1.00M HCl solution. What is the molarity of the NaOH solution?

  21. Monoprotic vs Polyprotic • Many acids are simple combinations of one H+ ion and one -1 ion. With these, the acid can donate no more than one H+ to another substance. These are called monoprotic acids • Other acids have more than one H+ ion bonded to the negative ion (meaning -2 or -3 anion). These can give up one H+ or more than one H+ ions to another substance. These are called polyprotic acids.

  22. Polyprotic acid dissociation • Dissociation – process by which ionic compounds separate into smaller ions • Step 1 H3PO4 + H2O  H2PO4- + H3O+ • Step 2 H2PO4- + H2O  HPO42- + H3O+ • Step 3 HPO42- + H2O PO43- + H3O+ complete dissociation • Net Reaction H3PO4 + 3H2O  PO43- + 3H3O+

  23. Acid-Base Reactions • When acids react with metals, metal salts and hydrogen gas are formed. • 2HCl + Mg  MgCl2 + H2 • H2 SO4 + 2Na  Na2 SO4 + H2 • In general: HX + M  MX + H2 • When acids and bases react, we simply observe a transfer of a H+ ion • HNO3 + NaOH  NaNO3 + H2O • HCH3CO2 + NaHCO3 NaCH3CO2 + H2CO3 • When we see a substance that can potentially act as an acid AND a base, it is called amphoteric. • H2O + H2O  H3O+ + OH-

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