Presentation on theme: " electrolytes electrolytes turn litmus red sour taste react with metals to form H 2 gas slippery feel turn litmus blue bitter taste "— Presentation transcript:
electrolytes electrolytes turn litmus red sour taste react with metals to form H 2 gas slippery feel turn litmus blue bitter taste vinegar, soda, apples, citrus fruits ammonia, lye, antacid, baking soda
Arhenius Definition HCl + H 2 O H 3 O + + Cl – In aqueous solutions AcidsIn aqueous solutions Acids form hydronium ions (H 3 O + ) H HHHH H Cl OO – + Acid Produce Hydronium
Arhenius Definition In aqueous solutions BasesIn aqueous solutions Bases form hydroxide ions (OH - ) NH 3 + H 2 O NH 4 + + OH - H H H H H H N NO O – + H H H H Base Produce Hydroxide
Exceptions? This did not encompass all of the compounds that we knew were basic and acidic. Oh what to do?
Bronsted-Lowry Definition HCl (aq) + H 2 O (l) Cl – (aq) + H 3 O + (aq) AcidsBases baseacid are proton (H + ) donors. are proton (H + ) acceptors. conjugate base conjugate acid
Example (Acid) H 2 O + HNO 3 H 3 O + + NO 3 – C onjugate B ase C onjugate A cid A cid B ase Hydronium C onjugate B ase
Example(Base) NH 3 + H 2 O NH 4 + + OH - CACBBA Hydroxide C onjugate A cid
H 2 O…Acid or Base? Amphiprotic: So water is an acid? So water is a base. BOTH A chemical species that can act as EITHER, an acid or a base.
HF HBr HI H3O+H3O+ Give the conjugate base for each of the following: Practice Activity
Partner Up! Partner 1, write the following bases on the back of a cue card. (one acid per card (3 cards). Partner 2, write the conjugate acid for the acid on the other side of the card.
Br – (aq) HSO 4 - (aq) CO 3 2- (aq) HBr (aq) H 2 SO 4(aq) HCO 3 - (aq) Give the conjugate acid for each of the following:
Practice Activity Partner Up! Partner 1, write the following acids on the back of a cue card. (one acid per card (3 cards). Partner 2, write the conjugate base for the acid on the other side of the card.
H 2 SO4 (aq) HCl (aq) HCO 3 - (aq) HSO 4 - (aq) H 2 SO 4(aq) H 2 CO 3 2- (aq) Give the conjugate base for each of the following: SO 4 2- (aq)
Auto-ionization of Water H 2 O + H 2 O H 3 O + + OH - K w = [H 3 O + ][OH - ] = 1.0 10 -14 Square Brackets indicate concentration Most water molecules do not ionize. Only 1 in 556 000 000 water molecules ionize! The other 555 999 999 remain H 2 O!
pH = -log[H 3 O + ] pH Scale 0 7 INCREASING ACIDITY NEUTRAL INCREASING BASICITY 14 pouvoir hydrogène (Fr.) “power of hydrogen”
pH of Common Substances Can go beyond 0 and 14
Super Acids/Super Bases A very concentrated (really hot) strong acid can have a pH below 0! (-0.5, -1) A very concentrated (really hot) strong base can have a pH above 14! (15, 16)
Relationship Between Hydronium Concentration [H 3 O + (aq) ] and pH [H 3 O + (aq) ] = 1.00 x10 -1 [H 3 O + (aq) ] = 1.00 x10 -3 [H 3 O + (aq) ] = 1.00 x10 -5 [H 3 O + (aq) ] = 1.00 x10 -9 [H 3 O + (aq) ] = 1.00 x10 -13 pH = 1 pH = 3 pH = 5 pH = 9 pH = 13 What’s the relationship? Relationship: [H30+(aq)] is related to pH by powers of 10.
Example: What is the pH of 0.050mol/L HNO 3 ? HNO 3(aq) +H 2 O (l) ↔H 3 O + (aq) +NO 3 - (aq) 1 1 1 1 C = 0.050 mol/L pH = -log[H 3 O + (aq) ] pH = -log[0.050] pH =
Example: What is the amount concentration of HBr in a solution that has a pOH of 9.6? pH + pOH = 14 HBr (aq) +H 2 O (l) ↔H 3 O + (aq) +Br - (aq) 11 1 1 C = ? [H 3 O + (aq) ]= 1 x 10 -pH pH = 14 - pOH pH = 14 – 9.6 = 4.4 [H 3 O + (aq) ]=1 x 10 -4.4 [H 3 O + (aq) ]=
Example: What is the amount concentration of HBr in a solution that has a pOH of 9.6? HBr (aq) +H 2 O (l) ↔H 3 O + (aq) +Br - (aq) 11 1 1 C =C = ? N G (N/G)(1/1) C = [HBr (aq) ] =
Why is the concentration of HBr the same as the hydronium ion concentration? Look at the dissociation/ionization equation HBr (aq) H + (aq) + Br - (aq) HBr (aq) + H 2 O (l) H 3 0 + (aq) + Br - (aq) There is a 1:1 relationship between HBr and the ions
Substances that change colour due to the acidity of a solution. Acid – Base Indicators:
They are a weak acid – conjugate base pair that exist in two forms (two different colours) due to presence or lack of a single proton (Hydrogen atom) in the chemical formula.
Because of the complex nature of the chemical formula of each indicator. Abbreviations are usually used to make using indicators less complex. Ex: HLt – Lt - are the acid and conjugate base of litumus with Hlt being the red form and Lt - being the blue form
Example Reactions Placing red litmus paper in a base: HLt (aq) + NaOH (aq) H 2 O (l) + Na + (aq) + Lt - (aq) Placing blue litmus in an acid: HCl (aq) + Lt - (aq) HLt (aq) + Cl - (aq)
Universal Indicators An indicator substance that changes a variety of different colours to indicate a more precise acidity of the solution being tested. ***most indicators DO NOT DO THIS*** **Usually only do two colours**
Uses of Indicators Mark the end point of a titration (chp. 8) to estimate the pH of a solution. ***We can use a series of indicators to get a fairly precise pH instead of using the more expensive pH meter.**** 0 7 14 Indicator Table (Pg. 10) 0-4.8 2.8-8.0 0-3.2 pH = 2.8 – 3.2
Example (You Try) 0-8.2 Phenolphthalein Bromothymol Blue 7.6-14 7.6-8.2 Phenol Red 8.0-14 pH = 7.6-8.0
Indicator Practice Activity
According to the modified Arrhenius theory Acids are substances that react with water (ionize in water) to produce hydronium ions. Alternately, according to Bronsted – Lowry theory Acids are proton donors that become basic (conjugate bases) once they donate their proton. Defining Acids
Example CH 3 COOH (aq) + H 2 O (l) H 3 O + (aq) + CH 3 COO - (aq) This reaction can be explained using either definition and requires the acid react with water. CH 3 COOH (aq) H 2 O (l) H 3 O + (aq) CH 3 COO - (aq)
Example How can we explain using either the modified arrhenius theory or the Bronsted-Lowry theory that a solution of NaHSO 4 will turn blue litmus paper red? NaHSO 4(aq) Arhenius Bronsted-Lowry Acidic Na + (aq) HSO 4 - (aq) H 2 O (l) + → H 3 O + (aq) + SO 4 2- (aq)
Example How can we explain using either the modified arrhenius theory or the Bronsted-Lowry theory that a solution of NaHSO 4 will turn blue litmus paper red? Bronsted-Lowry Acidic Na + (aq) HSO 4 - (aq) H 2 O (l) + → H 3 O + (aq) + SO 4 2- (aq) Acid Conjugate Base
Bases So far bases have been metal hydroxides that can be explained by simple dissociation to produce hydroxide ions according to the Arrhenius theory. Ex: Ca(OH) 2(aq) Ca 2+ (aq) + 2OH - (aq)
The original Arrhenius theory doesn’t account for the basic nature of ammonia or baking soda. The modified Arrhenius theory, that bases ionize in water (react with water) to produce hydroxide ions. Helps to explain why such substances are in fact basic. Modified Arhenius Theory
Example Na 2 CO 3(s) 2Na + (aq) + CO 3 2- (aq) Then, CO 3 2- (aq) + H 2 O (l) OH - (aq) + HCO 3 - (aq) The Bronsted – Lowry definition of a base as a proton acceptor can also help explain this reaction.
How did we know that the carbonate ion was going to produce a basic solution and that the sodium ion was a spectator? Bases are proton acceptors and usually have a negative charge (water and ammonia are exceptions) and acids are proton donors. Therefore, the sodium ion cannot act as an acid or a base. Na does not have H to give away and is + so it can accept an H.
A Special Case Nonmetal oxides in water will form acidic solutions. There is a two step process to explain how this occurs.
Ex: CO 2(g) + H 2 O (l) H 2 CO 3(aq) H 2 CO 3(aq) + H 2 O (l) H 3 O + (aq) + HCO 3 - (aq)
The overall reaction could be combined into one equation: CO 2(g) + 2H 2 O (l) H 3 O + (aq) + HCO 3 - (aq)
Neutralization Reactions An acid – base neutralization is a double replacement reaction that produces water (HOH) and a salt (an ionic compound) HCl (aq) + NaOH (aq) NaCl (aq) + H 2 O (l)
According to the modified Arrhenius definition acids produce hydronium and bases produce hydroxide in solution. Therefore, the reaction can be written as: H 3 O + (aq) + OH - (aq) 2 H 2 O (l) Neutralization can be defined as the reaction of hydronium and hydroxide to produce water.
How to get to that equation? The first equation is called the molecular equation. It shows everything present in a typical double replacement which you should be familiar with. HCl (aq) + NaOH (aq) NaCl (aq) + H 2 O (l)
Next we need to write out the total ionic equation. For this equation separate all soluble compounds into ions and strong acids into hydronium and the conjugate base (anion). Insoluble compounds and weak acids will remain the same. H 3 O + (aq) + Cl - (aq) + Na + (aq) + OH - (aq) 2 H 2 O (l) + Cl - (aq) + Na + (aq)
Now we cancel out spectator ions, ions that don’t change or react in the equation. This leaves us with the net ionic equation. H 3 O + (aq) + OH - (aq) 2 H 2 O (l)
Strong Acids: (more than 99%) React completely (more than 99%) with water to form hydronium ions. The more hydronium ions the greater the acidic properties such as conductivity and low pH. Weak Acids: (less than 50%) React incompletely (less than 50%) with water to form hydronium ions. Lower concentration of hydronium ions leads to less acidic properties. They have a higher pH and are poor conductors of electricity.
Strong Bases: Soluble ionic hydroxides that dissociate 100% in water to produce hydroxide ions. Weak Bases: Reacts partially with water (less than 50%) to produce fewer hydroxide ions.
Examples: Explain the weak base properties of baking soda (sodium bicarbonate/sodium hydrogen carbonate). Solid sodium acetate is dissolved in water. The final solution is tested and found to have a pH of about 8. Explain this evidence by writing balanced chemical equations.
Polyprotic Substances: Polyprotic Acids: Weak acids with multiple protons to donate and whose percent reaction with water decreases after each step. Polyprotic Bases: Weak bases that can accept multiple protons and whose percent reaction with water decreases after each step. BicarbonateCaCO3
H 2 PO 4 - (aq) H + (aq) H 3 PO 4(aq) H + (aq) HPO 4 2- (aq) H + (aq) PO 4 3- (aq) <50% <1% <0.00%