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Acidity is a measure of the hydrogen (hydronium) ion concentration of a solution (Acid). Alkalinity is a measure of the hydroxide ion concentration of a solution (Base). Arrhenius Acid is a substance that produces hydronium ions (H3O+) as the only positive ions when dissolved in water. Arrhenius Base is a substance that produces hydroxide ions (OH-) as the only negative ions when dissolved in water. Neutralization is a reaction between an acid and a base to produce salt and water. Acid + Base HNO3 + NaOH Salt + Water NaNO3 + H2O Electrolyte is a substance whose water solution conducts an electric current (H3O+ or OH-).BASE Salt is a product (other than water) of a neutralization reaction; an ionic substance. - gives the spectator ions; ex. Na+ and Cl- According to Bronsted and Lowry Acid: proton (H+) donor H+ proton – remove electron from H leaving one proton Acid is a substance that gives up H+ ions to another substance Base: proton (H+) acceptor Base is a substance that takes an H+ ion from another substance CONJUGATE ACID-BASE PAIRS: A conjugate pair refers to acids and bases with common features. These common features are the equal loss/gain of protons between the pairs. Conjugate acids and conjugate bases are characterized as the acids and bases that lose or gain protons. In an acid-base reaction, an acid plus a base reacts to form a conjugate base plus a conjugate acid Acid + Base→Conjugate Base + Conjugate Acid The conjugate acid of a base is formed when the base gains a proton. Refer to the following equation: Examples: HCl + H2O H3O+ + Cl- NH3 +H2O NH4++ OH- HC2H3O2 + H2O H3O+ + C2H3O2- Acid (H+ or H3O+) & Metal - to react with an acid, the metal must be higher on Table J than H2 Ca + 2HCl H2 + CaCl2 Cu + H3PO4 No Reaction ACIDS: 1. Sour Taste 2. Aqueous solution electrolytes 3. React with bases to produce water and salt 4. React with certain metals to produce H2 gas 5. Cause acid-base indicators to change color(turn blue litmus red) 6. pH less than 7 7. Acids can be formed by a reaction of gaseous oxides with water 1. 2. 3. 4. 5. 6. 7. BASES Bitter taste Aqueous solutions are electrolytes Slippery feel React with acids to form salt and water Cause acid-base indicators to change colors (turn red litmus blue) pH greater than 7 Formed when Group 1 and 2 metals react with H2O, H is released too HCl – dangerous acid Citric Acid – in fruits Boric Acid – eye washing These strong acids and bases are STRONG ELECTROLYTES – GOOD CONDUCTORS of electricity RELATIVE STRENGTHS OF ACIDS AND BASES Strength is determined by the position of the “dissociation” equilibrium Strong Acids (H+ or H3O+) - Found at the top of Table K - Low pH number (1, 2 or 3) - The more oxygen present in the polyatomic ion of an oxyacid (acid containing oxygen), the STRONGER its acid within that group is Strong Bases (OH-) - Found at the top of Table L - High pH number (12, 13 or 14) - Hydroxides or oxides of group 1 and 3 metals (except Mg and Be) - Those that are very soluble are very strong WEAK ACIDS AND BASES Majority of acids and bases are weak – they do not ionize much BINARY ACIDS: composed of H and one other elements (TABLE K) Ex) HCl - hydrogen chloride Name begins with hydro followed by the name of other element with modified ending – ic Ex) HCl – hydrochloric acid TERNARY ACIDS: Have polyatomic ions containing O – table E 1.Anion suffixes – ate and – ite usually replaced by –ic and –ous respectively -ate to –ic -ite to –ous Ex) HNO3 -> NITRIC ACID H2SO4 -> SULFURIC ACID BASES: - EX) Positive ion name is not modified and then of the base ends with hydroxide Ca(OH)2 Calcium hydroxide It was found that no matter how pure water is, it still conducts a minute current. This proves that water self-ionizes Since water will dissociate with itself to a slight extent only about 2 billion water molecules are ionized at any instant The equilibrium expression used here is referred to as the autoionization constant for water – Kw In pure water or dilute solutions, the concentration of water can be considered to be a constant value, so we include that with the equilibrium expression and write it as: Kw = [H+] [OH-] = 1.00 x 10-14 Knowing this value allows us to calculate the ion concentration for OH- and H+ for various situations [OH-] = [H+] solution is neutral (in pure water, each of these is 1.0 x 107 [OH-] > [H+] solution is basic [OH-] < [H+] solution is acidic Used to designate the [H+] in most aqueous solutions where [H+] is small pH = -log [H+] pOH = -log [OH+] pH + pOH = 14 0 = strongly acidic 7 = neutral 14 = strongly basic LOGARITHMIC SCALE Each change of a pH unit signifies a tenfold change in the concentration of the hydrogen ion - Ex) Concentration of [H+] is 10x greater in a solution with a pH = 5 than pH = 6 As concentration of H+ increases the concentration of OH- decreases - Acidity: strength of acid in H+ ions - Alkalinity: strength of base in OH- ions - 1. pH probes – electrodes that detect electrical conductivity 2. Acid –base indicators and narrowing down pH using multiple indicators (mixture of indicators gives great range of colors, pH paper) - Chemicals that have certain colors depending on pH pH paper to compare results to chart TABLE M Ex) A solution yields the following results when tested with various indicators Methyl Orange = yellow Phenolphthalein = clear Bromcresol green = blue Thymol blue = yellow AIM: HOW CAN WE DETERMINE THE CONCENTRATION OF AN ACID OR BASE? Controlled process of acid-base neutralization used to determine the concentration of an acid or a base Endpoint: the pH at which an indicator that has been added to a titration set up turns color Equivalence Point: the point at which the titrated solution has a pH of 7 [OH-] = [H+] – can be detected using probes 1 mole H+ neutralizes 1 mole of OHMoles of H+ = Moles of OHM = mol/L mol = M x L = M x V Macid x Vacid = Mbase x Vbase #H(MaVa) = #OH(MbVb) Want the number of moles #H(molesa) = #OH(molesb) 1. How many moles of LiOH are needed to exactly neutralize 2 moles of H2SO4? 2. How many moles of H2SO4 are needed to exactly neutralize 5.0 moles of NaOH? 3. How many moles of HCl are needed to neutralize 0.10L of 2.0M solution of NaOH? 4. How many moles of NaOH are needed to neutralize 0.10L of 0.20M H2SO4 solution? 5. If it takes 15.0mL of 0.40M NaOH to neutralize 5.0mL of HCl, what is the molar concentration of the HCl solution? 6. If it takes 10.0mL of 2.0M H2SO4 to neutralize 30.0mL of KOH, what is the molar concentration of KOH? 7. How many mL of 2.0M H2SO4 are required to neutralize 30.0mL of 1.0M NaOH solution? 8. How many mL of 0.10M Ca(OH)2 are required to neutralize 25.0mL of 0.50M HNO3 solution? A titration curve is drawn by plotting data attained during a titration, titrant volume on the x-axis and pH on the yaxis. The titration curve serves to profile the unknown solution. In the shape of the curve lies much chemistry and an interesting summary of what we have learned so far about acids and bases. The endpoint is when the indicator color changes – if you picked the right indicator the equivalence point and the endpoint will occur at the same time. WHAT IS A BUFFER? A buffer is a solution of a weak acid or base and its salt [which is its conjugate] WHAT DOES IT DO? A buffer resists a change in pH HOW DOES IT WORK? Since a buffer consists of both an acid or base and its conjugate, an acid and a base are present in all buffer solutions. If a small amount of strong acid is added to the buffer, there is a base component ready and waiting to neutralize the “invader” PREPARING BUFFER SOLUTIONS: Use 0.1M to 1.0M solutions of reagents and choose an acid whose Ka is near the [H+] or [H3O+] concentration we want. The pKa should be as close to the pH desired as possible. Adjust the ratio of weak A/B and its salt to fine tune the pH.