Speech Writing Elements
Essential elements of speech writing in this detailed content, covering aspects like audience engagement, structure, and key features. Learn how to craft impactful speeches that resonate with your target audience and make a lasting impression.
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Presentation Transcript
Book: Bruice Chapter: 2
Early chemists called any compound that tasted sour an acid (Lat.acidus=sour) Compounds that neutralize acids were called bases, or alkaline compounds
Introduction to acids and bases Br nsted and Lowry Acid - species that donates a proton Base - species that accepts a proton Water accepts accept a proton because it has two lone pairs, either of which can form a covalent bond with a proton Reaction of an acid with a base is called an acid base reaction
according to the BrnstedLowry definitions - any species that has a hydrogen can potentially act as an acid, and any compound that has a lone pair can potentially act as a base compound loses a proton - resulting species is called its conjugate base compound accepts a proton - resulting species is called its conjugate acid
Example Ammonia and water Ammonia is base it accepts a proton Water is acid it donates a proton HO- is conjugate base of H2O NH4+ is conjugate acid of NH3
Water can act as acid and base Acid - it has a proton that it can donate Base - it has a lone pair that can accept a proton Acidity - measure of the tendency of a compound to give up a proton Basicity- measure of a compound s affinity for a proton
Strong acid strong tendency to give up its proton Itsconjugate base weak because it has little affinity for the proton Weak acid - little tendency to give up its proton Itsconjugate base - strong because it has a high affinity for the proton -> the stronger the acid, the weaker is its conjugate base
pKa and pH Strong acid dissolves in water - all the molecules dissociate (break into ions) products are favored at equilibrium the equilibrium lies to the right Weak acid dissolves in water - very few molecules dissociate (break into ions) reactants are favored at equilibrium the equilibrium lies to the left
The degree to which an acid (HA) dissociates in an aqueous solution is indicated by the acid dissociation constant-Ka The larger the acid dissociation constant, the stronger is the acid Greater is its tendency to give up a proton
For convenience, the strength of an acid is generally indicated by its pKarather than Kavalue
concentration of positively charged hydrogen ions in the solution is indicated by pH The lower the pH the more acidic is the solution Acidic solutions pH less than 7 Basic solutions pH greater than 7
Important note!!!! Do not confuse pH and pKa!!!!!!! pH scale describe the acidity of a solution pKa characteristic of a particular compound; indicates the tendency of the compound to give up its proton
Organic acids and bases most common organic acids are carboxylic acids compounds that have a COOH group Examples: acetic acid and formic acid Carboxylic acids pKavalues: from 3 to 5 Weak acids
Alcohols have OH group much weaker acids than carboxylicacids pKavalues close to 16
Water can behave as acid and base Alcohol behaves similarly It can donate and accept proton
Protonated compound has gained additional proton Protonated alcohols and carboxylic acids very strong acids sp2 hydrogen is protonated
Amines can act as bases and acids Donate or accept the proton Have high pKa rarely behave as acids More likely to act as bases
Amines most common organic bases talk about the strength of its conjugate acid ->stronger the acid the weaker is conjugate base Example: protonated methylamine Stronger acid than protonated ethylamine Means that methylamine is a weaker base than ethylamine
Approximate pKa values of various classes of compounds important to know Easy to remember units of five (table) R- used when the particular carboxylic acid, alcohol, or amine is notspecified
How to determine the position of equilibrium Of an acid-base reaction Compare the pKa values of the acid on the left and on the right from the arrow The equilibrium favors reaction of the stronger acid and formation of the weaker acid equilibrium lies away from the stronger acid and toward the weaker acid Products are favored in first Reactants are favored in second
How the structure of an acid affects its pKa strength of an acid is determined by the stability of the conjugate base formed when the acid gives up its proton: the more stable the base, the stronger is its conjugate acid Stable base weak base; they do not share their electrons well -> the weaker the base, the stronger is its conjugate acid or, the more stable the base, the stronger is its conjugate acid
Relative electronegativity Factors taht affect the stability of a base: Size and electronegativity Example: elements in the second row of PSE Same size but very different electronegativities Increases across the row from left to right
Relative acidity Acids fromed by those elements - most acidic compound is the one that has its hydrogen attached to the most electronegativeatom
Relative stability Stability of the conjugate bases of formed acids Increase from left to right -> more electronegative the atom, the better it can bear its negative charge Strongest acid has the most stable conjugate base
Conclusion: when the atoms are similar in size, the strongest acid will have its hydrogen attached to the most electronegative atom Comparing pKa values of alcohols and amines appreciated the effect of electronegativity Oxygen more electronegative than nitrogen alcohol more acidic than amine (also in protonated form)
Size of an atom is more important than its electronegativity in determing how well it bears its negative charge Down the column in PSE - atoms get larger and their electronegativity decreases stability of the bases increases down the column - strength of their conjugate acid increases ->when atoms are very different in size, the strongest acid will have its hydrogen attached to the largest atom
Buffer solutions Buffer solution -solution of a weak acid and its conjugate base Will maintain constant pH when small amounts of acid or base are added to it
Example: Blood Transports oxygen to all the cells of the human body Normal pH 7.3-7.4 Death: pH below 6.8 and abouve 8.0 Hemoglobin (HbH+) oxygen carrier When binds to O2 loses a proton Would make blood more acidic if it did not contain a buffer to maintain its pH A carbonic acid/bicarbonate buffer controls the pH of blood Carbonic acid decomposes to CO2 and H2O
Lewis acids and bases 1923 new definition of acids and bases Acid - accepts a share in an electron pair Base - donates a share in an electron pair
Lewis acids are not limited to compounds that donate protons Compounds that have unfilled valence orbitals can accept a share in an electron pair (BH3 or AlCl3) React with a compund that has a lone pair Lewis definition of acid includes all proton-donating compounds and some additional compounds (without protons) Lewis acid : non-proton-donating acid
Lewis base All bases are Lewis bases All have a pair of electrons that they can share
Summary Acids/bases Acidity Lewis acid/base Acid dissociation constant pKa pH Acid-base reactions Buffer solutions
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