Common Acids: Binary & Ternary Acids H2SO4 = ________ HNO3 = ________ H2CO3 = _______ H3PO4 = _______

H2SO4 = Sulfuric acid HNO3 = Nitric acid H2CO3 = Carbonic acid H3PO4 = Phosphoric acid fyi: (don't use hydro when naming ternary acids) fyi : (adding H+ makes it an acid)

Common Bases: From Polyatomic ions NaOH = ______ KOH = _______ Mg(OH)2 = ______ Ca(OH)2 = _______ CaCO3 = _______ Li2CO3 = _______ NaHCO3 = _______ Ca(HCO3)2 = ______

NaOH = Sodium Hydroxide KOH = Potassium Hydroxide Mg(OH)2 = Magnesium Hydroxide (milk of mag) Ca(OH)2 = Calcium Hydroxide CaCO3 = Calcium Carbonate Li2CO3 = Lithium Bicarbonate NaHCO3 = Sodium Bicarbonate Ca(HCO3)2 = Calcium Bicarbonate

Bronsted-Lowry Acids & Bases: HCN + H2O ⇔ CN- + H3O+ Which are acids and which are bases? What are the conjugate pairs?

HCN + H2O ⇔ CN- + H3O+ HCN: acid H2O: base CN-: base H3O+= acid Conjugate pairs: HCN/CN- & H2O/H3O+

Bronsted-Lowry Aids & Bases: NH3 + H2O ⇔ NH4+ + OH- Which are acids and which are bases? What are the conjugate pairs?

NH3 + H2O ⇔ NH4+ + OH- NH3 = base H2O = acid NH4+ = acid OH- = base Conjugate pairs: NH3/NH4+ & H2O/OH-

Ka values for selected acids: Which of the following is stronger? Ka = 2.5 x 1010 OR Ka = 6.6 x 10-4 (Ka refers to acids, different than Keq)

Ka = 2.5 x 1010 is the stronger acid. fyi: (larger values are stronger acids)

Ch. 7 Acid, Bases & Equilibrium Flashcards by sandy hicks

Common Acids: Binary & Ternary Acids

HI = __________

HBr = _________

HCl = _________

HF = _________

HI = Hyroiodic acid

HBr = Hydrobromic acid

HCl = Hydrochloric acid

HF = Hydroflouric acid

fyi**: (adding H⁺makes it an acid)

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Common Acids: Binary & Ternary Acids

H₂SO₄= ________

HNO₃ = ________

H₂CO₃ = _______

H₃PO₄ = _______

H₂SO₄= Sulfuric acid

HNO₃ = Nitric acid

H₂CO₃ = Carbonic acid

H₃PO₄ = Phosphoric acid

fyi: (don’t* use hydro when naming ternary acids)
fyi**: (adding H⁺ makes it an acid)*

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Common Bases: From Polyatomic ions

NaOH = ______

KOH = _______

Mg(OH)₂= ______

Ca(OH)₂ = _______

CaCO₃ = _______

Li₂CO₃ = _______

NaHCO₃ = _______

Ca(HCO₃)₂ = ______

NaOH = Sodium Hydroxide

KOH = Potassium Hydroxide

Mg(OH)₂= Magnesium Hydroxide (milk of mag)

Ca(OH)₂ = Calcium Hydroxide

CaCO₃ = Calcium Carbonate

Li₂CO₃ = Lithium Bicarbonate

NaHCO₃ = Sodium Bicarbonate

Ca(HCO₃)₂ = Calcium Bicarbonate

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Acids Vs. Bases

Acids taste _______

Bases taste _______

Bases are also _______ to the touch.

Acids taste Sour

Bases taste Bitter

Bases are also Slippery to the touch.

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Bronsted-Lowry Acids & Bases:

Acids _______ H⁺

Bases _______ H⁺

Acids are proton _______.

Bases are proton _______.

Acids release H⁺

Bases accept H⁺

Acids are proton donors.

Bases are proton acceptors.

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Bronsted-Lowry Acids & Bases:

Define: Amphoteric

Amphoteric: Compounds that can act as acids or bases.

i.e: H₂O

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Bronsted-Lowry Acids & Bases:

Define: Conjugates

Conjugates: Compounds which differ only in the presence or absence of H⁺.

i.e: H₂O & H₃O⁺

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Bronsted-Lowry Acids & Bases:

HCN + H₂O ⇔ CN^-+ H₃O⁺

Which are acids and which are bases?

What are the conjugate pairs?

HCN + H₂O ⇔ CN^-+ H₃O⁺

HCN: acid

H₂O: base

CN^-: base

H₃O⁺= acid

Conjugate pairs: HCN/CN^-& H₂O/H₃O⁺

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Bronsted-Lowry Aids & Bases:

NH₃ + H₂O ⇔ NH₄⁺ + OH^-

Which are acids and which are bases?

What are the conjugate pairs?

NH₃ + H₂O ⇔ NH₄⁺ + OH^-

NH₃= base

H₂O = acid

NH₄⁺ = acid

OH^-= base

Conjugate pairs: NH₃/NH₄⁺& H₂O/OH^-

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Bronsted-Lowry Acids & Bases:

NH₃ + H₂O ⇔ NH₄⁺ + OH^-

What is the conjugate base of H₂O acid?

NH₃ + H₂O ⇔ NH₄⁺ + OH^-

OH^-

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Equilibrium:

Define: Equilibrium

Equilibrium: The rate at which forward and reverse reactions are equal.

i.e: N₂O_4(g) ⇔ 2NO_2(g)

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Equilibrium:

K_eq = _______

K_eq = Equilibrium constant

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Equilibrium:

Equilibrium constant values vary with _________.

Can a catalyst effect equilibrium or the value of K_eq?

Temperature(25°C)

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Equilibrium Constant:

______ & _____ are NOT included in equilibrium equations.

Solvents_(l) & Solids_(s) are NOT included in equilibrium constant equations(K_eq).

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Equilibrium Constant:

aA + bB ⇔ cC + dD

K_eq=

aA + bB ⇔ cC + dD

K_eq = [C]^cx [D]^d / [A]^{a<strong> </strong>}x [B]^b

aA + bB ⇔ cC + dD

fyi: (lowercase = coefficiants from balanced equation)

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Equilibrium Constant:

Balance the reaction and write the corresponding equilibrium constant expression.

CO_(g) + O_2(g) ⇔ CO_2(g)

K_eq =

2CO_(g) + O_2(g) ⇔ 2CO_2(g)

K_eq = [CO₂]²/ [CO]²[O₂]

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Equilibrium Constant:

Balance the reaction and write the corresponding equilibrium constant expression.

C_(s)+ H₂O_(g) ⇔ CO_(g) + H_2(g)

K_eq =

C_(s)+ H₂O_(g) ⇔ CO_(g) + H_2(g)

(balanced)

K_eq = [CO][H₂] / [H₂O]

fyi: (solvents_(l) & Solids_(s)are NOT included in K_eq.)

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Equilibrium Constant:

Balance the reaction and write the corresponding equilibrium constant expression.

H₂SO_4(aq)+ H₂O_(l) ⇔ HSO₄^-_(aq) + H₃O⁺_(aq)

K_eq=

H₂SO_4(aq)+ H₂O_(l) ⇔ HSO₄^-_(aq) + H₃O⁺_(aq)

(balanced)

K_eq = [HSO₄^-][H₃O⁺] / [H₂SO₄]

fyi: (solvents(l) & Solids(s) are NOT included in Keq.)

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Ka values for selected acids:

Which of the following is stronger?

K_a = 2.5 x 10¹⁰

K_a = 6.6 x 10^-4

(K_a refers to acids, different than K_eq)

K_a = 2.5 x 10¹⁰ is the stronger acid.

fyi: (larger values are stronger acids)

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K_a values for selected acids:

K_a values are always at ____°C

25°C

K_a values for selected acids:

Higher K_a value indicates a better _______ donor.

proton donor

Acid_(?)are ________

Aqueous

Acids_(aq)

Ionization of Water:

K_w = _________________

K_w= [OH^-][H₃O⁺] = 1.0 x 10^-14

Ionization of Water:

K_w is constant at ____°C.

25°C

Ionization of Water: [OH^-] **\>** [H3O⁺] = \_\_\_\_\_\_ [OH^-] **\<** [H3O⁺] = \_\_\_\_\_\_

[OH^-] **\>** [H3O⁺] = **Basic** [OH^-] **\<** [H3O⁺] = **Acidic**

Ionization of Water: [OH^-] = 8.4 x 10^-3M pH= 11.92 Calculate the [H₃O⁺] and pH. Acidic, basic, or neutral? ***(explain steps)***

1 x 10^-14 **/** [OH^-] = [H₃O⁺] [H₃O⁺] = 1.19 x 10^-12 pH = -Log([H₃O⁺]) = 11.92 11.92 = Basic

Ionization of Water: [H₃O⁺] = 9.1 x 10^-8M pH = 7.0 Calculate the [OH^-] and pH. Acidic, basic, or neutral? ***(explain steps)***

1 x 10^-14 **/** [H₃O⁺] = [OH^-] [OH^-] = 1.1 x 10^-7 pH = -Log([H₃O⁺]) = 7.0 7.0 = Neutral

Ionization of Water: How do you find the pH?

pH = -Log([H₃O⁺])

Ph Values: pH has ____ sig figs after the decimal

Ph Values: Acidic / Basic / Neutral? 0 → 6.9 = \_\_\_\_\_\_\_ 7. 0 = \_\_\_\_\_\_ 7. 1 → 14.0 = \_\_\_\_\_\_

0 → 6.9 = **Basic** 7. 0 = **Neutral** 7. 1 → 14.0 = **Acidic**

A ________ is a solution that resists large changed in pH when small amounts of acid or base are added.

Buffer

What makes a buffer solution?

A weak acid and its conjugate base H₂CO_3(aq)**and** HCO₃^-_(aq) H₂CO_3(aq)+ H₂O_(l) ⇔ HCO₃^-_(aq) + H₃O⁺_(aq)

Buffers are ______ acids.

Weak

Le Châtelier's Principle: What happens when a reversible reaction is pushed out of equilibrium?

The reaction responds to reestablish equilibrium.

Le Châtelier's Principle: What is one way to upset an equilibrium?

Varying the concentration of reactant or product.

Le Châtelier's Principle: H₂O_(l) + CO_2(g) ⇔ H₂CO_3(aq) Increasing CO₂ **upsets** the equilibrium and a net \_\_\_\_\_\_\_reaction takes place. ***fyi: (upsets = changing concentration)***

net forward reaction

Le Châtelier's Principle: H₂O_(l) + CO_2(g) ⇔ H₂CO_3(aq) Decreasing CO₂**upsets** the equilibrium and a net _________ reaction takes place. ***fyi: (upsets = changing concentration)***

net reverse reaction

Le Châtelier's Principle: H₂O_(l) + CO_2(g) ⇔ H₂CO_3(aq) Increasing H₂CO₃ **upsets** the equilibrium and a net ________ reaction takes place. ***fyi: (upsets = changing concentration)***

net reverse reaction

Le Châtelier's Principle: For the equilibrium reaction below, predict which direction *(forward or reverse)* will be the faster one until equilibrium is reestablished when: **2**SO_2(g) + O_2(g)⇔ **2**SO₃_(g) a. SO₂ is increased b. SO₃ is increased c. SO₂ is decreased d. O₂ is increased

**a.** Net forward reaction, SO₃ increases **b.** Net forward reaction **c.** Net reverse reaction, SO₃ decreases **d.** Net forward reaction

The pH of your blood normally ranges between _____ & \_\_\_\_\_. How is this narrow range maintained?

7.35 & 7.45 With the help of buffers

Blood pH below the normal range is called \_\_\_\_\_\_\_\_. Blood pH above normal range is called \_\_\_\_\_\_\_\_\_.

acidosis alkalosis

The most important buffer system in the blood is formed from _________ and its conjugate base, HCO₃^- .

H₂CO₃

Buffers in the blood: **Two important equations** CO₂_(g)+ H₂O_(l) ⇔ \_\_\_\_\_\_\_\_\_\_\_ CO₂_(g)+ H₂O_(l) ⇔ \_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_

CO₂_(g)+ H₂O_(l) ⇔ H₂CO_3(aq) CO₂_(g)+ H₂O_(l) ⇔ HCO₃^-_(aq)+ H₃O⁺_(l)

Acidosis & Alkalosis are potentially fatal! Why?

**Adding OH^- (base)** = Loss of salt bridge due to OH^- bonding w/ and forming H₂O. ***pH increases / H₃O decreases*** **Adding H⁺ (acid)** = Won't allow NH₃ and O^- to bond. Forms and Carboxylic acid(no longer a protein) ***pH decreases / H₃O increases***

The solution is alkaline **=** the solution is \_\_\_\_\_\_\_.

Basic

Common Causes / Symptoms of: Acidosis

1. Lung disease, asthma 2. Excessive alcohol consumption 3. Ketosis / Starvation 4. Holding your breath too long 5. Hy**po**ventilation **(*****not enough CO₂ is exhaled***)

Common Causes / Symptoms of: Alkalosis

1. Excessive use of antacids 2. Anxiety, fever 3. Hy**per**ventilation ***(CO₂ is blown off faster than it is produced in the cells.)***

Blood buffer (w/ help from organs):