chemistry pH and equilibrium chpters

Bronsted-Lowry

Acid: proton donor

Base: proton acceptor

Arrhenius acid-base

- Acid: produces H+

- Base: produces OH-

*limited to aqueous solutions

Weak acid parent

Kp

equilibrium constant for gas-phase reactions based on partial pressures of reactants and products (has to be in ATM)

Kp=

Kc(RT)^∆n

∆n

moles of gaseous products - moles of gaseous reactants

in equilibrium

-a reaction has not stopped

-happens indefinetly as long as conditions continue

-concentrations remain constant

Q>K

shift left towards reactants

Q

reaction goes towards products

Small K value

reactants are favored

K>1

product favored

Q=K

reaction is at equilibrium

T or F, K values is not temp dependent

false

Le Chatelier's Principle

when stress is applied to a system at equilibrium, the system changes to relieve the stress

What happens when pressure increases at equilibrium?

the reaction will go towards the side with less moles to maintain equilibrium

What happens when pressure increase in a reaction where the moles on the reactants are the same as the moles in the products?

The reaction wont do anything

What happends when we increase the temp in a reaction at equilibrium?

The reaction will shift wherever heat is not at.

If you decrease the temperature in a reaction at equilibrium...

reaction shifts toward the side where heat is located

Adding an inert gas

does not change \(K\) as long as the temperature stays constant.

when doing equilibrium expression always make sure that the states of matter are

(aq) or (g)

R value used for calculating Kp

0.0821

What happens to \(K_p\) when a reaction is reversed at the same temperature?

The new \(K_p\) is the reciprocal of the original:

if you need partial pressure and have molarity and temp

use PV= nRT

Write the equilibrium expression first.

Write \(K_p\) using the reaction exactly as given, regardless of what substances are initially present.

What happens to the reaction if only products are initially present?

The reaction must shift in reverse to produce reactants.

What happens to the reaction if only reactants are present?

The reaction shifts forward to produce products.

What is the purpose of setting up an ICE table in chemistry?

To track the changes in concentrations or partial pressures of reactants and products during a reaction.

What do the coefficients in a balanced chemical equation indicate in an ICE table?

They indicate how the partial pressures or concentrations change relative to one another.

How should changes be represented for substances being formed in an ICE table?

As positive changes.

How should changes be represented for substances being consumed in an ICE table?

As negative changes.

How do you find Ka if you know the Kb of a weak base?

Ka = Kw / Kb. This gives Ka for the conjugate acid.

A weak acid HA is neutralized by a strong base. What constant applies to A⁻ in water?

Kb, because A⁻ is the conjugate base and produces OH⁻.

5% rule

x/ [initial concentration] times 100 < 5%

How do coefficients affect the changes in an ICE table

Multiply x by each substance's coefficient. Coefficient 1 = ±x; coefficient 2 = ±2x; coefficient 3 = ±3x. The reaction direction determines the sign.

Example: H2 + Br2 ⇌ 2HBr. In the reverse direction: H2 = +x, Br2 = +x, HBr = −2x.

How do you predict which changes form more product in equilibrium when solids and liquids are present?

Decrease temperature favors the exothermic direction. 2. Increasing gas pressure shifts toward fewer gas moles. 3. Adding a gas reactant shifts toward products. 4. Adding or removing a pure solid or liquid does not shift equilibrium, as long as some remains.

T or false: removing a solid/ liquid does not affect the equilibrium of a concentration as long as there is still some

T

If a solid is left out of the equilibrium expression, can equilibrium shifts still produce more of it?

Yes. Being left out of the expression does not mean it cannot be formed or consumed. Temperature, gas pressure, and concentration changes can shift equilibrium and change the amount of solid produced.

Do catalyst change Kc

No

What is percentage dissociation, and how do you calculate it?

The percentage of an acid or base that ionizes in water. % dissociation = (x / initial concentration) × 100%, where x is the amount ionized from the ICE table.

What happens when a common ion is added to a weak acid solution?

Equilibrium shifts left, forming more undissociated acid and decreasing [H⁺]. The pH increases, while Ka stays constant at the same temperature.

Weak acid

: handles added base

Conjugate base:

handles added acid.

Why dont we use a strong acid for a buffer

essentially complete ionization, so very little original acid remain

In a buffer the conjugate base is like a

a sponge: when you add extra H+, the sponge absorbs it, so you have less conjugate base left and you are forming more acid

What determines the pH of a solution?

The hydrogen ion concentration, \([H_3O^+]\). More \(H_3O^+\) = lower pH; less \(H_3O^+\) = higher pH.

T or F. Forming more weak acid will Increase the pH for the buffer system

F. pH depends on \([H_3O^+]\), not the total amount of acid. More \(H_3O^+\) means lower pH; less means higher pH. Forming more weak acid doesn't automatically increase pH.

Lewis acid

accepts electrons

Lewis base

electron pair donor

What happens to the \([A^-]/[HA]\) ratio as pH increases?

More HA loses H⁺ and becomes A⁻, so the ratio increases.

How do you identify a buffer with high capacity?

Look for a high concentration of buffer components and a ratio of \([A^-]/[HA]\) close to 1.

What happens when strong acid is added to a buffer?

A⁻ reacts with H⁺ to form HA. The \([A^-]/[HA]\) ratio decreases.

What happens when strong base is added to a buffer?

HA reacts with OH⁻ to form A⁻. The \([A^-]/[HA]\) ratio increases.

How can you determine which buffer component to add to adjust pH?

Compare the \([A^-]/[HA]\) ratios. Add HA to lower the ratio or A⁻ to raise it.

What conditions give a buffer high capacity?

High total concentration of HA and A⁻.

A ratio of HA to A⁻ close to 1:1 for balanced buffering.

What is the difference between buffer capacity and buffer range?

Capacity: How much added acid or base the buffer can neutralize.

Range: The pH range where the buffer works effectively.

What is the effective pH range of a buffer?

pH = pKₐ ± 1

The buffer works most effectively within one pH unit above or below its pKa. Maximum buffering occurs near pH = pKa, where [HA] ≈ [A⁻].

when given the ratio of current concentrations use (not necessarily equilibrium)

Qc

Strenght of acid and its conjugate base are

inversely related

How to validate aprroximation?

X/ intital concentration

Why is Ka1 greater than Ka2, which is greater than Ka3?

Each successive H⁺ is harder to remove as the ion becomes more negatively charged. Stronger acid = larger Ka.

strenghts of oxyacids depends on

The number of oxygen atoms attached to the central atom.

The electronegativity of the central atom, when comparing acids with the same number of oxygen atoms.

what is bond energy?

How difficult is it to break the bond

what is bond polarity?

How unevenly are electrons shared

Lower bond energy

bond is easier to break = generally stronger acid.

Higher bond energy

= bond is harder to break = generally weaker acid.

Acid strength increases because atomic size increases, making the H-X bond weaker and easier to break.

Down a group

note

larger atom forms a weaker bond with hydrogen, making it easier for the acid to release its proton.

The stronger an acid is

the weaker its conjugate base is.

stronger base

the weaker the conjugate acid

What determines how acidic a hydrated metal ion is?

Higher positive charge → generally greater acidity.

Smaller ionic radius → greater charge density → generally greater acidity.

% ionization

[H3O+]/[HA] x 100 or [OH-]/[B-]x100

At the equivalence point, how do the moles of acid and base compare?

They react according to the balanced equation's mole ratio. For a 1:1 reaction, moles of acid = moles of base.

pH below pKa

acid form dominates.

pH equal to pKa

equal amounts.

pH above pKa

conjugate base dominates.

the point in which an indicator changes color

endpoint

when are kp and kc equal

delta n=0

what affects buffer capcity

the total concentration of the buffer components.

Why do we use \(+s\) in a common-ion Ksp ICE table?

The common ion is already dissolved. \(s\) measures how much additional solid dissolves to reach the new equilibrium.

True or False: The common ion affects how much dissolves, but the solid's Ksp determines the equilibrium.

True

T of F For equilibrium constant kc and kp. when the reaction is finished no more products are forming

FALSE

T of F For equilibrium constant kc and kp. When a system has reached equilibrium The concentrations of the reactans and equal

FALSE

T of F For equilibrium constant kc and kp. At equilibrium the concentrations are no longer changing

TRUE

T of F For equilibrium constant kc and kp. As equilibrium the reaction is not over, but will continue forever is isolated

T of F For equilibrium constant kc and kp, at equilibrium, the rate of the reaction at which products are formed equals the rate of the reation reactants are formed

True

What does equilibrium constant tell us

which direction is favored