What is the equilibrium constant expression for the following
reaction:
HCN(aq) +
H2O(l)↔CN−(aq) +
H3O+(aq)
Choose from the list below and enter the letters alphabetical
order. (e.g. For an equilibrium constant of
[H3O+]-1eq[HCN]eq
enter AH.)
A) [HCN]eq
E) [HCN]-1eq
I)
[Ru(NH3)62+]6eq
B) [H2O]eq
F) [H2O]-1eq
J) [Ru2+]-6eq
C) [CN-]eq
G) [CN-]-1eq
K) [NH3]-6eq
D) [H3O+]eq
H) [H3O+]-1eq
L)
[Ru(NH3)62+]-6eq
a) identify the conjugate acid-base pairs for the
following reaction: H2O (aq) + CN- (aq)--->HCN (aq)+OH- (aq) b)
give the conjugate acid of HSO4- and PO4 -2 c) give the
conjugate base of HCLO4 and H3O+
The reaction of the weak acid HCN with the strong base KOH is:
HCN(aq)+KOH(aq)-->HOH(l)+KCN(aq) To compute the pH of the
resulting solution if 54mL of 0.79M HCN is mixed with 2.0 × 10^1 mL
of 0.32 KOH we need to start with the stoichiometry. Let\'s do just
the stoich in steps:. a)How many moles of acid? b)How many moles of
base? c)What is the limiting reactant? d)How many moles of the
excess reagent after reaction? e)What is the concentration of...
HNO3 (aq) + CH3NH2
(aq) → CH3NH3+
(aq) +
NO3−(aq)
Kb of CH3NH2 = 4.4 x
10-4
CH3NH3+ (aq) + H2O
(l) ↔ H3O+ (aq) +
CH3NH2
(aq)
Ka of CH3NH3+ = 2.3 x
10-11
Exactly 100 mL of 0.10 M methylamine
(CH3NH2) solution are titrated with a 0.10 M
nitric acid (HNO3) solution. Calculate the pH for:
A:The initial solution
B:The point at which 30 mL of the acid has been added
C:At the half-equivalence point
D:The equivalence point...
For the following acid-base reaction H₂S(aq) + CN⁻(aq) ⇌ HS⁻(aq)
+ HCN(aq) ∆H° = -24.7 kJ/mol and ∆S° = -49.9 J/mol・K. If you mix
100 mL of 0.0150 M NaCN with 100 mL of 0.0150 M H₂S, after
equilibrium is established at 25°C what will be the molar
concentration of HCN?
Consider the following dissociation of the weak acid, HCN:
HCN(aq) ⇌ CN−(aq) + H+(aq) K = 6.2×10-10 A solution is made with an
initial concentration of 2.60 M HCN. At equilibrium, what is the
concentration of H+ ions in solution? (Hint: You may use the 5%
approximation.)
?M
The reaction HCN (aq) + 2H2O (l) > NH4HCO2 (aq) is
first order, and it's rate =k [HCN]. The rate constant k at 65°C is
8.06 ×10 -8 s-1. How long will it take for thr concentration of the
HCN solution to drop from an initial 0.0800M to 0.0600M at this
temperature? What is the half life of thr reaction
The enthalpy of reaction for HCN(aq) + OH-(aq) ->
CN-(aq) + H2O(l) is -12.1 kJ/mol. Assuming
that the correct enthalpy of reaction of the experiment is -54.0
kJ/mol, calculate the enthalpy change for the ionization of
HCN:
HCN(aq) -> H+(aq) + CN-(aq)
Be sure to report your answer to the correct number of
significant figures.
ΔH = ____ kJ/mol