chapter_4

Behavior of a strong electrolyte in aqueous solution

Undergoes nearly complete dissociation or ionization into ions.

Behavior of a weak electrolyte in aqueous solution

Undergoes only partial ionization, existing as a dynamic mixture of intact molecules and ions.

Behavior of a nonelectrolyte in aqueous solution

Dissolves as intact neutral molecules with essentially no ion formation.

Electrolyte classification of $KCl$ in water

Strong electrolyte; soluble ionic compound.

Electrolyte classification of $HNO_3$ in water

Strong electrolyte; strong acid.

Electrolyte classification of $HNO_3$ in water

Strong electrolyte; strong acid.

Electrolyte classification of $CH_3COOH$ in water

Weak electrolyte; weak acid.

Electrolyte classification of $NH_3$ in water

Weak electrolyte; weak base.

Electrolyte classification of sucrose ($C_{12}H_{22}O_{11}$) in water

Nonelectrolyte; dissolves as intact neutral molecules.

Orientation of water molecules during the hydration of $Na^+$

The partially negative oxygen ($O$) atoms point toward the cation.

Orientation of water molecules during the hydration of $Cl^-$

The partially positive hydrogen ($H$) atoms point toward the anion.

Four categories of ionic compounds that are always soluble

Salts containing Group 1 metal cations, $NH_4^+$, $NO_3^-$, and $ClO_3^-$.

Insoluble exceptions for chloride ($Cl^-$), bromide ($Br^-$), and iodide ($I^-$) salts

Salts containing $Ag^+$, $Hg_2^{2+}$, and $Pb^{2+}$.

Insoluble exceptions for sulfate ($SO_4^{2-}$) salts

Salts containing $Ag^+$, $Ca^{2+}$, $Sr^{2+}$, $Ba^{2+}$, $Hg_2^{2+}$, and $Pb^{2+}$.

Soluble exceptions for typically insoluble anions ($CO_3^{2-}$, $PO_4^{3-}$, $CrO_4^{2-}$, $S^{2-}$)

Salts paired with Group 1 alkali metal cations or $NH_4^+$.

Soluble exceptions for typically insoluble hydroxide ($OH^-$) salts

Hydroxides containing Group 1 metal cations and $Ba^{2+}$.

Solubility of calcium phosphate, $Ca_3(PO_4)_2$, in water

Insoluble; phosphates are generally insoluble except when paired with Group 1 cations or $NH_4^+$.

Solubility of silver chlorate, $AgClO_3$, in water

Soluble; chlorate salts are always soluble.

Solubility of potassium sulfide, $K_2S$, in water

Soluble; potassium ($K^+$) is an alkali metal cation.

Solubility of calcium carbonate, $CaCO_3$, in water

Insoluble; carbonates are insoluble except when paired with Group 1 cations or $NH_4^+$.

Precipitate formed upon mixing solutions containing $Ba^{2+}$ and $SO_4^{2-}$

$$BaSO_{4(s)}$$

Species split into separate ions in complete ionic equations

Strong soluble electrolytes in aqueous solution.

Species kept in intact molecular or formula unit form in ionic equations

Solids, liquids, gases, weak acids, weak bases, and nonelectrolytes.

Net ionic equation for $AgNO_{3(aq)} + NaCl_{(aq)}$

$$Ag^+ + Cl^- \rightarrow AgCl_{(s)}$$

Net ionic equation for $ZnCl_{2(aq)} + Na_2S_{(aq)}$

$$Zn^{2+} + S^{2-} \rightarrow ZnS_{(s)}$$

Net ionic equation for the reaction between $Sr^{2+}$ and $PO_4^{3-}$

$$3Sr^{2+} + 2PO_4^{3-} \rightarrow Sr_3(PO_4)_{2(s)}$$

Net ionic equation for the neutralization of a strong acid by a strong base

$$H^+ + OH^- \rightarrow H_2O_{(l)}$$

Net ionic equation for the reaction between $CH_3COOH$ and $OH^-$

$$CH_3COOH + OH^- \rightarrow CH_3COO^- + H_2O_{(l)}$$

Reason $CH_3COOH$ is written intact in net ionic equations

It is a weak acid that remains predominantly non-ionized in aqueous solution.

Net ionic equation for the reaction between $HBr$ and $NH_3$

$$H^+ + NH_3 \rightarrow NH_4^+$$

Representation when all potential products in an aqueous reaction remain soluble

NR (no reaction / no net ionic reaction).

Brønsted-Lowry acid definition

A substance that donates a proton ($H^+$).

Brønsted-Lowry base definition

A substance that accepts a proton ($H^+$).

Brønsted-Lowry classification of $NH_4^+$

Acid; it can donate a proton ($H^+$).

Brønsted-Lowry classification of $PO_4^{3-}$

Base; it can accept a proton ($H^+$).

Brønsted-Lowry classification of $HCO_3^-$

Amphiprotic (both acid and base); it can donate or accept a proton ($H^+$).

Brønsted-Lowry classification of $H_2PO_4^-$

Amphiprotic (both acid and base); it can donate or accept a proton ($H^+$).

Brønsted-Lowry classification of $HSO_4^-$

Amphiprotic (both acid and base); it can donate or accept a proton ($H^+$).

Brønsted-Lowry classification of $CH_3COO^-$

Base; it accepts a proton ($H^+$) to regenerate $CH_3COOH$.

Condition under which a metal displaces another metal ion according to the activity series

The solid elemental metal must be positioned higher on the activity series than the aqueous metal cation.

Outcome of combining $Cu_{(s)}$ with $HCl_{(aq)}$

NR (no reaction); copper lies below hydrogen on the activity series.

Net ionic equation for $Mg_{(s)} + CuSO_{4(aq)}$

$$Mg + Cu^{2+} \rightarrow Mg^{2+} + Cu$$

Redox outcome when $Cl_2$ is added to a solution containing $Br^-$

$Cl_2$ is reduced to $Cl^-$, while $Br^-$ is oxidized to $Br_2$.

Outcome of combining $I_2$ with $Br^-$

NR (no reaction); iodine is less reactive as an oxidizing agent than bromine.

Definition and calculation formula for molar concentration ($c$)

$c = \frac{n}{V}$, where $n$ is the amount in moles and $V$ is the solution volume in $dm^3$.

Dilution equation

$$M_1V_1 = M_2V_2$$

Quantity that remains constant during solution dilution

The total amount (moles) of solute.

Final concentration when $25.0\,cm^3$ of $0.866\,mol\,dm^{-3}\ KNO_3$ is diluted to $500\,cm^3$

$$0.0433\,mol\,dm^{-3}$$

Mass of $KI$ required to prepare $500\,cm^3$ of a $2.80\,mol\,dm^{-3}$ solution

$$232\,g\ KI$$

Procedure for finding concentration when mixing solutions of the same solute

Sum the moles of solute from each solution, then divide by the total combined volume.

Final concentration after mixing $35.2\,cm^3$ of $1.66\,mol\,dm^{-3}$ and $16.7\,cm^3$ of $0.892\,mol\,dm^{-3}\ KMnO_4$

$$1.41\,mol\,dm^{-3}$$

Total $[Na^+]$ after mixing $50.0\,cm^3$ of $0.20\,mol\,dm^{-3}\ NaClO_4$ and $25.0\,cm^3$ of $0.20\,mol\,dm^{-3}\ Na_2SO_4$

$$0.267\,mol\,dm^{-3}$$

Initial calculation step in gravimetric analysis after measuring precipitate mass

Convert the mass of the precipitate to moles using its molar mass.

Mass percentage of $Ba$ in a $0.6760\,g$ sample that produces $0.4105\,g\ BaSO_4$

$$35.73\%$$

Core workflow sequence for quantitative titration calculations

$$\text{Moles of titrant} \rightarrow \text{mole ratio} \rightarrow \text{moles of analyte} \rightarrow \text{divide by analyte volume in } dm^3$$

Concentration of $NaOH$ if $25.0\,cm^3$ neutralizes $17.4\,cm^3$ of $0.312\,mol\,dm^{-3}\ HCl$

$$0.217\,mol\,dm^{-3}$$

Volume of $0.500\,mol\,dm^{-3}\ HCl$ required to neutralize $10.0\,cm^3$ of $0.300\,mol\,dm^{-3}\ NaOH$

$$6.00\,cm^3$$

Volume of $0.500\,mol\,dm^{-3}\ HCl$ required to neutralize $10.0\,cm^3$ of $0.200\,mol\,dm^{-3}\ Ba(OH)_2$

$$8.00\,cm^3$$

Reason $Ba(OH)_2$ requires twice the stoichiometric amount of $H^+$ per mole compared to $NaOH$

Each formula unit of $Ba(OH)_2$ dissociates to yield two moles of hydroxide ions ($2\,OH^-$).

First requirement before using volume data in a permanganate/${H_2O_2}$ redox titration

Determine the stoichiometric mole ratio from the balanced redox reaction.

Concentration of $H_2O_2$ if $25.00\,cm^3$ requires $36.44\,cm^3$ of $0.01652\,mol\,dm^{-3}\ KMnO_4$

$$0.06020\,mol\,dm^{-3}$$

Definition of oxidation in terms of electron movement

The loss of electrons.

Definition of reduction in terms of electron movement

The gain of electrons.

Role and fate of an oxidizing agent in a chemical reaction

Undergoes reduction (gains electrons) and causes the oxidation of another species.

Role and fate of a reducing agent in a chemical reaction

Undergoes oxidation (loses electrons) and causes the reduction of another species.

Oxidation state rule for an uncombined element

$$0$$

Oxidation state rule for a monatomic ion

Equal to the ionic charge.

Stepwise sequence for balancing redox half-reactions in acidic solution

Balance non-$O$/$H$ atoms $\rightarrow$ balance $O$ using $H_2O$ $\rightarrow$ balance $H$ using $H^+$ $\rightarrow$ balance charge using $e^-$.

Standard technique for balancing redox reactions in basic solution

Balance completely under acidic conditions, then add equal amounts of $OH^-$ to both sides to neutralize $H^+$.

Step performed after neutralizing $H^+$ with $OH^-$ in basic redox balancing

Combine $H^+$ and $OH^-$ into $H_2O$, then cancel excess water molecules across both sides.

Two balance criteria verified in every finished redox reaction

Atom balance (conservation of mass) and electrical charge balance.

Permissibility of free $H^+$ ions in a finalized basic redox reaction

Not permissible; all $H^+$ must be neutralized to $H_2O$ using $OH^-$.

Number of electrons lost in the half-reaction $Zn \rightarrow Zn^{2+}$

$2\,e^-$ per zinc atom.

Oxidation state of manganese in permanganate ($MnO_4^-$)

$$+7$$

Number of electrons gained per manganese atom in $MnO_4^- \rightarrow Mn^{2+}$

$$5\,e^-$$

Oxidation state change of nitrogen in the transformation $NO_3^- \rightarrow NH_4^+$

Decreases from $+5$ to $-3$, corresponding to a gain of $8\,e^-$.

Total electrons transferred in $4Zn + 10H^+ + NO_3^- \rightarrow 4Zn^{2+} + NH_4^+ + 3H_2O$

$8\,e^-$ total.