Chem 101 Unit 2

BCC: Chem Unit 2

Solution

A homogeneous mixture of two or more substances; a solute dissolved in a solvent

Solute

The substance present in the lesser amount in a solution

Solvent

The substance present in the greater amount in a solution

Aqueous solution (aq)

A solution in which water is the solvent

Crystalline vs. amorphous solid

Crystalline solids have a highly organized, repeating arrangement at the atomic level; amorphous solids do not

Dissolution

Water pulls whole molecules away from a solid; the molecules stay intact (e.g., sucrose: C₁₂H₂₂O₁₁(s) → C₁₂H₂₂O₁₁(aq))

Dissociation

Water pulls individual ions off an ionic crystal, separating them in solution (e.g., NaCl(s) → Na⁺(aq) + Cl⁻(aq))

Complete dissociation

The ionic compound fully separates into ions; written with a one-way arrow; produces a strong electrolyte (e.g., NaCl)

Partial dissociation

Only some of the solute separates into ions; written with an equilibrium arrow (⇌); produces a weak electrolyte (e.g., NH₃ + H₂O ⇌ NH₄⁺ + OH⁻)

Nondissociation

The solute dissolves as intact neutral molecules with no ions; produces a nonelectrolyte (e.g., sucrose)

Equilibrium arrow (⇌)

Forward and reverse reactions continue at equal rates, so concentrations stop changing; reactants and products coexist

Electrolyte

A substance whose solution conducts electricity

Strong electrolyte

Result of complete dissociation; many ions in solution; bulb glows brightly

Weak electrolyte

Result of partial dissociation; few ions in solution; bulb glows dimly

Nonelectrolyte

Result of nondissociation; no ions in solution; bulb does not light

Conductivity test: what carries the current?

Ions carry charge through the solution to complete the circuit; electrons travel through the wires

Why might a sucrose solution light the bulb slightly?

The water was not deionized and contained dissolved ions

Acids and bases as electrolytes

They undergo complete or partial dissociation, so they are classified as electrolytes

Solubility: always soluble ions

Group 1 cations (alkali metals), NH₄⁺, nitrates (NO₃⁻), bicarbonates (HCO₃⁻), chlorates (ClO₃⁻), acetates (CH₃COO⁻)

Solubility: halides (Cl⁻, Br⁻, I⁻)

Soluble EXCEPT with Ag⁺, Hg₂²⁺, Pb²⁺, and Cu⁺

Solubility: sulfates (SO₄²⁻)

Soluble EXCEPT with Ba²⁺, Ag⁺, Ca²⁺, Sr²⁺, Hg₂²⁺, and Pb²⁺

Solubility: hydroxides (OH⁻)

Insoluble EXCEPT with Group 1 cations, Ca(OH)₂, Sr(OH)₂, and Ba(OH)₂

Solubility: sulfides (S²⁻)

Insoluble EXCEPT with Group 1 cations, NH₄⁺, CaS, SrS, and BaS

Solubility: carbonates (CO₃²⁻)

Insoluble EXCEPT with Group 1 cations and NH₄⁺

Solubility: phosphates (PO₄³⁻)

Insoluble EXCEPT with Group 1 cations and NH₄⁺

Solubility: chromates (CrO₄²⁻)

Insoluble EXCEPT with Group 1 cations and NH₄⁺

Solubility: fluorides

Most are insoluble, although not those of Group 1 cations or NH₄⁺

Is AgCl soluble?

No. It is a halide with Ag⁺, an exception to the halide rule

Is BaSO₄ soluble?

No. It is a sulfate with Ba²⁺, an exception to the sulfate rule

Is Na₃PO₄ soluble?

Yes. Phosphates with Group 1 cations are soluble

Chemical reaction

The result of bond-breaking or bond-making events that create new substances

Five signs of a chemical reaction

Color change; formation of a solid (precipitate); formation of a gas (bubbles, possibly an odor); temperature change; light emission

Why does temperature change in a reaction?

The net energy decides: if more energy is absorbed (breaking bonds) than released (making bonds) it is endothermic and cold; if more is released it is exothermic and hot

Balancing chemical equations

Done to obey the Law of Conservation of Mass

Metathesis (double displacement) reaction

A reaction between two aqueous salts that swap ion partners: AB(aq) + CD(aq) → AD(s) + CB(aq)

When does a metathesis reaction occur?

Only if at least one product is non-aqueous (a solid, a gas, or water); otherwise no reaction (NR)

Why does NaCl(aq) + KHCO₃(aq) not react?

The possible products KCl and NaHCO₃ are both aqueous, so there is no non-aqueous product

Acid-base (neutralization) reaction

HB + COH → HOH + CB; acid + base gives water (non-aqueous) + salt; involves the transfer of H⁺ ions

Gas-forming reaction

Metathesis produces an unstable product that immediately decomposes into a gas and water

Decomposition of H₂CO₃

H₂CO₃ → CO₂(g) + H₂O(l)

Decomposition of H₂SO₃

H₂SO₃ → SO₂(g) + H₂O(l)

Decomposition of NH₄OH

NH₄OH → NH₃(g) + H₂O(l)

Balanced: Na₂CO₃ + HCl

Na₂CO₃(aq) + 2 HCl(aq) → 2 NaCl(aq) + H₂O(l) + CO₂(g)

Balanced: AgNO₃ + NaCl

AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)

Balanced: Pb(NO₃)₂ + KI

Pb(NO₃)₂(aq) + 2 KI(aq) → PbI₂(s) + 2 KNO₃(aq)

Total ionic equation

Shows all dissolved strong electrolytes as separate ions (solids, liquids, and gases stay intact); shows how species actually exist in solution

Net ionic equation

Shows only the species that actually change in the reaction; spectator ions are removed

Spectator ions

Ions that appear unchanged on both sides of the equation (non-participants, crossed out)

Net ionic: AgNO₃ + NaCl

Ag⁺(aq) + Cl⁻(aq) → AgCl(s) (Na⁺ and NO₃⁻ are spectators)

Net ionic: strong acid + strong base

H⁺(aq) + OH⁻(aq) → H₂O(l)

Net ionic: Pb(NO₃)₂ + KI

Pb²⁺(aq) + 2 I⁻(aq) → PbI₂(s)

Net ionic: BaCl₂ + Na₂SO₄

Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s)

Net ionic: Na₂CO₃ + HCl

CO₃²⁻(aq) + 2 H⁺(aq) → H₂O(l) + CO₂(g)

Brønsted-Lowry acid

A proton (H⁺) donor

Brønsted-Lowry base

A proton (H⁺) acceptor

Arrhenius acid

Forms H₃O⁺ (hydronium) ions in water

Arrhenius base

Forms OH⁻ (hydroxide) ions in water

Why are Brønsted-Lowry definitions more inclusive?

They do not require OH⁻ or H₃O⁺ in solution; e.g., NH₃ is a Brønsted-Lowry base but not an Arrhenius base

The six common strong acids

HCl, HBr, HI, HNO₃, HClO₄, H₂SO₄

Strong acid behavior

Dissociates completely into ions in solution

H₂SO₄ as a diprotic strong acid

Loses its first H⁺ completely (H₂SO₄ + H₂O → H₃O⁺ + HSO₄⁻); the second is only partial (HSO₄⁻ + H₂O ⇌ H₃O⁺ + SO₄²⁻)

Strong bases

Hydroxides of Group 1 metals plus Ca(OH)₂, Sr(OH)₂, and Ba(OH)₂; strong electrolytes that dissociate completely

Weak acids: carboxylic acids

Contain a -COOH group and are always weak acids (e.g., acetic acid, CH₃COOH)

Weak bases

Most other bases; they are proton acceptors (e.g., NH₃ is a weak base)

Why is NH₃ a weak base?

It only partially reacts with water to form NH₄⁺ and OH⁻, so equilibrium leaves lots of unreacted NH₃

Amphiprotic substance

A substance that can act as either an acid or a base (donate or accept H⁺) depending on what it is with (e.g., H₂O)

Water acting as a base

HCl + H₂O → H₃O⁺ + Cl⁻ (water accepts H⁺)

Water acting as an acid

NH₃ + H₂O ⇌ NH₄⁺ + OH⁻ (water donates H⁺)

Redox reaction

A reaction involving a transfer of electrons (the key feature)

Oxidation

Species that donates/loses electrons (charge gets more positive)

Reduction

Species that accepts/gains electrons (charge gets more negative)

Half reactions

The two halves of a redox reaction: one is the oxidation component, one is the reduction component

Reducing agent

The species that is oxidized (it gives up electrons)

Oxidizing agent

The species that is reduced (it accepts electrons)

Redox example: Fe + Au³⁺

Fe(s) + Au³⁺(aq) → Au(s) + Fe³⁺(aq); Fe is oxidized (and is the reducing agent); Au³⁺ is reduced (Au(NO₃)₃ is the oxidizing agent)

Oxidation half-reaction (Fe + Au³⁺)

Fe(s) → Fe³⁺(aq) + 3 e⁻

Reduction half-reaction (Fe + Au³⁺)

Au³⁺(aq) + 3 e⁻ → Au(s)

Why can Fe donate electrons easily?

Its electrons are delocalized

Why use a net ionic equation for redox?

Determining the net ionic equation makes it easier to see the electron transfer

Oxidation number rule: sum

The oxidation numbers of all atoms in a species add up to the overall charge (NaCl = 0; PO₄³⁻ = -3)

Oxidation number: elemental form

0 (e.g., V, W, H₂, O₂, O₃)

Oxidation number: Group 1 metal ions

+1

Oxidation number: Group 2 metal ions

+2

Oxidation number: hydrogen

+1 in most compounds

Oxidation number: fluorine

Always -1 in compounds

Oxidation number: oxygen

-2 in most compounds

Oxidation number: halogens (Group 17) at the end of a formula

-1

Oxidation number of Mn in KMnO₄

+7

Oxidation number of P in PO₄³⁻

+5

Oxidation number of S in H₂SO₄

+6

Oxidation number of N in NO₃⁻

+5

Oxidation number of Cr in Cr₂O₇²⁻

+6

Single displacement reaction

A + BC → B + AC; usually a metal (A) reacting with an aqueous salt (BC); almost always redox

How to tell if a single displacement reaction occurs

Consult an activity series; the added metal must be more reactive than the metal ion in the salt

Single displacement example

Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)

Concentration (general)

Concentration = amount of solute / amount of solution

Units of concentration

Mass percent (% m/m); mass-volume percent (% m/v); volume percent (% v/v); parts per million (ppm); parts per billion (ppb); molality; molarity

Molarity (M)

Moles of solute / liters of solution

Meaning of 0.2 M NaCl

0.2 mol of NaCl per 1 L of solution ("0.2 molar NaCl")

Making a solution: step 1

Calculate and measure out the mass of solute (e.g., 1.58 g KMnO₄ for 100 mL of 0.100 M)