Ch. 3 Ionic Compounds, Formulas, Reactions

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Chemical Formula

$\ce{KMnO4}$, $\ce{Al(NO3)3}$, $\ce{CoCl2*6H2O}$

• Empirical formulas represent the simplest ratio of the atoms
  • used in ionic compunds
• Molecular formulas represent the actual number of each atom per molecule
  • used in covalent compounds
  • condensed formula: letters/numbers
  • structural formula: letters and bond lines
  • ball-and-stick model: literally balls and bond sticks, 3D
  • space-filling model: just balls, 3D

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Chemical Equations

Balance chemical equations using inspection method or ion-electron methodnoteCh. 12

Inspection Method

1. count number of each atom on each side
2. add coefficients to balance most complex molecule
3. repeat step 2 for each atom

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Combustion Example

$\ce{C6H6 + O2 -> CO2 + H2O}$
$\ce{C6H6}$ is most complex, so balance $\ce{C}$ and $\ce{H}$
$\ce{C6H6 + O2 -> \textcolor{#EE4266}{6}CO2 + \textcolor{#EE4266}{3}H2O}$
Finally, balance $\ce{O}$ and multiply to get whole numbers
$\ce{\textcolor{#EE4266}{2}C6H6 + \textcolor{#EE4266}{15}O2 -> \textcolor{#EE4266}{12}CO2 + \textcolor{#EE4266}{6}H2O}$

Double-Replacement Example

$\ce{Ba(OH)2 + H3PO4 -> H2O + Ba3(PO4)2}$
Start by balancing $\ce{Ba}$ from $\ce{Ba3(PO4)2}$
$\ce{\textcolor{#EE4266}{3}Ba(OH)2 + H3PO4 -> H2O + Ba3(PO4)2}$
Next balance $\ce{PO4}$ from $\ce{Ba3(PO4)2}$
$\ce{3Ba(OH)2 + \textcolor{#EE4266}{2}H3PO4 -> H2O + Ba3(PO4)2}$
Finally, we have 6 $\ce{H}$s and 6 $\ce{OH}$ groups so
$\ce{3Ba(OH)2 + 2H3PO4 -> \textcolor{#EE4266}{6}H2O + Ba3(PO4)2}$

Types of Chemical Equation

• Combustion: $\ce{Molecule + O2 -> CO2 + H2O}$
• Single-Replacement: $\ce{AB + C -> AC + B}$
• Double-Replacement: $\ce{AB + CD -> AD + CB}$
• Neutralization: $\ce{Acid + Base -> Salt + H2O}$
• Synthesis: $\ce{A + B -> AB}$
• Decomposition: $\ce{AB -> A + B}$
• Net Ionic: e.g. $\ce{Cl-(aq) + Ag+(aq) -> AgCl(s)}$
• Half-Reaction: e.g. $\ce{I2 + 2e- -> 2I-}$
• Oxidation-Reduction: one atom loses electron, another gains; combines 2 half-reactions

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Ionic Bonds/Salts

bond in which one atom takes the electron(s) of another

• cation (+ charge) attracted toward anion (- charge)
• representative (P 1, 2, 13-18) metals form cations; lose enough electrons to be isoelectrionic to noble gas
  • $\ce{M -> M^n+ + \textit{n}e-}$
  • period 4, 5, 6, metals may lose only p electrons
• representative nonmetals form anions; gain enough electrons to be isoelectronic to noble gas
  • $\ce{X + \textit{n}e- -> X^n-}$
• nonrepresentative transition metals form cations; can sometimes be predicted
  • may lose p electrons, s and p electrons, s and d electrons (half-filled d is stable)
• polyatomic ion: group of atoms that form an ion

Polyatomic Ions

$\ce{NH4^+}$: ammonium
$\ce{CO3^2-}$: carbonate
$\ce{HCO3^-}$: bicarbonate (hydrogen carbonate)
$\ce{ClO-}$: hypochlorite
$\ce{ClO2^-}$: chlorite
$\ce{ClO3^-}$: chlorate
$\ce{ClO4^-}$: perchlorate
$\ce{NO2^-}$: nitrite
$\ce{NO3^-}$: nitrate
$\ce{SO3^2-}$: sulfite
$\ce{HSO3^-}$: bisulfite (hydrogen sulfite)
$\ce{SO4^4-}$: sulfate
$\ce{HSO4^-}$: bisulfate (hydrogen sulfate)
$\ce{MnO4-}$: permanganate
$\ce{Cr2O7^2-}$: dichromate
$\ce{CrO4^2-}$: chromate
$\ce{S2O3^2-}$: thiosulfate
$\ce{PO4^3-}$: phosphate
$\ce{HPO4^2-}$: hydrogen phosphate
$\ce{H2PO4^-}$: dihydrogen phosphate

- enough of each ion is added to balance the compound charge

Naming

• cation name then anion name
• cations with multiple charges followed by parenthesis and roman numerals (eg $\ce{iron(II)}$)
• if monatomic anion, replace ending with -ide
• if polyatomic, refer to chart above

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Ionic Compounds in Solutions

• most dissolve in water, separate into aqueous ions
• Solubility Rules: soluble if containing any of the following
  • sodium cation
  • potassium cation
  • ammonium ion
  • $\ce{NO3^-}$ anion

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Predicting Reactions

Driving Forces

• double-replacement reaction likely occuring if
  • water forms (strongest)
  • precipitate (insoluble compound) forms
  • covalent compound (e.g. gas) forms
• stroner driving force allows prediction for whether reaction occurs or not

Net Ionic Equations

• if all the ions cancel out, no reaction occurs
  • i.e. there has to be a net ionic equation
• if there is none, there is probably no driving force
• in single-replacement, the activity series of the elements is what is important (see standard reduction potential table)