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Chapter 4

CHM 111 Chapter Notes - Chapter 4: Molar Mass, Zirconium, Disproportionation

Course Code
CHM 111

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Solution: homogeneous mixture of 2+ substances
Solute: substance present in smaller quantities
Solvent: substance present in larger quantities (eg. water = universal solvent)
Aqueous solutions: solute is initially liquid or solid, solvent is water
Electrolyte: substance dissolved in water becomes a solution that conducts electricity
Non-electrolyte: does not conduct electricity when dissolved in water
o Strong electrolyte: solute is assumed to be 100% dissociated in ions in the solution,
non reversible (eg. HCl, HNO3, HCLO4, H2SO4, NaOH, Ba(OH)2, ionic compounds
o Weak electrolyte: <100% dissociated, reversibility (eg. CH3COOH, HF, HNO2, NH3,
o Examples of non-electrolytes: urea, methanol, ethanol, glucose, sucrose
Hydration: process where ions are surrounded by H20 molecules arranged in a specific manner
→ stabilization of solution, prevent anions & cations from combining
Acids & bases are electrolytes
Ionization: separation of acids & base into ions
= reversible reaction - typical for weak electrolytes
Chemical equilibrium: state where no net change is observed → acid molecules ionize as fast
as ions recombine
Precipitation reactions: result in the formation of insoluble product/precipitate
o Precipitate: insoluble solid that separates from a solution (usually involves ionic
Metathesis reaction/double displacement reaction: involve exchanging of parts b/w 2
Solubility: maximum amount of solute that will dissolve in a given amount of solvent at a
specific temperature
o Used to help predict whether a precipitate will form when a compound is added to a
solution or when 2 solutions are mixed
Soluble: substance is able to dissolve a fair amount when added to H2O
Slightly soluble
Insoluble: dissolves only a tiny bit to a certain extent
Molecular equations: formulas of compounds written as though all species existed as
molecules/whole units
Ionic equation: shows dissolved species as free ions
o Spectator ions: ions that are not involved in the overall reaction, can be eliminated
from ionic equation → creates net ionic equation
Writing Ionic Equations
1. Write balanced molecular equation for the reaction, use table to determine solubility
of each molecule & see if there will be a precipitate
2. Write the ionic equation - compound that is not the precipitate should be shown as
free ions
3. Identify & cancel spectator ions, write net ionic equation
4. Check charges & # atoms for balanced equation
Acids: dissolve in water to form H+ ions
o Have sour taste, cause color changes in plant dyes, react w/ some metals to produce
H2(g), react w/ (bi)carbonates to produce CO2(g), aqueous acid solutions conduct
o Bronsted Acid: proton donor, does not require aqueous solution
Bases: dissolve in water to form OH- ions
o Have bitter taste, slippery (soaps), color changes in plant dyes, aqueous solutions
conduct electricity
o Bronsted Base: proton acceptor, does not require aqueous solution
H+ is simply a proton → cannot exist as separate entity in an aqueous solution due to
attraction to O atom in H2O
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o Exists instead as H3O+ (hydrated proton, hydronium ion)
Monoprotic acids: each unit yields 1 H+ upon ionization (HCl, HNO3, CH3COOH)
Diprotic acids: each unit yields 2 H+ in 2 steps (H2SO4)
Triprotic acids: yield 3 H+ ions (H3PO4)
Strong acids: HCl, HBr, HI, HNO3, H2SO4, HCLO4
Weak Acids: HF, HNO2, H3PO4, CH3COOH
Neutralization reactions b/w acids & bases
o Salt: ionic compound made up of cation other than H+ & an anion other than OH- or
acid + base → salt + water
o Some salts react w/ carbonates to form gases
Carbonates (CO32- ion)
Bicarbonates (HCO3- ion)
Sulfites (SO32- ion)
Sulfides (S2- ion)
Redox reactions: electron transfer reactions, not always aqueous
o OIL RIG (oxidation is losing electrons, reduction is gaining electrons)
o Half-reaction: shows electrons involved in the red or ox part of a redox reaction
o Element that is oxidized acts as the reducing agent
o Element that is reduced acts as the oxidizing agent
Oxidation #s: (oxidation state) # charges the atom would have in a molecule (or ionic
compound) if electrons were transferred completely
o Assigning Oxidation #s
1. Free elements - each atom has oxidation # of zero (eg. H2, Br2, K, P4 = 0)
2. Monatomic ions - oxidation # is the charge of the ion
a. Alkali metals = +1, Alkaline earth metals= +2, Al = +3
3. Oxidation # of O is -2, except for H2O2 peroxide ion (O22-) = -1
4. Oxidation # of H is +1, except when bonded to metals in binary compounds
(eg. LiH, NaH) = -1
5. F has oxidation # of -1 in all compounds, other halogens have negative #s
when in compounds as halides, but have positive #s when combined w/ O
(oxoacids, oxoanions)
6. Neutral molecule: sum of oxidation #s = 0
7. Polyatomic ion: sum of oxidation #s = net charge of ion
8. Not always whole #s - oxidation # of O in superoxide ion O2- is ½
Metallic elements = positive oxidation #s; nonmetallic = positive or negative oxidation #s
Highest possible oxidation # for Groups 1A - 7A is its group #
Transition metals have several possible oxidation #s
Combination Reactions: 2+ substances combine to form a single product
Decomposition Reactions: compound gets broken down into 2+ compounds
Combustion Reactions: substance reacts w/ oxygen, release of heat & light
Displacement Reactions: ion or atom in a compound is replaced by an ion or atom of another
o Hydrogen displacement: all alkali & some alkaline earth metals (Ca, Sr, Ba) will
displace H2(g) from cold water
o Metal Displacement: metals in a compound displace another metal in its elemental
o Activity series/electrochemical series - summary of results of many possible
displacement reactions
o Halogen displacement: F2>CL2>Br2>I2 - F2 is so reactive it attacks H2O - cannot be
aqueous reactions
Recovering halogens from halides requires oxidation
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