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Chemistry 205 > Ch. 3 > Flashcards

Ch. 3 Flashcards

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1
Q

Law of Conservation of Mass

who, what?

A
  • John Dalton: chemical change involves a reorg of the atoms in one or more substances
  • Antoine Lavoisier (18th c): Matter can neither be created nor destroyed.
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2
Q

Balancing Chemical Equations

steps

A
  1. Find the easiest element (one that’s in only one species on each side)
  2. Balance its partner
  3. Do the rest but save the toughest to the end
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3
Q

Equilibrium

def

A
  • Forward and reverse reactions occur at same rate
  • No net change
  • Reactions or products can be favored, but what’s important is the reaction RATE being equal
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4
Q

Polar molecules

def

A
  • A molecular compound with polar covalent bonds has a slightly more positive region and a slightly more negative one.
  • Bent geometry (V-shaped) like water
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5
Q

Polar Molecules vs Water

A
  • Polar molecular substances are soluble in water
  • LIKE DISSOLVES LIKE
  • Water molecules interact favourably with polar molecules due to electrostatic attraction between partial charges
  • Water molecules fully surround (hydrate) the solute
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6
Q

Dissolving Ionic Compounds

A
  • Bonds between ionic compounds break
  • Polyatomic ions remai intact
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7
Q

Dissolving Molecular Compounds

A
  • Intact molecules disperse
  • Covalent bonds do not break
  • ex: CO2(s) will simply become CO2(aq)
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8
Q

Molarity

A
  • Moles of solute per litre of solution
  • c=n(solute)÷v(solution)
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9
Q

Qualitatively describing concentration

A
  • Concentrated solution: large ratio of solute to solvent
  • Dilute solution: small ratio of solute to solvent
  • Stock solution: concentrated solution often used in a lab. diluted and used when needed
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10
Q

Preparing diluted solutions

A
  • If solid into a solution: c=n÷v
  • If pre-diluted stock solution into another more diluted solution: C1V1=C2V2
  • Can combine the two equations if we need moles
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11
Q

Water or Acid First when Adding?

A
  • Always add acid TO water
  • Heat is released, so if it splatters, at least it was diluted first (i.e. the stuff that will sputter is ultra-diluted, whereas if we added water to the acid, the stuff that would sputter is lots of acid with just a touch of water)
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12
Q

Electrolyte VS Solubility

A
  • Highly soluble = strong electrolyte = high concentration of ions
  • Slightly soluble = weak electrolyte = low concentration of ions
  • Insoluble = nonelectrolyte = contractration of ions is zero
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13
Q

Reasons for differences in ion solubility

A
  • Too complex to predict by theory alone
  • Complexe balance ion-ion vs ion-solvent interactions
  • Often, ions with high charge do not get sufficient charge compensation by being surrounded by water, so they have low solubility
  • If one ion has a low charge, it’s probably soluble.
  • If both are +2 or +3 or -2 or -3, it likely won’t dissolve (some exceptions)
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14
Q

Always soluble polyatomic ions

A
  • NH4+ (ammoniums)
  • Alkali metals, group 1A
  • Nitrates NO3-
  • Acetates CH3COO -
  • Chlorates ClO3-
  • Perchlorates ClO4-
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15
Q

Sometimes Soluble Polyatomic Ions

A
  • Chlorides Cl- , bromides Br- , iodides I- are soluble except when combined with:
  1. Ag+
  2. Hg2+
  3. Pb2+
  • Sulfides SO42- except CaSO4 , SrSO4 , BaSO4 , PbSO4
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16
Q

Always insoluble polyatomic ions

A
  • When both ions have charges 2 and above

The following are insoluble unless combines with a polyatomic ion that is always soluble, like from group 1A

  • Carbonates CO32-
  • Phosphates PO43-
  • Oxalates C2O42-
  • Chromates CrO42- (soluble with Mg2+)
  • Sulfides S2- (a bit soluble with group 2A
  • Oxides
  • Hydroxides (but the following two ARE soluble: Ba(OH)2 and Sr(OH)2)
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17
Q

Strong Electrolyte

def

A
  • Any compound that fully ionizes in a solution
  • Freely mobile solvated ions
  • Solution conducts electricity
    *
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18
Q

Nonelectrolyte

def

A
  • Any compound that does NOT dissociate into ions at all
  • No ions
  • Solution will not conduct electricity
  • Most molecular compounds
  • Compound reactant same as product but might go from s to aq
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19
Q

Predicting species present in aqueous solution

A
  1. Ionic: ions dissociate to strong electrolyte
  2. Molecular: nonelectrolyte if only solvated molecules
  3. Molecular: Weak electrolyte if molecules + some ions (weak acid or base) less than 100% reacts with H2O, so it’s a weak acid or base
  4. Molecular: Strong electrolyte if all solvated ions (strong acids or bases) 100% reacts with H2O, so it’s a strong acid or base
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20
Q

Electrolyte Solution vs Substance

A
  • Solution if contains ions
  • Substance if yields ions when dissolved
  • Weak ionic substance: ionic, but not much dissolves or ionizes partially when dissolved
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21
Q

Double Displacement

A
  • Mixing two substances that break apart and stick + to - of the other pair to form new products
  • AKA precipitation reactions
  • When a precipitate of a solid product is formed via counter-ion exchange upon mixing water-soluble ionic reactants

ex: KCl(aq) + AgNO3(aq) = KNO3(aq) + AgCl(s)

K reacts with NO3 and Ag reacts with Cl = double displacement

KNO3 dissolved in supernatant

AgCl precipitated (can filter out)

22
Q

Supernatant

def

A

denoting the liquid lying above a solid residue after crystallization, precipitation, centrifugation, or other process.

23
Q

What is IN the reactant solution?

A

Separated hydrated ions (cations and anions)

24
Q

What happens when we mix the separated hydrated ions in the reactants

A
  • Ions encounter new counter-ions via collisions
  • If new ion pair separates: soluble salt (stays in supernatant)
  • If new ion pair sticks together: insoluble salt (precipitate)
25
Q

How to collect soluble and insoluble salts?

A

Filtration!

  1. Collect insoluble salt on filter and rinse
  2. Evaporate solvent from supernatent/filtrate = collecting soluble salt
26
Q

Writing out Complete Ionic Equations

A
  • Write the strong electrolytes as ions on both sides
  • The insoluble ones get written as a polyatomic ion salt (precipitate)
  • The ones who don’t form into one and don’t even change phases and stay as floating ions are called spectator ions. They didn’t actually react.
  • Keep the spectator ions here.
27
Q

Writing out Net Ionic Equations

A
  • Only show species that changed phase
  • Spectator ions are not included
  • Must balance atoms AND charge
  • Summarizes the overall reaction that occurred

ex: Ag+(aq) + Cl-(aq) → AgCl(s)

28
Q

Writing out Molecular Equations

A
  • Called molecular even though ions… silly
  • Reactants and products written as compounds

ex: AgNO3(aq) + KCl(aq) → AgCl(s) + KNO3(aq)

29
Q

What do we know about ACIDS

A
  • Sour taste
  • Turns litmus paper red
  • Produce bubbles of CO2 when added to carbonate rocks
  • React with many metals to produce bubbles of H2
  • When dissolved in water, INCREASE the concentration of H+
30
Q

What do we know about BASES

A
  • Bitter taste (soap)
  • Slippery feel
  • Turns litmus paper blue
  • Counter-act the acidic nature of acids
  • When dissolved in water, increase the concentration of OH-
31
Q

Arrhenius theory on acids and bases

A
  • Late 19th century
  • Acids and bases yield ions when dissolved in water.
  • If completely dissolved, strong acid or base
  • If partially, it’s weak
32
Q

Bronsted-Lowry definition of Acids

A
  • Acids are H+ doners (AKA proton givers)
  • Strong acids: HCl, HBr, HI, H2SO4, HNO3 , HClO4
  • The first three are halogens with H (but HF is a weak acid)
  • Produce H3O+ (what makes the solution acidic. It’s the H3O that makes it sour)
33
Q

Bronsted-Lowry definition of bases

A
  • Bases are H+ acceptors (stealers)
  • Bases steal H+ ions. That’s how acids lose them.
  • Strong: OH- (ex NaOH and others that have an OH at the end)
  • Weak: CO32-, PO43-, NH3 , amines
  • They accept H+ by sharing a lone pair to form a bond.
  • They actually rip off the H+
  • H2 O can act as a base
34
Q

What happens in an acid reaction?

A
  • H is covalently bonded to H2O but often shown as just H+ for simplicity.
  • For a strong acid: This new bond is stronger than in HA, so H+ does not go back to the HA (which has a weak covalent bond). Single arrow in equation
  • For a weak acid, the original HA bond is moderately weak, so not all break apart to form H3O = low yield of A- and H+. New bond in H3 is slightly weaker than HA, so mostly stay bonded as HA. Double arrow in equation.
  • Equilibrium favours the products
  • Acids actually react with water, not just dissolve
35
Q

Strong Acids to memorize

A
  1. HCl
  2. HBr
  3. HI

1-2-3 are hydrohalic acids (but HF not a strong acid)

  1. HNO3 Nitric acid
  2. HClO4 Perchloric acid
36
Q

Common Weak Acids

A
  1. HF (hydrofluoric acid)
  2. H3PO4 (phosphoric acid)
  3. H2CO3 (carbonic acid)
  4. CH3COOH (acetic acid)
  5. H2C2O4 (oxalic acid)
37
Q

What happens in a Base Reaction?

A
  • Two types of bases:
  1. Molecule that reacts with water to steal H+ to make OH- (not in this course)
  2. Soluble hydroxide salts
  • Salts that are not very soluble will be weak bases (weak electrolyte)
  • Ionic yield 100% OH without reaction

B reacts to pull the H off H2O

Weak base = low yield of BH+ and OH- (double arrow)

ex: NH3(aq) + H2O(l) ⇔ NH4+(aq) + OH-(aq) WEAK
ex: NaOH(aq) → Na+(aq) + OH-(aq) STRONG

38
Q

Common bases

A

Strong: with OH

Weak: ammonia (N with 3 H) and amines (N with 3 bonds)

39
Q

Acids with no H in the formula

A
  • Oxides of nonmetals
  • Acidic oxides react with water to produce substances that release H+ (i.e. they don’t have any H+ to get robbed of, but they react with water and then do have them.

CO2 + H2O → H2CO3 (carbonic acid)

SO3 + H2O → H2SO4 (sulfuric acid)

NO2 + H2O → HNO3 (nitric acid)

40
Q

Bases with no OH in the formula

A
  • Oxides of metals
  • Basic oxides react with water and produce substances that release OH-

CaO(s) + H2O(l) → Ca(OH)2(s)

41
Q

Tips to recognize an acid

A
  • Formula begins with an H (Hcl, HBr, HCN)
  • Formula contains -COOH (CH3COOH, HCOOH)
  • Nonmetal hydroxides (OH) and oxides (O) like HNO3, H2SO4, H3PO4
  • Protonated amines (+NH4, +HN(CH3)3 )
42
Q

How to recognize a base

A
  • Anions (not if H+ form a strong acid) NaHCO3 , CaCO3 , Na2SO3
  • Metal hydroxides Mx(OH)y and oxides Mx(O)y like NaOH, Mg(OH)2 , CaO
  • Ammonia and amines ( :NH3, :N(CH3)3 )
  • contain OH- or CO32-
43
Q

pH

A
  • In water, more OH- means less H3O+
  • So more concentrated base solution has lower concentration of H3O
  • Concentration of H3O- = 10-pH
  • pH = -log[H3O]
  • If solution has a higher [H3O+], then acidic solution
44
Q

pOH

A
  • [OH-] = 10-pOH
  • pOH = -log[OH-]
  • If solution has a higher [OH-], then basic solution
  • Very basic solution will have a high pH
  • pH + pOH = 14
45
Q

What happens when we mix an acid and a base

A

Molecular equation: HCl + NaOH → NaCl + H2O

Ionic: H3O+ + Cl- + Na+ + OH- → 2H2O + Na+ + Cl-

Net ionic: H3O+ + OH- → 2H2O (true for any strong acid + strong base rxn)

  • Strongest acid will mix with strongest base: species produced will be less reactive than reactants consumed
  • Often called neutralization reactions (but pH not necessarily neutral, but def less acidic or basic as before)
  • strong + strong or strong + weak rxns are quantitative (→) because even a weak base is capable of ripping off an H+ from a strong acid
  • The H+ ends up covalently bonded to the base molecule, via a stronger bond than it had to the acid molecule
  • The stronger bond drives the rxn
  • Imagine the reaction as such:
  • acid(HA) + base(B) → new base(A) + new acid(BH)
  • product favoured if new acid is weaker (holds onto its H better)
46
Q

Which side is favoured in an acid-base rxn

A

TREND: products favoured in

  1. strong A + strong B
  2. strong A + weak B
  3. weak A + strong B

The side with the weaker acid is favoured because it holds onto its H so more stable

47
Q

How to know what type of reaction will occur

A
  • Precipitation reaction: involves ion exchange
  • Acid-base reactions: forming a new covalent bond to H+ (which moved from the acid to the base, so both must be there)
  1. Strong electrolytes: use ion exchange - if products are soluble and ionic, then no reaction
  2. If acid and base react, products likely soluble. New strong A or B
48
Q

Gas-forming reactions

A

Can be several types of reactions AND gas-forming (always say gas-formin AND they type of rxn)

  1. Gas-forming redox (e- transfer) (certain metals + acid → dissolved metals + gas
  2. Gas-forming acid-base (H+ transfer).
  • Sometimes acid/base rxns produce gas.
  • Products of some acid-base rxns decompose to create gas. For example, when you see CO32- + H+, it creates H2CO3 which decomposes into CO2 and H2O
49
Q

How to know if gas will be produced?

A
  1. H2CO3‍ : This will decompose into carbon dioxide gas and water.
  2. H2SO3‍ : This will decompose into sulfur dioxide gas and water.
  3. H‍2S : This is already a gaseous product. (smells like rotten eggs)
  4. If NH4+ and OH- ions are produced, they will form NH3(g) and water.
50
Q

Oxidation States

A
  • Hypothetical ionic charge
  • One element is more electronegative, so it attracts to covalently shared electrons more, therefore is more negative g-
51
Q

Oxidation numbers

AKA oxidation states

A

We assign oxidation numbers (ONs) to elements using these rules:

  • Rule 1: The ON of an element in its free state is zero — examples are Al, Zn, H₂, O₂, N₂.
  • Rule 2: The ON of a monatomic ion is the same as its charge — examples are Na⁺ = +1; S²⁻ = -2.
  • Rule 3: The sum of all ONs in a neutral compound is zero. The sum of all ONs in a polyatomic ion is equal to the charge on the ion.
  • Rule 4: The ON of Group 1 metal in a compound is +1; the ON of a Group 2 metal in a compound is +2.
  • Rule 5: The ON of O in a compound is usually –2, except in peroxides like H₂O₂ and Na₂O₂, where it is -1.
  • Rule 6: The ON of H in a compound is usually +1, except in metal hydrides such as NaH or CaH₂, where it is -1.
  • Rule 7: The ON of F in a compound is always –1. Cl, Br, and I usually have an ON of –1, unless they are combined with O or F.
52
Q

Acid or base redox reactions steps

A
  1. Write the half reactions
  2. Balance the atoms other than O and H
  3. Balance oxygen by adding H2O
  4. Balance H by adding H+
  5. Balance charges by adding electrons
  6. Make the number of electrons equal and add OH- if it’s a base
  7. Add half reactions and make sure atoms and charges balance

Chemistry 205 (3 decks)

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