Chapter 5 Types Of Reactions

CHEMISTRY What is the world made of? The material world that we live in is the macroscopic counter part to a microscopic world of atoms & molecules. We will become familiar with the various atoms & the molecules they form. Also, the states of matter & physical or chemical change & the tools we use to describe matter & its changes measurement & mathematical relationships We begin with a brief review of … __________________________________________________

The Periodic Table Mendeleev (1869) When the elements are ordered according to atomic mass, the chemical & physical properties vary in a periodic fashion eg. Li, 3

Na, 11 8

K, 19 8

Rb, 37 18

Cs 55 18

← atomic # ← spacing

are all similar ⇒ they form a “group”, the alkali metals - a column in the table

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The position (order) of an element in periodic table = atomic number = # of protons in nucleus of atom KNOW FORM OF TABLE KNOW THE FIRST 20 ELEMENTS → ATOMIC NUMBER, SYMBOL, NAME, & POSITION IN TABLE. Columns form groups → labeled by roman numerals Rows form periods

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Group I: Alkali metals Li, Na, K, Rb, Cs, Fr soft metals the most reactive metals react with most nonmetals to form ionic salts (not with noble gases though) egs. 2 Li(s) + H2(g) → 2 LiH(s) lithium hydride 2 Na(s) + Cl2(g) → 2 NaCl(s) sodium chloride - table salt __________________________________________________ Group II: Alkaline earth metals harder, higher melting points (than alkali metals) react more slowly with nonmetals egs. Ca(s) + H2(g) → CaH2(s) 2 Mg(s) + O2(g) → 2 MgO(s) Salts of group I and II metals with nonmetals are generally ionic (Be is an exception). → Electrical conductivity

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Electrical conductivity ionic & covalent solids do not conduct ionic liquids (high T) do conduct covalent liquids do not conduct eg. pure water has an extremely low conductivity However, dissolve a metal halide in H2O - the resulting solution is a good conductor. ⇒ metal halides are … Electrolytes → produce conducting solutions in water Covalent compounds are Non-electrolytes → produce non-conducting solutions in water Solubility & Precipitation Reactions A precipitation reaction results when an insoluble solid is formed from solvated ions brought together by mixing solutions of soluble substances eg.

AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq) ↑ solid particles precipitate out of solution

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Ag+(aq) + NO3−(aq) + Na+(aq) + Cl−(aq) → AgCl(s) + Na+(aq) + NO3−(aq) ↑ ↑ ↑ precipitates because AgCl(s) is an insoluble (in H2O) solid

these ions remain in solution

What solids are soluble? Solid classified as …

soluble insoluble

Solubility in water (@ 25° C)

≥ 0.1 mol L−1 sparingly soluble 0.01 to 0.1 mol L−1 ≤ 0.01 mol L−1

eg. most sulfates (SO42−) are soluble exceptions: Ca2+, Sr2+, Ba2+ & Pb2+ sulfates Learn: Table 4.1 on page 139 Solubility refers to equilibrium between solid & ions in saturated solution eg. Ag+(aq) + Cl−(aq) AgCl(s) net ionic equation corresponding to above rxn Equilibrium lies far to the right – replace

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by →

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Some Solubility Rules Table 4.1 on page 139 Almost all salts of NH4+ & the alkali metal cations are soluble. Otherwise solubility is classified according to anion … Anion

Soluble

Insoluble

____________________________________________________________________________________

NO3−, ClO4−, CH3COO−

all

Cl−, Br− & I−

all except …

SO42−

all except …

(more … eg. HCO3− & ClO3−)

Ag+, Hg22+, Cu2+ & Pb2+

Ca2+, Sr2+, Ba2+ & Pb2+ _____________________________

CO32, PO43−, S2−,

NH4+ & alkali metal cation salts

everything else

O2− & OH−

NH4+, alkali metal & larger alkaline earth (beginning with Ca2+) cation salts

everything else

_____________________________________________________________________________________

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ACID BASE REACTIONS Arrhenius definition of an acid: produces aqueous solution containing H+(aq) → actually H3O+(aq) is a better description of the H+ ion in acidic aqueous solution eg. Cl −H

HCl(g) + H2O(l) → H3O+(aq) + Cl−(aq) O−H | H

→

Cl−

+

H−O+−H | H

Aqueous acid solutions: → sour (eg. vinegar) → change color of indicator (eg. phenolphthalein: red → clear) → react with many metals to produce H2(g) i.e. H3O+ is an oxidizing agent eg. Zn(s) + 2 H+(aq) → Zn2+(aq) + H2(g) Fe(s) + 3 H+(aq) → Fe3+(aq) + 3/2 H2(g) In an acid electrons are pulled away from an H – making the H more likely to come off as H+ H−A egs H − Cl, H−S−H δ+ δ− A is more “electronegative” than H

δ+

δ−

δ+ 2δ− δ+

Contrast: Arrhenius bases form OH−(aq) in water eg. NaOH(s) + H2O(l) → Na+(aq) + OH−(aq) + H2O(l)

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BRONSTED-LOWRY DEFINITION Acid is a proton donor Base is a proton acceptor i.e. an acid-base reaction is a proton transfer reaction HA + H2O H3O+ + A− acid base acid base → an acid has an H which it can donate as H+ → a base has a lone pair of e−s which can accept the proton _____________________________________________________________________________________

Strong Acids (only ones)

Weak Acids (examples)

_____________________________________________________________________________________

HCl

HF

HOCl

HBr

H2CO3

H2SO3

HI

CH3COOH

H2SO4

H3PO4

HNO3

HNO2

HClO4

HClO2

HClO3

most acids are weak

H3BO3

_____________________________________________________________________________________

Learn the strong acids NOTE: strength is not related to concentration

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Strong acids HA(aq) + H2O(l)

H3O+(aq) + A−(aq)

→

→ completely ionized in H2O → high electrical conductivity Weak acids eg. → → →

H3O+(aq) + F−(aq)

HF(aq) + H2O(l)

lots of HF in solution smaller electrical conductivity weak acids are incompletely ionized in H2O Redox rxns of acids reduction

eg. Mg(s) + 2H3O+(aq) → Mg2+(aq) + H2(g) + 2H2O(l) oxidation Mg → Mg2+ + 2e− 2e− + 2H3O+ → H2 + 2H2O Bases → →

molecule or ion proton acceptor (must have lone e− pair)

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HCl(g) + NH3(g) → NH4Cl(s) (= NH4+Cl−)

eg. −

H δ | .. :N − H → :Cl: − + | .. H base weak base −

+

δ .. δ :Cl − H ..

+

acid In general,

H | H − N+− H | H weak acid

A− + HB+

HA + B

Weak Bases: eg. NH3(aq) + H2O(l) H | H − N: | H

+

NH4+(aq) + OH−(aq) H | H − N+− H | H

.. H − O: | H

+

.. :O:− | H

→ incomplete ionization → @ equilibrium - small concentration of NH4+ & OH− → larger conc. Of NH3 Strong Bases in H2O: eg. NaOH(s)

H2O(l)

→

NaOH(aq) = Na+(aq) + OH−(aq)

OH−(aq) + H2O(l)

H2O(l) + OH−(aq) 11

→

eg. NaH(s) + H2O(l) Na+ H− + H−O−H ↑ strong base & reducing agent

→

H2(g) + OH−(aq) + Na+(aq) H−H

+

−

O−H

+

Na+

Strong Bases (ionic) Hydroxides: Li+OH−, NaOH, KOH, RbOH, & CsOH Ca2+(OH−)2, Sr(OH)2 & Ba(OH)2 Not very soluble

Mg(OH)2 is even less soluble however, it dissolves in acidic solution, but not basic solution => it is a basic hydroxide like the above Corresponding oxides react with H2O to form OH− eg. Li2O, CaO Li2O(s) + H2O(l) → 2 LiOH(aq) O 2− + H−O−H Hydrides:

→

−

O−H +

Li+H−, NaH, KH, …

Amides (of group I): Na+NH2−, … Aqueous acid-base rxns

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−

O−H

egs. NaOH(aq) KOH(aq) base

+

HCl(aq)

+ HI(aq) → + acid →

→

NaCl(aq)

KI(aq) salt

+

H2O(l)

+ H2O(l) + water

Na+ + OH− + H3O+ + Cl− → Na+ + Cl− + 2H2O K+ + OH− + H3O+ + I− → K+ + I− + 2H2O OH− + H3O+ → 2H2O common rxn net ionic rxn Conjugate acid-base pairs: H3O+(aq) + CN−(aq) strong acid base

HCN(aq) + H2O(l) weak acid base Bronsted acids: Bronsted bases:

HCN & H3O+ stronger H2O & CN− Stronger

Acid-base "Half reactions": HCN → H+ + CN− acid 1 base 1

← conjugate acid-base pair

H2O + H+ → H3O+ base 2 acid 2

← conjugate acid-base pair

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Examples of Acid-base reactions HBr(aq) + H2O(l) → H3O+(aq) + Br−(aq) acid base acid base O2−(aq) + H2O(l) → OH−(aq) + OH−(aq) base acid base acid HSO4−(aq) + H2O(l) acid base

SO42− + H3O+(aq) base acid

HCO3−(aq) + H2O base acid

H2CO3(aq) + OH−(aq) acid base CH3COO− + H3O+

CH3COOH + H2O

Strong acid reacts with strong base to form (very) weak acid & (very) weak base. → Strong acids are conjugate to very weak bases → Strong bases are conjugate to very weak acids. Weak acids are conjugate to weak bases

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HYDROLYSIS OF SALTS Cations and anions can act as acids and bases … For example, NH4Cl forms an acidic solution when dissolved in water. NH4Cl(aq) consists of NH4+(aq) & Cl−(aq) ↑ ↑ no rxn with H2O reacts with H2O NH4+ + H2O ↑ weak acid

NH3 + H3O+ ↑ weak base

⇒ an acidic solution Contrast … NaCl(aq) is a neutral solution because neither Na+ or Cl− react with water Another example … LiOCl forms a basic solution when dissolved in water because of the acid-base reaction, OCl− + H2O

HOCl

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+

OH−

Carbon dioxide is (slightly) soluble in water. But it doesn’t exist as CO2(aq) CO2(g) + H2O(l) ↑ not very soluble

H2CO3(aq) ↑ carbonic acid

Carbonic acid is a weak acid … HCO3−(aq) + H3O+(aq)

H2CO3(aq) + H2O(l)

Soda-pop & beer are slightly acidic. Note that coca-cola has additional phosphoric acid. Adding acid to baking soda (NaHCO3) brings about the reverse of these two reactions. Bubbles of CO2(g) result.

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Oxidation Reduction Reactions REDOX eg.

2 Na(s) + Cl2(g) → 2NaCl(s) two “half reactions”

❶

Na → Na+ + e−

e− loss

❷

Cl2 + 2e− → 2Cl−

e− gain

Balance by balancing electrons (2× first reaction plus 1× the second reaction) … 2 Na → 2Na+ + 2e− 2e− + Cl2 → 2Cl−

____________________________

2Na + Cl2 → 2Cl− An e− transfer rxn Loss of Electron = Oxidation Gain of Electron = Reduction mnemonic: LEO the lion says GER

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An oxidizing agent (eqs. Cl2 & O2) can oxidize other substances → it takes e−s → it gets reduced A reducing agent (eqs. Na & NaH) can reduce other substances → it gives e−s → it gets oxidized Another example … Na(s) + H2O(l) → Na+(aq) + OH− (aq) + ½ H2(g) ~~~~~~~~~~~~~~~~~~~~~ Na(s) → Na+(aq) + e− H2O(l) + e− → OH− (aq) + ½ H2(g)

____________________________________________

Na(s) + H2O(l) → Na+(aq) + OH− (aq) + ½ H2(g)

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some … oxidizing agents (product) reducing agents (product) __________________________________________________ O2

O2−

H2

F2

F−

metals such as . . . . .

Cl2

Cl−

Na

Na+

Br2

Br−

K

K+

I2

I−

Fe2+

Fe3+

H2SO4

SO2

Al

Al3+

Cr2O72−

Cr3+ (in acid)

C

CO or CO2

MnO4−

H+

Mn2+ H2S SO2 or SO3 (in acid) MnO2 (in base) __________________________________________________

metals easily lose e−s → get oxidized → reducing agents → nonmetals easily gain e−s → get reduced Halogens are strong oxidizing agents →

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F2 > Cl2 > Br2 > I2 Cl2(g) + 2Br−(aq) → Br2(aq) + 2Cl−(aq)

eg.

2 Br− → Br2 + 2 e− 2 e− + Cl2 → 2 Cl− ⇒

Br− is oxidized to Br2 Cl2 is reduced to Cl−

Cl2 can oxidize Br− & I− Br2 can oxidize I− - not Cl− I2 will not oxidize Br− or Cl− F2 is so strong it can even oxidize H2O 2 F2(g) + 2 H2O(l) → 4 HF(aq) + O2(g) in contrast Cl2(g) + H2O(l)

HCl(aq) + HOCl(aq)

F2 is the strongest oxidizing agent ⇒ it is very difficult to oxidize F− to F2 F2(g) is produced by electrolysis of molten NaF →

order of halide reducing strength F− < Cl− < Br− < I−

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OXIDATION NUMBERS & OXIDATION STATES measures the extent of oxidation of an atom within a compound RULES for assigning oxidation numbers : 1.

Atom in elemental form has ON = O

2.

ON of monatomic ion = charge on ion egs. K+ ON = +1 Br− ON = −1

3.

ON = −1 for F in all it’s cmpds → F in most “electronegative” element

4.

ON = +1 for H in all its cmpds except metal hydrides for which ON(H) = −1 egs. H2O NaH

5.

ON(H) = +1 ON(H) = −1

ON(O) = −2 except in i) F2O where ON(O) = +2 & ii) cmpds with O−O bonds → X−O−O−Y ON(O) = −1 eg. H−O−O−H

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6. Other halogen atoms (Cl, Br & I) have ON = −1 unless they are bonded to more EN atom (i.e. O or halogen above in group) - in which case, they have +ve ON. 7. Sum of all ON = net charge → sum over all atoms of neutral molecule = O → sum over all atoms of ion = charge of ion egs. SO2 IV state of S ON(S) + 2 ON(O) = 0 ↑ -2

⇒

SO3 ON(S) + 3 ON(O) = 0 ↑ -2 H2SO4

ON(S) = +4

VI state of S ⇒

ON(S) = +6

VI state of S

2 ON(H) + ON(S) + 4 ON(O) = 0 ↑ ↑ +1 -2 HSO4−

⇒

ON(S) = +6

VI state of S

ON(H) + ON(S) + 4 ON(O) = -1 ↑ ↑ +1 -2

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⇒

ON(S) = +6

What does oxidation number (ON) represent? It is the effective number of electrons lost or gained relative to the elemental form of the atom. eg.

K

→

K+ + e−

ON(K) in K+ is +1 K (elemental state of potassium) loses one electron to get to K+ state Br2 + 2 e− →

2 Br−

ON(Br) in Br− is –1 Each atom of Br2 (elemental state of bromine) gains one electron to get to Br− state eg.

+δ −δ H − Cl

ON(Cl) = –1

there is a partial transfer of one electron of

H⋅

to

.. :Cl⋅ ..

when HCl is formed Oxidation number treats polar bonds as though there is a complete transfer of an electron (one for each bond formed)

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Balancing redox reactions in cases with solvent (specifically water) participation eg.

4 H+(aq) + SO42−(aq) + 2 Br−(aq) → 2 H2O(l) + SO2(g) + Br2(aq)

• •

This reaction consumes H+(aq) & produces H2O(l) How do we determine the coefficients?

•

The forward reaction is favored under conditions of high concentration of H+(aq) – i.e. high [H+] – this is why concentrated H2SO4(aq) is a strong oxidizing agent, whereas dilute H2SO4(aq) is not - concentrated H2SO4(aq) is required to oxidize Br−(aq), as shown here Balancing Redox Reactions in Acid or Base Medium

1. Identify oxidizing & reducing agents & write unbalanced reaction SO42−(aq) + Br−(aq) → SO2(g) + Br2(aq) +6

−1

+4

0

2. Separate half-reactions & balance atoms other than O & H 2− SO4 (aq) + 2 e− → SO2(g) & 2 Br−(aq) → Br2(aq) + 2 e− These same steps are used for all redox reactions. The next step is specific to cases of solvent participation … 3. Complete balancing the half-reactions – if not already balanced 24

If solution is acidic … • first balance O ’s by adding H2O ’s to the O deficient side of reaction • then balance H ’s by adding H+ ’s If solution is basic • balance as in acidic solution, then add OH− ’s to both side, neutralizing H+ ’s resulting from previous step • use H+ + OH− = H2O In above example, only the SO42−(aq) half-reaction needs balancing 4 H+(aq) + SO42−(aq) + 2 e− → SO2(g) + 2 H2O The last step applies to all redox reactions with or without solvent participation. 4. Add the half-reactions, balancing electrons. In some cases, we will need a multiple of one or both half-reactions in order to cancel the electrons. 4 H+(aq) + SO42−(aq) + 2 e− → SO2(g) + 2 H2O 2 Br−(aq) → Br2(aq) + 2 e−

4 H+(aq) + SO42−(aq) + 2 Br−(aq) → 2 H2O(l) + SO2(g) + Br2(aq)

Figures denoted by ♣ are courtesy of …

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