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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, Na, K, Rb, Cs 3 11 19 37 55 ← atomic # 8 8 18 18 ← 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 these ions because AgCl(s) remain in solution is an insoluble (in H2O) solid

What solids are soluble? Solid classified as … Solubility in water (@ 25° C) soluble ≥ 0.1 mol L−1

sparingly soluble 0.01 to 0.1 mol L−1

insoluble ≤ 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 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 (more … eg. HCO3

− & ClO3−)

Cl−, Br− & I− all except … Ag+, Hg2

2+, Cu2+ &

Pb2+

SO4

2− all except … Ca2+, Sr2+, Ba2+ & Pb2+

_____________________________ CO3

2, NH4+ & everything else

PO43−, alkali metal

S2−, cation salts O2− & OH− NH4

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

_____________________________________________________________________________________

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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. HCl(g) + H2O(l) → H3O+(aq) + Cl−(aq) Cl −H O−H → Cl− + H−O+−H | | 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− δ+ δ− δ+ 2δ− δ+

A is more “electronegative” than H 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 Weak Acids (only ones) (examples) _____________________________________________________________________________________ HCl HF HOCl HBr H2CO3 H2SO3 HI CH3COOH H2SO4 H3PO4 HNO3 HNO2 HClO4 HClO2 H3BO3 HClO3 most acids are weak _____________________________________________________________________________________

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. HF(aq) + H2O(l) H3O+(aq) + F−(aq) → 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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eg. HCl(g) + NH3(g) → NH4Cl(s) (= NH4+Cl−)

H H δ− .. δ+ δ− | .. | :Cl − H + :N − H → :Cl: − + H − N+− H .. | .. | H H acid base weak base weak acid In general, HA + B A− + HB+

Weak Bases:

eg. NH3(aq) + H2O(l) NH4+(aq) + OH−(aq)

H H | .. | .. H − N: + H − O: H − N+− H + :O:− | | | | H H H H → incomplete ionization → @ equilibrium - small concentration of NH4

+ & OH− → larger conc. Of NH3

Strong Bases in H2O:

eg. NaOH(s) → NaOH(aq) = Na+(aq) + OH−(aq) OH−(aq) + H2O(l) H2O(l) + OH−(aq)

H2O(l)

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eg. NaH(s) + H2O(l) → H2(g) + OH−(aq) + Na+(aq) Na+ H− + H−O−H → H−H + −O−H + Na+ ↑ strong base & reducing agent

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 → −O−H + −O−H Hydrides: Li+H−, NaH, KH, … Amides (of group I): Na+NH2

−, … Aqueous acid-base rxns

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egs. NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l) KOH(aq) + HI(aq) → KI(aq) + H2O(l) base + acid → salt + 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:

HCN(aq) + H2O(l) H3O+(aq) + CN−(aq) weak acid base strong acid base

Bronsted acids: HCN & H3O+ stronger Bronsted bases: 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) SO4

2− + H3O+(aq) acid base base acid

HCO3

−(aq) + H2O H2CO3(aq) + OH−(aq) base acid acid base

CH3COOH + H2O CH3COO− + H3O+

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) ↑ ↑ reacts with H2O no rxn with H2O

NH4+ + H2O NH3 + H3O+

↑ ↑ weak acid 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 + OH−

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Carbon dioxide is (slightly) soluble in water. But it doesn’t exist as CO2(aq)

CO2(g) + H2O(l) H2CO3(aq) ↑ ↑

not very soluble carbonic acid

Carbonic acid is a weak acid …

H2CO3(aq) + H2O(l) HCO3−(aq) + H3O+(aq)

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 H+ F2 F− metals such as . . . . . Cl2

Cl− Na Na+ Br2 Br− K K+ I2 I− Fe2+ Fe3+ H2SO4 SO2 Al Al3+ Cr2O7

2− Cr3+ C CO or CO2 (in acid) MnO4

− 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 eg. Cl2(g) + 2Br−(aq) → Br2(aq) + 2Cl−(aq)

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 ON(H) = +1 NaH ON(H) = −1 5. 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 ⇒ ON(S) = +4 ↑ -2 SO3 VI state of S ON(S) + 3 ON(O) = 0 ⇒ ON(S) = +6 ↑ -2 H2SO4 VI state of S 2 ON(H) + ON(S) + 4 ON(O) = 0 ⇒ ON(S) = +6 ↑ ↑ +1 -2 HSO4

− VI state of S ON(H) + ON(S) + 4 ON(O) = -1 ⇒ ON(S) = +6 ↑ ↑ +1 -2

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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 +δ −δ eg. 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) + SO4

2−(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 SO4

2−(aq) + Br−(aq) → SO2(g) + Br2(aq) 2. Separate half-reactions &

balance atoms other than O & H SO4

2−(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

+6 −1 +4 0

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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 SO4

2−(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 …