Download Descriptive Chemistry for Midterm Exam #2

Chemistry 1051
Descriptive Chemistry
Intersession
Material on Midterm #2
Note: This handout is meant to highlight important facts and reactions. Exam
questions can be based on anything in the assigned text pages. You must read
each section carefully.
Alkali Metals:
PHHM (9th Edition): pp 870-874
PHH (8th Edition): pp 873-876
Common properties:
1.
2.
3.
4.
The most reactive (active) group of metals.
Stored under oil to prevent rapid reaction with O2 or H2O.
Low melting and soft.
Characteristic colors in Bunsen flame.
lithium: red sodium: yellow potassium: purple
5. React by losing one electron to form M+ ion (isoelectronic
with nearest noble gas). Powerful reducing agents.
6. Reactivity increases down group as attraction for ns1
valence electron decreases with distance from nucleus.
Production: by electrolysis of melts (molten i.e. liquid salts)
e.g.
2 NaCl(l)
∆
⎯
⎯→
2 Na(l) + Cl2(g) in Down’s cell
oxidation of anode: 2 Cl⎯ → Cl2 + 2 e⎯
reduction of cathode: 2 Na+ + 2e⎯ → 2 Na
Uses Na
1.
to produce other metals by reduction of their compounds
KCl(l) + Na(l)
Li
∆
⎯
⎯→
NaCl(l) + K(g) T≥ 850°C
2.
3.
as a heat transfer medium in nuclear reactors
in sodium vapour lamps for outdoor lighting.
1.
as an anode in batteries due to its ease of oxidation and high
electron output relative to its low molar mass.
as an alloying agent to make high-strength, low density alloys with
Al and with Mg for use in aerospace and aircraft industries.
2.
Some Reactions: with water to form H2(g) and a strong hydroxide base
e.g. 2 K(s) + 2 H2O(l) → H2(g) + 2 KOH(aq)
(H reduced from +1 oxidation state to zero by K)
with carboxylic acids to form H2(g) and a salt
e.g. 2 Na(s) + 2 HC2H3O2(aq) → H2(g) + 2 NaC2H3O2(aq)
(H reduced from +1 oxidation state to zero by Na)
Further explanation of reaction of Alkali metals with H2O and carboxylic acids:
The metal atoms (M) donate an electron to either H atom of H2O or the acidic hydrogen
of the carboxylic acid (the one bonded to oxygen and listed first in the formula). In
either case, the hydrogen leaves as an H atom leaving a pair of electrons on the oxygen
and a negative charge. H atoms combine to form H2. The Alkali metal cation M+
combines with OH⎯ to form a strong base (reaction with H2O) or with the anion of the
carboxylic acid to form a salt.
Page 1 of 6
Questions:
PHHM (9th Edition): Chapter 21: 1a, 2a
PHH (8th Edition): Chapter 22: 15(a), 16(a).
Indicate the colours formed by Na, K and Li in a Bunsen flame.
Give two uses each for sodium and lithium.
Alkaline Earth Metals:
PHHM (9th Edition): pp 879-882
PHH (8th Edition): pp 880-884
Properties:
1. Less reactive than the alkali metal in same period but undergo similar
reactions e.g. with H2O and acids to yield H2(g).
2. Higher melting and harder than alkali metals
3. React by losing 2 electrons to form M2+ ion (isoelectronic with nearest
noble gas). Powerful reducing agents.
Production: 1. Mg, Ca, Sr obtained by electrolysis of melts
CaCl2(l)
∆
⎯
⎯→
Ca(l) + Cl2(g) in electrolysis cell
oxidation of anode: 2 Cl⎯ → Cl2 + 2 e⎯
reduction of cathode: Ca2+ + 2 e⎯ → Ca
2. Ca, Sr, Ba obtained by reduction with Al
∆
⎯
⎯→
3 BaO(l) + 2 Al(l)
Al2O3(s) + 3 Ba(l)
The production of Mg from seawater.
1.
Removal of Mg2+ ions from seawater by precipitation with strong base.
Mg2+(aq) + 2 OH⎯(aq) → Mg(OH)2(s)
2.
Formation of chloride salt solution by acid/base neutralization
Mg(OH)2(s) + 2 HCl(aq) → MgCl2(aq) + 2 H2O(l)
3.
Isolation of solid chloride salt by evaporation of water.
MgCl2(aq) → MgCl2(s)
4.
Reduction of Mg2+ to Mg metal by electrolysis of molten MgCl2.
MgCl2(l)
Uses
∆
⎯
⎯→
Mg(s) + Cl2(g)
Mg is a lightweight construction material.
Important Group 2 Compounds
PHHM (9th Edition) pp 882-884
PHH (8th Edition) pp 884-886
CaO: quick lime
Ca(OH)2: slaked lime
CaCO3: limestone
Page 2 of 6
Important reactions of Group 1 and Group 2 compounds:
1.
Thermal decomposition of carbonates:
CaCO3(s)
calcination
∆
⎯
⎯→
CaO(s) + CO2(g)
Note: Alkali metal carbonates are thermally stable except Li2CO3.
2.
Thermal decomposition of hydrogen carbonates:
2 NaHCO3(s)
∆
⎯
⎯→
Na2CO3(s) + CO2(g) + H2O(g)
∆
⎯→ MgO(s) + 2 CO2(g) + H2O(g)
Mg(HCO3)2(s) ⎯
(the MgCO3 initially formed, decomposes to yield MgO and CO2)
3.
Formation of plaster of Paris (dry wall, casts) from gypsum
CaSO4 · 2 H2O(s)
4.
∆
⎯
⎯→
CaSO4 · ½ H2O(s) + 3/2 H2O(g)
Carbonates as bases
react with acids to form CO2 due to decomposition of carbonic acid (H2CO3)
CaCO3(s) + 2 HCl(aq) → CaCl2(aq) + CO2(g) + H2O(l)
5.
Hydrogen carbonates can react with both acids and bases
NaHCO3(s) + NaOH(aq) → Na2CO3(aq) + H2O(l)
NaHCO3(s) + HCl(aq) → NaCl(aq) + CO2(g) + H2O(l)
6.
CO2 is an acidic oxide
CO2(g) + H2O(l) → H2CO3(aq) ⇌ H+(aq) + HCO3⎯(aq)
Carbonation: CO2(g) + Ca(OH)2(aq) → CaCO3(s) + H2O(l).
This is used as a lab test for CO2, i.e. formation of white solid when CO2 is
bubbled into an aqueous solution of Ca(OH)2 (lime water).
Further explanatory notes on carbonation:
Think of this reaction as occurring in 2 steps:
7.
Step 1:
CO2(g) + H2O(l) → H2CO3(aq)
Step 2:
H2CO3(aq) + Ca(OH)2(aq) → CaCO3(s) + 2 H2O(l)
Overall:
CO2(g) + Ca(OH)2(aq) → CaCO3(s) + H2O(l)
Group 1 and Group 2 oxides are basic oxides. They form ionic hydroxides when
they react with water.
Slaking:
CaO(s) + H2O(l) → Ca(OH)2(s)
Page 3 of 6
PHHM (9th Edition) Chapter 21: 15, 16
PHH (8th Edition) Chapter 22: 29, 30
Questions:
Write 4 chemical equations to describe the production of Mg from seawater.
Describe the lab test for CO2. Write a chemical equation for the reaction.
Identify the formulas for limestone, quicklime, slaked lime and lime water.
The Noble Gases:
PHHM (9th Edition): Section 22.1 pp 918-920
PHH (8th Edition): Section 8.4, pp 281-282 and 23.1 pp 907-908
Some uses of the noble gases.
1.
2.
3.
4.
5.
Argon/nitrogen mixtures are used in long-life light bulbs
Krypton and xenon are used in lasers
Helium and argon are used to protect materials which are sensitive to oxygen
and nitrogen from exposure to air.
Helium/oxygen mixtures are used in deep-sea diving to prevent the bends.
Liquid helium, whose temperature is just above zero Kelvin, is used to cool
metals to become superconductors and cool the electromagnets in NMR (nuclear
magnetic resonance) instruments used in research and MRI (magnetic
resonance imaging) devices in hospitals.
Xenon compounds:
First compound made in 1962 by Bartlett and Lohman
XePtF6 a compound thought to contain Xe+ and PtF6⎯ ions.
Xenon forms compounds with highly electronegative fluorine and oxygen. It has
the second lowest ionization energy of the Noble gases.
Xe oxidation state
example
+2
XeF2
+4
XeF4, XeOF2
+6
XeF6, XeO3
+8
XeO4, H4XeO6
Properties of xenon compounds
1.
2.
3.
4.
All are powerful oxidizing agents.
All are unstable in water, oxidizing it to O2(g).
2 XeF2(aq) + 2 H2O(l) → 2 Xe(g) + O2(g) + 4 HF(aq)
Fluorides are stable solids.
Oxides are explosively unstable
PHH (9th Edition): Chapter 22 # 8, 9, 15, 16
PHH (8th Edition): Chapter 23 # 17, 18, 19
Identify the gas mixtures used in long-life light bulbs and by deep-sea divers to
prevent the bends.
Write formulas for (a) a compound of Xe, F and O in which Xe has an oxidation
state of +4 (b) a compound of Xe and O in which Xe has an oxidation state of +8.
Which noble gases are used to protect sensitive materials from air?
Give the formula of the first Noble gas compound synthesized.
Which noble gases are used in lasers?
Which liquid noble gas is used to cool the electromagnets in hospital MRI
machines?
Questions:
1.
2.
3.
4.
5.
6.
Page 4 of 6
Hydrogen:
PHHM (9th Edition) pp 948 - 950
PHH (8th Edition) pp 286 - 288
Occurrence: found in more compounds than any other element on earth. It is the
most abundant element in universe.
Oxidation states: 0 in H2, +1 in compounds with other non-metals, −1 in metal
hydrides.
Industrial Preparation of H2:
This is carried out through the reduction of +1 oxidation state in H2O using electrolysis
or at high temperature where C(s), CO(g), and hydrocarbons act as strong reducing
agents, combining with the O atom of H2O. Electrolysis uses electrical energy to force
non-spontaneous redox reactions to occur at the cell electrodes.
Note: The anode and cathode half reactions which result in the overall cell reaction are
given for each important electrolysis reaction to provide insight into the actual
process involved.
Electrolysis of water:
2 H2O(l) → 2 H2(g) + O2(g) in dilute acid or base solutions.
2 H2O(l) → O2(g) + 4 H+(aq) + 4 e⎯
4 H2O(l) + 4 e⎯ → 2 H2(g) + 4 OH⎯(aq)
cell reaction:
oxidation of anode:
reduction of cathode:
∆
⎯
⎯→
CO(g) + H2(g) ∆ means at high T
∆
⎯→
CO(g) + H2O(g) ⎯
CO2(g) + H2(g)
∆
⎯→
CO(g) + 3 H2(g)
CH4(g) + H2O(g) ⎯
C(s) + H2O(g)
Water gas reactions:
Reforming of methane:
Lab preparation of H2:
by reaction of active metals with acids.
e.g. Zn(s) + 2 HCl (aq) → ZnCl2(aq) + H2(g)
Reactions of Hydrogen
1.
With non-metals to form covalent hydrides
(a)
with halogens to form hydrogen halides which readily dissolve in water to
give corresponding acids
(b)
H2(g) + Cl2(g) → 2 HCl(g)
HCl(g)
(b)
2O
⎯H⎯
⎯
→
HCl(aq)
with nitrogen to form ammonia by the Haber process
3 H2(g) + N2(g) ⇌ 2NH3(g) T> 400°C, Fe catalyst, total P > 100 atm
2.
With metals to form ionic hydrides
e.g. 2 Na(s) + H2(g) → 2 NaH(s)
3.
To form methanol, an alternate fuel by reduction of CO(g).
CO(g) + 2 H2(g)
4.
⎯catalyst
⎯⎯→
CH3OH(g)
Hydrogen is a good reducing agent at high temperature and is used to reduce
metal oxides to metals, combining with oxygen to form water.
WO3(s) + 3H2(g)
∆
⎯
⎯→
Page 5 of 6
W(s) + 3 H2O(g)
Questions:
PHHM (9th Edition): Chapter 22: 59, 60
PHH (8th Edition): Chapter 8: 51, 52
Hints on identifying reaction types and predicting products
1.
2.
3.
4.
5.
Pure metals are reducing agents, often combining with oxygen at high
temperature.
C, CO(g), H2(g) and hydrocarbons are good reducing agents at high
temperature.
Electrolysis involves oxidation of anions at anode (+) and reduction of cations at
the cathode(−) of an electrochemical cell.
Heat(∆) usually means a thermal decomposition of the reactant forming one or
more gases and other products. Heat is required to melt salts for electrolysis.
Acids and bases react to form a salt and usually water.
Additional questions (See answers on Dr. Flinn’s web page.)
Complete and balance the following chemical reactions including physical states of the
products:
____ Na(s) + ____ HCHO2(aq) →
____ CO(g) + ____ H2O(g)
∆
⎯
⎯→
____ Fe(s) + ____ HCl(aq) →
∆
⎯
⎯→
∆
⎯
⎯→
____ Fe2O3(s) + ____ H2(g)
____ SrO(l) + ____ Al(l)
____ K(s) + ____ H2O(l) →
∆
⎯
⎯→
∆
⎯→
____ KHCO3(s) ⎯
∆
⎯→
____ CaSO4 · 2 H2O(s) ⎯
____ MgCO3(s)
____ CO2(g) + ____ Ca(OH)2(aq) →
____ KHCO3(s) + ____ HCl(aq) →
____ KHCO3(s) + ____ KOH(aq) →
____ Mg(s) + ____ HC2H3O2(aq) →
____ Ca(s) + ____ H2O(l) →
____ C(s) + ____ H2O(g) →
What is the most abundant element in the universe?
Write complete chemical reactions for the following:
The thermal decomposition of calcium carbonate (calcination).
The Haber process for the manufacture of ammonia.
The electrolysis of H2O, molten NaCl, molten CaCl2.
Two water gas reactions.
The reforming of methane.
The formation of plaster of Paris from gypsum.
Developed by Dr. Chris Flinn
Page 6 of 6