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CHAPTER 13 -- CRYSTALS, IONS, SOLUTIONS
PLUS ACIDS, BASES, AND SALTS
Solids -- consist of atoms, ions, or molecules packed closely together and held
together by electric forces
Crystalline -- particles arranged in regular, repeated, three-dimensional patterns
Amorphous -- irregularly arranged particles arranged with no definite pattern
Ionic Crystal Classes -- although there are 230 different ways to arrange
particles, all crystals are assumed to fall into 6 crystal classes -- two examples
of one class(cubic) are described below
Face centered cubic -- Na+Cl-
Ions of either kind can be thought of as being located at the corners and at
the faces of an assembly of cubes [each ion has six closest neighbors of
the other kind]
side view Na+---- Cl- ----- Na+
9 ions per | | |
face; one in - +
Cl ----- Na ---- Cl-
center of | | |
face + -
Na ---- Cl ----- Na+
Body centered cubic -- Cs+Cl-
each ion is located at the center of a cube whose corners are ions of the
other kind [each ion has eight nearest neighbors of the other kind]
Covalent Crystals -- cohesive forces arise from the presence of electrons between
adjacent atoms. Example: diamond -- tetrahedral carbon system
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purely covalent crystals are rare . . .
C (diamond)
Si (silicon)
Ge (germanium)
SiC (carborundum)
They are usually very hard and have very high melting points
The “metallic” bond -- the electron “gas” that bonds metals makes them good
electrical and heat conductors.
Characteristic of metals: only a few electrons are present in the outer shells
and these electrons are not securely attached.
a. Outer electrons that are free to move help hold the positively charged metal
ions together.
b. Since adjacent atoms in a metal are not linked by specific bonds, most
alloys(mixtures of different metals) do not obey the Law of Definite
Proportions.
Molecular crystals -- produced by Van der Waal’s forces
Major example is hydrogen bonded materials.
Polar-polar attraction [H-bonding]
Time average electrical attraction
Polar-non-polar attraction -- one dipole induces the electrons to separate in
another species [electrostatic induction]
Non-polar/non-polar attraction -- caused by instantaneous dipoles, sometimes
called London or dispersion forces; reason for liquefaction of inert gases
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Characteristics of Crystal Types
Type Bond Example Properties
Ionic Electrical Na+Cl- hard; high
Attraction E(cohesive)=3.3eV mp; soluble
atom in polar liquid
Covalent Shared C (diamond) very hard,
Electrons E(cohesive)=7.4eV high mp;
atom insoluble in
most solvents
Metallic Electron Sodium ductile,
Gas E(cohesive)=1.1eV luster,
atom conductive
Molecular Van der Waal’s Methane soft, low
forces E(cohesive)=0.1eV mp, low bp,
atom soluble in
non-polar
SOLUTIONS -- intimate mixtures of two or more substances
EXAMPLES: alloys, air, sea water, gases in liquids
Homogenous -- same composition throughout
EXAMPLES: sugar in water, liquor
Heterogeneous -- composition varies
EXAMPLES: air, sea, water
Solvent versus Solute
(1) When two liquids are in solution, the one in greatest amount is the
solvent; the lesser amount is solute.
(2) When solids or gases dissolve in liquids, the liquid is the solvent; the
solid/gas is solute.
Solubility -- usually there is a limit to how much solute can “dissolve” in a
solvent
DEFINITION: the solubility of a substance is the maximum amount
that can be dissolved in a given quantity of solvent at a given
temperature, usually expressed in grams of solute per 100 mL of solvent.
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Saturated -- maximum quantity dissolved at temperature specified
Solubility of solids (in liquids) usually increase with increase in temperature.
EXAMPLES: salt in water, sugar in water
IMPORTANT EXCEPTION: calcium carbonate produces “limeout” in
hot water heaters because it is less soluble in hot water than in cold water
Solubility of gases (in liquids) usually decreases with increase in temperature;
[thermal pollution] Example: CO2 in water (soda water).
Solubility of gases (in liquids) usually increases with increase in pressure.
[cause of “Bends”] Examples: CO2 in water (soda water), N2 in blood
Polar/non-polar liquids as solvents
polar liquids usually dissolve polar solutes [ionic species]
non-polar liquids usually dissolve non-polar solutes [covalent species]
“Like Dissolves Like”
DISSOCIATION -- separation of a compound into ions as it dissolves
Electrolytes -- substances that separate into free ions of solution in water.
Includes all ionic compounds that are soluble in water and some covalent
compounds containing hydrogen (called acids). [Electrolytes in solution
conduct electricity by the motion of ions]
Ionic compounds conduct electricity when in the molten state or in aqueous
solution; whereas acids do not conduct electricity in the molten state, but do in
aqueous solution.
Non-Electrolytes -- soluble covalent compounds that do not dissociate in
solution
Strong electrolytes vs weak electrolytes vs non-electrolytes.
COLLIGATIVE PROPERTIES -- characteristics of solution that depend on
the # of particles of solute and not on the kind of solute
( 1 ) Vapor-pressure lowering -- addition of non-volatile solute causes a
decrease in vapor pressure of solvent
( 2 ) Boiling point elevation -- addition of non-volatile solute causes a
increase in boiling point of solvent
( 3 ) Freezing point depression -- addition of a non-volatile solute causes a
decrease in the freezing point of the solvent [antifreeze]
( 4 ) Osmotic pressure -- addition of a non-volatile solute causes change in the
internal pressure of the solvent
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SPECIAL NOTES:
(1) Since ionic compounds consist of multiple particles, electrolytes cause
larger changes in colligative properties than non-electrolytes.
sugar NaCl CaCl2
In equimolar amounts, calcium chloride has a highest effect of the list;
sugar has the lowest effect.
(2) The major use of colligative properties is to determine the Gram-
Molecular-Weights of compounds.
Molecular vs Ionic [Cl-]
Chlorine [Cl2] Chloride Ion
Greenish-yellow color colorless
Strong, irritation taste and odor mild, pleasant taste
Combines with all metals does not react w/metals
Combines readily with hydrogen no reaction with hydrogen
Does not react with Ag forms AgCl (s) with Ag+
Very soluble in CCl4 insoluble in CCl4
ARRHENIUS THEORY/HYPOTHESIS
That many substances exist as ions in solution was proposed in 1887 by a
young Swedish chemist named Svante Arrhenius.
Conflict with the already well-known Faraday on the concept of electrolysis:
(1) Faraday said that passage of a current caused a substance in solution
to break up into ions.
(2) Arrhenius said electrolytes set ions free when they dissolve.
Arrhenius Proof
(1) Reactions between electrolytes are very fast in solution; whereas
when dry, the reactions are extremely slow or do not occur at all.
(2) Electrolytes yield unexpectedly greater freezing point depressions
than non-electrolytes.
Na+Cl twice as much as sugar
Arrhenius is also famous for his initial, simple theory of acids, bases, and
salts.
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ACIDS, BASES, AND SALTS
Acids -- by the simplest definition are substances which yield H+ in aqueous
solution; i.e., dissociates into H+ and its anion (or produces the hydronium ion
when reacted with water)
Ex. HCl + H2O H3O+ + Cl-
HCl(aq) is called hydrochloric acid.
HNO3 + H2O H3O+ + NO3-
HNO3(aq) is called nitric acid.
H2SO4 + H2O H3O+ + HSO4-
H2SO4(aq) is called sulfuric acid.
Bases -- by the simplest definition are substances which yield OH
(hydroxide ions) in aqueous solutions; i.e., dissociate into cations and OH in
water solution
NaOH(aq) Na+ + OH
Mg(OH)2(aq) Mg++ + 2 OH
NH4OH(aq) NH4+ + OH
Salts -- by the simplest definition are substances which are the products,
along with water, of neutralization reactions (acid + base salt + water)
Acids which can yield only one proton(H+) in solution are called monoprotic.
Examples: HCl, HNO3, HI, HC2H3O2 (acetic acid)
If two H+ diprotic Examples: H2SO4 , H2CO3
If three H+ triprotic Example: H3PO4
If more than three polyprotic EDTA or
(ethylenediaminetetraacetic acid)
Bases which yield only one OH- in solution are called monobasic.
Examples: KOH, NaOH, LiOH, CsOH
If two OH dibasic Examples: Mg (OH)2 , Ca(OH)2
If three OH tribasic Example: Fe(OH)3
If more than three polybasic Example: Ti(OH)4
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ACIDIC SALTS -- salts formed from strong acids and weak bases
BASIC SALTS -- salts formed from strong vases and weak acids
NEUTRAL SALTS -- salts formed from either strong acids and strong
bases OR weak acids and weak bases
ACID(or BASE) STRENGTH depends only on the ease of dissociation in
water solution.
Strong Acids Weak Acids Strong Bases Weak Bases
HCl HF Any Group NH4OH
HNO3 H2S I A or II A organic
HClO4 HC2H3O2 Metal bases
H2SO4 H2CO3 Hydroxide (amines)
HBr H3PO4 if soluble
HI HCN (in water)
+
“Mineral” Organic
Acids Acids
ACIDIC SALTS BASIC SALTS NEUTRAL SALTS
NH4NO3 NaF NaCl
NH4Cl K2CO3 CaSO4
(NH4)2SO4 Na2HPO4 NH4C2H3O2
Mg(CN)2 KClO4
Ca(C2H3O2)3 BaI2
PHYSICAL PROPERTIES of ACIDS
(1) contain H+
(2) water solutions taste sour
(3) turn blue litmus paper red
(4) liberate hydrogen gas when reacted with active metals
PHYSICAL PROPERTIES of BASES
(1) contain OH
(2) water solutions taste bitter
(3) turn red litmus paper blue
(4) slippery to the touch
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NEUTRALIZATION
NaOH (aq) + HCl (aq) NaCl (aq) + H2O (aq)
base + acid salt + water
colorless + colorless colorless + colorless
How do we know a reaction occurs?
(1) mixture becomes warm (*exothermic* rxn)
(2) taste of the acid gets less and less sour
(3) solution of base (if adding acid to the base) becomes less and
less slippery
*gives off heat exothermic
takes in heat endothermic
Base destroys acid properties!! Acid destroys base properties!!
ACIDS NEUTRALIZE BASES!
How can we determine the acidity or basicity of solutions?
(1) use litmus paper [blue in base, red in acid] or use pHydrion paper
(indicates pH value by color)
(2) use *indicators* different color in acid than base
(3) use a pH meter (electrical device)
0 7 exactly 7 7 14
acidic neutral basic
Indicators COLOR
Name of Indicator acid base
Phenolphthalein colorless pink
methyl orange yellow salmon pink
cherry juice reddish blue
bromocresol green yellow blue
bromothymol blue yellow blue
Amphiprotic compounds -- materials which can behave as either an acid
or base depending on what it reacts with; water since it reacts with itself to
produce both acid and base species is called autoprotolytic.
Notice that one of the HOH’s H+; notice that one of the HOH”s OH.
H2O + H2O H3O+ + OH
acid base acid! base!
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If the species starts out as a metalloid hydroxide, it is called amphoteric!
base acid salt
Sn(OH)2 + 2 HCl SnCl2 + 2 H2O
acid base salt
H2SnO2 + 2 NaOH Na2SnO2 + 2 H2O
H3BO3 / B(OH)3 boric acid/boric hydroxide
H3AlO3 / Al (OH)3 aluminum hydroxide/aluminic acid
acid base
HCO3 + OH HOH + CO3=
base acid
HCO3 + H3O+ HOH + H2CO3
Therefore, sodium bicarbonate or baking soda is amphiprotic!!!
NaHCO3 can be used to clean up both acid and base spills!!!
ACID BASE NOMENCLATURE
(1) Bases are named as hydroxides.
NaOH (aq) sodium hydroxide
Ca(OH)2 (aq) calcium hydroxide
(2) Binary acids [hydro + non-metal stem + ic + acid]
HCl (aq) hydrochloric acid
H2S (aq) hydrosulfuric acid
(3) Polyatomic acids [element “stem” + ic + acid];
oxygen does not enter into naming.
H2SO4 (aq) sulfuric acid
H3PO4 (aq) phosphoric acid
HNO3 (aq) nitric acid
However; polyatomic (-ate) -ic
polyatomic (-ite) -ous
HClO4 (aq) perchloric acid
HClO2 (aq) chlorous acid
HNO2 (aq) nitrous acid
HClO (aq) hypochlorous acid
Again, PRACTICE! PRACTICE! PRACTICE!
