Chem 1000 Midterm 2

malleable, ductile, good conductor of heat and electricity, shiny, resistivity increases with increasing T

brittle when solid, poor conductor of heat and electricity (insulator)

intermediate between metal and nonmetal, often semiconducting

electrical conductivity is between that of a conductor and insulator, resistivity decreases with increasing T

Explains metals
-heat and electrical conductivity
-deformability (ductility, malleability)

repeating arrangement of points in space

a material composed of many microscopic crystals (grains) stuck together in different orientations

surface where two grains meet

a material composed of a single, (nearly) perfectly ordered crystalline material without grain boundaries

square
rectangle
hexagon

cubic unit cell in which lattice points are situated only at the corners
8x(1/8)=1 atom per unit cell

A common crystal structure found in some elemental metals. Within the cubic unit cell, atoms are located at corner and cell center positions
8x(1/8)+1=2 atoms per unit cell

A crystal structure found in some common elemental metals. Within the cubic unit cell, atoms are located at all corner and face-centered positions.
8x(1/8)+6(1/2)=4 atoms per unit cell

shared by 8 unit cells
1/8 of an atom is inside any given cell

shared between 2 unit cells
1/2 of an atom is inside any given cell

minimum empty space
74% of the space is occupied by atoms
ABCABC...

an identical packing fraction is obtained
ABAB...

from this experiment we get:
-crystal structure
-positions of atoms within unit cell
-dimensions of unit cell

group 1 except H
soft metals
lowest ionization energies and electronegativities in periodic table, low melting point and boiling points
have among the most negative reduction potentials
very powerful reducing agents
in nature these elements only appear as their +1 cations
almost universally ionic
relatively large, negative enthalpies of hydration
carry a single charge, the forces holding their crystals together, while significant, are less strong than those holding together crystals of more highly charged ions
almost all are extremely soluble in water (Exception, Li with highly charged anions)

metal hydroxide + hydrogen
M(s) + H2O(l)-> M+(aq) + OH-(aq) + 1/2H2(g)

metal halide
M(s) + 1/2 X2 -> MX(s)

Lithium oxide (Li2O)
2Li(s) + 1/2O2(g) -> Li2O(s)

half cell potential for gaining electrons

an ion in solution is surrounded by water molecules

enthalpy change for the transfer of an ion from the gas phase to solution
M+(g) -> M+(aq)

Metal ions are often identified by precipitation
Alkali metals compounds are extremely soluble so that wont work. Instead we use flame tests
-putting a sample into a flame puts energy into it. This energy can put ions in excited electronic states
-when the ions return to their ground state (possibly in multiple hops), they emit light
-the emission spectrum depends on a number of factors (including the flame temperature), but is most strongly dependent on the energy levels of the emitter, leading to characteristic colours

crimson

yellow

purple

blue

LiCl(l) -> Li(l) + 1/2 Cl2(g)
NaCl(l) -> Na(l) + 1/2 Cl2(g)

commercial electrochemical cell used to obtain sodium metal by the electrolysis of molten NaCl
melting point of NaCl= 804

Na+(l) + e- -> Na(l)

Cl-(l) -> 1/2 Cl2 (g) + e-

NA+(l) + Cl-(l) -> Na(l) + 1/2 Cl2(g)

electrolysis of an aqueous solution of NaCl
Half Reactions:
2Cl- (aq) -> Cl2(g) + 2e-
2H2O (l) + 2e- -> H2(g) + 2OH-(aq)
Overall:
2Cl-(aq) + 2H2O(l) -> Cl2(g) +H2(g) +2OH-(aq)

enthalpy change when an ionic compound is broken up into the corresponding gaseous ions

An enthalpy cycle that allows you to calculate the lattice enthalpy change of formation for system.

Summing the forces between the ions, assuming that they are point charges occupying their respective lattice sites

Group 2, except maybe Be
soft metals
form 2+ cations
very negative reduction potentials
M2+(aq) + 2e- -> M(s)
tend to be ionic (other than compounds of Be)

M(s) + 2H2O -> M(OH)2 + H2

M(s) + 1/2 O2(g) -> MO(s)

3M(s) + N2(g) -> M3N2(s)

M(s) + X2 -> MX2(s)

term applied by early chemists to nonmetallic substances which are insoluble in water and remain stable when heated. The alkaline earth metal oxides have these properties

means basic. The alkaline earth metal oxides, to the extent that they dissolve at all in water, make basic solutions
MO(s) + H2O(l) -> M(OH)2

with monovalent anions (NO3-,Cl-) tend to be soluble in water, except hydroxides and fluorides, which are at best slightly soluble
with divalent anions tend to be insoluble (MCO3, MSO4) or only slightly soluble (MS) in water

Brick red

crimson

green

no colour

When heated, alkaline earth carbonates decompose to the corresponding oxide and CO2:
MCO3(s) -> MO(s) + CO2(g)

We can regenerate the carbonate in 2 steps:
1. MO(s) + H2O(l) -> M(OH)2
2. M(OH)2 + CO2(g) -> MCO3(s) + H2O(l)

Reaction with excess acid:
MCO3(s) + 2H+(aq) -> M2+(aq) + CO2(g) + H2O(l)

blocks potassium channels
when ingested, Ba2+ released from barium carbonate when it is combined with stomach acid, leading to life-threatening symptoms
Barium sulfate on the other hand is essentially insoluble, so it is safe to ingest. It is used as a contrast agent in x-ray imaging of the GI tract

contains divalent cations, esp. Mg2+, Ca2+, and Fe2+

has the counterion HCO3-
Called temporary because alkaline earth metal (and other high-valent metal ions) carbonates can be removed by boiling:
2HCO3- (aq) + heat -> CO3 2-(aq) + CO2 +H2O(l)
M2+ (aq) + CO3 2-(aq) -> MCO3(s)
If untreated, leads to scale in hot water heaters, pipes, kettles, etc.

contains counterions other than bicarbonate (chloride, sulfate, nitrate) whose alkaline earth metal compounds cant be removed by boiling
Divalent cations tend to precipitate with soaps and detergents= bathtub ring, greying of clothes
water can have both temporary and permanent hardness

Be tends to make covalent bonds instead of ionic bonds
Forms a thin layer of BeO on contact with air or water
BeO is nonporous, so it prevents further oxidation. This is called passivation
BeO doesn't react with water
BeF2 and BeCl2 are covalent (molecular) compounds, not ionic
BeO is amphitaric (both an acid and a base)
Reaction with acid: BeO(s) + 2H3O+(aq) +H2O(l) -> [Be(OH2)4]2+ (aq)
[Be(OH2)4]2+ contains covalent bonds between O and Be
[Be(OH2)4]2+ is acidic:
[Be(OH2)4]2+ + H2O(l) <-> [Be(OH2)3(OH)]+(aq) + H3O+(aq)
Reaction with base:
BeO(s) + 2OH-(aq) + H2O(l) -> [Be(OH)4]2-(aq)

similarities between pairs of elements in different groups and periods of the periodic table
Li-Mg
Be-Al
B-Si

Li is the only alkali metal to form a nitride, whereas this is a typical reaction for the alkaline earth metals
A number of lithium compounds are much less soluble in water than the corresponding compounds of other alkali metals (e.g. lithium phosphate)
Li2CO3 decomposes to Li2O and CO2 like the alkaline earth metal carbonates, but the other alkali metal carbonates are thermally stable

both form a tough and chemically resistant passivation layer
both have amphoteric oxides. Contrast: reaction of alkaline earth metals with water to make alkaline earth hydroxides
Both form covalent halides. Contrast: Ionic halides of alkaline earth metals (Note: ionic vs. covalent character determined from electrical conductivity of melt)

All the elements in group 13 except boron are metals
common ions: Al3+, Ga3+, In+, TI+, also sometimes In3+ and TI3+
Under typical reactions, these metals don't react with water
at room temperature, the oxide formed is M2O3
Al(OH)3, Ga(OH)3 and In(OH3) are all insoluble in water, but TIOH is very soluble
Aluminum is a good reducing agent

Thallium has many compounds of its +1 ion
TI+ and K+ are very similar in size: 164 and 153pm, respectively. In the body, TI+ can go wherever K+ can go, which is pretty much everywhere
Most thallium compounds are toxic. TISO4 was once used as a rat and ant poison
Thallium (I) sulfate is tasteless, so it was once a popular poison, nicknamed "Inheritance powder"
Unlike K+, TI+ reacts with sullfur ligands, disrupting proteins containing cysteine in particular

Like Be, Al is passivated by an oxide layer which forms extremely rapidly when the metal is exposed to air
Aluminum oxide:
is hard
has a high melting point
is non-porous

deliberate, electrolytic thickening of the oxide layer:
2Al(s) + 3H2O(l) -> Al2O3 (s) + 6H+ (aq) + 6e-
Has the following effects:
makes Al more chemically resistant
makes Al harder
may be used to pigment the oxide layer with appropriate additives

Al2O3 is amphoteric
reaction with acid:
Al2O3(s) + 6H+(aq) -> 2Al 3+(aq) + 3H2O(l)
reaction with base:
Al2O3(s) + 2OH-(aq) + 3H2O(l) -> 2[Al(OH)4]-(aq)

HNO3 is an oxidizing acid because it typically attacks metals through the reduction of NO3-:
NO3- + 4H+ +3e- -> NO + 2 H2O
For Al, this reaction, just tends to make the oxide layer thicker, i.e. makes the metal more passive:
2Al(s) + 2NO3- + 2H+ -> Al2O3(s) + 2NO(g) + H2O(l)
Al reacts with non-oxidizing acids (e.g. HCl) in the typical way for a metal:
Al(s) + 3H+ (aq) -> Al3+(aq) + 3/2 H2(g)

Aluminum metal can be oxidized in aqueous base:
Al(s) + OH-(aq) + 3 H2O(l) -> [Al(oH)4]-(aq) + 3/2 H2(g)
aluminum can react with acid or base, so it is amphoteric

Al(OH)3 is amphoteric
Reaction with metal (typical for metal hydroxides):
Al(OH)3 (s) + 3 H3O+(aq) -> [Al(H2O)6]3+(aq)
Reaction with base (unusual for a metal hydroxide):
Al(OH)3(s) + OH-(aq) -> [Al(OH)4]- (aq)

aluminum ore (bauxite) contains Al2O3, AlO(OH) and Fe2O3 (among other things)
Al2O3 and AlO(OH) dissolve in base, but FeO3 doesnt:
Al2O3(s) + 2OH- (aq) + 3H2O(l) -> 2[Al(OH)4]-(aq)
AlO(OH)(s) + OH-(aq) + H2O(l) -> [Al(OH)4-](aq)
filter off solid Fe2O3, then precipitate Al(OH)3:
[Al(OH)4]-(aq) + H3O+(aq) -> Al(OH)3(s) + 2H2O(l)
Recover aluminum oxide by heating:
2Al(OH)3(s) + heat -> Al2O3(s) + 3H2O(g)

Al2O3 has a very high melting point (2072)
A 15:85 mixture of Al2O3, and cryolite (Na3AlF6) has a melting point of about 1000
this molten mixture is electrolyzed using a graphite anode
Anode:C(s) + 2 O2- (l) -> CO2(g) + 4e-
Cathode: Al3+(l) + 3e- -> Al(l)
Overall: 3C(s) + 6 O2-(l) + 4 Al3+(l) -> 4 Al(s) +3 CO2(g)
liquid aluminum is denser than the ionic liquid from which it is electrolyzed, so it is collected from the bottom of the electrolysis vat

a chemical bond in which electrons are shared between two atoms
typically formed between two nonmetallic elements

material held together by electrostatic force between ions (no sharing of electrons)
Electronegativities differ by 2 units or more

electrons shared among all the metal atoms

a convenient graphical accounting system for valence electrons in molecules
a shared electron pair is represented by a line between the two atoms
a non bonding valence electron is represented by a dot
dots are paired to represent their occupation of orbitals

number of shared electron pairs in a bond

a pair of nonbonding electrons in an "orbital"

typically, main-group elements in chemical compounds (except H) are surrounded by eight valence electrons (including any shared in covalent bonds).
H an only have/ share two electrons
Some elements, notably B, sometimes form electron deficient compounds in which there is less than a full octet around the atom, with important implications for their chemistry

1. Count up the number of valence electrons in all the atoms
If the species is cation, subtract the charge
If the species is an anion, add the absolute value of the charge
This is the total number of valence electrons to place in your Lewis diagram
2. Draw a skeleton with one bond between each pair of atoms which are covalently bonded
3. Complete the octets for any group 17 elements with lone pairs. Exception: compounds where a halogen is the central atom
4. placing the remaining electrons becomes a puzzle
complete as many octets as possible with lone pairs or multiple bonds
double bonds are common for C, group 15 and group 16
triple bonds are common for C and group 15
never exceed an octet in period 2
you can exceed an octet in periods 3 and up if necessary
5. the best lewis structure for a compound minimizes the formal charges even if this means violating the octet rule for atoms in the 3rd period and beyond

when there is more than one reasonable lewis diagram

a calculation of the charge of an atom in a lewis diagram assuming electrons are equally shared
Calculation: Draw the Lewis diagram, then count the electrons on a particular atom, counting half of any electrons shared with a neighbour
the formal charge is the number of valence electrons of the neutral atom minus the number of electrons on the atom in the Lewis structure
*Sum of formal charges= charge of molecule

shorter bond length

enthalpy change for the breaking of a covalent bond in the gas phase
depends on the particular molecule in which the bond occurs, not just on the atoms involved
increases with bond order

we can estimate reaction enthalpies from bond dissociation enthalpies by counting up the number of bonds of each type made and broken

but its Carbon-Carbon double and triple bonds are very reactive

electronegativity is a measure of the ability of an atom to attract electrons to itself
two atoms in a covalent bond may not share the electrons in the bond equally, depending on the difference in electronegativity

VSEPR
Key ideas:
electrons repel
valence electrons (bonding and nonbonding) are particularly close together at the end of a bond near a central atom
electrons can be thought of as "living in groups" of nonbonding electron lone pairs and covalent bonds. A bond of any order (single, double, triple) is a single group
The electronic geometry around an atom is therefore dictated by the mutual repulsion of electron groups near that atom
electrons organize themselves so as to minimize the repulsion between groups

"fatter" than bonding pairs
Consequences:
1. lone pairs locate themselves where there is the most space
2. Lone pairs push the bonding electrons away, distorting their geometry away from the ideal geometry for n identical electron groups

a pair of equal and opposite charges separated by a distance d

polar bonds can be thought of as little dipole vectors by convention in chemistry, these vectors point toward the negative (more electronegative) end of a bond

u=qd
measures the strength of a dipole
the overall dipole moment of a molecule is the vector sum of the bond dipoles
Nonzero dipole moment=polar
zero dipole moment=nonpolar