Download COLLEGE ENTRANCE TEST REVIEW CHEMISTRY

COLLEGE ENTRANCE TEST
REVIEW
CHEMISTRY
CVAndaya
• The Atomic Theory of MATTER
Can matter be subdivided into fundamental particles?
• •
Democritus (460–370 BC): All matter can be divided
into indivisible atomos.
John Dalton: Proposed atomic theory with the following
postulates:
•
Elements are composed of atoms.
•
All atoms of an element are identical.
•
In chemical reactions atoms are not changed into
different types of atoms. Atoms are neither created nor
destroyed.
•
Compounds are formed when atoms of elements
combine.
LAWS OF CHEMICAL
COMBINATION
• Law of constant composition: The relative kinds and numbers of
atoms are constant for a given compound.
• Law of conservation of mass (matter): During a chemical
reaction, the total mass before reaction is equal to the total
mass after reaction.
• Law of multiple proportions: If two elements A and B combine to
form more than one compound, then the mass of B which
combines with the mass of A is a ratio of small whole numbers.
e.g.: CO and CO2 (mass C = 12, O = 16, O2 = 32; ratio of O to C =
4:3, ratio of O2 to C = 8:3)
• •
Dalton’s theory predicted the law of multiple proportions.
• A cathode ray tube (CRT)
is a hollow vessel with an
electrode at either end.
• A high voltage is applied
across the electrodes.
• The voltage causes
negative particles to
move from the negative
electrode to the positive
electrode.
• The path of the electrons
can be altered by the
presence of a magnetic
field.
– The amount of deflection of the cathode rays
depends on the applied magnetic and electric
fields.
– In turn, the amount of deflection also depends
on the charge to mass ratio of the electron.
• In 1897, Thomson determined the
charge to mass ratio of an electron to
1.76  108 C/g.
Consider the following
experiment:
• Oil drops are sprayed above a
positively charged plate
containing a small hole.
• As the oil drops fall through the
hole, they are given a negative
charge.
• Gravity forces the drops
downward. The applied
electric field forces the drops
upward.
• When a drop is perfectly
balanced, the weight of the
drop is equal to the
electrostatic force of attraction
between the drop and the
positive plate.
• Using this experiment, Millikan determined
the charge on the electron to be
1.60  10-19 C.
• Knowing the charge to mass ratio,
1.76  108 C/g, Millikan calculated the
mass of the electron: 9.10  10-28 g.
• With more accurate measurements, we
get the mass of the electron to be
9.10939  10-28 g.
Consider the following
experiment:
• A radioactive substance is
placed in a shield containing a
small hole so that a beam of
radiation is emitted from the
hole.
• The radiation is passed
between two electrically
charged plates and detected.
• Three spots were noted on the
detector:
– a spot in the direction of
the positive plate,
– a spot which is not
affected by the electric
field,
– a spot in the direction of
the negative plate.
• A high deflection towards the positive plate
corresponds to radiation which is
negatively charged and of low mass. This
is called β-radiation (consists of electrons).
• No deflection corresponds to neutral
radiation. This is called gamma radiation.
• Small deflection towards the negatively
charged plate corresponds to high mass,
positively charged radiation. This is called
α-radiation.
• From the separation
of radiation we
conclude that the
atom consists of
neutral, positively,
and negatively
charged entities.
• Thomson assumed all
these charged
species were found in
a sphere.
• Rutherford carried out the following experiment:
• A source of a-particles was placed at the mouth
of a circular detector.
• The a -particles were shot through a piece of
gold foil.
• Most of the a-particles went straight through the
foil without deflection.
• Some a-particles were deflected at high angles.
• If the Thomson model of the atom was correct,
then Rutherford’s result was impossible.
• In order to get the majority of -particles
through a piece of foil to be undeflected,
the majority of the atom must consist of a
low mass, diffuse negative charge - the
electron.
• To account for the small number of high
deflections of the -particles, the center or
nucleus of the atom must consist of a
dense positive charge.
The Nuclear Atom
Rutherford modified
Thomson’s model as
follows:
– assume the atom is
spherical but the
positive charge must
be located at the
center, with a diffuse
negative charge
surrounding it.
• The atom consists of positive, negative, and
neutral entities (protons, electrons, and
neutrons).
• Protons and neutrons are located in the nucleus
of the atom, which is small. Most of the mass of
the atom is due to the nucleus.
– There can be a variable number of neutrons for the
same number of protons. Isotopes have the same
number of protons but different numbers of neutrons.
• Electrons are located outside of the nucleus.
Most of the volume of the atom is due to
electrons.
The Periodic Table of Elements
• By 2002, there were 115 elements known.
• The majority of the elements were
discovered between 1735 and 1843.
• How do we organize 115 different
elements in a meaningful way that will
allow us to make predictions about
undiscovered elements?
• First attempt
Dmitri Mendeleev and
Lothar Meyer
arranged the
elements in order of
increasing atomic
weight.
Certain elements were
missing from this
scheme.
In 1871, Mendeleev
noted that As, at. no.
33, belonged properly
underneath P (15) and
not Si (14), which left
a missing element
underneath Si. He
called this missing
element eka-silicon. He
predicted a number of
properties for this
element. In 1886, Ge
(32) was discovered.
The properties of Ge
match Mendeleev’s
predictions well.
Atomic Structure
Proton
Electron
Neutron
Atomic Number
Atomic mass
Ions
Isotopes
Given:
56
26Fe
++
How many protons,
electrons, and
neutrons does the ion
have?
p=
e=
n=
• Some of the groups in the periodic table
are given special names.
• These names indicate the similarities
between group members:
– Group 1A: Alkali metals.
– Group 2A: Alkaline earth metals.
– Group 7A: Halogens.
– Group 8A: Noble gases.
Periodic Table
Effective Nuclear
Charge: the charge
experienced by an
electron on a manyelectron atom.
• The effective
nuclear charge is not
the same as the
charge on the
nucleus because of
the effect of the
inner electrons.
• Electrons are attracted to the
nucleus, but repelled by the
electrons that screen it from
the nuclear charge.
• The nuclear charge
experienced by an electron
depends on its distance from
the nucleus and the number of
core electrons.
• As the average number of
screening electrons (S)
increases, the effective nuclear
charge (Zeff) decreases.
• As the distance from the
nucleus increases, S increases
and Zeff decreases.
Moseley revised
Mendeleev’s PT:
• The Modern periodic
table: elements were
arranged in order of
increasing atomic
number.
Periodic Trends in
Atomic Radii
• As the principle
quantum number
increases (i.e., we
move down a group),
the distance of the
outermost electron
from the nucleus
becomes larger.
Hence, the atomic
radius increases.
Size of Atoms and Ions
Consider a simple diatomic
molecule:
• The distance between the
two nuclei is called the
bond distance.
• If the two atoms which
make up the molecule are
the same, then half the
bond distance is called
the covalent radius of the
atom.
• As the principal
quantum number
increases, the size
of the orbital
increases.
Consider the s
orbitals:
• All s orbitals are
spherical and
increase in size as n
increases.
• Periodic Trends in Atomic
Radii
• As a consequence of the
ordering in the periodic
table, properties of elements
vary periodically.
• Atomic size varies
consistently through the
periodic table.
• As we move down a group,
the atoms become larger.
• As we move across a period,
atoms become smaller.
• There are two factors at
work:
– principal quantum number,
n, and
– the effective nuclear
charge, Zeff.
• As the principal quantum number (n) increases
(i.e., we move down a group), the distance of
the outermost electron from the nucleus
becomes larger. Hence, the atomic radius
increases.
• As we move across the periodic table, the
number of core electrons remains constant.
However, the nuclear charge increases.
Therefore, there is an increased attraction
between the nucleus and the outermost
electrons. This attraction causes the atomic
radius to decrease.
Trends in the Sizes of Ions
• Ion size is the distance between ions in an ionic
compound.
• Ion size also depends on nuclear charge,
number of electrons, and orbitals that contain
the valence electrons.
• Cations vacate the most spatially extended
orbital and are smaller than the parent atom.
• Anions add electrons to the most spatially
extended orbital and are larger than the parent
atom.
Trends in the Sizes of Ions
• For ions of the same charge, ion size
increases down a group.
• All the members of an isoelectronic
series have the same number of
electrons.
• As nuclear charge increases in an
isoelectronic series the ions become
smaller:
O2- > F- > Na+ > Mg 2+ > Al 3+
Ionization Energy
First ionization energy is the
amount of energy required to
remove an electron from a
gaseous atom:
• Na(g)  Na+(g) + e-.
The second ionization energy is
the energy required to remove
an electron from a gaseous
ion:
• Na+(g)  Na 2+ (g) + e-.
• The larger the ionization
energy, the more difficult it is to
remove the electron.
• Variations in Successive Ionization
Energies
• There is a sharp increase in ionization
energy when a core electron is removed.
• Ionization energy decreases down a group.
• This means that the outermost electron is more
readily removed as we go down a group.
• As the atom gets bigger, it becomes easier to
remove an electron from the most spatially
extended orbital.
– Ionization energy generally increases across a period.
– As we move across a period, Zeff increases.
Therefore, it becomes more difficult to remove an
electron.
• Two exceptions: removing the first p electron
and removing the fourth p electron.
• The s electrons are more effective at
shielding than p electrons. Therefore,
forming the s2p0 becomes more
favorable.
• When a second electron is placed in a p
orbital, the electron-electron repulsion
increases. When this electron is
removed, the resulting s2p3 is more
stable than the starting s2p4
configuration, thus, a lower I.E. for the
4th e- in the p orbital.
ELECTRON AFFINITY
• Electron affinity is
the energy change
when a gaseous atom
gains an electron to
form a gaseous ion:
• Cl(g) ) + e-  Cl-(g) +
energy
or
• Ar(g) + e- + energy
 Ar-(g)
Metals and Nonmetals
• Metallic character refers to the
properties of metals (shiny or lustrous,
malleable and ductile, oxides form basic
ionic solids, and tend to form cations in
aqueous solution).
• Metallic character increases down a
group.
• Metallic character decreases across a
period.
• Metals have low ionization energies.
• Most neutral metals are oxidized rather
than reduced.
• When metals are oxidized they
tend to form characteristic cations.
• All group 1A metals form M+ ions.
• All group 2A metals form M2+ ions.
• Most transition metals have
variable charges.
Most metal oxides are basic:
• Metal oxide + water  metal hydroxide
• Na2O(s) + H2O(l)  2NaOH(a
Nonmetals are more diverse in their
behavior than metals.
• When nonmetals react with metals,
nonmetals tend to gain electrons:
• metal + nonmetal  salt
• 2Al(s) + 3Br2(l)  2AlBr3(s)
Alkali Metals
• Alkali metal produce different oxides when
reacting with O2:
• 4Li(s) + O2(g)  2Li2O(s) (oxide)
• 2Na(s) + O2(g)  Na2O2(s) (peroxide)
• K(s) + O2(g)  KO2(s)
(superoxide)
• Alkali metals emit characteristic colors
when placed in a high temperature flame.
• The s electron is excited by the flame and
emits energy when it returns to the ground
state.
Li line: 2p  2s
transition
Na line (589 nm):
3p  3s transition
K line: 4p  4s
transition
• Group 6A: The Oxygen Group
• Oxygen (or dioxygen, O2) is a potent oxidizing
agent since the O2- ion has a noble gas
configuration.
• There are three oxidation states for oxygen: 2(e.g. H2O) and 1- (e.g. H2O2), ½-(e.g. KO2)
• Sulfur is another important member of this
group.
• Most common form of sulfur is yellow S8.
• Sulfur tends to form S2- in compounds (sulfides).
The Halogens
• Group 7A: The HalogensThe chemistry of the
halogens is dominated by gaining an electron to
form an anion:
X2 + 2e-  2X--.
• Fluorine is one of the most reactive substances
known:
2F2(g) + 2H2O(l)  4HF(aq) + O2(g)
H = -758.7 kJ.
• All halogens are diatomic molecules, X2.
• Chlorine is the most industrially useful halogen.
It is produced by the electrolysis of brine (NaCl):
2NaCl(aq) + 2H2O(l)  2NaOH(aq) + H2(g) +
Cl2(g).
• The reaction between chorine and water
produces hypochlorous acid (HOCl) which
disinfects pool water:
Cl2(g) + H2O(l)  HCl(aq) + HOCl(aq).
• Hydrogen compounds of the halogens are all
strong acids with the exception of HF.
As we go down Gp VIIA, the acid formed becomes
stronger.
CHLOROFLUOROCARBONS
First CFC was synthesized
in 1928 by
THOMAS MIDGLEY, Jr.
(who also synthesized
tetraethyllead, a gasoline
additive)
as alternative refrigerant for
NH3. It was a miracle
substance.
Trade name--Freon
CFCs have varied uses:
• Refrigerant
• Foam blowing agent
• Cleaning solvents for
electronic parts
• As metal degreaser
• Sterilants for medical
instruments
• Propellants for
pressurized chem’ls
• Fire extinguisher
Ozone Depleting Substances
• O3 molecules are dispersed in the stratosphere,
and are formed from
3 O2 + uv → 2 O3
• When brought all together to form a layer, the result amount
to a thickness of 3 to 5mm.
• In air, there are only 3 O3 molecules for every 10million air
molecules.
• Ozone is depleted where it is needed the most:
90% of all the O3 in the world is in the
stratosphere.
• CFCs break up into ions
and atoms when they
absorb UV.
• CF2Cl2 + uv → CF2Cl +
Cl
2Cl + 2O3 →2ClO + 2 O2
2ClO → ClOOCl
ClOOCl + uv →ClOO + Cl
ClOO →Cl + O2
Net eq’n: 2O3 → 3 O2
1 Cl atom could “destroy” as
many as 100,000 O3
molecules
Other ODS:
Halons (w/ Br or F,
substituting for Cl), used
as fire extinguisher
Methyl Bromide (more
dangerous ODS), as soil
fumigant
Carbon Tetrachloride, a
cheap, toxic solvent
Methyl chloroform, a dry
cleaning agent
Hydrogen Chloride (from US
space shuttles)
Group 8A: The Noble Gases
• These are all nonmetals and monatomic.
• They are notoriously unreactive because
they have completely filled s and p subshells.
• In 1962 the first compound of the noble
gases was prepared: XeF2, XeF4, and
XeF6.
• To date the only other noble gas
compounds known are KrF2 and HArF.
Chemical Bonds
.
• Covalent bond results from sharing
electrons between the atoms. Usually
found between nonmetals.
• Ionic bond results from the transfer of
electrons from a metal to a nonmetal.
• Metallic bond: attractive force holding
pure metals together.
Lewis structures are used to
represent covalent bonds
Cl Cl
H F
H O
H
H N H
H
H
H C H
H
• It is possible for more than one pair of electrons
to be shared between two atoms (multiple
bonds):
– One shared pair of electrons = single bond (e.g. H2);
– Two shared pairs of electrons = double bond (e.g.
O2);
– Three shared pairs of electrons = triple bond (e.g.
N2).
• Generally, bond distances decrease as we move
from single through double to triple bonds.
Multiple bonds
• Multiple Bonds
• It is possible for more than one pair of electrons
to be shared between two atoms (multiple
bonds):
– One shared pair of electrons = single bond (e.g. H2);
– Two shared pairs of electrons = double bond (e.g.
O2);
– Three shared pairs of electrons = triple bond (e.g. N2).
• Generally, bond distances decrease as we move
from single through double to triple bonds.
• In a covalent bond, electrons are shared.
• Sharing of electrons to form a covalent
bond does not imply equal sharing of
those electrons.
• Unequal sharing of electrons results in
polar bonds.
• Electronegativity: The ability of one
atoms in a molecule to attract electrons to
itself.
• Pauling set electronegativities on a scale
from 0.7 (Cs) to 4.0 (F).
• Electronegativity increases
– across a period and
– down a group.
• Difference in electronegativity is a gauge
of bond polarity:
– electronegativity differences around 0 result in
non-polar covalent bonds (equal or almost
equal sharing of electrons);
– electronegativity differences around 2 result in
polar covalent bonds (unequal sharing of
electrons);
– electronegativity differences around 3 result in
ionic bonds (transfer of electrons).
Electronegativity and Bond
Polarity
• There is no sharp
distinction between
bonding types.
• The positive end (or
pole) in a polar bond
is represented + and
the negative pole -.
• In general, the least
electronegative
element is named
first.
Ionic Bonding
• The formation of Na+(g) and Cl-(g) from
Na(g) and Cl(g) is endothermic.
• NaCl(s) Na+(g) + Cl-(g) endothermic
(H = +788 kJ/mol).
• The formation of a crystal lattice from the
ions in the gas phase is exothermic:
• Na+(g) + Cl-(g)  NaCl(s) H = -788
kJ/mol
Electron Configurations of Ions
of the Representative Elements
• These are derived from the electron
configuration of elements with the required
number of electrons added or removed from the
most accessible orbital.
• Electron configurations can predict stable ion
formation:
–
–
–
–
–
Mg: [Ne]3s2
Mg+: [Ne]3s1
not stable
Mg2+: [Ne]
stable
Cl: [Ne]3s23p5
Cl-: [Ne]3s23p6 = [Ar]
stable
Covalent Bond
• When two similar atoms bond, none of
them wants to lose or gain an electron to
form an octet.
• When similar atoms bond, they share pairs
of electrons to each obtain an octet.
• Each pair of shared electrons constitutes
one chemical bond.
• Example: H + H  H2 has electrons on a
line connecting the two H nuclei.
Ionic
MgH2
Molecular
Magnesium
hydride
FeF2 Iron(II) fluoride
Mn2O3, Manganese(III)
oxide
H2S Hydrogen
sulfide
OF2 Oxygen
difluoride
Cl2O3
Dichlorine
trioxide
Resonance structures are
attempts to represent a real
structure that is a mix between
several extreme possibilities
The ozone molecule
O
O
O
O
O
O
Resonance in Benzene
• Resonance in Benzene
• We write resonance structures
for benzene in which there are
single bonds between each
pair of C atoms and the 6
additional electrons are
delocalized over the entire
ring:
• Benzene belongs to a category
of organic molecules called
aromatic compounds (due to
their odor).
or
The Chemistry of Global Warming
Greenhouse Gases: those atmospheric gases
which absorb infrared radiation, and help keep
the earth warm. These are:
H2O
SO2
N2O (nitrous
CO2
CH4
oxide)
Global Warming Potential (GWP)
CO2 = 1 N2O = 296
CH4 = 23
Abundance (ppm)
CO2 = 375, N2O = 0.31
CH4 = 1.8
Anthropogenic Greenhouse Gases
CF3CH2F
SF6
N 2O
GWP
1300
5700