Early Models of the Atom
The Early Greeks and Medieval Alchemists
Democritus (400 BC) stated that matter
was composed of small particles called
"atoms" that could be divided no further.
These atoms were all composed of the
same primary matter with the only
differences between them being their size,
shape and weight.
The differences in these characteristics explained the
differences in the properties of the matter around us.
In the early 4th century, Aristotle
stated that “matter” had four possible
properties:
moist, dry, hot, and cold
These properties were contained in
various proportions by four major
elements which made up everything
our senses could detect:
water, air, fire, and earth
Medieval chemists developed the idea of “corpuscles” that
were subject to attractive and repulsive forces.
Any ideas by the Greeks or Medieval chemists were largely
philosophical in nature (not based on scientific
observations).
Dalton’s Atomic Theory
In 1808, John Dalton based his model
of the atom on solid experimental
discoveries.
He stated that:
- Elements are made up of extremely small particles
called atoms.
-All atoms making up an element are the same.
-Each compound is unique and consists of particular
atoms arranged in a particular way.
-Chemical reactions involve reshuffling of atoms to form
new compounds made from the old atoms.
He also developed three laws:
1. The Law of Definite Proportions
compounds
2. The Law of Multiple Proportions
mole ratios
3. The Law of Conservation of Mass
The Thomson Model of the Atom
By the middle of the 19th century,
there was extensive evidence of
atoms, but also smaller parts
which make up the atom.
J.J. Thomson discovered and
showed that atoms have
negatively charged “electrons”
and positively charged “protons”.
His model was
called the
“Plum Pudding”
model due to it’s
resemblance to
the English desert.
The Rutherford-Bohr Model of the Atom
Rutherford proposed that the atom consists
of a tiny, positively charged nucleus,
surrounded by a cloud of negatively
charged electrons.
The protons are equal in number to the
electrons thus making an electrically neutral
entity.
Niels Bohr took Rutherford’s
model a bit further…
He stated that electrons are restricted to
certain paths called an “orbital” a fixed
distance from the nucleus.
Electrons could only emit or absorb
energy when they moved from one orbital
to another.
It was not until 1932 that James Chadwick discovered the
neutron that was predicted by Rutherford.
Year
Milestone
Scientist(s)
~400
BC
The first coherent atomic
theory
Democritus
1804
First "modern" atomic theory
J. Dalton
1869
First periodic table
D. Mendeleev
Radioactivity discovered and
identified
H. Becquerel, M.
Curie, E. Rutherford
Electron discovered
J.J. Thomson
Identification of an atomic
nucleus
E. Rutherford, H.
Geiger
1913
Mass of electron determined
R. A. Millikan
1913
Positive charge in the nucleus
and naturally occuring isotopes J.J. Thomson
discovered
1913
Atomic number determined,
periodic table reorganized
H. Moseley
1913
Bohr Model of the Atom
N. Bohr
1919
Proton existance confirmed,
neutron proposed
E. Rutherford
1931
Neutron Identified
J. Chadwick
1896-9
1897
1909-11
Exercise p. 144 # 7,8,10,12 & define Dalton’s 3 Laws
Atomic Number and Atomic Mass
The elements are differentiated from one another by the
numbers of protons in the nucleus.
Atomic Number = the number of protons in the nucleus.
A neutral atom has no charge, therefore:
In a neutral atom,
# of Electrons = # of Protons
There are three pieces of information are usually shown for
each element on the periodic table:
NOTE:
If electrons are added or subtracted, the resulting
particle is an ION.
For Example:
How many electrons are possessed by the following?
a) N3b) Ca2+
c) At
Since both neutrons and protons have a mass of 1.0 amu
(atomic mass unit), the total atomic mass of an atom will be
found by their combined totals.
What about the electrons?
1 / 1836 amu. …… so we consider them to have no mass
For Example:
a)
27
Al
13
protons =
neutrons =
electrons =
b)
75
As
33
protons =
neutrons =
electrons =
c)
52
Al+3
24
protons =
neutrons =
electrons =
d)
19
F-1
9
protons =
neutrons =
electrons =
Isotope
Species having the same atomic number, but different
atomic masses (same # of protons, different number of
neutrons).
Eg:
1
1H
= H = Ordinary Hydrogen (called “protium”)
2
1H
= D = Deuterium (sometimes call “heavy” hydrogen)
3
1H
= T = Tritium (called “radioactive” hydrogen)
The molar masses given on the periodic table are found by
calculating the average mass of a sample containing a
mixture of isotopes.
For Example:
Experiments show that chlorine is a mixture which is
75.77% Cl-35, and 24.23% Cl-37. If the precise molar mass
of Cl-35 is 34.968853 g/mol and of Cl-37 is 36.965903 g/mol,
what is the average molar mass of the chlorine atoms in
such a mixture
You may also use the atomic mass to calculate the average.
The average mass will be less exact, but still satisfactory.
Exercises p. 146 #13-17; p. 147 #19
p.149 #22a-f; p. 150 #23a,b
The Electronic Structure of the Atom
When a hydrogen atom is irradiated by
energy, some of the energy is absorbed
then reemitted as light.
If the light is passed through a prism, a “line spectrum” is
observed.
In 1913, Niels Bohr proposed a model that explained why
the observed line spectrum for hydrogen looks the way it
does.
He proposed that:
• The electron in hydrogen can
only exist in specific energy
states.
• These energy states are
associated with specific circular
orbits which the electron can
occupy around the atom.
• When an electron absorbs
energy, it instantaneously
moves from one orbit to
another.
• The greater the energy, the
farther the orbit is from the
nucleus.
ENERGY LEVEL:
A specific amount of energy which an electron in an
atom can possess.
The energy levels of hydrogen have the pattern below (“n” is
the number of the energy level).
The observed spectrum represents energy level
differences occurring when an electron gives off energy
and drops from a higher energy level.
The energy difference between two different energy levels is
called the QUANTUM energy associated with the transition
between the two levels.
A few years after Bohr published his theories, several
changes were made to his ideas.
The idea of electrons orbiting along a specific path in a well
defined orbit had to be abandoned.
Instead, different electrons, depending on their energies,
occupy particular regions of space called “orbitals”.
ORBITAL:
The actual region of space occupied by an electron in a
particular energy level.
The Energy Level Diagram for Hydrogen
The lowest sets of energy levels for hydrogen are as follows:
Each dash represents the energy possessed by a particular
orbital in the atom.
The letter s, p, d, and f refer to the four “types” of orbitals.
SHELL:
The set of all orbitals having the same n value.
Eg: The 3rd shell consists of the 3s, 3p, and 3d
orbitals.
SUBSHELL:
A set of orbitals of the same type.
Eg: The set of five 3d orbitals in the 3rd shell is a
subshell.
- All the shells for a hydrogen atom with a given value
of n have the same energy (not true for atoms with
more than one electron).
Rules governing which types of orbitals can occur:
o For a given value of “n”, n different types of orbitals
are possible
For n=1: only the s-type is possible
For n=2: the s- and p-types are possible
For n=3: the s-, p-, and d-types are possible
For n=4: the s-, p-, d-, and f-types are
possible.
o An s-type subshell consists of ONE s-orbital
o A p-type subshell consists of THREE p-orbitals.
o A d-type subshell consists of FIVE d-orbitals.
o An f-type subshell consists of SEVEN f-orbitals.
The Energy Level Diagram for Polyelectronic Atoms
The energy level diagram must be modified to describe
any other atom.
The below diagram applies to ALL polyelectronic atoms
(atoms having more than one electron).
ELECTRON CONFIGERATIONS
The addition of electrons to the orbitals follows three
simple rules:
Aufbau Principal:
As atomic number increases, electrons are added to the
available orbitals. To ensure LOWEST POSSIBLE ENERGY
for the atom, electrons are added to the orbitals having the
lowest energy first.
Pauli Exclusion Principal:
A maximum of TWO electrons can be placed in each orbital.
Hunds Rule:
When electrons occupy subshells of equal energy, they must
be singly occupied with electrons having parallel spins. 2nd
electrons are then added to each subshell so each electron
has opposite spin.
Writing Electron Configurations for Neutral Atoms
ELECTRON CONFIGERATION:
Describes which orbitals in an atom contain electrons and
how many electrons are in each orbit.
The electron configuration of most elements can be easily
determined by using the “orbital version” of the periodic
table:
2 Simple Rules:
1. As atomic number increases, electrons are added to the
orbitals having the lowest energy first ... this is to ensure the
LOWEST POSSIBLE ENERGY.
2. A maximum of 2 electrons can be placed in each orbital.
Notice that 4s fills before 3d, which fills before 4p orbital –
refer to the periodic table above.
Ex.
H = (1s1 )
He = (1s2 )
Li = (1s2 2s1 )
Be = (1s2 2s2 )
B = (1s2 2s2 2p1 )
O = (1s2 2s2 2p4 )
F = (1s2 2s2 2p5 )
Ne = (1s2 2s2 2p6 )
Fe = (1s2 2s2 2p6 3s2 3p6 4s2 3d6 )
Se = (1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p4 )
After neon, the 2nd shell has been filled and is said to be
CLOSED.
For Example:
Predict the electron configuration of the following:
a) Si
b) Tc
c) Ca
d) Zr
e) Ga
CORE NOTATION
The electrons belonging to an atom can be broken into two
subsets:
1. The CORE electrons.
2. The OUTER electrons.
The CORE of an atom is the set of electrons with the
configuration of the nearest noble gas having an atomic
number LESS than that of the atom being considered.
The OUTER electrons are all those outside the core. Since
the core electrons are not involved in chemical reactions,
they are excluded from the electron configuration.
Rules:
1. Locate atom and note the noble gas at the end of row
above the element.
2. Place noble gas into square brackets [ ] and write the rest
of the electron configuration.
3. Exception: If you asked to write the core notation for a
noble gas such as Kr, you must show the electron
configuration as Kr ([Ar] 4s23d104p6) Do not write Kr ([Kr])
For example:
Write the following using core notation.
a. Al
b. Zr
c. Ga
d. Co
Writing Electron Configurations for Ions
Negative Ions: (anions)
Add electrons to the last unfilled subshell, starting where the
neutral atom left off.
For Example:
O([He] 2s2 2p4) + 2e- 
S([Ne] 3s2 3p4) + e- 
Positive Ions: (cations)
2 Rules:
1. Electrons in the outermost shells (largest n value) are
removed first.
2. If there are electrons in both the s- and p- orbitals of the
outermost shell, the electrons in the p- orbitals are removed
first.
p-electrons BEFORE s-electrons BEFORE d-electrons
Outermost electrons are removed preferentially. Also, e- in
the highest energy outermost orbital require the least
amount of energy to be completely removed from the atom.
For Example:
Sn([Kr] 5s2 4d10 5p2)  2e- +
Sn([Kr] 5s2 4d10 5p2)  4e- +
V([Ar] 4s2 3d3)  2e- +
Ru3+
Sb3+
S2N3-
Electron Configuration Exceptions
2 exceptions to the configurations of elements up to Kr:
Cr ([Ar] 4s2 3d4) - “3d4” is one e- short of a half filled subshell.
Cu ([Ar] 4s2 3d9) - “3d9” is one e- short of a filled subshell.
The actual configurations for Cr and Cu are found to be:
Cr ([Ar] 4s1 3d5) - “4s1” and “3d5” are two half filled subshells.
Cu ([Ar] 4s1 3d10) - “4s1” is a half filled subshell, and “3d10” is
a filled subshell.
Therefore:
A filled or exactly half filled d subshell is especially stable.
Because of this extra stability, an atom or ion that is one eshort of a “d5” or “d10” configuration will shift an e- from the ssubshell having the highest energy into the unfilled dsubshell.
Predicting the Number of Valence Electrons
Valence Electrons:
• Electrons that can take place in chemical reactions.
• Are all the electrons in the atom EXCEPT:
o Core electrons.
o In filled d- or f- subshells.
For Example:
Al([Ne] 3s2 3p1) has 3 valence electrons:
 “3s2 3p1”
Ga([Ar] 4s2 3d10 4p1) has 3 valence electrons:
 Omit “3d10” b/c filled
Pb([Xe] 6s2 4f14 5d10 6p2) has 4 valence electrons:
 Omit “4f14” and “5d10” b/c filled
Xe([Kr] 5s2 4d10 5p6) has ZERO valence electrons:
 Noble gas configuration
Exercises p. 156 #27a-g; p.157 #28; p.158 #29a-g
The Periodic Table
By 1817, chemists had discovered 52 elements; by 1863
that number had risen to 62.
William Odling (1857):
- Elements could be divided into 13 groups.
- Based in chemical and physical properties.
John Newlands (1863-1866):
- Assigned hydrogen an arbitrary mass of 1 and ordering
elements by their mass.
- Every eighth element shared a common set of
properties (the law of octaves).
- Very elegant, but did not allow for prediction of new
elements.
- Every time new elements were discovered, the table
had to be rearranged.
Dimitri Mendeleev (Russian) (1869):
- Organized the elements according to their masses and
properties.
- When he did this, he discovered that certain properties
recur PERIODICALLY.
- He left gaps in his table for elements he proposed had
not yet been discovered.
- This new periodic table allowed chemists to:
o Organize and understand their data.
o Predict new properties.
Mendleev’s work in Progress Periodic Table
Henry Mosley (1887-1915):
- Mosley developed the modern table in 1911.
Periodic Law:
The physical and chemical properties of the elements
are periodic functions of their atomic numbers.
Plain English:
When elements are arranged in order of increasing
atomic number, similar properties appear at regular
intervals.
- Elements are still grouped by properties.
- Similar properties are in the same column.
- Order is by increasing atomic number.
- Added a column of elements Mendeleev didn’t know
about.
The Modern Periodic Table:
- The table we see today wasn’t actually finished until
the early 1900’s.
- Strutt and Dorn added the Noble Gases in 1894.
- The Lanthanides and Actinides were added in the
middle of the 1900’s.
- Still changes being made today, there are always
discoveries yet to be made.
Metals, Non-Metals, and Semiconductors
In addition to the previous method, we can also classify
elements according to their metallic character.
The Properties of Metals:
• Reflect light when polished.
• Are opaque.
• Are good conductors of electricity or heat.
• Generally, but not always, flexible when in thin sheets.
• Generally malleable (can be hammered into thin
sheets, and ductile (can be stretched into wires).
• Usually solid at room temperatures (Hg is an
exception).
The Properties of Non-Metals:
• Are gasses, liquids, or brittle solids at room
temperature.
• Are poor heat and electricity conductors.
• If solids, are dull to lustrous in appearance and
opaque to translucent.
Some elements share properties from both metals and
non-metals.
Therefore, we have two subgroups:
1. Non-metals with very low electrical conductivities.
2. Non-metals with fair to moderate electrical conductivity.
Semiconductor:
A non-metal that has good electrical conductivity
A couple of important notes:
- Semiconductors were formally called “Metalloids”.
- As temperature increases
Metals
Conduc. decreases
Metalloids
Conduc. Increases
Major Divisions Within the Periodic Table
Period:
The set of elements in a given row
going across the table.
Group or Family:
The set of elements in a given column
going up and down the table.
Two important trends:
1) Elements become more metallic (or better metals)
going down a family on the table.
2) The properties of the elements change from metallic
to non-metallic going from left to right across the table.
There are several groups, rows, and “blocks” of elements:
Exercises p.160 #30; p. 162 #31-33
p.162 #34; p. 164 #35-39
Chemical Bonding
Atomic Radius:
Two factors must be taken into consideration in explaining
this periodic trend:
1. Increasing nuclear charge
2. Increasing shell
Along a Period (left to right):
• The atomic number increases while the valence electrons
remain in the same shell.
• Due to the increasing nuclear charge (pulling electrons
closer to the nucleus):
The atomic radius decreases from left to right.
Along a Group (top to bottom):
• The atomic number continues to increase.
• However, the shell increases from shell 1 to shell 2 etc.
• The atomic orbitals for each successive shell get larger and
larger.
The result is:
Atomic radii increase top to bottom along a group.
Things change when you take ions into consideration:
Cations:
Get smaller than the atomic species because you lose
outer electrons and the net positive charge draws in
remaining electrons.
Anions:
You gain electrons and they repulse each other.
The ionic radius expands out to accommodate the
repulsive forces.
Ionization Energy:
The energy needed to remove an electron from the atom.
For example:
X(g)  X+ (g) + eI1 = First Ionization Energy
It is a minimum for the alkali metals which have a single
electron outside a closed shell.
It generally increases across a row on the periodic
maximum for the noble gases which have closed shells.
The representative elements tend to gain or lose
electrons until they become isoelectronic with the
nearest noble gas.
Electron Sheilding:
Imagine a campfire.
The same happens for electrons:
When a new orbital is started, every orbital of lower energy
shields these electrons from feeling the full nuclear charge.
Electron Affinity / Electronegativity:
The energy change when an electron is accepted by
an atom
The Pauling scale is the most commonly used.
Fluorine (the most electronegative element) is assigned a
value of 4.0.
Values range down to cesium and francium which are the
least electronegative at 0.7.
If an atom has a HIGH electronegativity, it strongly attracts
electrons from a neighboring atom and may completely
remove it.
Atoms with high electronegativies also STRONGLY
ATTRACT THEIR OWN VALENCE ELECTRONS.
As a result, these valence electrons are difficult to remove
and the atom has HIGH ionization energy.
Lewis Theory
1916-1919 - Lewis, Kossel, and Langmuir made several
important proposals on bonding which lead to the
development of Lewis Bonding Theory.
Elements of the theory:
1. Valence electrons play a fundamental role in chemical
bonding.
2. Sometimes bonding involves the transfer of one or more
electrons from one atom to another. This leads to the ion
formation and IONIC BONDS.
3. Sometimes bonding involves sharing electrons between
atoms, this leads to COVALENT BONDS.
4. Electrons are transferred or shared such that each atom
gains a more stable electron configuration.
Usually this is that of a noble gas.
o Eg: having 8 outer shell electrons.
This arrangement is called an OCTET.
LEWIS SYMBOLS:
A common chemical symbol surrounded by up to 8 dots.
The symbol represents the nucleus and the electrons of the
filled inner shell orbitals.
The dots represent the valence electrons.
Elements in the same group have similar Lewis symbols:
Note:
This only works well for the representative elements.
Transition metals, actinides and lanthanides have
incompletely filled inner shells - we can't write simple Lewis
structures for them.
Exercises p. 165 #40; p. 168 #50-51;
p. 170 #53, 55; p.173 #60-61
Types of Bonding
Intramolecular Bonds- between atoms
Ionic Bonds
forces that hold ionic compounds together.
Forming the ionic bond:
Step 1:
A positive ion forms by the LOSS of 1 or more e-.
Representative elements become cations, isoelectronic
with the nearest Noble gas.
For others (transition metals), not necessarily.
Step 2:
A negative ion forms by GAINING sufficient e- to
become isoelectronic with the nearest Noble gas.
The Lewis symbol will show an OCTET (8) of electrons
for the anion.
Step 3:
The oppositely charged ions come together to form an
ionic compound.
The electrostatic attraction between the ions forms the
bond.
In the solid state, each anion surrounds itself with
cations, and each cation with anions, forming an ionic
crystal.
Step 4:
A formula unit of an ionic compound is the smallest
collection of ions that would be electrically NEUTRAL.
The formula unit is automatically obtained when the
Lewis structure of the compound is written.
The ionic crystal then consists of each constituent ion
bound together in the crystal, not of individual
molecules.
When an ionic compound forms, its constituent atoms
change size. The overall process can be visualized in
the following figure for NaCl:
ATOMS  COMPOUNDS: Properties change…
- Their reactivity decreases considerably.
- They become neutral overall.
- There are no unique molecules in many ionic solids.
- Ionic compounds become electrically conductive when
melted or dissolved.
IONIC BONDS are Very STRONG
- compounds held together by ionic bonds have HIGH
MELTING TEMPERATURES.
Covalent Bonds
Covalent bonds arise from the sharing of electrons between
atoms (generally of groups IVA, VA, VIA, and VIIA).
Each electron in a shared pair is attracted to both nuclei
involved in the bond.
The valence electrons involved in the bond are called the
BONDING ELECTRONS or the BOND PAIR.
Those not involved in the bond are called the
NONBONDING ELECTRONS or the LONE PAIRS.
The pairs repel each other and thus tend to stay as far away
as possible.
OCTET RULE :
An atom other than hydrogen tends to form bonds until it is
surrounded by eight valence e-.
COVALENT BONDS are VERY STRONG
- compounds held together by covalent bonds can have
LOW or HIGH MELTING TEMPERATURES.
Multiple Bonds:
The Bond order describes the bond:
1 = single, 2 = double, 3 = triple
A single bond is formed when 1 pair of e- is shared; double 2
pairs, triple 3 pairs.
Polar Covalent Bonds
Assume one atom with a somewhat higher electronegativity
than the other:
1.
This will cause the electrons to be shared unevenly,
such that the shared electrons will spend more time on
average closer to the atom that has the higher
Electronegativity.
2.
The greater the difference in electronegativity in the
bonding atoms, the greater the polarity of the bond.
Atoms with widely different electronegativity values
( ≥ 2.0) tend to form IONIC BONDS.
True NON-POLAR COVALENT BONDS form only
when diatomic molecules are formed with two identical
atoms. ( ≤ 0.4)
Everything else will form a POLAR COVALENT BOND.
(0.5 – 1.9)
Ionic Vs Covalent Compound Properties
Ionic
1. High melting and boiling
points
2. Conduct electricity when
melted
3. Many soluble in water
but not in nonpolar liquid
4. Generally are odorless
Covalent
1. Low melting and boiling
points
2. Poor electrical
conductors in all phases
3. Many soluble in nonpolar
liquids but not in water
4. Generally have an odor
Intermolecular Bonds – between molecules
Hydrogen Bonding
Polar molecules, such as water molecules, have a weak,
partial negative charge at one region of the molecule (the
oxygen atom in water) and a partial positive charge
elsewhere (the hydrogen atoms in water).
When water molecules are close together, their positive and
negative regions are attracted to the oppositely-charged
regions of nearby molecules.
The force of attraction, shown here as a dotted line, is called
a hydrogen bond.
Each water molecule is hydrogen bonded to four others
HYDROGEN BONDS are WEAK
London Dispersion Forces
The London dispersion force is the weakest intermolecular
force.
The London dispersion force is a temporary attractive force
that results when the electrons in two adjacent atoms occupy
positions that make the atoms form temporary dipoles.
Because of the constant motion of the electrons, an atom or
molecule can develop a temporary (instantaneous) dipole
when its electrons are distributed unsymmetrically about the
nucleus.
Dispersion forces are present between any two molecules
(even polar molecules) when they are almost touching.
LONDON DISPERSION FORCES are the WEAKEST
Exercises p. 172 #57a-d; p.175 #62-63; p.177 #68; p.180
#73; p. 181 # 74-76; p.182 # 79,82,83a-d
Writing Lewis Structures
SIMPLE IONIC COMPOUNDS
Example 1:
Draw the following Lewis structures:
a) MgCl2
b) KBr
c) Li2S
d) MgO
THE LEWIS STRUCTURES OF COVALENT
COMPOUNDS THAT OBEY THE OCTET RULE
Lewis structures show how the VALENCE electrons are
distributed in a molecule.
For Example:
H2O
For most chemistry problems, you’ll find the configuration for
a compound of type AXn, where A is the central atom, and X
is the external or terminal atom.
1. Find the total number of valence e-.
Go by the column or group that it’s in.
2. After connecting your central atom to the terminal atoms
with single bonds, begin adding the remaining valence e- as
lone pairs.
First around the terminal atoms.
Then around the central (if you have any left over).
3. Satisfy the octet rule for your central atom by either:
Replacing a lone pair on your terminal atoms with a
bond (to make a double bond).
OR
By replacing two lone pairs with two bonds (to make a
triple bond).
REMEMBER:
You can’t just add double bonds without first removing a lone
pair.
Not only are you adding more electrons than you started
with, but you’re probably breaking the octet rule for the
terminal atoms.
Examples:
NH4+
CHO2-
HOPO
THE LEWIS STRUCTURES OF COVALENT COMPOUNDS
THAT VIOLATE THE OCTET RULE
Electron Deficient Molecules: one or more atoms (other than
H) does not possess a full octet of e-’s
In addition to H, Be, B, and Al are exceptions to the octet
rule.
These atoms have such low electronegativities that they gain
1e- for every 1 they contribute
For Example:
BF3
Atoms Having an Expanded Octet of Valence Electrons:
Elements in the third and fourth periods frequently attain
more than an octet of valence e- when they form covalent
compounds (the electrons are placed in d-orbitals).
Other than the fact that the central atom will end up with
more than eight valence electrons, the same rules as above
apply.
Two common examples include P and S.
For Example:
PCl5
SF4
Exercises p. 188 # 86 a, b, d, e, f, k, m, n,
p, q, bb, cc
Chemical Families
Noble Gases
Known for their extremely slow reactivity, these were once
thought to never react; neon, one of the noble gases, is used
to make bright signs.
Alkali Metals
The most reactive metal family, these must be stored under
oil because they react violently with water! They dissolve
and create an alkaline, or basic, solution, hence their name.
Alkaline Earths
These also are reactive metals, but they don’t explode in
water; pastes of these are used in batteries.
Halogens
Known as the “salt formers,” they are used in lighting and
always exist as diatomic molecules in their elemental form.
Chemical Reactivity
Chemical Reactivity in a given period:
In the same period, it is the number of electrons that chiefly
determines reactivity. The smaller the number of electrons
transferred between reacting atoms, the more vigorous the
reaction.
Chemical Reactivity in a given group:
In the same group, elements have the same number of
outershell electrons and it is the atomic radius which largely
determines reactivity. The larger metals loose outer shell
electrons more easily and smaller nonmetals (whose
attraction for electrons by the nucleus is greater) are more
likely to take electrons away from other metals (or share with
other nonmetals).
Example:
Cs, Li
Ti, Ca
N, F
What is on the test?
History of the Atom
o Early Greeks and Medieval Chemists
o John Dalton
o J.J. Thomson
o Rutherford-Bohr
The Atom
o Atomic number and atomic mass
o # of protons, neutrons, and electrons
o Isotopes
The Electronic Structure of the Atom
o Theory
o Electron configurations for neutral atoms, ions.
o Core notation
o Exceptions
o Valence electrons
I,F
The Periodic Table
o Development
o Divisions within the Table
o Metals, non-metals, and semiconductors.
Periodic Trends
o Atomic and Ionic Radii
o Ionization Energy
o Electron Affinity and Electronegativity
Lewis Theory
o Elements of the Theory
o Ionic Bonds
o Covalent Bonds
- Multiple Bonds
- Polarity
o Hydrogen Bonding
o London Dispersion Forces
o Writing Lewis Structures
- Simple Ionic Compounds
- Structures that Obey the Octet Rule
- Structures that Violate the Octet Rule