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UNIT 2
Atomic Structure
and
The Periodic Table
(Chapters 4,5,6)
The smallest particle of an
element that retains the
properties of that element.
Atoms consist of electrons,
protons and neutrons.
These are called subatomic particles.
Sub Atomic Particles
Electrons (e-):
 negatively charged
 revolve around the nucleus in
electron clouds
 mass much less than protons or
neutrons (mass = 0 a.m.u)
 Up to 2 electrons on 1st energy
level, up to 8 on each level after
1st.
Protons (p+):
 positive charge
 inside the nucleus
 Mass equals 1 a.m.u
 Mass is 1840 times more than eNeutrons (n0):
 Neutral (no) charge
 inside the nucleus
 mass equals 1 a.m.u.
Nucleus
 Central
core of atom
 Composed of protons and neutrons
 Tiny (compared to atom as a whole)
 Very Dense
 Contains most of the atoms mass
ATOMIC NUMBER
Number of protons (p+) in the nucleus
 p+ = e- (in a neutral atom)
 Each element contains a different # of
protons (what makes each element
different)
 “the element’s address on the
periodic table”
MASS NUMBER
 Total
number of protons & neutrons in
an atom
 n0 + p+ = Mass #
 Remember (atomic # = # of p+)
 Mass # - atomic # = n0
ISOTOPES
Atoms of the same element that have:
• same number of protons
• different number of neutrons
• different masses
How many neutrons does this atom have?
Mass #
Atomic #
Chemical
Symbol
NOTE: Some periodic tables list the mass # on top
and atomic # on the bottom while other periodic
tables do the reverse. Remember, the smaller of
the two numbers is always the atomic number!!
For about 50 years past the time of
John Dalton, the atom was
considered a solid indivisible mass.
The later discovery of subatomic
particles shattered this theory.
Dalton’s Atomic Theory
1.
2.
3.
4.
5.
Elements are composed of tiny particles
called atoms.
Atoms of the same element are identical.
Different elements have different atoms.
Atoms of different elements can
physically or chemically combine.
Chemical Reactions (RXNs) occur when
atoms are separated, joined, or
rearranged.
Atoms of one element never change into
atoms of another element as a result of a
chemical reaction.
He discovered the electron and
proposed the plum-pudding model of
the atom. This model said nothing
about the number of protons and
neutrons.
He discovered the nucleus and proposed a
nuclear atom in which electrons surround
a dense nucleus. He thought of the rest of
the atom as empty space.
He had nothing to do with
discovering a model of the atom but
really likes teaching chemistry.
(Just checking to see if you are
paying attention!)
A student of Rutherford’s, he proposed that
electrons are arranged in orbits around
the nucleus. He explained that the electrons
do not fall into the nucleus because they
have fixed energy levels.
Development of Periodic
Table
 70
elements known by 1800s
 Dmitri Mendeleev discovered a
way to systematically & logically
group elements
 What are some ways elements
are grouped?
Mendeleev’s Work
 Arranged
Elements in columns, so
elements w/ similar properties were side
by side
 Arranged in columns of increasing Atomic
Mass
 Left blank spaces in table because there
were no known elements w/ the correct
properties and mass for that space.
Moseley’s Periodic Table
Corrections
 Put
table according to atomic number
instead of atomic mass
PERIODIC LAW
1)
2)
3)
Elements arranged in order of
increasing atomic number
Periodic Repetition of physical &
chemical properties
Elements with same properties in
columns
Modern Periodic Table
– horizontal rows going
across the periodic table
 7 periods (each has its own energy
level)
 Properties of elements change as you
move across
 Pattern of properties repeats when
you move from one period to the next.
 Periods
 Groups/Families–
elements in the
periodic table going down in columns.
 Similar
chemical & physical
properties
 Classified by number and/or letter

Ex: 1A, 2A, 3A…
Representative Elements =
Group A
Transition Metals = Group B
METALS
a)
b)
c)
d)
e)
Left side of periodic table (except H)
High electrical conductivity
High Luster (very shiny when clean)
Ductile (drawn into wires)
Solid @ room temperature (except
Hg)
Alkali Metals: Group 1A (not H)
1)

Very Reactive
Alkaline Earth Metals: Group 2A
3) Transition Metals: Group B
2)

Middle of the periodic table
NON-METALS
a)
b)
c)
d)
Upper right corner of periodic table
Non-lustrous (not shiny)
Poor conductors of electricity
Some are gases at room temp.(O,N)
--others brittle solids (S)
1)
Halogens – group 7A
1)
2)
2)
Very Reactive
7 e- on outer energy level
Noble Gases -- group 8A
1)
2)
Unreactive – Few to no chemical
RXNs
Full outer energy level
Semi-Metals (metalloids)
 Divides
metals from nonmetals
 properties intermediate between metals
and nonmetals
 Ex: Si, Ge, ...
 “stair case” on right
Side of periodic table
The Modern Atom
 Quantum
Mechanical Model:
 The Energy level of an electron is
the region around the nucleus
where the electron is likely to be
moving.
 Each
period has its own principle
energy level
 principle energy levels are assigned
values in order of increasing energy:
n = 1, 2, 3, 4, and so forth
One Important Rule
 Aufbau
Principle: Electrons enter
orbitals of lowest energy first.
Energy Sublevels
 each
principal energy level has
energy sublevels.
 sublevels are called atomic orbitals
and are represented by the letters:
s, p, d, f.
Principle
Number of
NRG Level Sublevels
Orbital
Types
n =1
1
s
n=2
2
s, p
n=3
3
s, p, d
n=4
4
s, p, d, f
Each sub-orbital can hold a different
number of electrons
An s orbital can contain 2 electrons.
A p orbital can contain 6 electrons
A d orbital can contain 10 electrons.
A f orbital can contain 14 electrons.
p
s
d
f
3 QUESTIONS TO ASK
 What

(principle energy level)
 What

section?
(type of sub-orbital)
 What

Row?
seat?
(how many electrons in that
sub-orbital)
Example 1:
Write the electron configuration
for nitrogen.
7N
2
2
3
1s 2s 2p
Example 2:
Write the electron configuration
for Fe.
26Fe
2
2
6
2
6
2
6
1s 2s 2p 3s 3p 4s 3d
p
s
d
f
S orbitals
in 1st NRG level
 1 orbital per energy level
 Each can contain up to 2 electrons
 Alkali and Alkaline Earth metals
 Spherical in nature
 Starts
P orbitals
in 2nd NRG level
 3 orbitals per NRG level
 Each orbital can contain up to 2 e(6 e- total per sublevel)
 Groups 3A – 8A
 Dumbbell shaped
 Starts
D orbitals
at 3rd NRG level
 5 orbitals per NRG level
 2 e- per orbital (10 total e- per
sublevel)
 Clover shape
 Transition Metals
 Starts
F orbitals
in 4th NRG level
 7 orbitals per NRG level
 2 e- per orbital (14 e- total per
sublevel)
 Complicated, unexplainable shape
 Lanthanide & Actinide series
 Starts
2 Important Rules
 Pauli
Exclusion Principle: An
atomic orbital may contain a
maximum of two electrons. The
electrons must have opposite spins.
 Hund’s Rule: When electrons occupy
orbitals of equal energy, one electron
enters each orbital before they pair.
For Example:
2s
2p
After the s sublevel gets two
electrons, three electrons enter
the p orbitals before they pair.
What do we know about the table so far?
Trends in Atomic Size
 radius
of an atom cannot be
measured directly
 atomic radius is estimated as half of
the distance between the nuclei of
two like atoms in a diatomic molecule.
Atomic size generally
increases as you move
down a group of the
periodic table.
More principal energy levels
are added.
Atomic size generally decreases as you
move from left to right across a period.
Increasing Nuclear charge pulls
electrons in closer to nucleus
Trends in Ionization Energy
 When
an atom gains or loses an
electron, it becomes an ion.
 Ionization Energy is the amount of
energy required to remove an
electron from a neutral atom.
 The smaller the atom , the more
ionization energy required.
Decreases
Increases
Trends in Ionic Size
 Nonmetallic
elements:
 gain electrons to form negative ions.
 Negative ions (anions) are always
larger than the neutral atoms from
which they form.
 Metallic elements:
 lose electrons to form positive ions.
 Positive ions (cations) are always
smaller than the neutral atoms from
which they form.
Trends in Electronegativity
 tendency
for an atom to attract
electrons when it is chemically
combined with another atom.
 decreases as you move down a
group
 increases as you go across a period
from left to right.