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Transcript
ATOMS & THE
PERIODIC TABLE
Subatomic Particles
• Protons and electrons are the only particles that
have a charge.
• Protons and neutrons have essentially the same
mass.
• The mass of an electron is so small we ignore it.
Atoms and the Periodic Table
Subatomic Location
Particle
Charge
Size
Mass
Proton
p+
Nucleus
Positive
1+
About same
as neutron
D = 10-15 m
1 amu
1.6726 x 10-24 g
Neutron
n0
Nucleus
No charge
Neutral
0
About same
as proton
D = 10-15 m
1 amu
1.6749 x 10-24 g
Electron
e-
Orbital
cloud
Negative
1-
Tiny
compared to
proton &
neutron
D = 10-18 m
1/1840 amu
9.11 x 10-28 g
Size of Atoms
Particle Diameter (meters)
atom 10-10
nucleus 10-14
proton 10-15
neutron 10-15
electron 10-18
1 Angstrom = 10-10 m
History of the Atomic Model
Democritus – 400 BC
atoms make up all substances
John Dalton – 1766-1844
atom is a solid hard sphere
Joseph John Thomson – 1856-1940
discovered the electron in 1897
plum pudding model of atom
positive sphere with negative e- embedded
• Lived from (1766-1844)
• All elements are composed of atoms
• Atoms of the same element are identical.
Each element has unique properties .
• Atoms of different elements can be
chemically combined in simple whole
number ratios to form compounds.
• Law of Conservation of Matter/Mass
The Electron
• Streams of negatively charged particles were
found from cathode tubes.
• J. J. Thompson is credited with their discovery
(1897).
The Atom, circa 1900:
• “Plum pudding” model,
put forward by Thompson.
• Positive sphere of matter
with negative electrons
imbedded in it.
Ernest Rutherford – 1871-1937
gold foil experiment
most of atom empty space
positive nucleus contains most of the mass
discovered protons in 1919
James Chadwick – 1891-1974
discovered the neutron in 1932.
Niels Bohr – 1885-1962 (Danish)
electrons move around the nucleus in fixed
orbits that have a set amount of energy
Electron Cloud Model of Atom
1. came to be used to estimate the positions of
electrons in an atom
2. uncertainty principle, which states that it is
not possible to obtain precise values of both
position and momentum of a particle at the
same time
3. probability of finding an electron
Protons
• The number of protons distinguished 1 atom
from another
• Most atoms are very stable
• It takes a lot of energy to add or remove a
proton from an atom
• Atomic number = number of protons
• The Periodic Table is arranged by number of
protons
Symbols of Elements
Elements are symbolized by one or two letters.
Atomic Number
Number of protons = The atomic number
Atomic Mass
The mass of an atom in atomic mass units (amu)
is the total number of protons and neutrons in
the atom.
Mass Number
• The number of protons plus neutrons in an
atom.
• Always a whole number.
• Written or indicated like this:
H
C
Si
Cu
K
Calculating Number of Neutrons
• Subtract:
Mass Number - Atomic Number =
Neutrons
Mass Number - # of protons =
Neutrons
Notes on Finding Numbers of Protons,
Neutrons, Electrons
Isotopes
• Atoms of the same element with different
numbers of neutrons
Isotopes
• To distinguish between isotopes of an
element
• Ex: Neon has 3 main isotopes
Neon
Protons
Neutrons
Ne-20
Ne-21
10
10
10
11
Mass
number
20
21
Ne-22
10
12
22
Average Atomic Mass
• Atomic mass unit – 1/12 the mass of a C-12
atom
• To calculate avg. atomic mass
Cesium
Natural %
Mass
Abundance Number
Avg.
Atomic
Mass
Cs-133
75%
133
99.75
Cs-132
20%
132
26.4
Cs-134
5%
134
6.7
Example – Avg. Atomic
Mass
Chemical Goals
• To be Chemically
Stable
Unstable atoms are
radioactive: their
nuclei change or
decay by spitting out
radiation, in the form
of particles or
electromagnetic
waves.
• To be Electronically
neutral
To have no charge
on the atom.
To have the same #
of protons as
electrons.
Why does an atom stay together?
The strong nuclear force keeps protons and
neutrons together in the nucleus in spite of the
repulsion of the protons for each other.
The strong nuclear force acts only over a very
short distance. It doesn’t work outside the
nucleus.
The strong nuclear force is stronger than the
electromagnetic force.
Valence Electrons
• Electrons in outer energy level.
• Can only have 8 or less. This is the Octet
Rule.
• These electrons are the ones involved in
bonding with other atoms and the ones with
the most energy.
• Looking at the Periodic Table, you can tell the
number of valence electrons for the A
Families from the Roman Numeral
designation
Electrons
1.
e- located far from nucleus in a series of energy levels.
2.
Each e- has a certain amount of energy.
3.
The further the e- gets from the nucleus the more energy it has. Valence ehave the most energy. Those closest to the nucleus have the least amount
of energy.
Energy Levels
1. Each energy level can only hold a certain number of e2. Electrons always fill low energy orbitals (closest to
the nucleus) before filling higher energy ones.
3. The high the energy level occupied by the
e-, the easier it is for the e- to escape from
the atom.
4. Quantum of energy – amount of energy needed to move an e- from its
present energy level to the next one
5. Ground State – the lowest energy level for
an e-.
Electron Placement on the Periodic Table
2 e-
8 e18 e32 e-
Energy Levels & Orbitals
Energy
Level
1
2
3
4
Maximu
# of
m
Orbitals
Number
of
Electrons
2
8
18
32
1
4
9
16
Sublevel
Names
s
s, p
s, p, d
s, p, d, f
Type of
Sublevel
Number of
orbitals in
Sublevel
Number of
electrons in
orbital
s
1
2
p
3
6
d
5
10
f
7
14
Electron Configuration Chart
Principal
Energy Level
Number of
Sublevels
Type of
Sublevel
(Atomic
Orbitals)
Number of
Electrons
1
1
s
2
2
2
s, p
s-2, p-6 = 8
3
3
s, p, d
4
4
s, p, d, f
s-2, p-6,
d-10 = 18
s-2, p-6,
d-10, f-14
= 32
Aufbau /Orbital Diagram
EX: Electron Configuration for Potassium
EX: Electron Configuration for Arsenic
EX: Electron Configuration for Silver
19 electrons
33 electrons
47 electrons
Pauli Exclusion Principle
• An orbital can hold 0, 1, or 2 electrons and if
there are 2 electrons in the orbital they must
have opposite spin.
Aufbau Principle
• Rules for Orbital Filling
• Lower-energy orbitals fill first.
• An orbital can hold only 2 e- with opposite
spins.
• The most stable arrangement for e- is one with
the maximum number of unpaired e-. This
minimized e- to e- repulsion and stabilizes the
atom. – Hund’s Rule
Hund’s Rule
• When filling up sublevels other than s,
electrons are placed in individual orbitals
before they are paired up.
Increasing Energy in Electron
Sublevels
THE DIAGONAL RULE
MUST GO IN THIS
ORDER:
1s, 2s, 2p, 3s, 3p, 4s, 3d,
4p, 5s, 4d, 5p, 6s, 4f, 5d,
6p, 7s, 5f, 6d
and 7p. These orbitals will
account for all the
elements now known.
This diagonal rule can help
account for the octet rule,
too.
The Energy Flow
• The order of increasing energy of the orbitals is then
read off by following these arrows, starting at the top
of the first line and then proceeding on to the
second, third, fourth lines, and so on. This diagram
predicts the following order of increasing energy for
atomic orbitals.
• 1s < 2s < 2p < 3s < 3p <4s < 3d <4p < 5s < 4d < 5p < 6s
< 4f < 5d < 6p < 7s < 5f < 6d < 7p < 8s ...
Bohr Model of Atoms
only represents energy levels, not orbitals
Lithium
Nobel Gases
Neon
Krypton
Argon
Drawing Atoms
The Bohr Model
11 p+
e- 3 p+
Lewis Structures
Electron Dot Diagrams
• Describes e- arrangement in atoms
• Describes bond arrangement in molecules.
• Uses dots to represent valence e- around an
atom
• EX:
Li
Ne
O
Si
Dimitri Mendeleev
• In the late 1860's, Mendeleev
began working on his great
achievement: the periodic table
of the elements. By arranging all
of the 63 elements then known
by their atomic weights, he
managed to organize them into
groups possessing similar
properties. Where a gap existed
in the table, he predicted a new
element would one day be found
and deduced its properties. And
he was right. Three of those
elements were found during his
lifetime: gallium, scandium, and
germanium.
Mendeleev’s Periodic Table
Moseley’s Periodic Table
In 1913 Henry Moseley came up with
this Periodic Table. The elements are
arranged by increasing atomic number.
• A group (also known as a family) is a vertical
column in the periodic table of the chemical
elements. There are 18 groups in the standard
periodic table.
• Elements in a group have similar configurations of
the outermost electron of their atoms – same
number of valence e• This gives the groups of elements similar physical
and chemical characteristics.
• With each group across a period, the elements
have one more proton and electron and become
less metallic.
• Rows of elements are called periods. The period
number of an element signifies the highest
unexcited energy level for an electron in that
element.
Physical Properties
Metals
• Good electrical conductors
and heat conductors.
• Malleable - can be beaten
into thin sheets.
• Ductile - can be stretched
into wire.
• Possess metallic luster.
• Opaque as thin sheet.
• Solid at room temperature
(except Hg).
Nonmetals
• Poor conductors of heat and
electricity.
• Brittle - if a solid.
• Nonductile.
• Do not possess metallic
luster.
• Transparent as a thin sheet.
• Solids, liquids or gases at
room temperature.
Chemical Properties
Metals
• Usually have 1-3 electrons
in their outer shell.
• Lose their valence electrons
easily.
• Have lower
electronegativities.
Nonmetals
• Usually have 4-8 electrons
in their outer shell.
• Gain or share valence
electrons easily.
• Have higher
electronegativities.
Metalloids
• Electronegativities and ionization energies
between those of metals and nonmetals
• Possess some characteristics of metals/some of
nonmetals
• Reactivity depends on properties of other
elements in reaction
• Often make good semiconductors
• Boron, Silicon, Germanium, Arsenic, Antimony,
Tellurium, Polonium
Ionic Bonding
Between Metals & Nonmetals
Metals
• Make Ionic Compounds
• Lose Electrons
• Have positive oxidation
numbers
• Are first in a formula
Ex: Na2S
• When naming, write the
name just as it is
• Ex: Sodium sulfide
Nonmetals
• Make Ionic Compounds
• Gain Electrons
• Have negative oxidation
number
• Are second in the formula
Ex: MgO
• When naming, drop the
ending of the name and add
IDE
• Ex: Magnesium oxide
More Ionic Bonding
• Strong attractions or forces between atoms in
these compounds
• High melting and boiling points, good
conductors
• Have a lattice structure
Sodium chloride Lattice
Covalent Bonding
Covalent Molecules
•
•
•
•
Between 2 or more nonmetals
Share electrons
Still try to obey the Octet Rule
Weak bonds between molecules but strong
bonds between atoms
• Low melting and boiling points, usually non
conductors
• Simple molecules or giant structures can form
More Covalent Bonding
• Diamond and graphite are held together by
this type of bond (allotropes)
• Ex: H2O, H2O2, CH4, CO2
• Ex: Diatomics – H2, N2, O2, F2, Cl2, Br2, I2
Water Molecule
Moles
•The mole is the SI unit for amount of
substance
•A mole (abbreviated mol) is the amount
of a substance that contains as many
particles as there are atoms in exactly
12 g of carbon-12.
Avogadro’s Number
• 6.02 x 1023 – is the number of particles in exactly
one mole of a pure substance.
• Conversion factor!
• 6.02 x 1023 particles = 1 mole
If I have 3.45 moles of hydrogen, how
many particles do I have?
6.02 x 1023 = 1 Mole