Transcript Atoms and the Periodic Table

Atoms and the Periodic
Table
Unit 6
Standards
• SPS1. Students will investigate our current
understanding of the atom.
• SPS1a. Examine the structure of the atom in
terms of proton, electron, and neutron
locations. atomic mass and atomic number,
atoms with different numbers of neutrons
(isotopes), and explain the relationship of the
proton number to the elementʼs identity.
Standards
• SPS4. Students will investigate the arrangement of
the Periodic Table.
• SPS4a. Determine the trends of the following:
number of valence electrons; Types of ions formed
by representative elements; location of metals,
nonmetals, and metalloids; Phases at room
temperature.
• SPS4b. Use the Periodic Table to predict the above
properties for representative elements
Democritus (400 B.C.)
• Proposed that matter was
composed of tiny indivisible
particles
• Not based on experimental
data
• Greek: atomos
John Dalton (1807)
• British schoolteacher; based
his atomic theory on others’
experimental data
• Billiard Ball Model (atom is a
uniform, solid sphere)
John Dalton
Dalton’s Four Postulates
1. Elements are composed of small indivisible
particles called atoms.
2. Atoms of the same element are identical. Atoms
of different elements are different.
3. Atoms of different elements combine together in
simple proportions to create a compound.
4. In a chemical reaction, atoms are rearranged, but
not changed.
John Dalton
• Law of definite proportions: a chemical
compound always contains the same elements
in exactly the same proportions
• (while his theory had flaws,) Dalton’s theories
are the foundation for modern atomic theory
Henri Becquerel (1896)
• Discovered radioactivity
(spontaneous emission of
radiation from the nucleus)
• Three types:
– alpha () - positive
– beta () - negative
– gamma () - neutral
J. J. Thomson (1903)
• Discovered electrons
(negative particles within
the atom)
• Plum-Pudding Model of
atomic structure
Ernest Rutherford (1911)
• Discovered the nucleus
– dense, positive charge in
the center of the atom
• Nuclear Model
Niels Bohr (1913)
• Energy Levels
– electrons can only exist in
specific energy states
• Planetary Model
-electrons move in circular
orbits
Erwin Schrödinger (1926)
• Quantum mechanics
– electrons can only exist in
specified energy states
• Electron Cloud Model
– orbital: region around
the nucleus where e- are
likely to be found
Erwin Schrödinger (1926)
Electron Cloud Model (orbital)
• dots represent probability of finding an enot actual electron
James Chadwick (1932)
• Discovered neutrons
– neutral particles in the
nucleus of an atom
The Structure of Atoms
• protons: + charge
• neutron: no charge
• electrons: - charge
-very small in mass
-electric force holds
holds atom together
James Chadwick (1932)
Neutron Model
• revision of Rutherford’s Nuclear Model
Atomic Structure
3 main subatomic particles:
• Proton: +1; in the nucleus
• Neutron: 0; in the nucleus
• Electron: -1; cloud around nucleus
Atomic Structure
• Each element has a unique # of protons
• Unreacted atoms have no overall charge
(# +protons = # -electrons)
• Electric force (opposite charges) holds atom
together
Elements
• Atomic Number = # protons (never, ever
changes!)
• Mass Number = # protons + # neutrons
• Mass number – atomic number = # neutrons
Isotopes
• Isotopes: vary in mass because the number of
neutrons is different
• Radioisotopes: emit radiation and decay over
time
Atomic Mass
• Atomic mass unit (u): expresses tiny atomic
mass
ex. 1 u = 1/12 the mass of a 12C atom
• 1 proton = 1 u; neutron = 1 u
• Average atomic mass = weighted average
(mass # )(# of atoms)  (mass # )(# of atoms)

total # of atoms
Mole
• Mole (mol): SI counting unit for small particles
• (1 mol = 602,213,670,000,000,000,000,000
• 6.022 x 10³²- Avogadro’s number)
Modern Atomic Theory
• Modern models: electrons are in certain
energy levels (cannot be predicted- Bohr)
-orbital: region where electron is likely to be
found
• Electron energy levels: number of energy
levels depend on number of electrons
• Electron transitions: Electrons jump between
levels when atom gains or loses electrons
Electron Energy Levels
Valence electrons:
• outer level e• Determine chemical properties of atom
3
4 Types of Orbitals
Energy Level
#/ Type of
Orbitals
s
p
Total # Orbitals
#
Electrons
d f
1
1
1=1
2
1
3
1+3=4
8
3
1
3 5
1+3+5=9
18
4
1
3 5 7 1 + 3 + 5 + 7 =16
(x2)
2
32
Bohr Model Diagrams
• Simplified energy levels using Bohr’s idea of
circular orbits.
Lithium
Atomic #: 3
Mass:
7
# of p:
# of e:
# of n:
3
3
4
Can replace with:
e-
e-n
p
p
n
n
n
p
e-
3p
4n
Maximum eLevel 1 2eLevel 2 8eLevel 3 18eLevel 4 32e-
Bohr Model Activity
• Choose a number between 2 & 35.
• Find your element by the atomic number you picked.
Record the name, atm # and mass # on your paper.
• Draw a Bohr Model diagram for the element.
– Round off the mass listed on the table and subtract the
atomic # to find the # of neutrons.
– Abbreviate the # of ‘p’ and ‘n’ in the nucleus
• Repeat with a new element (total of two)
The Periodic Table
• Dmitri Mendeleev
(1869, Russian)
• Organized elements by
increasing atomic mass
• Predicted the existence
of undiscovered
elements.
The Periodic Table
• Henry Mosely (1913, British): created
modern periodic table.
• Organized elements by increasing atomic
number.
• Fixed problems in Mendeleev’s arrangement.
Modern Periodic Table
• Periodic Law: elements with similar
properties appear at regular intervals;
elements change periodically with atomic #
• Elements are organized by:
- period (row): ↑ atomic #
- group (column): result of valence electrons
Groups and Periods
1
2
3
4
5
6
7
• Group (Family)
• Period
Modern Periodic Table
• Valence Electrons
– e- in the outermost energy level
• Atomic Radius:
• First ionization energy:
-energy required to remove an e- from a
neutral atom
-cation: + ion (loses electrons)
-anion: - ion (gains electrons)
Form cations
Form anions
Valence Electrons
• Group # = # of valence e- (except He)
• Period # = # of energy levels
1
1
2
3
4
5
6
7
18
2
13 14 15 16 17
Charges
+1
1
2
3
4
5
6
7
18
+2
+13 -+14 -15 - 16 -17
Classification of Elements
• Metals: alkali, alkaline-earth metals,
transition metals
• Nonmetals: halogens, noble (inert) gases, C,
N, O, P, S, Se
• Semiconductors
Alkali Metals
+1
VERY REACTIVE!
Rarely found alone in nature
Alkaline Earth Metals
+2 React less violently than
Alkali Metals
Transition Metals
Many are hard, high melting and
boiling points
+3
+12
F- Transition Metals/Lanthanides
↓Shiny, reactive, many emit light
Actinides
↗ Primarily radioactive- used in fuel, bombs
Halogens
Combine readily with metals to form
salts; diatomic (X₂)
+17
Noble (Inert) Gases
Almost entirely unreactive (stable)
18
Octet Rule
• All elements gain or lose electrons so they end
up with the same electron configuration as
the nearest noble gas
- completely filled outermost energy levels
make an element stable ()
(elements with 1 electron away from a noble
gas are extremely active)
Metals and Nonmetals
Metals
Nonmetals
good conductors of electricity
poor conductors of electricity
good conductors of heat
poor conductors of heat
ductile (easily made into
wires)
malleable (easily shaped or
formed)
not ductile
usually shiny
dull, not shiny
not malleable
Metalloids and Semiconductors
• Metalloids/semiconductors have properties
of both metals and nonmetals
-able to conduct heat, electricity under certain
conditions
-silicon is used in computers
-boron increases steel’s hardness
-arsenic treats wood to prevent decay
Last, But Not Least…
• Hydrogen: single proton, electron
-unique behavior
-most abundant element in universe
-reacts with many other elements