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Transcript
```II. Atom. Periodic Table
and Trends
Prepared by PhD Halina Falfushynska
Rutherford model (planetary model)
of atom
Draw a diagram showing the location of
each part of the atom.
Components of Atoms
The components that
make up the atom are
known as sub-atomic
particles. Some
particles have charges
which are integer
multiples of the
elementary charge.
For comparison:
• Human hair ~106 atoms wide
• Pencil line ~106 atoms wide
• HIV virus ~800 atoms wide (~108 atoms total)
• E. coli ~1011 atoms total
Line Spectra and the Bohr Model
Bohr Model
• Since the energy states are quantized, the light emitted
from excited atoms must be quantized and appear as
line spectra.
• After lots of math, Bohr showed that
 1 
En   2.178  10 18 J  2 
n 
where n is the principal quantum number (i.e., n = 1, 2,
3, … and nothing else).
Line Spectra and the Bohr Model
Bohr Model
E  E f  Ei  h
En   2.178  10
• We can show that
Ei  f  h 
hc

  2.178  10
• When ni > nf, energy is emitted.
• When nf > ni, energy is absorbed
18
18
1 1
J  2  2 
 n f ni 
 1 
J  2 
n 
Line Spectra and the
Bohr Model
Bohr Model
Quantum Mechanics and Atomic Orbitals
• Schrödinger proposed an equation that contains both
wave and particle terms.
^
H  E 
• Solving the equation leads to wave functions.
• The wave function gives the shape of the electronic
orbital. [“Shape” really refers to density of electronic
charges.]
• The square of the wave function, gives the probability of
finding the electron ( electron density ).
Quantum Mechanics and Atomic Orbitals
Solving Schrodinger’s
Equation gives rise to
‘Orbitals.’
These orbitals
provide the electron
density distributed
Orbitals are
described by
quantum numbers.
Quantum Mechanics and Atomic Orbitals
Orbitals and Quantum Numbers
• Schrödinger’s equation requires 3 quantum numbers:
1. Principal Quantum Number, n. This is the same as Bohr’s n.
As n becomes larger, the atom becomes larger and the
electron is further from the nucleus. ( n = 1 , 2 , 3 , 4 , …. )
2. Angular Momentum Quantum Number, . This quantum
number depends on the value of n. The values of begin at
0 and increase to (n - 1). We usually use letters for (s, p, d
and f for = 0, 1, 2, and 3). Usually we refer to the s, p, d
and f-orbitals.
3. Magnetic Quantum Number, m . This quantum number
depends on . The magnetic quantum number has integral
values between - and + . Magnetic quantum numbers
give the 3D orientation of each orbital.
Quantum Numbers of Wavefuntions
Quantum #
Symbol
Values
Description
Principal
n
1,2,3,4,…
Size & Energy of orbital
Angular
Momentum

0,1,2,…(n-1)
for each n
Shape of orbital
Magnetic
m
-…,0,…+ 
for each 
Relative orientation of orbitals within
same 
Spin
ms
+1/2 or –1/2
Spin up or Spin down
Angular Momentum Quantum #
Name of Orbital
()
0
s (sharp)
1
p (principal)
2
d (diffuse)
3
f (fundamental)
4
g
Representations of Orbitals
The s-Orbitals
Representations of Orbitals
The p-Orbitals
d-orbitals
Orbitals and Their Energies
Many-Electron Atoms
Many-Electron Atoms
Electron Spin and the Pauli Exclusion
Principle
Many-Electron Atoms
Electron Spin and the Pauli Exclusion
Principle
• Since electron spin is quantized, we define ms =
spin quantum number =  ½.
• Pauli’s Exclusions Principle: no two electrons
can have the same set of 4 quantum numbers.
• Therefore, two electrons in the same orbital must have
opposite spins.
Orbitals CD
Figure 6.27
Atomic and Mass Numbers, Isotopes
• Using the periodic
table, the atomic
number (number of
protons; used to
identify the element),
elemental symbol,
and atomic weight can
be identified.
Nuclides
• Atoms may exist as more than one isotope –
atoms with the same number of protons but
have a different number of neutrons. The
mass number is used to distinguish isotopes.
• One of two or more atoms whose nuclei have
the same number of neutrons but different
numbers of protons are called isotones.
• One of two or more atoms or elements having
the same atomic weights or mass numbers
but different atomic numbers are called
isobars
Average atomic weight
• The average atomic weight is a weighted
average of naturally occurring isotopes and
fractional abundance in nature.
Example: Carbon has two appreciably present
isotopes, 12C and 13C.
Respectively, the abundances are 98.9% and 1.1%.
• average atomic weight = (0.989)(12 amu) +
(0.011)(13 amu)
= 12.011 amu (on
periodic table)
av. atomic wt. = (mass isotope A)(% A) + (mass isotope B)(% B)
Molecules and Molecular Compounds
• Sharing electrons molecular (covalent) compound
• Trading electrons  ionic compound
Charges
• A cation is formed when
more protons than
electrons are present in an
atom or molecule
• An anion is formed when
more electrons than
protons are present in an
atom or molecule.
Examples
19
9
F
9 protons
10 neutrons
9 electrons
+
Na
23
11
11 protons
12 neutrons
10 electrons
2-
O
18
8
8 protons
10 neutrons
10 electrons
Predicting Ionic Charges
• Metals form cations.
•
Nonmetals form anions.
Naming Inorganic Compounds
• Positive ions:
• • Cations formed from metal (main group or transition)
atoms have the same name.
• • If a metal can form different cations, the positive charge
is indicated by a Roman numeral in parenthesis
following the name of the metal. Older names using
–ous and –ic are still seen but their use is fading.
• • Cations formed from nonmetals have names that end in
–ium (i.e. hydronium ion and ammonium ion).
The Periodic Law says:
• When elements are arranged in order
of increasing atomic number, there is a
periodic repetition of their physical and
chemical properties.
• Horizontal rows = periods
–There are 7 periods
• Vertical column = group (or family)
–Similar physical & chemical prop.
–Identified by number & letter (IA, IIA)
Areas of the periodic table
• Three classes of elements are:
1) Metals: electrical conductors, have luster,
ductile, malleable
2) Nonmetals: generally brittle and nonlustrous, poor conductors of heat and
electricity
3) Metalloids: border the line-2 sides
– Properties are intermediate between
metals and nonmetals
Electron Configurations in Groups
• Elements can be sorted into 4
different groupings based on their
electron configurations:
1) Noble gases
2) Representative elements
3) Transition metals
4) Inner transition metals
Classify elements based on electron
configuration.
• Group IA – alkali metals
–Forms a “base” (or alkali) when
reacting with water (not just dissolved!)
• Group 2A – alkaline earth metals
–Also form bases with water; do not
dissolve well, hence “earth metals”
• Group 7A – halogens
–Means “salt-forming”
Electron Configurations in Groups
1) Noble gases are the elements in
Group 8A (also called Group18 or 0)
– Previously called “inert gases”
because they rarely take part in a
reaction; very stable = don’t react
– Noble gases have an electron
configuration that has the outer s
and p sublevels completely full
Electron Configurations in Groups
2) Representative Elements are in
Groups 1A through 7A
– Display wide range of properties,
thus a good “representative”
– Some are metals, or nonmetals, or
metalloids; some are solid, others
are gases or liquids
– Their outer s and p electron
configurations are NOT filled
Electron Configurations in Groups
3) Transition metals are in the “B”
columns of the periodic table
– Electron configuration has the
outer s sublevel full, and is now
filling the “d” sublevel
– A “transition” between the metal
area and the nonmetal area
– Examples are gold, copper, silver
Electron Configurations in Groups
4) Inner Transition Metals are located
below the main body of the table,
in two horizontal rows
– Electron configuration has the outer
s sublevel full, and is now filling the
“f” sublevel
– Formerly called “rare-earth”
elements, but this is not true
because some are very abundant
H
1
Li
3
Na
11
1s1
1s22s1
Do you notice any similarity
in these configurations of
the alkali metals?
1s22s22p63s1
1s22s22p63s23p64s1
K
19
Rb
37
Cs
55
Fr
87
1s22s22p63s23p64s23d104p65s1
1s22s22p63s23p64s23d104p65s24d10 5p66s1
1s22s22p63s23p64s23d104p65s24d105p66s24f
145d106p67s1
ALL Periodic Table Trends
• Influenced by three factors:
1. Energy Level
–Higher energy levels are further
away from the nucleus.
2. Charge on nucleus (# protons)
–More charge pulls electrons in
closer. (+ and – attract each other)
• 3. Shielding effect
(blocking effect?)
What do they influence?
Energy levels and Shielding have
an effect on the GROUP (  )
Nuclear charge has an effect on a
PERIOD (  )
#1. Atomic Size - Group trends
• As we increase the
atomic number (or
go down a group).
• each atom has
another energy
level,
• so the atoms get
bigger.
H
Li
Na
K
Rb
#1. Atomic Size - Period Trends
• Going from left to right across a period,
the size gets smaller.
• Electrons are in the same energy level.
• But, there is more nuclear charge.
• Outermost electrons are pulled closer.
Na
Mg
Al
Si
P
S Cl Ar
Ions
• Some compounds are composed of
particles called “ions”
–An ion is an atom (or group of atoms)
that has a positive or negative charge
• Atoms are neutral because the number
of protons equals electrons
–Positive and negative ions are formed
when electrons are transferred (lost or
gained) between atoms
Ions
•
Metals tend to LOSE
electrons, from their outer energy level,
and thus a positively charged particle is
formed = “cation”. Cations are smaller
than the atom they came from
• Nonmetals tend to GAIN one or more
electrons. Negative ions are called
“anions”. Anions are bigger than the
atom they came from
Ion Group trends
• Each step down a
energy level
• Ions therefore get
bigger as you go
down, because of the
level.
Li1+
Na1+
K1+
Rb1+
Cs1+
Ion Period Trends
• Across the period from left to right, the
nuclear charge increases - so they get
smaller.
• Notice the energy level changes between
anions and cations. (more protons would
pull the same # of electrons in closer)
Li1+
B3+
Be2+
C4+
N3-
O2-
F1-
#2. Trends in Ionization Energy
• Ionization energy is the amount of
energy required to completely remove
an electron (from a gaseous atom).
• Removing one electron makes a 1+ ion.
• The energy required to remove only
the first electron is called the first
ionization energy.
Ionization Energy
• The second ionization energy is the
energy required to remove the second
electron.
–Always greater than first IE.
• The third IE is the energy required to
remove a third electron.
–Greater than 1st or 2nd IE.
Ionization Energy - Group trends
• As you go down a group, the
first IE decreases because...
–The electron is further away
from the attraction of the
nucleus, and
–There is more shielding.
Ionization Energy - Period trends
• All the atoms in the same period have
the same energy level.
• Same shielding.
• But, increasing nuclear charge
• So IE generally increases from left to
right.
• Exceptions at full and 1/2 full orbitals.
Electron Affinities
• Electron affinity is the opposite of ionization
energy.
• Electron affinity: the energy change when a
gaseous atom gains an electron to form a gaseous
ion:
Cl(g) + e-  Cl-(g)
• Electron affinity can either be exothermic (as the
above example) or endothermic:
Ar(g) + e-  Ar-(g)
```
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