Download The Bohr Atom The Bohr Atom Electronic Transitions 2.3 Light, Atom

The Bohr Atom
Electrons exist in fixed
energy levels
surrounding the nucleus
Promotion of
electron occurs as
it absorbs energy
Energy is released as
the electron travels
back to lower levels
Quantization of energy
Excited State
Relaxation
2.3 Light, Atomic Structure,
and the Bohr Atom
The Bohr Atom
Electronic Transitions
• Amount of energy absorbed in jumping
from one energy level to a higher energy
level is a precise quantity
• Energy of that jump is the energy
difference between the orbits involved
• Orbit - what Bohr called the fixed energy
levels
• Ground state - the lowest possible energy
state
1
2.3 Light, Atomic Structure,
and the Bohr Atom
2.3 Light, Atomic Structure,
and the Bohr Atom
Bohr Theory
• Allowed levels are quantized energy levels,
orbits
• Electrons are found only in these energy levels
• Highest-energy orbits are farthest from the
nucleus
• Atoms
– absorb energy by excitation of electrons to higher
energy levels
– release energy by relaxation of electrons to lower
energy levels
• Energy differences may be calculated from the
wavelength of light emitted
Modern Atomic Theory
• Bohr’s model of the atom when applied to
atoms with more than one electron failed to
explain their line spectra
• One major change from Bohr’s model is that
electrons do not move in orbits
• Atomic orbitals - regions in space with a
high probability of finding an electron
• Electrons move rapidly within the orbital
giving a high electron density
The Periodic Law and the
Periodic Table
• Dmitri Mendeleev and Lothar Meyer - two
scientists working independently developed
the precursor to our modern periodic table.
• They noticed that elements have distinct
regular variation of their properties when
listed in order of atomic mass.
• Periodic law - the physical and chemical
properties of the elements are periodic
functions of their atomic numbers.
2
Classification of the Elements
2.4 The Periodic Law
and the Periodic Table
2.4 The Periodic Law
and the Periodic Table
Periods
Parts of the Periodic Table
• Period – a horizontal row of elements in
the periodic table. They contain 2, 8, 8,
18, 18, and 32 elements,
• Group – also called families are columns
of elements in the periodic table.
• Elements in a particular group or family
share many similarities.
Category Classification of
Elements
• Metals - elements that tend to lose
electrons during chemical change,
forming positive ions.
• Nonmetals - a substance whose atoms
tend to gain electrons during chemical
change, forming negative ions.
• Metalloids - have properties intermediate
between metals and nonmetals.
3
Classification of the Elements
Periods
2.4 The Periodic Law
and the Periodic Table
2.4 The Periodic Law
and the Periodic Table
Classification of Elements
Metals
• Metals:
– A substance whose atoms tend to lose
electrons during chemical change
– Elements found primarily in the left 2/3 of
the periodic table
• Properties:
–
–
–
–
High thermal and electrical conductivities
High malleability and ductility
Metallic luster
Solid at room temperature
Classification of Elements
Nonmetals
• Nonmetals:
– A substance whose atoms may gain
electrons, forming negative ions
– Elements found in the right 1/3 of the
periodic table
• Properties:
– Brittle
– Powdery solids or gases
– Opposite of metal properties
4
2.4 The Periodic Law
and the Periodic Table
2.4 The Periodic Law
and the Periodic Table
Atomic Number and Atomic Mass
• Atomic Number:
– The number of protons in the nucleus of
an atom of an element
– Nuclear charge or positive charge from
the nucleus
• Most periodic tables give the element
symbol, atomic number and atomic
mass
Element Information in the
Periodic Table
20
Ca
Calcium
40.08
atomic number
symbol
name
atomic mass
Electron Arrangement
and the Periodic Table
• The electron arrangement is the primary
factor in understanding how atoms join
together to form compounds
• Electron configuration - describes the
arrangement of electrons in atoms
• Valence electrons - outermost electrons
– The electrons involved in chemical bonding
5
Valence Electrons
• The number of valence electrons is the
group number for the representative
elements
• The period number gives the energy level
(n) of the valence shell for all elements
Valence Electrons and Energy
Level
• How many valence electrons does fluorine
have?
– 7 valence electrons
• What is the energy level of these electrons?
– Energy level is n = 2
Valence Electrons - Detail
• What is the total number of electrons in
fluorine?
– Atomic number = 9
– 9 protons and 9 electrons
• 7 electrons in the valence shell, (n = 2 energy
level), so where are the other two electrons?
– In n = 1 energy level
– Level n=1 holds only two electrons
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Determining Electron Arrangement
List the total number of electrons, total number of valence
electrons, and energy level of the valence electrons for
silicon.
1. Find silicon in the periodic table
•
•
•
Group IVA
Period 3
Atomic number = 14
2. Atomic number = number of electrons in an atom
•
Silicon has 14 electrons
Determining Electron Arrangement #2
List the total number of electrons, total number of valence electrons,
and energy level of the valence electrons for silicon.
3. As silicon is in Group IV, only 4 of its 14
electrons are valence electrons
•
Group IVA = number of valence electrons
4. Energy levels:
•
•
•
n = 1 holds 2 electrons
n = 2 holds 8 electrons (total of 10)
n = 3 holds remaining 4 electrons (total = 14)
Determining Electron Arrangement
Practice
List the total number of electrons, total
number of valence electrons, and energy
level of the valence electrons for:
• Na
• Mg
• S
• Cl
• Ar
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The Quantum Mechanical Atom
• Bohr’s model of the hydrogen atom didn’t
clearly explain the electron structure of other
atoms
– Electrons in very specific locations, principal energy
levels
– Wave properties of electrons conflict with specific
location
• Schröedinger developed equations that took into
account the particle nature and the wave nature
of the electrons
Schröedinger’s equations
• Equations that determine the probability of
finding an electron in specific region in space,
quantum mechanics
– Principal energy levels (n = 1,2,3…)
– Each energy level has one or more sublevels or
subshells (s, p, d, f)
– Each sublevel contains one or more atomic
orbitals
Energy Levels and Sublevels
PRINCIPAL ENERGY LEVELS
• n = 1, 2, 3, …
• The larger the value of n, the higher the energy
level and the farther away from the nucleus the
electrons are
• The number of sublevels in a principal energy
level is equal to n
– in n = 1, there is one sublevel
– in n = 2, there are two sublevels
8
Principal Energy Levels
• The electron capacity of a principal
energy level (or total electrons it can hold) is
2(n)2
– n = 1 can hold 2(1)2 = 2 electrons
– n = 2 can hold 2(2)2 = 8 electrons
• How many electrons can be in the n = 3
level?
– 2(3)2 = 18
• Compare the formula with periodic table…..
n=1, 2(1)2=2
n=2, 2(2)2=8
n=3, 2(3)2=18
n=4, 2(4)2=32
Sublevels
• Sublevel: a set of energy-equal orbitals within
a principal energy level
• Subshells increase in energy:
s<p<d<f
[sharp, principal, diffuse, and fundamental]
• Electrons in 3d subshell have more energy than
electrons in the 3p subshell
• Specify both the principal energy level and a subshell
when describing the location of an electron
9
Sublevels in Each Energy Level
Principal energy
level (n)
Possible
subshells
1
1s
2
2s, 2p
3
3s, 3p, 3d
4
4s, 4p, 4d, 4f
Orbitals
• Orbital - a specific region of a sublevel
containing a maximum of two electrons
• Orbitals are named by their sublevel and
principal energy level
– 1s, 2s, 3s, 2p, etc.
• Each type of orbital has a characteristic
shape
– s is spherically symmetrical
– p has a shape much like a dumbbell
Orbital Shapes
• s is spherically
symmetrical
10
Orbital Shapes
• Each p has a shape much like a dumbbell,
differing in the direction extending into space
Orbital Shapes
Electron Arrangement and
the Periodic Table
• There are five different d shapes.
• The f orbitals have seven different shapes, too
complicated and therefore seldom shown.
Subshell
Number of
orbitals
s
1
p
3
d
5
f
7
• How many electrons can be in the
4d subshell?
•10
11
Shell 4
• Each orbital within a
sublevel contains a
maximum of 2
electrons
• Energy increases as n,
shell number
increases, but ALSO
increases as move
from s to p to d to f
sublevels
4f •• •• •• •• •• •• ••
Increasing Energy
Electron Arrangement and
the Periodic Table
Electron Arrangement and
the Periodic Table
Electron Arrangement and
the Periodic Table
Quantum Mechanical Model
4d
•• •• •• •• ••
Sublevel
4p
•• •• ••
Orbital
4s
••
Electron
Electron Spin
• Electron Configuration - the
arrangement of electrons in atomic orbitals
• Aufbau Principle - or building up
principle helps determine the electron
configuration
– Electrons fill the lowest-energy orbital that is
available first
– Remember s<p<d<f in energy
– When the orbital contains two electrons, the
electrons are said to be paired
Rules for Writing Electron
Configurations
• Obtain the total number of electrons in the atom
from the atomic number
• Electrons in atoms occupy the lowest energy
orbitals that are available – 1s first
• Each principal energy level, n contains only n
sublevels
• Each sublevel is composed of orbitals
• No more than 2 electrons in any orbital
• Maximum number of electrons in any principal
energy level is 2(n)2
12
Electron Arrangement and
the Periodic Table
Electron Distribution
• This table lists the number of electrons in each
shell for the first 20 elements
• Note that 3rd shell stops filling at 8 electrons even though if
could hold more
Electron Distribution
Electron Arrangement and
the Periodic Table
Writing Electron Configurations
• H
– Hydrogen has
only 1 electron
– It is in the
lowest energy
level & lowest
orbital
– Indicate
number of
electrons with a
superscript
– 1s1
• Li
– Lithium has 3
electrons
– First two have
configuration
of helium – 1s2
– 3rd is in the
orbital of
lowest energy
in n=2
– 1s2 2s1
13
Electron Arrangement and
the Periodic Table
Electron Arrangement and
the Periodic Table
Electron Configuration Examples
• Give the complete electron
configuration of each element
– Be
–N
– Na
– Cl
– Ag
Shorthand Electron
Configurations
• Uses noble gas symbols to represent the
inner shell and the outer shell or valance
shell is written after
• Aluminum- full electron configuration is:
1s22s22p63s23p1
What noble gas configuration is this?
•Neon
•Configuration is written: [Ne]3s23p1
• Remember:
– How many subshells are in each principal energy
level?
– There are n subshells in the n principal energy level.
– How many orbitals are in each subshell?
– s has 1, p has 3, d has 5, and f has 7
– How many electrons fit in each orbital?
– 2
– Hence: s can have 2, p 6, d 10 and f 14 electrons
14
• Remember:
– How many electrons is subshells?
– There are n subshells in the n principal energy level.
s up to 2, p up to 6, d up to 10, f up to 14 electrons
Therefore:
for n=1: s subshell - up to 2 electrons
for n=2: s and p subshells - up to 8 electrons
for n=3: s, p, and d subshells - up to 18 electrons
Electron Arrangement and
the Periodic Table
Electron Arrangement and
the Periodic Table
for n=4: s, p, d and f subshell, up to 32 electrons
Shorthand Electron
Configuration Examples
• N
• S
• Ti
• Sn
Classification of Elements
According to the Type of
Subshells Being Filled
15
The Octet Rule
[eight in Latin of Greek - octo, οκτώ …ocho]
• The noble gases are extremely stable
– Called inert as they don’t readily bond to other
elements
• The stability is due to a full complement of
valence electrons in the outermost s and p
sublevels:
– 2 electrons in the 1s of helium
– the s and p subshells full in the outermost shell of
the other noble gases (eight electrons)
The Octet Rule
Octet of Electrons
• Elements in families other than the noble
gases are more reactive
– Strive to achieve a more stable electron
configuration
– Change the number of electrons in the atom to
result in full s and p sublevels
• Stable electron configuration is called the
“noble gas” configuration
The Octet Rule
The Octet Rule
• Octet Rule - elements usually react in such a way
as to attain the electron configuration of the noble
gas closest to them in the periodic table
– Elements on the right side of the table move right to the
next noble gas
– Elements on the left side move “backwards” to the
noble gas of the previous row
• Atoms will gain, lose or share electrons in
chemical reactions to attain this more stable
energy state
16
The Octet Rule
Ion Formation and the Octet Rule
• Metallic elements tend to form positively
charged ions called cations
• Metals tend to lose all their valence
electrons to obtain a configuration of the
noble gas
Na
Na+ + e-
Sodium atom
11e-, 1 valence e[Ne]3s1
Sodium ion
10e[Ne]
Cation Formation
Na
Na+ + e-
Sodium atom
11e-, 1 valence e[Ne]3s1
Sodium ion
10e[Ne]
Energy required to remove an electron from
an isolated atom, to form a cation is called
the ionization energy
The Octet Rule
Ion Formation and the Octet Rule
• All atoms of a group lose the same number of
electrons
• Resulting ion has the same number of electrons as
the nearest (previous) noble gas atom
Al
Al3+ + 3e-
Aluminum atom
13e-, 3 valence e[Ne]3s23p1
Aluminum ion
10e[Ne]
Formation of a cation, positively charged species with a
missing electron(s)
17
The Octet Rule
Using the Octet Rule
• The octet rule is very helpful in predicting
the charges of ions in the representative
elements
• Transition metals still tend to lose electrons
to become cations but predicting the charge
is not as easy
• Transition metals often form more than one
stable ion
– Iron forming Fe2+ and Fe3+ is a common example
Composition of the Atom
The Octet Rule
Examples Using the Octet Rule
• Give the charge of the
most probable ion
resulting from these
elements
–
–
–
–
Ca
Sr
S
P
• Which of the
following pairs of
atoms and ions are
isoelectronic?
–
–
–
–
Cl-, Ar
Na+, Ne
Mg2+, Na+
O2-, F-
Calculating Subatomic Particles
in Ions
• How many protons, neutrons and electrons
are in the following ions?
39
19
K
32
16
S2-
24
12
Mg 2
18
Trends in the Periodic Table
• Many atomic properties correlate with
electronic structure and so also with their
position in the periodic table
Trends in the Periodic
Table
–
–
–
–
atomic size
ion size
ionization energy
electron affinity
Atomic Size
• The size of an element increases moving
down from top to bottom of a group
• The valence shell is higher in energy and
farther from the nucleus traveling down the
group
Trends in the Periodic
Table
• The size of an element decreases from left
to right across a period
• The increase in magnitude of positive charge in
nucleus pulls the electrons closer to the nucleus
Variation in Size of Atoms
19
Trends in the Periodic
Table
Trends in the Periodic
Table
Trends in the Periodic
Table
Cation Size
Cations are smaller than their parent atom
• More protons than electrons creates an increased
nuclear charge
• Extra protons pulls the remaining electrons
closer to the nucleus
• Ions with multiple positive charges are even
smaller than the corresponding monopositive
ions
– Which would be smaller, Fe2+ or Fe3+?
Fe3+
• When a cation is formed isoelectronic with a
noble gas the valence shell is lost decreasing the
diameter of the ion relative to the parent atom
Anion Size
Anions are larger than their parent
atom.
• Anions have more electrons than protons
• Excess negative charge reduces the pull
of the nucleus on each individual electron
• Ions with multiple negative charges are
even larger than the corresponding
monopositive ions
Relative Size of Select Ions and
Their Parent Atoms
20
Trends in the Periodic
Table
Ionization Energy
• Ionization energy - The energy required to
remove an electron from an isolated atom
• The magnitude of ionization energy
correlates with the strength of the attractive
force between the nucleus and the
outermost electron
• The lower the ionization energy, the easier
it is to form a cation
ionization energy + Na  Na+ + e-
Trends in the Periodic
Table
Ionization Energy of Select Elements
• Ionization decreases down a family as the
outermost electrons are farther from the nucleus
• Ionization increases across a period because the
outermost electrons are more tightly held
• Why would the noble gases be so unreactive?
Trends in the Periodic
Table
Electron Affinity
• Electron Affinity - The energy released
when a single electron is added to an
isolated atom
• Electron affinity gives information about
the ease of anion formation
– Large electron affinity indicates an atom
becomes more stable as it forms an anion
-
Br + e  Br + energy
Formation of an anion, negatively charged species with an
extra electron
21
Trends in the Periodic
Table
Periodic Trends in Electron
Affinity
• Electron affinity
generally
decreases down a
group
• Electron affinity
generally increases
across a period
22