Download Periodic Relationships

Electrons in Atoms &
Periodic Table
1
Where Are the
Electrons?
Quantum Theory or Wave Theory:
 a description of the
electron “configuration.”
One of the greatest
achievements of mankind.
2
Arrangement of
e in
Atoms
Determines:
 chemical reactivity
 bonding between atoms
 Periodic Table
 many physical properties
3
Atomic Models: History
Each atomic model was
eventually replaced because of
new experimental evidence.
1
2
3
4
Dalton: 1803
Concept of the atom
as smallest unit of an
element.
Indivisible particle
5
Thomson: 1897
 Discovered the e Atom has parts!!
electron
positive charge
“Plum pudding” model
6
Rutherford: 1911
 Au foil experiment
 Dense nucleus with
positive charge
 Most of atom is empty space
+
Nuclear model
7
Nuclear Model: Problem
Opposite charges attract, so what
keeps the electrons and nucleus apart?
+
8
Bohr: 1913
e- held in “orbits”
Motion of e- keeps them from
“falling” into nucleus
Similar to planets around sun
9
Bohr: “Planetary”
Model
 e- move in circular
orbits around nucleus,
and each orbit has a
certain energy.
10
Bohr: “Planetary”
Model
E3
E2
E1
+
“Quantized”
energy levels
11
“Bright Line Spectrum”
of Hydrogen
12
Stair Analogy: H spectrum due
to e- transition between orbits.
energy
E5
Stairs are quantized.
E4
E3
E2
Not a ramp
E1
13
e
energy
E5
E4
in Ground State
Ground state is lowest
energy of the e .
E3
E2
E1
14
e
in Excited State
e absorbs energy to move
energy
E5
to a higher energy level.
E4
E3
E2
E1
15
e
energy
E5
E4
in Excited State
E3
E2
E1
16
e
Returning to Ground
e gives off energy as light
energy
E5
photon
E4
E3
E2
Elight=Eexcited-Eground
E1
17
Elight=Eexcited-Eground
The energy of the light is the
difference between the higher and
lower energy level of the electron.
Each energy of light corresponds
to a unique color of light.
18
e
energy
E5
Returning to Ground
E4
lower energy
photon
E3
E2
E1
19
Bohr: Hydrogen Emission
Spectrum
e- absorbs energy
(heat, elec.)
E3
E2
E1
+
e- falls to lower E
and gives off
energy as light
Elight=E3-E1
20
Bohr Theory: Failings
• Why do e only have
certain orbit energies?
• Only explains the
hydrogen atom exactly.
21
Quantum Mechanics
(Wave Theory)
1926: E. Schrodinger
Currently accepted theory
Can not determine exact location
of an electron! Wow!
22
Quantum Mechanics
e
in “atomic orbitals”
Can only determine
the probability of
locating an electron.
electron cloud
23
Models
+
Dalton Thomson Rutherford
+
Bohr
+
Quantum
24
Atomic Orbital
A region in space around the
nucleus with high probability
of finding an electron.
Each atomic orbital
can hold 2 e
-
-
Analogy: student in a desk
25
Wave Model
e
Each is arranged in an atom
according to its energy.
+
3rd energy level
(higher shell)
2nd energy level
1st energy level
(lowest shell)
26
Overview
Bohr
Quantum
https://www.youtube.com/watch?feature=player
_embedded&v=8ROHpZ0A70I#t=4
27
Regents
e
Notation
Regents Periodic Table gives
the number of electrons in each
energy level or shell.
st
1
shell -
[C] = 2
-
nd
2
shell -
rd
3
shell…
4
28
Regents
What
e
is the e
“Configuration”
configuration for:
1. sodium
2. argon
3. calcium
4. copper
7. lead
6. radium (Ra)
What is the maximum electrons in:
Shell 1?
Shell 2?
Shell 3?
29
Regents
e
“Configuration”
What happens to the number of
electrons in each shell going from:
Ca to Sc?
Zn to Ga?
30
Noble Gases
At end of each row in Periodic
Table are the noble or inert gases
with 8 e in the highest shell.
Stable (not reactive) elements
31
Periodic Table by Shell
1
1
2
3
4
5
6
7
4
5
Transition elements
2
3
4
5
6
3
4
5
6
7
Inner transition elements
32
Valence Electrons
Electrons that are in the highest energy
level are called “valence electrons.”
These are the most important electrons
when atoms bond. Why?
How many “valence electrons” in:
Li
Fe
Cu
33
Valence Electrons
Note all elements in a Group have
same number of valence electrons.
Group 1: Li, Na, K, Rb, Cs, Fr
Group 16: O, S, Se, Te, Po
This is why elements in the same
group have similar properties.
34
Excited State
Remember “excited state”?
(e- have absorbed energy to move
to a higher energy level)
[Al] is 2-8-3 [Al]* could be 2-8-2-1
ground state Al
excited state Al
What element is 2-7-6-1?
35
Flame Test (Lab)
Adding energy can cause e- to
“jump” from ground state (as
written in Regents table) to
“excited state”. When e- falls
back, it emits light.
36
Flame Test for Copper
Cu atom in excited state:
2-8-17-2
Cu atom in ground state:
2-8-18-1
Can return to ground state
by emitting energy
37
Flame Test for Copper
Which photon has greater energy:
When an e the falls from E5 to E3 or
When an e the falls from E to E ?
5
2
38
Emission Spectrum
Flame test
Neon signs
Fireworks
Fireplace colors
39
Bonding
•Electron configurations are
the key to bonding.
•Atoms become ions to achieve
Noble gas electron configuration.
40
Atoms vs. Ions
F atom: [F] = 2-7
What does F need in order to have
the e configuration of a Noble gas?
F ion: [F-] = 2-8 = [Ne]
Na atom: [Na] = 2-8-1
+
+
Na ion: [Na ] = 2-8 = [Ne]
41
Practice

Write the e- configuration for
Fe & Cu.
 Based on e configuration,
predict the charge of:
Mg ion
S ion
 Write two e- configurations for
excited states of calcium.
42
Periodic Relationships
43
Early chemists describe the first element.
44
Tabulation of Elements
Mendeleev (1869)
•Tabulated by chem. &
physical properties
•Arranged by mass
•Predicted missing elements
and properties
45
Modern Periodic Table
Argon vs. potassium problem.
Now ordered by atomic number,
not mass.
Element 101 (Md)
46
Periodic Table
Most important tool in chemistry
Key to understanding chemical
and physical properties
Each group has same electron
configuration for outer shell.
47
Regents Periodic Table
Elements arranged by atomic no. (#p+)
Symbol
Atomic number & atomic mass
Electron configuration
“Charges”
48
Representative Elements
 Groups 1, 2 & 13-17
 Last digit of group number gives the
number of valence electrons.
 Examples: oxygen Group 16
(6 VE)
sodium Group 1
(1 VE)
49
Representative Elements
Some groups have special names
Group 1: alkali metals
Group 2: alkaline earth metals
Group 17: halogens
50
Noble Gases
Last element in each Period
8 VE, except He
very stable
(non-reactive)
51
Transition Elements
e.g. Iron
Regents: 2-8-14-2
e
filling
Compounds with these elements
have colored solutions.
52
Inner Transition Elements
At bottom of Periodic Table for
convenience.
53
Trends in Atomic Size
Atomic size is
measured by radius.
R
Table ‘S’
For chlorine:
radius = 100. pm
What is its radius
in meters?
54
Atomic Radius: Trends
?????????
OK
(model)
55
Atomic Radius
Down a Group: size increases
due to adding electrons to
higher energy levels (shells)
further from the nucleus.
56
Atomic Size:
Across a Period
Electrons added to same shell
+
Nuclear charge increases (more p )
Greater inward pull on the electrons
Atoms get smaller
p+ =11
2-8-1
12
13
2-8-3
14
15
2-8-5
16
17 18
2-8-8
57
Atomic Size:
Across a Period
smaller
+5
+6
Boron (2-3) vs. Carbon (2-4)
58
Atomic Radius
larger
smaller
Row: greater
nuclear charge
Column: e in
higher shell
59
Atomic Radius
Try It:
Arrange these atoms in order of
increasing size.
N, O, P, S
O<N<S<P
60
Ionization Energy (I)
Chemical properties determined
by valence electrons.
First ionization energy (I):
energy (kJ/mol) to remove a
valence e from an atom.
If ionization energy is high,
e held tightly.
61
Ionization Energy
I is endothermic (need to put
energy in to pull off an e )
+
I + X(g)  X (g) + e
ionization
energy
62
Ionization Energy:
Table ‘S’
I1 across Period
I1 down Group
II11
Atomic Number
63
Trends in I (due to size)
I1 decreases going down a Group.
The e are farther from the nucleus.
I1 increases going across a Period.
The e are closer to the nucleus.
Which corner of Periodic Table has:
-highest I1?
-lowest I1?
64
I Predicts Ionic Charges
Element
Na
Mg
I1
I2
I3
(kJ/mol) (kJ/mol) (kJ/mol)
496
4565
6912
738
1450
7732
Na atom 2-8-1
Mg atom 2-8-2
lose 1 elose 2 e-
+
Na
ion 2-8
+2
Mg ion
2-8
65
Ionization Energy
Which has smaller I and why?
O or S
Ge or Br
66
Trends in Ionic Size
Cation is smaller than its atom.
(less e with same # protons)
Na
160 pm
-1e
+
Na
95 pm
Al
124 pm
-3e
+3
Al
50 pm
67
Trends in Ionic Size
Anion is larger than its atom.
(more e with same # protons)
Cl +1e100 pm
Cl
181 pm
F
60 pm
+1e
F
136 pm
68
(model)
Ionic Radii
cations
anions
69
Ionic Radii
Place in order of increasing size.
Fe,
2+
Fe
and
3+
Fe
70
Try It !!!
e
1. Use configuration to
predict the charge of Ca ion.
2. Is this ion larger or smaller
than its atom?
71
Electronegativity
 The tendency of an atom to
attract bonding electrons.
Water: which atom
“wins the battle” for
the bonding e ?
O
H
H
72
Electronegativity
An arbitrary scale from 0 to 4.
0
Least EN
Fr (0.7)
Low attraction
for e in bond
4
Most EN
F (4.0)
High attraction
for e in bond
73
Electronegativity
Why don’t the Noble gases
have electronegativity values?
74
Electronegativity
slightly
Example: Water
d-
3.4
O
H
O
H
H
d+
H
d+
2.1
2.1
Water is a “polar” molecule.
75
Electronegativity
Group Trend: EN decreases going
down a group. Atoms get larger, so
bonding e are farther from the nucleus.
Period Trend: EN increases going
across a period. Atoms get smaller, so
bonding e are closer to the nucleus.
(Same trend as ionization energy.)
76
Metallic Character
Metals lose e to become cations.
Which element is the most metallic?
(smallest ionization energy)
Nonmetals gain e- to become anions.
Which element is the least metallic?
(largest ionization energy)
77
“Diagonal
Relationships”
Largest R
Smallest I1
Smallest EN
Most metallic
Smallest R
Largest I1
Largest EN
Least metallic
78
79
Warm-up
What did Rutherford’s gold foil experiment
show about the structure of the atom?
How did Bohr’s model of the atom differ
from the prior model of the atom?
80
Warm-up
What was Bohr’s explanation
for the emission or bright-line
spectrum of hydrogen?
+
81
Warm-up
What is the name of the region
outside the nucleus where electrons
are most probably found?
82
Warm-up
What is the name of the region outside
the nucleus where electrons are most
probably found?
Write the Regents electron
configuration for arsenic.
What does each of the numbers mean?
How many valence electrons does
arsenic have?
83
Warm-up
•What is 2-8-7?
•What is 2-7-8?
•What is the e configuration of gold?
•How many valence electrons does
manganese have?
•What is the electron configuration
of the nitride ion?
84
Warm-up
What are the names of
Groups: 1, 2, 17, and 18?
How many valence e in Co?
What is the trend size:
-down a group?
Why?
-across a row?
What is e config. of Al+3?
85
Warm-up
Which element, P or S, is bigger
(larger radius)? Explain.
Define first ionization energy, I1.
What is the trend in I1 across a row and
down a group? Explain.
Place the following elements in order
of increasing I1: P, Cl, As
86
Warm-up
What is “metallic character”?
How is metallic character related to
ionization energy?
What happens to metallic character
going down Group 15?
Which has greater metallic
character: Fe or Na?
87
Warm-up
Define each term, state the trend, and
explain why:
•Atomic radius across a row
•Ionization Energy down a group
•Electronegativity across a row
•Metallic character down a group
88
Element Song
http://www.privatehand.com/flash
/elements.html
89