electrons in atoms & periodic table - maine-endwell … test (lab) adding energy can cause e-to...
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1
Electrons in Atoms & Periodic Table
2
Quantum Theory or Wave Theory:
a description of the
electron “configuration.”
Where Are the Electrons?
3
Quantum Theory
One of the greatest
achievements of mankind.
4
Arrangement of e- in Atoms
chemical reactivity
bonding between atoms
Periodic Table
many physical properties
Determines:
5
Atomic Models: History
Each atomic model was
eventually replaced because of
new experimental evidence.
1 2 3
6
Dalton: 1803
Concept of the atom
as smallest unit of an
element.
Indivisible particle
7
Thomson: 1897
Atom has parts!!
Discovered the e-
“Plum pudding” model
electron
positive charge
8
Rutherford: 1911
Dense nucleus with
positive charge
Au foil experiment
Nuclear model+
Most of atom is empty space
9
Nuclear Model: Problem
What keeps the electrons
and nucleus apart?
+
10
Bohr: 1913
e- held in “orbits”
Motion of e- keeps them from
“falling” into nucleus
Similar to planets around sun
11
Bohr: “Planetary” Model
e- move in circular
orbits around nucleus,
and each orbit has a
certain energy.
12
Bohr: “Planetary” Model
+
E1
E2
E3
“Quantized”
energy levels
13
“Bright Line Spectrum” of Hydrogen
Stair Analogy: H spectrum due to e- transition between orbits.
14
E1
E2
E3
E4
E5
ener
gy
Stairs are quantized.
Not a ramp
e- in Ground State
15
E1
E2
E3
E4
E5
ener
gy
Ground state is lowest
energy of the e-.
e- in Excited State
16
E1
E2
E3
E4
E5
ener
gy
e- absorbs energy to move
to a higher energy level.
e- in Excited State
17
E1
E2
E3
E4
E5
ener
gy
e- Returning to Ground
18
E1
E2
E3
E4
E5
ener
gy
photon
Elight=Eexcited-Eground
e- gives off energy as light
19
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.
e- Returning to Ground
20
E1
E2
E3
E4
E5
ener
gy
lower energy
photon
21
Bohr: Hydrogen Emission Spectrum
+
E3
E2
E1
e- absorbs energy
(heat, elec.)
e- falls to lower E
and gives off
energy as light
Elight=E3-E1
22
Bohr Theory: Failings
• Why do e- only have
certain orbit energies?
• Only explains the
hydrogen atom exactly.
23
Quantum Mechanics
(Wave Theory)
1926: E. Schrodinger
Can not determine exact location
of an electron! Wow!
Currently accepted theory
24
Quantum Mechanics
e- in “atomic orbitals”
Can only determine
the probability of
locating an electron.
electron cloud
25
Models
Dalton Thomson
+
Rutherford
Bohr
++
Quantum
26
Atomic Orbital
Each atomic orbital
can hold 2 e-- -
A region in space around the
nucleus with high probability
of finding an electron.
Analogy: student in a desk
Wave Model
27
Each e- is arranged in an atom
according to its energy.
+1st energy level
(lowest shell)
2nd energy level
3rd energy level
(higher shell)
Overview
28
https://www.youtube.com/watch?feature=player
_embedded&v=8ROHpZ0A70I#t=4
Bohr Quantum
29
Regents e- Notation
Regents Periodic Table gives
the number of electrons in each
energy level or shell.
1st shell - 2nd shell - 3rd shell…
[C] = 2 - 4
30
Regents e- “Configuration”
What is the e- 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?
Regents e- “Configuration”
31
What happens to the electrons in
each shell going from
Ca to Sc? Zn to Ga?
32
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
33
1 1
2 2
3 3
4 3 4
5 4 5
6 5 6
7 6 7
4
5
Periodic Table by Shell
Inner transition elements
Transition elements
Valence Electrons
34
How many “valence electrons” in:
Li Fe Cu
Electrons that are in the highest energy
level are called “valence electrons.”
These are the most important electrons
when atoms bond. Why?
Valence Electrons
35
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
Excited State
36
(e- have absorbed energy to
move to a higher energy level)
Remember “excited state”?
What element is 2-8-2-1?
What element is 2-7-6-1?
37
Flame Test (Lab)
Adding energy can cause e-
to “jump” from ground
state (as written in Regents
table) to “excited state”.
When e- fall back, they
emit light.
38
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
39
Flame Test for Copper
Which photon has greater energy:
An e- the falls from E5 to E3 or
An e- the falls from E5 to E2 ?
40
Emission Spectrum
Flame test
Neon signs
Fireworks
Fireplace colors
41
Bonding
•Electron configurations are
the key to bonding.
•Atoms become ions to achieve
Noble gas electron configuration.
42
F atom: [F] = 2-7
= [Ne]
Na atom: [Na] = 2-8-1
= [Ne]
Atoms vs. Ions
F- ion: [F-] = 2-8
Na+ ion: [Na+] = 2-8
What does F need to have the e-
configuration of a Noble gas?
43
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.
44
Periodic Relationships
45Early chemists describe the first element.
46
Tabulation of Elements
•Tabulated by chem. &
physical properties
•Arranged by mass
•Predicted missing elements
and properties
Mendeleev (1869)
47
Modern Periodic Table
Now ordered by atomic number,
not mass.
Element 101 (Md)
Argon vs. potassium problem.
48
Periodic Table
Key to understanding chemical
and physical properties
Each group has same electron
configuration for outer shell.
Most important tool in
chemistry
49
Regents Periodic Table
Symbol
Atomic number & atomic mass
Electron configuration
“Charges”
Elements arranged by
atomic no. (#p+)
50
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)
Representative Elements
51
Representative Elements
Group 1: alkali metals
Group 2: alkaline earth metals
Group 17: halogens
Some groups have special names
52
Noble Gases
Last element in each Period
8 VE, except He
very stable
(non-reactive)
53
Transition Elements
e.g. Iron
Regents: 2-8-14-2
Compounds with these
elements have colored solutions.
e- filling
54
Inner Transition Elements
At bottom of
Periodic Table
for convenience.
55
Trends in Atomic Size
Atomic size is
measured by radius. Table ‘S’
For chlorine:
radius = 100. pm
What is its radius
in meters?R
56
Atomic Radius: Trends?????????
OK
(model)
57
Atomic Radius
Down a Group: size
increases due to adding
electrons to higher energy
levels (shells) further from
the nucleus.
58
Atomic Size: Across a Period
Electrons added to same shell
Nuclear charge increases (more p+)
Greater inward pull on the electrons
Atoms get smaller
2-8-1 2-8-3 2-8-5 2-8-8
p+ =11 12 13 14 15 16 17 18
59
Atomic Size: Across a Period
+5 +6
Boron (2-3) vs. Carbon (2-4)
smaller
60
Atomic Radius
smallerRow: greater
nuclear charge
larg
er Column: e- in
higher shell
61
Atomic Radius
Try It:
Arrange these atoms in order of
increasing size.
N, O, P, S
O < N < S < P
62
Ionization Energy (I)
Chemical properties determined
by valence electrons.First ionization energy (I): energy
(kJ/mol) to remove an e-
from an atom.
If I high, e- held tightly.
63
I + X(g) X+(g) + e-
I is endothermic (need to put
energy in to pull off an e-)
Ionization Energy
ionization
energy
64Atomic Number
I1
Atomic Number
I1
Ionization Energy: Table ‘S’
I1 across Period
I1 down Group
65
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?
66
I Predicts Ionic Charges
Element I1(kJ/mol)
I2(kJ/mol)
I3(kJ/mol)
Na 496 4565 6912
Mg 738 1450 7732
Na atom 2-8-1 lose 1 e-
Na+ ion 2-8
Mg atom 2-8-2 lose 2 e-
Mg+2 ion 2-8
67
Ionization Energy
Which has smaller I and why?
O or S
Ge or Br
68
Trends in Ionic Size
Cation is smaller than its atom.
(less e- with same # protons)
Na-1e- Na+
160 pm 95 pm Al-3e- Al+3
124 pm 50 pm
69
Trends in Ionic Size
Anion is larger than its atom.
(more e- with same # protons)
Cl-+1e-Cl
100 pm 181 pmF-
+1e-F
60 pm 136 pm
70
Ionic Radii
cations anions
(model)
71
Ionic Radii
Place in order of increasing size.
Fe, Fe2+ and Fe3+
72
Try It !!!
1. Use e- configuration to
predict the charge of Ca ion.
2. Is this ion larger or smaller
than its atom?
73
Electronegativity
The tendency of an atom to
attract bonding electrons.
HH
OWater: which atom
“wins the battle” for
the bonding e-?
74
Electronegativity
Low attraction High attraction
for e- in bond for e- in bond
Least EN Most EN
An arbitrary scale from 0 to 4.
0 4
Fr (0.7) F (4.0)
75
Electronegativity
Why don’t the Noble gases
have electronegativity values?
76
Electronegativity
Example: Water
HH
O
2.1 2.1
3.4
HH
O
d+ d+
d-
Water is a “polar” molecule.
slightly
77
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.)
78
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)
79
“Diagonal
Relationships”
Smallest R
Largest I1Largest EN
Least metallic
Largest R
Smallest I1Smallest EN
Most metallic
80
Warm-up
81
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?
Warm-up
82
What was Bohr’s explanation
for the emission or bright-line
spectrum of hydrogen?
+
Warm-up
83
What is the name of the region
outside the nucleus where electrons
are most probably found?
Warm-up
84
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?
Warm-up
85
•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?
Warm-up
86
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?
-across a row?
What is e- config. of Al+3?
Why?
Warm-up
87
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
Which element, P or S, is bigger
(larger radius)? Explain.
Warm-up
88
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?
Warm-up
89
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
Element Song
90
http://www.privatehand.com/flash
/elements.html
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