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Chapter 4
Atomic Structure
History: Early 1800’s
Lots of introductory work with electricity
Atoms are electrical in nature
Have parts with + charge
Have parts with – charge
Technology was limited until late 1800’s,
early 1900’s
Cathode Ray Tube
Tube containing two
metal electrodes/plates
Under VERY low
pressure (vacuum)
Glows / green beam
(cathode ray)
So what are
Cathode Rays???
John Thompson (1897)
Able to deflect a cathode
ray using an electric field
Beam attracted to + plate,
deflected from – plate
THUS- cathode ray is
composed of – charged
particles
DISCOVERY OF
ELECTRONS!!!
John Thompson, cont’d
DISCOVERED ELECTRONS
Also used magnetic fields to deflect cathode
ray
Allowed him to calculate ratio of electron mass
to charge
1906 Nobel Prize
Canal Rays
Eugen Goldstein (1886)
Used a discharge tube with holes in cathode
+ ions shot in opposite direction of cathode rays
(electrons) == CANAL RAYS
Some went through the holes
1907- repeated in presence of magnetic field--found the + particles had varying massess
Millikan Oil-Drop Experiment
Robert A. Millikan-(1909)
Discovered charge of an
electron
Placed electrically charged
oil drops in an electric
field
Oil drops sprayed in upper
chamber
Some gain – charge (static
electricity)
Calculate mass by
measuring how fast oil
drops fall though hole
(gravity)
Millikan Oil-Drop Experiment,
cont’d
Then Millikan adjusted
voltage on plates in
bottom chamber
- particles attracted to upper
plate to offset pull of
gravity---suspended drop in
space
Calculate mass: charge ratio
Know mass from first exps,
so can calculate charge!
Smallest change between
two drops taken to be
charge of an electron
1923 Nobel Prize
Millikan Oil-Drop Experiment,
cont’d
Mass of an electron:
9.1X10-28 g
Charge of an electron:
-1
Radioactivity
1895- Antione Bequerel
Placed a sample of uranium on covered photographic
paper (NO light exposure)
Fogged the photographic paper!!
Student = Marie (Sklodowska) Curie called it
“radioactivity”
1903- Bequerel, Pierre Curie, Marie Curie shared Nobel
Prize in Physics
1911- Marie Curie won Nobel Prize in chemistry for
additional work on radioactivity
Radioactivity, cont’d
3 types of radioactivity (Ernest Rutherford)
Alpha (α)
Beta (β)
Gamma (γ)
Radioactivity, cont’d
Alpha Particles:
Deflected magnetic
field like + particles
Mass = 4 X that of
hydrogen atom
Charge: 2 X that of an
electron, but + charged
Radioactivity, cont’d
Beta particles:
Electrons / cathode rays
Charge = -1
Mass = 1/1837 of a
hydrogen atom
Gamma Particles:
NOT deflected in a
magnetic fields
Very high energy
Radioactivity, cont’d
particle
alpha
Mass (u)
4
charge
Penetrating
ability
+2
Low (paper
shield)
beta
1/1837
-1
Intermediate
(plastic shield)
gamma
0
0
High (lead
shield)
Gold Foil Experiment
Ernest Rutherford
Bombard gold foil
with alpha particles
Gold Foil Experiment
Many α particles
passed through gold
foil as expected, but
SOME DEFLECTED
Concluded that +
charge and most of
mass of an atom is
contained in a tiny
core (NUCLEUS)
Models of Atomic Structure
By early 1910’s, scientists knew about
electrons and protons, and that the majority
of the mass (an + charge) was in the
“nucleus”
Began to develop models of atomic
structure
Rutherford’s “Plum Pudding
Model”
Atoms are like plum pudding
Lots of empty space (pudding), but a central
core (plum) every so often, where most of the
mass is located
Nucleus: concentrated, tiny volume where vast
majority of the mass is located, + charges there
Electrons: surround nucleus, but occupy almost
the entire volume of an atom
Rutherford’s “Plum Pudding
Model”or “Football Analogy”
Football Stadium
Subatomic Particles
3 main ones (there are others, but that’s for
physics..)
Neutrons
Protons
Electrons
Electrons
JJ Thompson, RA Millikan
1st discovered late 1800’s, early 1900’s
“cathode rays”
Negatively (-) charged particles
Mass = 9.109389 X 10-28 g
Charge = “-1” = - 1.60217733X10-19 coulombs
Occupy majority of space in an atom
Protons
Eugen Goldstein (1886)
“canal rays”
Mass = 1.672623X10-24 g
1800 X as heavy as an electron!
Charge = “+1” = + 1.60217733X10-19
coulombs
Located in tiny, dense nucleus
Neutrons
James Chadwick (1932)
Provide mass, but no charge
Mass = 1.6749286 X 10-24 g
(nearly the same as a proton)
Located in the nucleus
Nuclear Symbols
X = element symbol
Z = mass #
= # protons + #
neutrons
A = atomic # = # protons
(gives an element its
identity!!)
So… # neutrons = Z-A
Z
X
A
charge
Atomic #
# protons
Gives an element it’s identity
If 6 P…….C
7 P…….N
1 P…….H
53 P……I
NOTE: # above element symbol in periodic
table
= # protons = Atomic #
Electrons and atom charge
If # P = # electrons
(e-)
Then atom is NOT
CHARGED
If #P > #e Then + charged
(cation)
Charge = #P - #e-
If # P < # e Then – charged
(anion)
ISOTOPES
Atoms with same #
Protons, but different
# neutrons
Example– 1H, 2H, 3H
All are H (have 1
proton)
Z is different!!!
Isotope
#P
1H
1
1 = (1+0)
2H
1
2 = (1+1)
3H
1
3 = (1 +2)
#N
0
1
2
How many P, N, e- ??
197 Au
79
# P = 79
# N = (197-79)
= 118
# e- = 79
23 Na+1
11
# P = 11
# N = (23-11)
= 12
#e- = 10
+1 = #p - #e+1 = 11- # e-
Write the nuclear symbol for…
An atom containing 18 e-,
16 P, and 17 N.
What element is it??
Atomic # = 16 P = Sulfur
(S)
Z = #N + # P = 16+17=33
Charge = #p -#e= 16 – (18) = -2
33
S
16
2-
Which of the following represent
isotopes of the same element?
16 X
8
16 X
7
14 X
7
14 X
6
12 X
6
Look for same atomic #: (same # protons)
16 X and 14 X both have 7 protons, so are isotopes
7
7
of Nitrogen.
14 X and 12 X both have 6 protons, so are isotopes
6
6
of C.
The Periodic Table
LOTS on INFORMATION!!!
Columns = Groups = Families = up and down,
# 1-18
Main groups : 1, 2, 13-18
Transition metals: 3-10
Inner Transition metals: at bottom of chart
Members of the same group have similar physical
and chemical properties
Rows = “Periods”
Main groups
Transition
metals
Inner transition metals
Metals
Gps 1,2,3-10, some in
14 and 15 (blue)
Maleable, ductile,
conductive,
Usually form cations
Metaloids
Green in picture
Elements on “step”
between metals and
non-metals
Intermediate
properties
(semi-conductors)
Non-metals
Rest of chart (yellow)
Nonconductive
usually form anions
Gp 1- alkalai metals
Very soft
not found in nature in
uncharged atom form
(find as cation)
VERY reactive with
nonmetals and water
Gp 2- alkali earth metals
Harder, more dense vs
gp 1
Less reactive than gp 1
All react with oxygen
to form metal oxides
Gp 17- Halogens
“salt-producing”
React with metals to
form salts
Exist as diatomic
molecules in nature
(F2, Cl2, etc.)
Gp 18- Noble Gases
Colorless gases
Exist as single atoms
Very Unreactive
Some limited
reactivity of Xe, Kr,
Rn
What is it about elements that
explain the reactivity and other
properties?
-arranged by atomic # (# protons), for
UNCHARGED elements (#p = #e-)
So as #P increases, #e- increases
Thus– reactivity based on
ARRANGEMENT OF ELECTRONS
around the nucleus
Electron Organization
Look at how electrons are positioned
around the nucleus
Not in any one place, but fill different
regions of space around the nucleus
Continuous Line Spectra
When white light
(sun) is separated
through a prism, see
that colors blend into
one another
Spectrum is
“continuous”
Each color represents
a different wavelength
of light
Emission Line Spectra
Take pure element in
gas phase and put in
cathode ray tube,
apply high voltage
Atoms “excited” as
they absorb energy
Atoms emit energy as
LIGHT
Emission Line Spectra
When pass light
through prism, get
bands of color at
specific wavelengths–
NOT continuous like
white light!
Each element has a
“signature” ELS
Bohr Model of Atomic Structure
(1913)
Based on Emission line spectra of
Hydrogen
Electrons can have only a specific amount
of energy = QUANTA
Energy values = energy levels
Quanta = difference between any two
energy levels
Bohr Model- Cont’d
6 to 2, 5 to 2, 4 to 2 and 3 to 2
transitions emit energy of visible
wavelengths
Analogy- you can stand on the
rungs of a ladder, but not in
between
Bohr Model- “Planetary” Model
Energy levels are in rings/
“orbitals” around the
nucleus
Each orbital (ring) is
assigned an integer value
“n” = principle quantum #
Radius of orbit increases as
n increases
n=1 is the lowest possible
level, where e- in H usually
resides
Bohr Model- “Planetary” Model
Ground state = lowest energy level an e- occupies
in the normal state
Excited state = when an e- is in an energy level
greater than the ground state
Takes energy to move e- from ground to excited state
(move away from + pull of nucleus)
Excited state is unstable, e- wants to go back to normal
state
e- releases energy (light) as if falls back to the ground state
This is the light in the Emission Line Spectra
Bohr Model- “Planetary” Model
Bohr won 1922 Nobel Prize for model
For atoms with more than 1 electron, the
Bohr model suggests that the energy levels
can hold only a certain # of e #e- = 2n2 (n= orbital #)
N=
1
2
3
#e- =
1
8
18
4
32
5
50
Bohr Model- “Planetary” Model
BUT…the outermost energy level can only
hold 8 e-….
1st level now filled
1st
row
on PT
2nd
row on
PT
3rd
row
on PT
2n2 = 2(1)2 = 2
2nd level now
filled
2(2)2 = 8e3rd level
maxed with
8e-; but can
hold 18
Bohr Model- “Planetary” Model
After this, things get more complicated…
Bohr’s model begins to fail….
Erwin Schrodinger- 1920’s
Treat electrons as both waves and particles (Bohr
treated only as particles..)
Devised UGLY mathematical equations that state
the PROBABILITY of finding an electron in a
given region around the nucleus
Cannot tell exact location , just the likelihood of
where an electron will be located!
“Quantum Mechanics”
Schrodinger model subdivides Bohrs energy
levels
Orbitals = regions of space
(replaced Bohr’s planetary orbits)
Camera analogy- if you could look at a timelaspe photo of an electron, the area exposed is
the orbital
“Quantum Mechanics”
Each orbital can hold 2 electrons
Some energy levels hold more than 1 orbital
# orbitals = n (row #, principle quantum #)
n
# orbitals
“names”
1
1
s
2
2
s, p
3
3
s, p, d
4
4
s, p, d, f
“Quantum
Mechanics”
Each orbital has a distinct
shape
S = sphere
P=dumb bell
D, f = very
complicated
“Quantum Mechanics”
Inverted Triangle Apartment
Building
Floor = n
Apt = orbital (s, p, d, f)
s = 1 bedroom = 2 ep = 3 bedroom = 6 ed = 5 bedroom = 10 ef = 7 bedroom = 14 e(2 e- per room)
Fits Periodic Table!
Order of Filling
14 e-
Follow atomic # on
Periodic Table
10 e-
6e2 e-
Order of Filling
Practice: Give the electron
configurations for each of the
following elements
Na
O
Mn
Still have only 8e- max in outer most shell??
Don’t forget the f-block!
Ba
La
Ce
Lu
Hf
Hg
Identify the following elements
given the electron configuration:
1s2 2s2 2p6 3s2 3p6 4s2 3d7
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5
1s2 2s2 2p6 3s2 3p4
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2
4f14 5d3
Octet Rule
No more than 8 e- in the outermost shell!!
outershell e- = valence electrons
Outershell = greatest n!!!
General Rule: Atoms combine with other
atoms (form bonds) in order to have a
COMPLETE octet (8 outershell electrons)
Share e- (covalent bond)
Gain/lose electrons (forming IONS)
Want to have electron configuration of
Nobel gases! (fill a p6 sublevel!)
IONS
Atom that has either a + or a – charge,
obtained by either by
gaining an e- (extra – vs +)
so – charge…anion
Lose an e- (less – vs +)
So + charge….cation
Anions
Gained an e-….now more e- than
protons(+)
Net – charge
Ex– Oxygen
Atom : 1s2 2s2 2p4
Needs 2 e- to fill up p level (be 2p6), like Neon
So will GAIN 2 e-
Ion: O 2- : 1s2 2s2 2p6
(put charge in upper right)
Anions, cont’d
Chlorine:
atom: 1s2 2s2 2p6 3s2 3p5
needs 1 e- to get configuration like Argon
(..3p6) (isoelectric)
Picks up 1 e-
Cl - : 1s2 2s2 2p6 3s2 3p6
Cations
Lose an e-….now fewer e- than protons(+)
Net + charge
Ex– Sodium
Atom : 1s2 2s2 2p2 3s1
One electron, all alone on 3rd floor…jumps off
So will LOSE 1 e-
Ion: Na + : 1s2 2s2 2p6
(put charge in upper right)
Cations, cont’d
Calcium:
Atom: 1s2 2s2 2p6 3s2 3p6 4s2
Lose the two outer shell (n=4) electrons
Ion: Ca 2+ : 1s2 2s2 2p6 3s2 3p6
(like Argon)
Ions in Summary
Gp 1: (ns1) : lose 1 eGp 2: (ns2) : lose 2 eGp 16: (ns2 np4) gain 2 eGp 17: (ns2 np5) gain 1 e-
+1 cations
+2 cations
-2 anions
-1 anion
Periodic Trends
Certain trends based on electron
configurations / organizations
Atom size
Parent atom vs Ion size
Reactivity
Electronegativity
Atom Size
Atom Size
What is happening?
1. More orbitals (energy levels, apt floors), the
larger the atom
So…INCREASE in size DOWN a group
Li < Na < K < Rb < Cs < Fr
n=1 < 2 < 3 <4 < 5 <6
Atom Size
Atom Size
2. As e- fill in the same energy level /floor
(across from left to right)
get more protons (+) in the nucleus while efilling same level
So more + to attract e- in towards nucleus
Trend: DECREASE atom size ACROSS (L
to R) a row
Li > Be > B > C > N > O > F > Ne
Atom Size
Example
Arrange the atoms by increasing size
(smallest to largest):
Mg Al Na S Ar
Find location in PT: note all in n = 3
so Left = largest and right = smallest
Ar < S < Al < Mg < Na
Ion size vs. Atom size
What happens to radius when an atom gains
an e-?
When it loses an e-?
Ion size vs Parent Atom
Anions:
add e-: now # e- > # P (+)
each electron now has less than a full +1 pull
from the nucleus
e- not held as tightly
Radius increases / larger in anion vs parent
atom
Ion size vs Parent Atom
Cations:
lose e-: now # e- < # P (+)
each electron now has more than a full +1 pull
from the nucleus
e- held more tightly
Radius decreases / smaller in cation vs parent
atom
Trends in Reactivity
Metals:
the larger the atom, the farther away the
valence e- is from the + nucleus
Lose the electron more easily (more reactive)
as size increases
Ex: Li vs Na vs K : which is more reactive
with water?
K is largest, so more reactive
Trends in Reactivity
Non-metals:
- want to gain e- (Opposite trend of metals)
- Closer the new e- will be to the nucleus, the
greater the attraction or pull to get that
electron to the new atom
- So Reactivity of non-metals increases as
atomic radius decreases
Electronegativity
“desire” to gain an electron
F is most EN atom
All increases to F
Fr is Least EN atom
Why is F the most EN atom?
F REALLY wants an additional electron to
fill its p level
The F p-level (n=2) is the closest p-level to
the nucleus, so there is a VERY strong,
unshielded nuclear pull
So an F atom will pull as many electrons to it
as it can (more on this in the next
chapter…)
Chapter Objectives
1.
Know how electricity played an
important role in unraveling the
structure of the atom. Define: cathode
ray tube, cathode rays, canal rays. Be
able to describe the role cathode ray
tubes played in the discovery of
electrons and protons.
2. Be familiar with the key experiments
performed by Thomson (properties of
cathode rays), Goldstein (canal rays),
Milliken (charge and mass of electrons), and
Chadwick (neutrons). Be able to describe
Thomson’s “plum pudding model” of the
atom.
3.
Know the brief history of radioactivity
and how work with radioactive elements led
to an understanding of the existence of
subatomic particles. Be familiar with the
accidental discovery of radioactivity by
Becquerel and the contributions of the
Curies.
4.
Know three types of radioactivity:
alpha, beta, and gamma emissions as
defined by Rutherford. Be able to distinguish
them by mass and charge.
5.
Discuss Rutherford’s gold foil
experiment, which led to the discovery of the
atom’s nucleus. Contrast Rutherford’s
atomic model with Thomson’s model and
current models.
6.
Be able to define and understand the
following terms: atomic number, mass
number, and isotope. Know how to find the
atomic number on the periodic table and
know its significance in identifying an
element. Understand that the mass number
of an isotope can not be found on the
periodic table.
7.
Be able to use the atomic number and
mass number of any element to determine
the number of electrons, protons, and
neutrons in a neutral atom. Also, given the
number of protons and neutrons, be able to
determine atomic number and mass number
and be able to draw the notation for an
isotope.
8 8.
Know the difference between
continuous spectra and line spectra. Be
able to explain how the line spectra of
hydrogen can be explained by the Bohr
model of the atom. Define the terms ground
state and excited state of an electron as
defined by the Bohr model.
9.
Be able to explain the Schroedinger
model of the atom based on probabilities of
electron location and pathway. Define the
term electron orbital and know the shape of
the s and p orbitals. Know how electron
configurations for neutral atoms are derived
according to the Schroedinger model using s,
p, d, and f orbitals.
10. Be familiar with the organization of the
periodic table. Know the difference between
a group (column) and a period (row). Know
and locate the Group A elements. Be able
to derive electron configurations for main
group elements. Given an electron
configuration, be able to locate the
element on the periodic table.
11. Using the periodic table be able to point
out the alkali metals, alkaline earth metals,
halogens, and noble gases. Be able to find
transition metals and inner transition
metals.
12. Know area on the periodic table that
correspond to the filling of s, p, d, and f
orbitals. Be able to use the periodic table to
do electron configurations.
13. Define valence electrons. Be able to
relate group A numbers to number of
valence electrons as derived from electron
configurations. Know the significance of the
number of valence electron s of group VIIIA
elements (ie, noble gases) and how this
might relate to chemical reactivity of other
main group elements.
14. Be able to distinguish general physical
properties of metals and nonmetals; eg,
malleability, ductility, ability to conduct
heat and current. Be familiar with general
chemical properties of metals and
nonmetals; eg, metals tend to give up
electrons to form positive ions when
reacting with nonmetals, nonmetals tend
to gain electrons when reacting with
metals. Understand why this might occur.
15. Predict trends in sizes of atoms and
trends in chemical reactivity going down
groups and going across periods of the
periodic table. Be able to understand reasons
for these trends.