Download Early Atomic Theory and Structure Empedocle (440 BC): all matter

Chapter 5
Early Atomic Theory
and Structure
Empedocle (440 BC): all matter consists
of 4 “elements” earth, fire, water and air.
Democritus (470-370 BC): matter is composed of
indivisible particles (Greek atomos - not cuttable).
Aristotle (384-322 BC): endorsed Empedoclean
theory, so that it dominated until 17th century.
John Dalton (1766-1844), 2000 years after
Democritus, revived concept of atoms:
Dalton’s rules:
1. elements are composed of indivisible particles-atoms.
2. atoms of the same element are alike in mass and size.
3. atoms of different elements are not alike – they have
different masses and sizes.
4. compounds are union of two or more atoms of
different elements.
5. compounds contain simple numerical ratios of atoms.
6. more than one compound can be formed of atoms of
two elements.
Dalton rules (cont.)
Consequence of simple numerical ratio
of atoms in compound:
Law of definite composition – a
compound contains two or more elements
in a definite proportion by mass.
For instance, 2g of hydrogen and 16 g of
oxygen combine to form 18g of water.
Law of multiple proportions – atoms of two
or more elements may combine in different
ratios to produce more than one compound.
H2O and H2O2 are composed of two types
of atoms, but different relative number.
compound
water
hydrogen peroxide
Copper (I) chloride
Copper(II) chloride
Methyl alcohol
Ethyl alcohol
formula
H2O
H2O2
CuCl
CuCl2
CH4O
C2H6O
percent composition
11.2 % H, 88.8% O
5.9% H, 94.1% O
64.2% Cu, 35.8% Cl
47.3% Cu, 52.7% Cl
37.5% C, 12.6% H, 49.9% O
52.1%C, 13.1% H, 34.7% O
All atoms of one
type have the
same mass, so
they can be
measured either
by counting or by
weighing.
Water
Hydrogen
peroxide
While most of the
Dalton’s theory
still holds, some
modifications
were necessary:
1. An atom is composed of subatomic particles,
2. Most elements contain atoms have isotopes - that have different masses.
Electric Charge
Matter may have electric charge
1. charge may be of two types, positive and negative
2. unlike charges attract, like repel
3. can be transferred from one object to the other by contact or induction
4. the smaller the distance, the stronger the force between two charges
x q
q
x
1
2
(F – force, k - constant, q1, q2 - charges, r - distance)
F=k
2
r
G.J. Stoney (1826-1911) named the unit of electricity electron.
Joseph Thompson (1856-1940) experimentally showed the existence of
electron, which is: negative in charge, deflected by magnetic and electric fields,
and capable of moving small paddle wheel.
Michael Faraday (1791-1867)
1. Some compounds dissolved in water conduct electricity, others decompose
2. Particles of some elements are attracted to the positive or negative electrode.
3. Concluded that they are charged and called them ions (Greek-wanderer).
Svante Arhenius (1859-1927)
1. Melted NaCl also conduct electricity. NaCl breaks up to Na+ and Cl- ions.
2. Ions move to oppositely charged electrode. Ion that goes towards positively
charged cathode is called cation. The other goes to positive anode is called anion.
Charge and Periodic Table
IA
H+
Li+
Na+
K+
Rb+
Cs+
IIA
Invented by Dmitri Mendeleev, the
Periodic Table shows recurring trends
in properties of elements and their
charges.
VIIIA
IIIA
Be2+
Mg2+
Ca2+
IIIB IVB
VB
Cr2+
Cr3+
Sr2+
VIB
VIIB
Fe2+
Fe3+
------ VIIIB -------
Al3+
IB
IIB
Cu2+
Cu+
Zn2+
Ag+
Cd2+
IVA
VA
VIA
N3-
O2-
P3-
S2-
VIIA
F-
ClBr-
I-
Ba2+
Group IA: Group IIA:
always 1+ always 2+
Charge = group number
Elements of B groups
(transition metals) form
ions of different charges
Group VA: Group VIA: Group VIIA:
often 3usually 2- always 1Charge = 8 - group number
Metals form positive charges, nonmetals negative charges. Group VIIIA:
no charge: noble gases
Subatomic particles
Atom is extremely small, cannot be observed by optical microscope.
Atomic diameters range from 0.1 to 0.5 nm.
Consider this dot on the right.
If its diameter is 1 mm, 10 million H atoms could be arranged in it.
Yet, there are particles smaller than H atom. They are called subatomic particles.
Three subatomic particles are sufficiently long-lived to be used in experiments:
electron, proton and neutron.
Proton and electron have charges opposite in sign, identical in absolute value.
Proton and neutron have almost equal masses.
Number of other short-lived subatomic particles are detected in recent years.
quarks, neutrinos, antineutrinos, positrons and antiprotons ...
Their existence is still debated.
The Evolution of Atomic Theory
Becquerel in 1896 discovered radioactivity. Radioactive elements spontaneously
emit alpha particles, beta particles and gamma rays from their nuclei.
By 1907 Rutherford found that alpha particles emitted by
certain radioactive elements were helium nuclei.
J.J. Thomson discovered electron in 1897. Shortly after,
he proposed the first atomic ‘Plum-Pudding’ model where
negatively-charged electrons (‘plums’) are embedded into
a positively-charged sphere (‘pudding’).
+
Thomson and Goldstein discovered proton in 1907,
and Chadwick discovered neutron in 1932.
α particles should pass through the
Rutherford’s
experiment
showed
that the
plumpudding
model is
incorrect.
sphere because it is effectively neutral.
α particle velocity ~ 1.4 x 107 m/s
(~5% speed of light)
The Evolution of Atomic Theory (cont.) Indeed, most of them did, but a few (1 in
20,000) were deflected, some in great
If Thomson’s model was correct, αangles. It was “as if you were firing 15”
particles would go straight through the foil.
shell at a tissue paper, and it came back
and hit you”.
Rutherford’s Conclusions:
1. Atom consists of a nucleus and electrons.
2. Size of an atom is about 100,000 times
bigger than that of its nucleus.
Size of an atom: If the
nucleus is as big as a
marble, the atom is as
big as football field.
Thus, atoms are
mostly empty space!
3. Atoms’ positive
charge is concentrated
in the nucleus.
4. Electron has negative
charge, and nucleus
has positive charge.
Nucleus has both protons and neutrons.
Mass of p is 1840 x mass of
e- (1.67 x 10-24 g).
In order to avoid falling onto the nucleus,
electron must be constantly rotating with
great speed, as if it was everywhere at once.
Problem: classical physics says that the
moving charge (electron) radiates energy.
Creation of Ions
An neutral atom has equal
number of protons and electrons.
Recall that the proton has
equal charge to that of an
electron, but opposite in sign.
When atom loses one or
more electrons, a positive
cation is formed.
When one or more electrons are
added to a neutral atom, an
anion is formed.
The word ‘cation’ comes from the
fact that positive ion (cation)
moves towards negativelycharged (cathode) in an
electrolyte solution. The negative
ion (anion) moves towards
positively-charged electrode
(anode).
Atomic Numbers of the Elements
The atomic number of an element is equal to the number of protons in the
nucleus of that element. It gives the identity of the element.
Atomic number (Z) = number of protons in nucleus
Mass number (A) = number of protons + number of neutrons
= atomic number (Z) + number of neutrons
Isotopes: atoms of the same element with different numbers of neutrons in their
nuclei. Neutrons can be envisioned as glue that holds nuclei together.
Charge: the difference in protons and electrons.
The mass of 12C is 12.000 g.
Mass Number
Atomic Number
A
ZX
There exists 1.11 % 13C and
Charge
traces of 14C. The mass of
Element Symbol C atom (12.01 g) is weighted
average of all atoms.
Two isotopes of uranium
235
92
238
92
U
U
Protons: 92
Neutrons:
235-92=143
Protons: 92
Neutrons:
238-92=146
Three isotopes of hydrogen
Hydrogen
(Protium)
Deuterium
Tritium
(radioactive)
Masses of Elements and Isotopes
To overcome the problem in
measuring extremely small masses, a
system of relative atomic masses
using “atomic mass units” was devised
to express the masses of elements
using simple numbers.
A mass of exactly 12 atomic mass
units (amu) was assigned to 126 C.
Element masses are
average numbers,
because of isotopes.
Mass of H atom is
1.008 amu.
The mass of a single
atom is too small to
measure on a balance.
Using a mass spectrometer, the
mass of the hydrogen atom was
determined = 1.673 x 10-24 g.
The standard to which the masses
of all other atoms are compared to
was chosen to be the most abundant
isotope of carbon.
1 amu is defined
as exactly equal to
1/12 the mass of a
carbon - 12 atom:
1 amu = 1.6606 x 10-24 g
12
6
C
Masses of elements and Isotopes (cont.)
Isotopes of the same element have different masses.
The listed atomic mass of an element is the average relative mass of the isotopes
of that element compared to the mass of carbon-12 (exactly 12.0000…amu).
To calculate the atomic mass, multiply the isotopic mass of each isotope
by its abundance and add the results.
Percent abundances must
If the abundance is given as percent, divide it by 100.
be divided by 100!
Example: calculate average atomic mass of copper.
Isotope
Average
Isotopic
Abundance
atomic
mass (amu)
(%)
mass (amu)
63
29
Cu
62.9298
69.09
65
29
Cu
64.9278
30.91
(62.9298 amu)
x
0.6909 =
(64.9278 amu)
x
0.3091 =
63.55
+
43.48 amu
20.07 amu
63.55 amu
Finding the Number of Protons, Neutrons and Electrons
Element
Chlorine
Gold
Bromine
Symbol
36Cl
Au
Z
Mass# #p
17
79
56
36
197
Br -
35
Fill in the table below.
1. From Periodic table, find: Element,
symbol, atomic number and #protons.
At least one of those must be given, the
other three are read from Periodic Table.
Symbol Element
Z = #p.
2. Find mass number and #n
Mass# = #p + #n
3. Find charge or #electrons
(charge) = #p – #e
#e = #p – (charge)
80
17
79
35
#n
#e
charge
19
118
79
16
17
79
18
0
0
+2
-2
45
36
-1
Charge = 0
Locate Cl in periodic table
Look up Atomic number (Z) for Cl
Mass # is 36 = #p + #n
#n = 36 – #p = 19
4. You can get #e and charge. Use step 3
to find #p, and from that identify element.
See if you can solve the other
two examples. The solutions are:
135Ba2+, and 32S2-.
56
16
Chapter 10
The Modern Atomic Theory and Periodic Table
Rutherford model, based on classical
(Newtonian) physics could not explain
periodicity of the elements, nor how
electrons stay in orbit around nucleus.
A new model was needed.
Light moves with speed of 300,000 km/s.
It is characterized with wavelength λ.
When passed through a prism, a familiar
set of colors are seen. Spectrum is
continuous as colors blend.
Visible spectrum is just a tiny fraction of the electromagnetic radiation.
Bohr’s Theory
Niels Bohr in 1913 postulated that an
electron inside an atom can possess
only certain values of energy. Those
“packs” of energies are called quanta
(pl. of quantum).
Atom is in the lowest energy when its
electrons are closest to the nucleus.
When atom absorbs quantum of energy,
electrons are moved to higher orbits.
Electrons can only have allowed
energies, or the distances from the
nucleus, not those between the allowed
ones. It is like a person on a ladder who
can stand on a 1st, 2nd, 3rd rung, but not
between the rungs.
Allowed orbits are called shells. They are
characterized by the principal quantum
number, n that goes from n = 1 (shell
closest to the nucleus), to infinity.
The maximum number of electrons a
shell can hold increases with n as the
radius of the shell increases.
Number of electrons in a
shell equals to 2n2.
Li and Na have a single electron in
the valence shell.
Periodicity
Elements
have similar
chemical
properties
because their
valence shell
configurations F and Cl have 7 electrons
are similar.
in the valence shell.
The number of electrons in the
valence shell is equal to the
Roman-numeral group number
for the representative elements.
The dependence of chemical on
valence shell occupancy holds for
each group of elements.
Helium is an exception. It behaves as
a noble gas although it has only 2 ein the valence shell.
Group IA has one electron, Group IIA has two…
The lowest reactivity is found in
the group VIIIA (noble gases).
Elements of other groups can
attain the valence shell of group
VIIIA by producing compounds.
Line Spectra
Ground state
s: 2 e- p: 6 ed: 10 e- f: 14 e-
Excited state
New, refined model
includes subshells
that are very close in
energy and size.
Every subshell is
designated with a
letter (s, p, d, f).
s subshell can hold
two electrons. Each
subsequent subshell
holds 4 e- more than
the previous.
Transition n=3 n=2 produces 13.1
eV – 11.2 eV = 1.9 eV, i.e. red color.
Similarly, n=4 n=2 produces 13.8 –
11.3 = 2.6, green color; other colors
in the visible spectrum are also found.
Bohr’s model correctly predicted all
transitions of H atom; transition into
n=1 produces a line in UV, into n=2
(visible) and into n≥3 (infrared). But
the model failed to predict spectrum
for any atom with more than one e-.
n
s
p
d
f
1
2
2
2 6
3
2 6 10
4
2 6 10 14
Each ‘box’ (orbital)
contains 2 e- with
paired spins.
Filling of Subshells
The number of subshells equals
to n. They increase in energy in
a shell: s < p < d < f.
Subshells may cross each other;
the crossing is more pronounced
when spacing between shells
reduces as the principle quantum
number increases.
35Br
This crossing is the
reason for the socalled ‘potassium
problem’.
Present the electron
configuration of 19K.
1s2 2s2 2p6 3s2 3p6 4s1
Valence
electron
Orbital
diagram
is
slightly more
complicated.
1s2 2s2 2p6
3s2 3p6 4s2
3d10 4p5
Rearrange:
1s2 2s2 2p6
3s2 3p6 3d10
4s2 4p5
Valence electrons
Move He next to H to get four blocks, each with different width
(s:2, p:6, d:10, f:14).
Subshell fill up
16S
S
Try 26Fe:
+
16 e
Rule of filling p, d, f subshells:
Each new electron goes in its own orbital,
assuming the same spin as the others in
that subshell. When all orbitals in a
subshell have at least one electron, the
next coming electron pairs with another in
a half-filled orbital.
1s2 2s2 2p6 3s2 3p4
1s2 2s2 2p6 3s2 3p6
4s2 3d6 long e config.
…first to leave are its
or: [Ar] 4s2 3d6
electrons in the s orbital
short e config.
When a transition of the valence shell !
metal forms cation… Valence electrons: e- in unfilled shells + e- in the outermost shell.
Valence electron jumps from Na to Cl
The best way for Mg to
achieve octet is to lose two
electrons from its 3rd shell.
Theoretically, Mg could also gain 6 e-,
but the 6 e- excess is too much for Mg
nucleus to hold on to.
The best way for F to achieve
octet is to gain one e-.
Each e- from Mg 3rd shell jumps
into 2nd shell of an F atom.
Mg atom becomes Mg2+ cation.
Each F atom becomes F- anion.
All three ions have 8 valence e-.
Compound Formation
NaCl
MgF2
Neutral atoms: Val.eNa: 1s2 2s2 2p6 3s1
Cl: 1s2 2s2 2p6 3s2 3p5
Note the size of atoms!
Ions: no electrons in
n = 3 shell of Na.
Na+ has 8e- in n=2 shell.
Cl- has 8e- in n=3 shell.
Both ions satisfy octet rule.
Note the size of ions!
Atomic Size
Understanding of atomic
size with Bohr’s model.
One consequence: shells
get larger as n increases.
A combination of same
valence shell and increased
nuclear charge shrinks the
atom in a group.
Quantum Mechanical Model
Electron has wavelike characteristics, it can tunnel
to appear in ‘forbidden’ places. Uncertainty principle
states that it is impossible to know at any given time
where the electron is, or where it is going. Instead, etravels around nucleus in an electron cloud or orbital.
The QM model also explains why e- stays in
orbit around the nucleus: An e- orbital is allowed
only if the electron wave closes in on itself.
Allowed
orbit must
close in on
itself.
Nonallowed
orbit does not.
Electron cloud exhibits wavelike motion, so that
different orbitals have different shapes. It has a
particular shape and energy determined by the
principle quantum number n.
HW Chapter 5 (p.95). 12. List similarities and
differences in the three isotopes of hydrogen.
15. Write isotopic notation symbols for each of
the following: Z=29, A=65; Z=20, A=45;
Z=36, A=84
23. An element exists in two stable isotopic
forms. One isotope has a mass of 62.9296 amu
(69.17 % abundance). The other isotope has a
mass of 64.9278 amu. Calculate the average
mass of the element and determine its identity.
27. What experimental evidence supports these
statements? (a) The nucleus of an atom is
small; (b) The atom consists of both positive and
negative charges; (c) The nucleus of the atom is
positive.
39. Complete the following table with the
appropriate data for each isotope given Assume
all are neutral atoms.
Element Symbol At.# Mass # #p #n #e
134Xe
silver
107
9
92
143 92
41
19
HW, Chapter 10 (p.208)
3. Write the complete (long) and
abbreviated (short) electronic
configurations for the elements: Sc,
Rb, Br, S.
17. Which elements have these
configurations?
1s2 2s2 2p5.
1s2 2s2 2p6 3s1
[Ne] 3s2 3p4
[Ar]4s2 3d8.
21 (extra credit) For each of the
element, write the corresponding
orbital diagrams: F, S, Co, Kr, Ru.
35. Pick the electron structures that
represent elements in the same
chemical family:
1s2 2s1
1s2 2s2 2p6 3s2 3p6
1s2 2s2 2p4
1s2 2s2 2p6 3s2 3p6 4s2
1s2 2s2 2p2
1s2 2s2 2p6 3s2 3p6 4s1
1s2 2s2 2p6 3s2 3p4
1s2 2s2 2p6 3s2 3p6 4s2 3d1.
41. In which period and group does an
electron first appear in an f orbital?