Atoms
And what they form…
Element Information
• Webelements…
http://www.webelements.com
The Atom
• ~400 BCE Democritus: small
particle (“atomos” is Greek for
indivisible)
• 1803
John Dalton:
atomic model with several
postulates
• 1897-now Modern
Investigations…
Dalton’s Atomic Theory
The Postulates
1. All matter consists of atoms which are indivisible
and indestructible.
2. Atoms of one element cannot be converted into
atoms of another element.
3. Atoms of an element are identical in mass and
other properties and are different from
atoms of any other element.
4. Compounds result from the chemical combination of
a specific ratio of atoms of different elements.
From Silberberg, Principles of Chemistry
Atomic Structure
• ~1870-1935 saw experimentation that
showed that Dalton was right and wrong.
• 1874 – Stoney: electricity is made of
individual particles with negative charge
called electrons
• 1879 – Crookes: discovered “cathode
rays” have distinct properties like
electrons
Figure 2.4 Silberberg, Principles of Chemistry
Experiments to determine the properties of cathode rays.
Atomic Structure
• 1896 – Becquerel discovers emissions
from materials (radioactivity)
• 1898 – Rutherford uncovers properties
of the some Becquerel emissions and
names them alpha and beta
www.lbl.gov/abc/graphics/magnet.gif
Atomic Structure
• 1897 – Thomson discovers “canal rays”
are the same as positive hydrogen
atoms
• Model of atom is “plum
pudding” with + and
– particles mixed together
http://commons.wikimedia.org/wiki/File:Plum_pudding_atom.svg
• 1911 – Rutherford conducts “gold foil
experiment”
Rutherford’s Gold Foil Experiment
Atomic Structure
• After gold foil experiment, atomic
model changes to one with + charges
(protons) in dense center with – charges
(electrons) surrounding the center
• 1932 – Chadwick: discovers the missing
mass in the atom comes from neutral
particles named neutrons
Properties of the Three Key Subatomic Particles
Charge
Mass
Name(Symbol) Relative Absolute(C)* Relative(amu)†
Proton (p+)
Neutron (n0)
Electron (e-)
Location
Absolute(g) in the Atom
1+ +1.60218x10-19
1.00727
1.67262x10-24 Nucleus
0
0
1.00866
1.67493x10-24 Nucleus
1-
-1.60218x10-19
0.00054858
9.10939x10-28
* The coulomb (C) is the SI unit of charge.
†
The atomic mass unit (amu) equals 1.66054x10-24 g.
Outside
Nucleus
Atomic Symbols, Isotopes, Numbers
A
X
Z
The Symbol of the Atom or Isotope
X = Atomic symbol of the element
A = mass number; A = Z + N
Z = atomic number
(the number of protons in the nucleus)
N = number of neutrons in the nucleus
Isotope = atoms of an element with the same
number of protons, but a different number
of neutrons
Figure 2.8 Silberberg, Principles of Chemistry
Figure 2.10
Silberberg, Principles of Chemistry
The modern periodic table.
The Modern Reassessment of Dalton’s Atomic Theory
1. All matter is composed of atoms that are indivisible and
indestructible. The atom is the smallest body that retains the unique
identity of the element. However, it can, under unusual
circumstances, be destroyed (converted to energy) and it can be
divided into smaller parts.
2. Atoms of one element cannot be converted into atoms of another
element in a chemical reaction. Elements can only be converted into
other elements in nuclear reactions.
3. All atoms of an element have the same number of protons and
electrons, which determines the chemical behavior of the element.
Isotopes of an element differ in the number of neutrons, and thus in
mass number. A sample of the element is treated as though its
atoms have an average mass.
4. Compounds are formed by the chemical combination of two or more
elements in specific ratios.
• To understand the electronic structure
of the atom we need to review the
properties of electromagnetic radiation.
Spectra Site
• http://jersey.uoregon.edu/vlab/element
s/Elements.html
Absorption and emission spectra for
element arranged on the periodic table
Metals and Color
Metal ions:
• provide the color in fireworks
and flares (and when you burn colored
newspapers)
• cause glass to have different colors
(stained glass)
• are responsible for the colors of many
precipitates (like the purple Co3(PO4)2)
The Wave Nature
of Light
Frequency and
Wavelength
c=ln
l = wavelength
n = frequency
C = speed of light
Silberberg, Principles of Chemistry
λ, ν, and Energy
• As λ decreases and ν increases, what
happened to the energy of the
radiation?
E=hn =
hc
l
where h = Planck’s constant
(6.626 × 10-34 m2 kg/s)
The infinite number of wavelengths of
electromagnetic radiation have been classified
into groups as shown below.
Regions of the electromagnetic spectrum.
Silberberg, Principles of Chemistry
Light is a wave…right?
• Light falling on alkali
metals causes electrons
to be released from the
metal.
• The # of electrons
depends on the intensity
of light.
• There are specific
wavelengths of light that
cause the release of e-.
• This is called the
photoelectric effect.
Light is a wave…right?
• Einstein’s interpretation of the
photoelectric effect (1905) was that
light is quantized in packets of set
energy called photons. (He won the
Nobel Prize for this.)
• This meant that light had
characteristics of particles!
Electrons are particles…right?
• In 1925, de Broglie stated that all
particles have a wavelength described
by the equation:
λ = h/p where p= momentum
• Electrons show diffraction pattern like
light when passing through a slit.
• So light and particles have a dual
nature.
Back to atomic structure…
• Bohr theorized that the emission
spectra of atoms described by
Rydberg’s equation were caused by the
transition of electrons between specific
energy levels (orbits).
• http://www.upscale.utoronto.ca/General
Interest/Harrison/BohrModel/Flash/Bo
hrModel.html
Figure 7.10
The Bohr explanation of the three series of spectral lines.
Silberberg, Principles of Chemistry
Electron locations
• When an electron occupies its usual
energy level it is in the ground state.
• When an electron absorbs a photon and
moves to a higher energy level it is in an
excited state.
• The energy levels are “quantized”.
Atoms can only transition between set
levels.
• Why are the levels set where they are?
More on electrons as waves
• Since electrons have wave motion
Schrödinger applied the classic wave
equations to the motion of a hydrogen
electron. Certain wavelengths
reinforced each other and were allowed.
• This generated regions occupied by an
electron of set energy termed orbitals.
More on electrons as waves
• Heisenberg stated that in measuring
the electron there is uncertainty so we
can only calculate a probable location
for the electron. This is called the
Heisenberg Uncertainty Principle.
Figure 7.16
Electron probability in the
ground-state H atom.
Silberberg, Principles of Chemistry
Figure 7.19
Silberberg, Principles of Chemistry
The 2p orbitals.
Figure 7.20
Silberberg, Principles of Chemistry
The 3d orbitals.
Figure 7.21
Silberberg, Principles of Chemistry
One of the seven
possible 4f orbitals.
Atomic Shape
• So…when you get all the electrons in
their orbitals surrounding the nucleus,
the shape is basically spherical. This is
why you often see atoms represented as
balls.
Ions…
• Electrons are arranged on levels or
“shells”. Atoms are most stable with 8
electrons on their outermost shell
(except for the small atoms). This is
often referred to as the octet rule.
• Number of electrons on the outer shell
= Group number for elements in Groups
I-VIII A (using the US convention
labeling on periodic table)
Figure 2.10
US Convention
IUPAC
Convention
Silberberg, Principles of Chemistry
The modern periodic table.
Ions…
• To achieve the octet, atoms with more
than 4 electrons on the outer shell will
gain enough electrons to reach 8.
• Cl
Group 7
7 electrons on
outer shell
Cl will gain 1 e- .
• Now #e- = 18 but #p+ = 17 so the
chlorine is an ion with the formula Cl1-.
Ions…
• O Group 6
6 electrons on outer
shell
O will gain 2 e- .
• Now #e- = 10 but #p+ = 8 so the oxygen
is an ion with the formula O2-.
• P
Group 5 5 electrons on outer shell
P will gain 3 e• Now #e- = 18 but #p+ = 15 so the
phosphorus is an ion with the formula P3-
Ions…
• Na Group 1 1 electrons on outer shell
Na will lose 1 e- to expose the complete
inside shell.
• Now #e- = 10 but #p+ = 11 so the sodium is
an ion with the formula Na1+.
• Mg Group 2 2 electrons on outer shell
Mg will lose 2 e• Now #e- = 10 but #p+ = 12 so the
magnesium is an ion with the formula Mg2+
--
1+
2+
3+ -Generally metals form cations
and non-metals form anions.
CHM 1010
Barbara Gage
PGCC
3-
2-
1-
Figure 2.11
From Silberberg, Principles of Chemistry
The formation of an ionic compound.
Transferring electrons from the atoms of one
element to those of another results in an ionic
compound.
Binary Ionic Compounds
• Cations (positively charged ions) and
anions (negatively charged ions) will
associate with each other and form a
neutral binary compound to reduce
energy.
• 1 Na+
1 ClNaCl
• 1 Ca2+
1 O2CaO
• 3 K+
1 N3K3N
• 2 Al3+
3 S2Al2S3
Binary Ionic Compounds
• The “B” group elements may form
cations with more than one charge.
• To specify which cation forms a
compound, a roman numeral equal to the
charge is added to the cation name.
Fe2+ = iron (II) Fe3+ = iron (III)
Polyatomic Ions
• Some atoms form bonds that hold atoms
together in a structure that has an
overall charge (rather than as a neutral
compound. These ions are called
polyatomic ions.
• CO32- carbonate SO42- sulfate
• Compounds with polyatomic ions are
named with the cation and anion name.
• Na2SO4 – sodium sulfate
Common Polyatomic Ions
NH4+
H3O+
NO2NO3SO32SO42S2O32HSO4OHCNPO43HPO42H2PO4CO32HCO3ClOClO2ClO3ClO4C2H3O2-
(Silberberg pg 54 – most common)
ammonium
MnO4hydronium
CrO42nitrite
Cr2O72nitrate
O22sulfite
sulfate
thiosulfate
hydrogen sulfate or bisulfate
hydroxide
cyanide
phosphate
hydrogen phosphate
dihydrogen phosphate
carbonate
hydrogen carbonate or bicarbonate
hypochlorite
chlorite
chlorate
perchlorate
acetate (or CH3COO- or CH3CO2-)
permanganate
chromate
dichromate
peroxide
Figure 9.9
Solid ionic
compound
Silberberg, Principles of Chemistry
Electrical conductance and ion mobility.
Molten ionic
compound
Ionic compound
dissolved in water
Covalent Compounds
• Non-metals form bonds by sharing
electrons rather than transferring them
to achieve the octet.
• The resulting bond is referred to as a
covalent bond. Each pair of shared
electrons = 1 bond
• Shared electrons move around the
nuclei of both atoms in the bond so both
atoms have possession of the shared
electrons.
Covalent Binary Compounds
CO2
carbon dioxide
N2 O
dinitrogen oxide
P2S5
diphosphorus pentasulfide
SiCl4
silicon tetrachloride
For compounds where two vowels occur
together when the prefix is added, the
vowel from the prefix can be dropped
(except for iodine).
• N2O5
dinitrogen pentaoxide
•
•
•
•
•
Electron Distribution in a
Covalent Bond
• Are electrons shared equally in a
covalent bond?
• If not, why not?
• Distance of electrons from nucleus and
number of protons in the nucleus
• Electronegativity – attraction of one
atom in a bond for the electrons in that
bond
Figure 9.20
The Pauling electronegativity (EN) scale.
Silberberg, Principles of Chemistry
Polarity
• When atoms in a bond have different
electronegativities, the electron sharing
is unequal.
• As the ΔEN increases, the electron
distribution becomes more uneven and
the molecule becomes polar.
Polarity
• HCl
• ENH = 2.1
ENCl = 3.0
ΔEN = 0.9
• The end with the higher EN will be
slightly negative and the other will be
slightly positive
δ+H
– Clδ-
H – Cl
Figure 9.22
Boundary ranges for classifying ionic character
of chemical bonds.
3.0
DEN
2.0
0.0
Silberberg, Principles of Chemistry