Atomic structure is composed of protons, neutrons, and electrons. Rutherford's gold foil experiment showed that atoms are mostly empty space with a small, dense nucleus at the center containing protons and neutrons. Electrons orbit the nucleus in specific energy levels. The Bohr model explained electrons in hydrogen having discrete energy levels. Atoms bond via ionic bonds between metals and nonmetals by transferring electrons, covalent bonds between nonmetals by sharing electrons, and metallic bonds between metals by delocalized electrons.

1. ATOMIC STRUCTURE

2. Atomic Structure
All matter is composed of atoms.
Understanding the structure of atoms is
critical to understanding the properties
of matter

3. Subatomic Particles
Particle
Mass
(g)
Charge
(Coulombs)
Charge
(units)
Electron (e-
) 9.1 x 10-28
-1.6 x 10-19
-1
Proton (p) 1.67 x 10-24
+1.6 x 10-19
+1
Neutron (n) 1.67 x 10-24
0 0
mass p = mass n = 1840 x mass e-

4. Rutherford’s experiment.

5. Actual Results.

7. Rutherford observed that most of the alpha-particles
passed through the foil without any deflection from their
path and struck the screen at its centre, causing
illuminations.
A few of them were deflected at some angles (90 0 or
wider angles) after passing through the foil.
Very few (not more than one in 10,000) turned back to
their original path.

8. Postulates of Rutherford's Model
1. Atom contains a massive (heavy) and positively charged
part at its centre. This central part of the atom is called
nucleus.
2. The volume occupied by the nucleus is only a minute
fraction of the total volume of the atom.
3. Atom is not all solid, as was earlier suggested by
Dalton, but is
extraordinarily hollow, since it consists of a lot of empty
space round the nucleus.
4. Electrons are revolving round the nucleus in closed
orbits with a fast speed and hence almost all the space
round the nucleus is occupied by the revolving electrons.

10. Atomic Structure
Atoms are composed of
-protons – positively charged particles
-neutrons – neutral particles
-electrons – negatively charged particles
Protons and neutrons are located in the
nucleus. Electrons are found in orbitals
surrounding the nucleus.

11. Atomic Structure
Every different atom has a characteristic
number of protons in the nucleus.
atomic number = number of protons
Atoms with the same atomic number
have the same chemical properties and
belong to the same element.

12. Atomic Structure
Each proton and neutron has a mass of
approximately 1 dalton.
The sum of protons and neutrons is the atom’s
atomic mass.
Isotopes – atoms of the same element that
have different atomic mass numbers due to
different numbers of neutrons.

13. 1. e- can have only specific
(quantized) energy values
2. light is emitted as e- moves
from one energy level to a
lower energy level
Bohr’s Model of
the Atom (1913)
En = -RH ( )
1
n2
n (principal quantum number) = 1,2,3,…
RH (Rydberg constant) = 2.18 x 10-18J

14. The Bohr Model of the Atom

15. The Bohr Model of the Atom:
Ground and Excited States
• In the Bohr model of hydrogen, the lowest amount
of energy hydrogen’s one electron can have
corresponds to being in the n = 1 orbit. We call this
its ground state.
• When the atom gains energy, the electron leaps to a
higher energy orbit. We call this an excited state.
• The atom is less stable in an excited state and so it
will release the extra energy to return to the ground
state.
– Either all at once or in several steps.

17. Barbara A. Gage PGCC CHM
1010
Why Do Atoms Bond?
• Chemical bonds form because they lower
the potential energy between the charged
particles that compose atoms
• A chemical bond forms when the potential
energy of the bonded atoms is less than
the potential energy of the separate
atoms

18. Barbara A. Gage PGCC CHM
1010
Types of Bonds
Types of Atoms Type of Bond
Bond
Characteristic
metals to
nonmetals
Ionic
electrons
transferred
nonmetals to
nonmetals
Covalent
electrons
shared
metals to
metals
Metallic
electrons
pooled
• We can classify bonds based on the kinds
of atoms that are bonded together

19. Barbara A. Gage PGCC CHM
1010
Ionic Bonds
• When a metal atom loses electrons it becomes a
cation
– metals have low ionization energy, making it
relatively easy to remove electrons from
them
• When a nonmetal atom gains electrons it
becomes an anion
– nonmetals have high electron affinities,
making it advantageous to add electrons to
these atoms
• The oppositely charged ions are then attracted
to each other, resulting in an ionic bond

20. Barbara A. Gage PGCC CHM
1010
Covalent Bonds
• Nonmetal atoms have relatively high ionization
energies, so it is difficult to remove electrons
from them
• When nonmetals bond together, it is better in
terms of potential energy for the atoms to
share valence electrons
– potential energy lowest when the electrons are
between the nuclei
• Shared electrons hold the atoms together by
attracting nuclei of both atoms

22. Barbara A. Gage PGCC CHM
1010
Metallic Bonds
• The relatively low ionization energy of metals
allows them to lose electrons easily
• The simplest theory of metallic bonding
involves the metal atoms releasing their
valence electrons to be shared as a pool by all
the atoms/ions in the metal
– an organization of metal cation islands in a
sea of electrons
– electrons delocalized throughout the metal
structure
• Bonding results from attraction of cation for
the delocalized electrons

23. Hydrogen Bonding
These three bonds all have;
• A strong permanent dipole
• A hydrogen atom
• An atom with lone pair electrons
The three types of bonds which give
molecules significant hydrogen bonding
are; (i) N – H (ii) O – H (iii) F – H

24. 0
50
100
150
200
250
300
350
400
Boiling
Point
(K)
CH4
SiH4
GeH4
SnH4
H2O
H2S H2Se
H2Te

25. Hydrogen Bonding
Hydrogen bonding in water results in some unusual
properties;
• Higher than expected boiling point
• High specific heat capacity
(absorbs a lot of heat energy with
only a small change in temperature)
• Ice is less dense than water

34. Valence Bond (VB) Theory
The basic principle of VB theory:
A covalent bond forms when the orbitals of two atoms
overlap and a pair of electrons occupy the overlap region.
The space formed by the overlapping orbitals can
accommodate a maximum of two electrons and these
electrons must have opposite (paired) spins.
The greater the orbital overlap, the stronger the bond.
Extent of orbital overlap depends on orbital shape and direction.

35. Figure 11.1 Orbital overlap and spin pairing in H2.
A covalent bond results from the overlap of orbitals from two atoms.
The shared space is occupied by two electrons, which have opposite spins.

36. Figure 11.2 Orbital orientation and maximum overlap.
Hydrogen fluoride, HF. Fluorine, F2.
The greater the extent of orbital overlap, the stronger the bond.

37. VB Theory and Orbital Hybridization
The orbitals that form when bonding occurs are different
from the atomic orbitals in the isolated atoms.
If no change occurred, we could not account for the molecular shapes
that are observed.
Atomic orbitals “mix” or hybridize when bonding occurs
to form hybrid orbitals.
The spatial orientation of these hybrid orbitals correspond with
observed molecular shapes.

38. Features of Hybrid Orbitals
The number of hybrid orbitals formed equals the number
of atomic orbitals mixed.
The type of hybrid orbitals formed varies with the types of
atomic orbitals mixed.
The shape and orientation of a hybrid orbital maximizes
overlap with the other atom in the bond.

39. Figure 11.3 Formation and orientation of sp hybrid orbitals
and the bonding in BeCl2.
orbital box diagrams
atomic
orbitals
hybrid
orbitals
One 2s and one 2p atomic orbital mix to form two sp hybrid orbitals.

40. Figure 11.3 continued
box diagram with orbital contours
Overlap of Be and Cl orbitals to form BeCl2.

41. Figure 11.4 The sp2 hybrid orbitals in BF3.
Mixing one s and two p orbitals gives three sp2 hybrid orbitals.
The third 2p orbital remains unhybridized.

42. Figure 11.4 continued
The three sp2 orbitals point to the corners of an equilateral triangle,
their axes 120° apart.
Each half-filled sp2 orbital overlaps with the half-filled 2p orbital of a
F atom.

43. Figure 11.5 The sp3 hybrid orbitals in CH4.
The four sp3 orbitals adopt a
tetrahedral shape.

44. Figure 11.6 The sp3 hybrid orbitals in NH3.
The N lone pair occupies an sp3
hybrid orbital, giving a trigonal
pyramidal shape.

46. Molecular Orbital (MO) Theory
The combination of orbitals to form bonds is viewed as the
combination of wave functions.
Atomic wave functions (AOs) combine to form molecular
wave functions (MOs).
Addition of AOs forms a bonding MO, which has a region
of high electron density between the nuclei.
Subtraction of AOs forms an antibonding MO, which has
a node, or region of zero electron density, between the
nuclei.

47. Molecular Orbital Diagrams
An MO diagram, just like an atomic orbital diagram,
shows the relative energy and number of electrons in
each MO.
The MO diagram also shows the AOs from which each
MO is formed.
Bond order is calculated as follows:
½[(# of e- in bonding MO) – (# of e- in antibonding MO)]

48. Figure 11.17 MO diagram for H2.
H2 bond order = ½ (2 − 0) = 1

49. Electrons in Molecular Orbitals
• MOs are filled in order of increasing energy.
• An MO can hold a maximum of 2 e- with opposite spins.
• Orbitals of equal energy are half-filled, with spins
parallel, before pairing spins.
Electrons are placed in MOs just as they are in AOs.
A molecular electron configuration shows the type of
MO and the number of e- each contains. For H2 the
configuration is (σ1s)2.