Summary of Atomic Models

SUMMARY OF ATOMIC MODELS

NAME / REPRESENTATION / MAIN IDEAS
Dalton Model
“Billiard Ball” model / /

1. All matter is made up of tiny indivisible particles called .
2. All atoms of the same element are the same in , mass, etc.
3. Atoms of different elements are different in and .
4. Atoms of different elements combine in small whole number ratios to form

.

1. In chemical reactions, atoms are not created or . They are

or .
Thomson Model
“ ______”
Model / /

1. The atom is a spherical cloud of positive charge with negatively charged

embedded in it.

1. The amount of positive charge the amount of negative charge, so the entire atom is neutral.

Rutherford Model
“Solar system Model” / /

1. Discovered using the gold . experiment.
2. The atom is mostly empty space.
3. Most of the mass is concentrated in a small, dense .
4. Electrons move around the nucleus like planets around the .

______Model
“Energy level” Model / /

1. Electrons travel around the nucleus in specific .
2. If an atom is excited, electrons absorb a certain amount of and are boosted to a energy level.
3. Electrons later drop back to a lower state releasing of a colour equal to the amount of energy released.

4.  The first energy levels can hold a maximum of electrons. The others have a valence shell of electrons.

PROBABILITY

OBJECTIVE:

The purpose of this activity is to illustrate how probability and charge density are related to distance from the nucleus for the 1s orbital of a hydrogen atom. An analogy of throwing stones at a target will be used. The target is marked off in circles with radii increasing in 5 cm increments.

PROCEDURE:

1. Place the target on the desk. Stand back just beyond reach of the table edges and toss 100 stones at the target one at a time, always aiming for the centre.

2. Count the number of stones in each ring. Any stones which land on a line should be counted as landing in the inner ring.

RESULTS:

Ring Radius
(cm) / Number of Hits (Your Data) / Number of Hits (Class Data)
0 - 5
5 - 10
10 - 15
15 - 20
20 - 25
25 - 30
30 - 35
missed target
total throws

DATA ANALYSIS:

1.  Plot the number of hits vs. the ring radius as a bar graph. Use the totals for the class.

Graph 1
#
o
f
H
i
t
s
0 - 5 / 5 – 10 / 10 – 15
Radius / 15 - 20
(cm) / 20 - 25 / 25 - 30 / 30 - 35

2.  What is the most probable distance from the center for the next stone to land?

3.  Predict the probability of a hit in a ring 35 - 40 cm from the center. This can be done by extrapolation of the graph. The probability is the estimated number of hits divided by the total number of throws.

4.  In the following chart:

a) calculate the area of each circle

b) calculate the area of each ring by taking the difference in area between successive circles

c) calculate the density of hits for each ring; that is, the number of hits per square centimeter

Ring Radius
(cm) / Area of Circle
(cm2) / Area of Ring
(cm2) / Density of Hits
(hits/cm2)
0
5
10
15
20
25
30
35

5.  Plot the density of hits vs. the ring radius as a bar graph.

Graph 2
D o
e f
n
s H
i i
t t
y s
0 - 5 / 5 – 10 / 10 - 15 / 15 - 20
Radius (cm) / 20 - 25 / 25 - 30 / 30 - 35

6.  How does the shape of this graph differ from the shape of the first graph?

7.  Which graph most clearly shows that you were aiming at the centre of the target?

8.  The shape of which graph combines the effect of aiming for the centre with the fact that the area decreases as the ring radius decreases?

9.  The following graph shows the probability of finding the electron at a distance r from the nucleus for the 1s orbital of hydrogen. Draw a vertical line showing the electron’s most probable distance from the nucleus.

10.  There are two factors affecting the shape of this curve. These factors are analogous, but not identical to those affecting the shape of the bar graph in the stone throwing experiment.

(a) What factors in the atom would result in the electron tending to be close to the nucleus?

(b) In spite of the previous factor, probability decreases as the distance from the nucleus decreases below that of the peak on the graph. Why is this so?

11.  Imagine that the electron charge could be “smeared out” over the entire volume of the atom with more of the charge concentrated in places where it is more likely to be found. Sketch a graph of charge density (charge per unit volume) vs. distance r for the 1s orbital of a hydrogen atom using the graphs from the stone throwing activity as a guide.

SHAPES OF THE ORBITALS

s orbital

p orbitals

d orbitals

QUANTUM NUMBERS

PRINCIPAL QUANTUM NUMBER, n
The main electron energy level or shell.
n = 1, 2, 3, 4...... / SECONDARY OR ORBITAL QUANTUM NUMBER, L
The electron sublevel or sub shell or orbital.
L = 0 to L = n-1 / MAGNETIC QUANTUM NUMBER, mL
The spatial orientation of the orbital.
+L ³ mL ³ -L / SPIN QUANTUM NUMBER, mS
The spin angular momentum.
+1/2 or - 1/2 / ORBITAL NAME
4 / 0 / 0 / +1/2, - 1/2 / s for sharp
1 / -1 / +1/2, - 1/2 / p for principal
0 / +1/2, - 1/2
+1 / +1/2, - 1/2
2 / -2 / +1/2, - 1/2 / d for diffuse
-1 / +1/2, - 1/2
0 / +1/2, - 1/2
+1 / +1/2, - 1/2
+2 / +1/2, - 1/2
3 / -3 / +1/2, - 1/2 / f for fundamental
-2 / +1/2, - 1/2
-1 / +1/2, - 1/2
0 / +1/2, - 1/2
+1 / +1/2, - 1/2
+2 / +1/2, - 1/2
+3 / +1/2, - 1/2

PERIODIC PROPERTIES OF THE ATOM/ELECTRON

FACTORS TO CONSIDER

1.  Coulomb’s Law: where k is a constant, q is electrical charge (+ proton, - electron), r is inter-particle distance (proton in nucleus to electron, electron to electron) and F is electrical force of attraction (opposite charges) or repulsion (same charge).

2.  Nuclear charge, i.e. Z (Atomic Number, # of protons). As Z­ then q­ and F­.

3.  Principal quantum number, i.e. principle energy level, “n”. The greater the principle energy level of the electron, the greater the radius. As n­, r­and F¯¯ because r2.

4.  Shielding, i.e. the greater the number of inner principle energy level electrons (core electrons), the greater the outward repulsions on the valence electrons and the weaker the net force of attraction from the nucleus.

5.  Sublevel or orbital type, eg. “s” orbitals have the lowest potential energy (lower “r”) and the “p” have a higher potential energy (higher “r”) in the same principal energy level, same “n”.

6.  Charge distribution. Symmetrical distributions of electrons in the same orbital type: empty, 1/2 filled or fully filled are more energetically favorable (more stable) than non-spherically symmetrical distributions.

7.  Note: ∆IE is the change in Ionization Energy.

Explain the data in each of the following tables using all or some of the above factors.

GROUP I SUCCESSIVE IONIZATION ENERGIES

ELEMENT / LEVEL / IE
(kJ/mol) / Z / n, r / shielding / orbital
shielding / symmetry / Therefore, \
Li
(relative to H) / 2 / 520
Na / 3 / 491
K / 4 / 414
Rb / 5 / 404
Cs / 6 / 376

INCREASING Z and IONIZATION ENERGY

Element and Z / IE
(kJ/mol) / Valence Orbital Box Diagram / n / shielding / r / orbital shielding / e- - e- / symmetry change / Therefore, \
H / 1310
He / 2396
Li / 520
Be / 901
B / 805
C / 1088
N / 1396
O / 1310
F / 1676
Ne / 2080
Na / 491
Mg / 738
Al / 641
Si / 786
P / 1060
S / 1002
Cl / 1252
Ar / 1522
K / 414
Ca / 587

CHEMICAL BONDING

The energy change that occurs when two hydrogen atoms approach each other is shown below.

1. Why does the potential energy drop as distant atoms approach each other?

2. Why does the potential energy increase when the atoms are very close?

3. Mark on the graph the equilibrium bond distance.

4. When the atoms are at the equilibrium bond distance, how does their potential energy compare to what they had when the distance of separation was large?

5. When the atoms are at the equilibrium bond distance, how does their kinetic energy compare to what they had when the distance of separation was large?

6. When the atoms are at the equilibrium bond distance, what must happen to prevent them from moving apart; that is, what happens if a bond forms?

7. If a bond forms, what must be done in order to break it and separate the atoms once again?

8. Sketch a graph similar to that above showing how it would look for two helium atoms as they approach.

9. Explain why the two graphs are different. Refer specifically to the attractions and repulsions which occur in each case.

10. What is there about the atoms that determines whether a chemical bond will form?

# e-densities / Basic Geometry
0 lone pair / 1 lone pair / 2 lone pairs / 3 lone pairs / 4 lone pairs
2 /
Linear
AX2
3 /
Trigonal planar
AX3 /
Bent
AX2E1
4 /
Tetrahedral
AX4 /
Trigonal pyramid
AX3E1 /
Bent
AX2E2
5 /
Trigonal bipyramid
AX5 /
Seesaw
AX4E1 /
T-shaped
AX3E2 /
Linear
AX2E3
6 /
Octahedral
AX6 /
Square pyramid
AX5E1 /
Square planar
AX4E2 /
T-shaped
AX3E3 /
Linear
AX2E4

FIVE BASIC SHAPES OF MOLECULES

# e- densities / FORMULA / Base Shape / Base Name / Hybridization
2 / AX2 / / Linear
3 / AX3 / / Trigonal planar
4 / AX4 / / Tetrahedral
5 / AX5 / / Trigonal bipyramid
6 / AX6 / / Octahedral

MOLECULAR SHAPES & THEIR REPRESENTATIONS

1.  For this activity, you will complete the charts giving the Lewis dot diagram, the VSEPR formula and Name, hybridization type, the perspective 3D diagram and hydridization development for several molecules.

2.  The following guidelines will apply:
Lewis dot diagram: must include all bonds and non-bonding electron pairs.

VSEPR formula and Name:

a) 1 of the 5 basic shapes

® use basic formula and name

b) modifications of the basic shapes

® use basic formula and name and also specific formula and name

c) 2 central atoms

use a) and b) above for each central atom.

Hybridization type: state for each central atom

Perspective 3D diagram: use the following conventions

® a bond in the plane of the paper is shown as a solid line

® a bond projecting out of the plane towards the viewer is drawn as a wedge

® a bond projecting out of the plane away from the viewer is drawn as a dashed line

Hydridization development: for each unique central atom show the orbital box diagrams for the ground state, valence state (if needed) and the hybridized state.

3.  For example, using methane, CH4:
- Draw the Lewis Dot diagram.

- Using the MolyMod Molecular kits construct a model of methane, CH4. For single bonds use the rigid bond sticks.

- Create a perspective drawing using the conventions listed above.

- Include the angle for at least 1 of each different angle in the structure.

ie. 109.5°, 120°, 180° or 90°.

- for each central atom, state the basic VSEPR name and formula for that atom.
ie. Linear–AX2, trigonal planar–AX3, tetrahedral–AX4, trigonal bipyramidal–AX5 or octahedral-AX6.

- if the central atom symmetry is not a basic VSEPR shape, state also the specific name and formula.

ie. Angular–AX2E2.

- state the hydrization type for each central atom and its development as instructed above

4.  In methane there is only one carbon, The central atom, which is bonded to several other atoms. In more complex molecules there may be several central atoms. The correct bond angles must be represented around each atom of the model. A perspective drawing will be simpler if as many atoms as possible are arranged in the plane of the paper.

5.  It is easy to forget the unbonded pairs of electrons surrounding some atoms in a molecule. In order to predict the shapes of molecules correctly these electrons must also be taken into account. Since unbonded pairs of electrons crowd closer to the central atom, they repel other pairs more strongly, causing the bond angles to be somewhat smaller than the usual 109.5°. This should be noted in your perspective drawings.

6.  For some compounds, the octet rule does not apply.

7.  The shape of molecules with multiple bonds is correctly predicted by using the flexible bond sticks.

Molecular Models and The Shapes of Molecules

Type of Molecule / Lewis Structure
(Electron Dot Diagram) / VSEPR formula, Name
& Hybridization / 3-D Structural Formula / Hybridization Development
CH4 /
Ð =109.5° / AX4
Tetrahedral
\ sp3 / / C [He] / ­¯ / ­ / ­
2s2 / 2p2
C [He] / ­ / ­ / ­ / ­
2s1 / 2p3
C [He] / ­ / ­ / ­ / ­
sp3
C2H6
PCl3
SCl2
CH3OH
H2O2
HONH2
H3COCH3
H3CCOOH
PH5
SF4
IF3
XeF2
SF6
IF5
XeF6
[I3 –]
C2H4
C2H2
N2O
HCN
N2H2
C4H6
1,3 -butadiene

INTERMOLECULAR FORCES

1.  Complete the following table indicating the type of intermolecular bonding found in each compound.

2.  Drawing the Electron Dot diagram may help in the answer to some of the questions.