COMPETENCY STANDART :

1. To understand the atom structure for predicting the element periodic characteristics, molecule structure and compound characteristics

THE BASE COMPETENCY :

1.1 To explain about Bohr atom theory and quantum mechanics for writing electron configuration, orbital diagram and determining the element location in periodic table

INDICATOR :

1. To explain about quantum mechanics theory

2. To determine quantum number (electron is possible)

3. To apply Aufbau principle, arrangement Hund and prohibition Pauli

4. To connect the element electron configuration with the location in periodic table

Electron orbitals and the periodic table

In the planetary model of the atom, electrons orbit the nucleus, but if this actually happened, they should lose energy in a dying scream of electromagnetic radiation and crash into the nucleus (the gravitational equivalent is what happened to the space station Mir). This is because the negatively charged electron is inherently attracted to the positively charged nucleus, and crashing into it represents a lower energy, and therefore more favourable state.

The solution to this problem is the quantum model of the atom. In quantum theory, the energy of an electron is only allowed to be gained or lost in discrete chunks called quanta. No 'scream' of energy is allowed, only discrete jumps between energy states. This makes the the allowed energy levels in an atom like the discrete steps of a staircase, rather than a continuous gentle slope. Furthermore, transitions are only allowed to unoccupied energy states (this is called the Pauli exclusion principle). Hence, energies higher than the lowest energy state are stable, because there is nowhere unoccupied further down the staircase of energy for the electrons to fall to. However, this still leaves the 'ground state' (the lowest energy state). Why is this stable? To account for this, we invoke Heisenburg's uncertainty principle, an important maxim of quantum theory, which states that an electron's position and velocity cannot be simultaneously known with certainty. One way we can tell where an electron is in space is by hitting it with light waves. However, the more precisely we need to know the position of the electron, the smaller the wavelength of light we will need: a dinghy will leave an obvious, observable shadow behind short-wavelength water waves, but a long-wavelength tsunami will not even notice it's there.

λ = h ⁄ p , or equivalently, E = m c² = hν

(λ is wavelength, ν frequency, E energy, p momentum, h Planck's constant, m mass, c speed of light)
However, by de Broglie's and Planck's relationships (above), short wavelengths carry more energy: X-rays (short wavelength) are more energetic than radio waves (long wavelength). These energetic light waves will kick the electron sufficiently to change its movement, so although we know exactly where the electron is, we will have no idea in which direction it is now travelling. From this, we can see that if an electron tries to get nearer the nucleus, its position will become more certain. Hence its velocity must become less certain, and therefore higher on average. This means its energy of motion is higher, and therefore there is no net energy advantage in crashing into the nucleus, because the increase in kinetic energy would exceed any loss in potential energy. Hence the ground state is also stable.
Quantum physics states that all particles have wave-like properties (like wavelength) and that all waves have particle-like properties (like momentum). In fact, all subatomic particles can be thought of as a sort of hybrid between a wave and a particle. For example, electrons, previously considered as particles (i.e. tiny billiard balls), also have wave like properties. The wave we're talking about here is a wave of probability of finding an electron at a given point. This is like a crime wave: the probability of getting mugged is higher when a crime wave runs through your town, hence when an electron probability wave runs through your equipment, you stand a good chance of seeing an electron particle somewhere. In an atom, the electron probability waves form 'standing waves', much like those you get when you pluck a guitar string. In a similar way to a guitar string, the atom allows various modes of vibration of the standing wave:

Schrödinger's equation describes these modes of vibration, and the solutions are generally graphically displayed as orbitals, which are the space where there's a better than 99% chance of finding an electron.
There are many sorts of orbitals, as we shall see. The various orbitals available to an atom are described by four quantum numbers, which can take certain values to create differently sized and shaped orbital of various energies:

Principal (n).
Azimuthal (l).
Magnetic (m_l).
Spin (m_r).

Th principal quantum number (n) describes the:

Largest energy differences ('shells').
Size of orbital.
Numbered 1, 2, 3, 4, 5, etc (although the shells it describes are often lettered K, L, M, etc.).
Low numbers are closest to the nucleus, and have the lowest energies.
Higher values of n indicate larger orbitals

The 2p_x, 2p_y and 2p_z orbitals all have the same energy, but different orientations in space.
Summary of Orbitals:

n=1 shell

One s-orbital: 1s

n=2 shell

One s-orbital: 2s
Three p-orbitals: 2p_x 2p_y 2p_z

n=3 shell

One s-orbital: 3s
Three p-orbitals: 3p_x, 3p_y, 3p_z
Five d-orbitals: 3d_xy, 3d_xz, 3d_yz, 3d_x²_−y², 3d_z²

n=4 shell

One s-orbital: 4s
Three p-orbitals: 4p_x, 4p_y, 4p_z
Five d-orbitals: 4d_xy, 4d_xz, 4d_yz, 4d_x²_−y², 4d_z²
Seven f-orbitals: 4f_z³, 4f_xz² 4f_z(x²_−y²₎, 4f_xy², 4f_xyz, 4f_x(x²_−3y²₎, 4f_y(3x²_−y²₎

The final quantum number is the spin quantum number (m_r). Each orbital (e.g. a 2p_x) can hold two electrons. The Pauli exclusion principle states that no two electrons in a single atom can have the same quantum numbers, hence they must differ in their spin quantum number, which takes the values +½ or -½. Two electrons in the same orbital have paired (opposite) spins. The number of electrons a set of degenerate orbitals can contain is therefore just 2 for an s orbital, 6 for a set of p orbitals (3p_x 3p_y 3p_z can hold 2 electrons each), 10 for a set of five d orbitals, 14 for a set of f, etc.
The periodic table is a list of the elements by their atomic number (i.e. by the number of electrons a neutral atom of the element possesses). Hence, as we traverse the periodic table, we are filling up electron orbitals from lowest energy to highest energy. We can write electronic configurations for atoms based on the arrangement of their electrons. We just need to write the (non-degenerate) orbitals out in energy order, then fill them up, writing the number of electrons each set of orbitals contains as a superscript
Hydrogen (1 electron) has the electronic configuration 1s¹
Boron (5 electrons) has the electronic configuration 1s² 2s² 2p¹
Sodium (11 electrons) has the electronic configuration 1s² 2s² 2p⁶ 2s¹

Worksheet 1-Atomic structure

1. Use your periodic table to complete the following table:

Name of element	Atomic Number	Atomic Symbol	Atomic Weight	Ionic Charge	No. of electrons in atom	No. of protons in ion	No. of electrons in ion
Lithium	3	Li	6.94 gm	+1	3	3	2
		H
			24.3 gm
	19
Barium
			55.85 gm	+2
						29
Oxygen
			*20.2 gm*	0	10	-	-
			200.6 gm

2.     You are an atom with 6 electrons in the outer orbital and 16 protons.
        A.    What is your name?
        B.    What is your weight?
        C.    What will your most likely charge be when you ionize?
        D.    How many sodium ions will you react with to produce a net charge of zero?
        E.    What will your formula be?
3.     You have a -1 charge and weigh 35.45 gm.
        A.    What is your element name?
        B.    What is your ionic symbol?
        C.    How many protons do you have as an atom?
        D.    How many electrons are in the outer orbital?
        E.    How many total electrons do you have as an ion?
        F.    How many of chloride ions can react with an Aluminum ion?
        G.    What will your formula be?
4.     How many electrons can fill an s suborbital?
        How many electrons can fill a p suborbital?
        How many electrons can fill a d suborbital
5. Which of the following is a correct electron structure?
        A. 1s²1p⁶2s²
        B. 1s²1p⁶2s³2p⁶3s²3p⁶4s²
        C.1s²2s²2p⁶3s²3p⁶4s²
        D. 1s²2s²2p⁶3s²3p⁶4s²3d⁴
        E. 1s²2s²2p⁶3s²3p⁵4s²
        F. 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p¹
6. Name the element whose uncharged atoms have the following electronic structure.
        A.    1s²2s²2p⁴
        B.    1s²2s²2p⁶3s²3p⁶4s¹
        C.    1s²2s²2p⁶3s²3p⁶4s²3d⁴
        D.    1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p³
        E.    1s²2s²2p⁶3s²3p³
7. What is the weight of 1 proton?
    What is the charge on an electron?
    What is the charge on a neutron?
    What is the weight of an electron?
    What is the charge on a proton?
8. Which bond is the most difficult to break?
    Which bond has overlapping orbitals?
    Which bond is the result of electrochemical attraction?

evaluasi

I. MULTIPLE CHOISE

Choose the correct answer in the following questions by give cross symbol (X) on the alphabet a, b, c, d or e in the answer sheet!

1. Which one he following statement is not the mistake of Dalton’s atomic model?

a. Atom are empty structure

b. The atom of the same type elements are totality identical

c. It is not true thet atoms are indivisible

d. Atoms form compounds by combining with a definite number ratio

e. Atoms show the characteristics of the element

2. Which one of the following electron structures belongs to the ion ₁₂ Mg²⁺?

a. [Ne]3s¹

b. [Ne]3s²

c. [Ne]

d. 1s²2s² 2p⁴

e. 1s²2s² 2p⁵

3. Which one the following electron distributions does not Hund’s rule?

a. 1s²2s² 2px² 2py¹ 2pz¹

b. 1s²2s² 2px¹ 2py¹

c. 1s²2s²2px² 2py² 2pz¹

d. 1s²2s² 2px² 2py¹ 2pz⁰

e. 1s²2s² 2px¹ 2py¹ 2pz¹

4. Calculete the ΔH₃!

a. ΔH₁+ΔH₂ - ΔH₄

b. ΔH₄+ΔH₂ - ΔH₁

c. ΔH₁–ΔH₃ + ΔH₄

d. ΔH₁–ΔH₂ - ΔH₄

e. ΔH₁+ΔH₄ – ΔH₂

ΔH₂

MgO

ΔH₁ ΔH₃

Mg(OH)₂

ΔH₄

5. After mixing 40 mL of 2.0 M HCl with 100 mL of 1 M NaOH, the temperature increases from 18.9ºC to 26.7ºC. What is q for the reaction? c = 4.18 J/g K

a. -4.6 kJ

b. -4.8 kJ

c. -2.3 kJ

d. 2.3 kJ

e. 4.6 kJ

6. When 50.0 mL of 1.00 M HCl is mixed with 50.0 mL of 1.00 M NaOH in a calorimeter, a neutralization reaction takes place in which the temperature of the mixture increases from 25oC become 31.6oC. How much heat (in kJ) is released or absorbed by this reaction, assuming the calorimeter absorbs very little heat ?C = 4.18 J/goC

a. +2.76

b. -23.6

c. -2.76

d. -13.2

e. -1.38

7. What is the bond angle between hydrogens in NH₃?

a)	120°
b)	109.5°
c)	107°
d) e)	90° 180^o

8. What is the hybridization of XeF₂?

a)	sp
b)	sp²
c)	d²sp³
d) e)	dsp³ sp³

9. Cl₂ is a….

a)	polar covalent molecule
b)	non polar covalent molecule
c)	ionic molecule
d)	metallic molecule
e)	covanlent molecule

10. At equilibrium: A (g) + B (g) ↔ C (g) + D (g). The value of K = 1, [W] = 2 Y, [X] = ….

a. 4 x [Z]

b. 2 x [Z]

c. [Z]

d. ½ x [Z]

e. ¼ x [Z]

kimia

Kamis, 03 Januari 2013

ATOMIC STRUCTURE

Contents

Electron orbitals and the periodic table