Electrons in Atoms (Chapter 4) Notes Name: __________________________________________ I. Properties of Light-Different types of electromagnetic radiation (x-rays, radio waves, microwaves, etc…) SEEM to be very different from one another. Yet they share certain fundamental characteristics. All types of electromagnetic radiation, also 8 called radiant energy, move through a vacuum at a speed of 3.00 x l0 meters per second. A. Wavelength – distance between identical points on successive waves; may be measured in any length unit but is usually dependent on how long the wave is (X-rays are usually measured in nanometers or Angstroms while the very long radio waves might be measured in meter. The Greek letter lambda, , is used to depict wavelength (pg 92)

B. Frequency – the number of complete wave cycles that pass a given point in one second: the unit is cycles/second but is -1 written as sec , or Hertz. The letter f or the Greek letter nu,, is used to depict frequency. If the frequency and wavelength are known then the product of the two (wavelength x frequency) is always equal to the same speed. It is known as the speed of light or c. 8

c = speed of light = 3.00 x l0 m/s

c =f  = wavelength (in m) f = frequency (in Hz)

Examples: 13 -1 l. What is the wavelength of radiation whose frequency is 6.24 x l0 sec ?

-6

2. What is the frequency of radiation whose wavelength is 2.20 x l0 nm? (1 m = 1,000,000,000 nm)

3. A television relay transmitter operates at a frequency of 927.9 MHz. Calculate the wavelength of the signal in nanometers. Is this visible light? (1 MHz = 1,000,000 Hz)

Energy relationship to frequaency:

E = hf h = Planck’s constant = 6.63 x l0 E = energy (in Joules) f = frequency (in Hz)

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Joules  sec

Examples: 19 1. If a certain light has 7.18 x l0- J of energy what is the frequency of this light?

b. What is the wavelength, in m, of this light?

2. If the frequency of a certain light is 3.8 x l0

3. The energy of a certain light is 3.9 x l0

-19

14

Hz, what is the energy of this light?

J. What is the wavelength, in nm, of this light? Is it visible?

4. Calculate the smallest increment of energy that an object can absorb from yellow light, the color given off when sodium atoms are heated in a flame. The wavelength of this light is 589 nm.

II. Quantum Mechanics Model of the Atom Periodic Table with predicted ending electron configurations. A. Energy Levels: energy levels (represented by the letter n) are assigned values in order of increasing energy: n=1,2,3,4, and so forth…. which correspond to the periods in the periodic table. The principle q. n. is related to the size and energy of the orbital. n=1, n=2, n=3, n=4, n=5, etc… Which energy level is furthest away from the nucleus and has electrons with the highest energy - 1, 2,3, or 4? B. Sublevels: Within each energy level, the electrons are located in various sublevels – there are 4 different sublevels s, p, d, and f. The sublevel defines the shape of the orbital (s, p, d, & f).

1s

2s

2p 3p

C. Orbitals: Where are the electrons in the various sublevels located in relation to the nucleus? Electrons are NOT confined to a fixed circular path, they are, however, found in definite regions of the atoms – these regions are called atomic orbital’s! Each orbital can only hold 2 electrons at a time (Pauli exclusion principle).

Within the s sublevel there is only 1 orbital (which is spherical) it is called the s orbital. Within the p sublevel there are 3 orbital’s (which are dumbbell shaped) called the px, py, pz orbital’s. 2

2

2

Within the d sublevel there are 5 orbital’s (4 of which are cloverleaf shaped) called the dxy, dxz, dyz, dx -y , dz orbital’s. Within the f sublevel there are 7 orbital’s - which are too complex to draw

D. How many electrons can go into each energy level? 2

Each orbital can hold two electrons. (2n = number of electrons per energy level) st

The 1 energy level (n=1) only has 1 sublevel called 1s. s only has 1 orbital called the s orbital, so only 2 electrons will st 2 be found in the 1 energy level. (2n = ) nd

The 2 energy level (n=2) has 2 sublevels called 2s and 2p. s only has 1 orbital called the s orbital, p has 3 orbital’s nd 2 called px, py, and pz orbital’s, so 8 electrons will be found in the 2 energy level. (2n = ) rd

The 3 energy level (n=3) has 3 sublevels called 3s, 3p, and 3d. s only has 1 orbital called the s orbital, p has 3 orbital’s rd 2 called px, py, and pz orbital’s, and d has 5 orbital’s, so 18 electrons will be found in the 3 energy level. (2n = ) th

How about the 4 energy level? It has 4 sublevels called 4s, 4p, 4d, and 4f. s only has 1 orbital, p has 3 orbital’s, d has 5 orbital’s, and f has 7 orbital’s, th 2 so 32 electrons will be found in the 4 energy level. (2n = ) Lets put it all together - Example of neon atom:

E. Spin of an electron within an orbital. Each orbital can hold two electrons, both with different spins. Clockwise spin and counterclockwise spin. Electrons fill the orbital’s one at a time with the same spin, then fill up the orbital(s) with electrons of the opposite spin.

Part 3 Electron Configurations I. Electron Configuration: It should be obvious to you now that it is very difficult to draw a representation or model of atom showing where the electrons are located, so what we do instead is write electron configurations for elements. A. Definition of electron configuration: An electron configuration is a written representation of the arrangement of electrons in an atom. When constructing orbital diagrams and electron configurations, keep the following in mind: l.

Aufbau Principle – electrons fill in order from lowest to highest energy.

2. The Pauli Exclusion Principle – An orbital can only hold two electrons. Two electrons in the same orbital must have opposite spins.

3. Hund’s rule – the lowest energy configuration for an atom is the one having the maximum number of unpaired electrons for a set of degenerate orbital’s. By convention, all unpaired electrons are represented as having parallel spins with spin “up”. How do we write an electron configuration? st A. 1 rule - electrons occupy orbital’s that require the least amount of energy for the electron to stay there. So always follow the vertical rule (Aufbau Principle):

You notice, for example, that the 4s sublevel requires less energy than the 3d sublevel; therefore, the 4s orbital is filled with electrons before any electrons enter the 3d orbital!!!! So just follow the above chart and you can’t go wrong!!!!) nd

B. 2 rule – only 2 electrons can go into any orbital, however, you must place one electron into each orbital in a nd sublevel before a 2 electron can occupy an orbital. Orbital’s with only 1 electron in the orbital are said to have an unpaired electron in them.

III. Writing Electron Configurations (3 ways): A. Orbital Notation: an unoccupied orbital is represented by a line______, with the orbital’s name written underneath the line. An orbital containing one electron is written as _____, an orbital with two electrons is written as ____. The lines are labeled with the principal quantum number and the sublevel letter.

Example: Hydrogen

____ 1s

Helium

__ 1s

Lithium ___ 1s

____ 2s

Carbon ____ ____ ____ ____ _____ 1s 2s 2px 2py 2pz Titanium

B. Electron Configuration Notation: eliminates the lines and arrows of orbital notation. Instead, the number of electrons in a sublevel is shown by adding a superscript to the sublevel designation. The superscript indicates the number of electrons present in that sublevel. Ex: Hydrogen: 1s

1

Helium: 1s

2

2

Lithium: 1s 2s

1

2

2

2

Carbon: 1s 2s 2p

Titanium

C. Short Hand or Noble Gas Notation: Use the noble gases that have complete inner energy levels and an outer energy level with complete s and p orbital’s. Use the noble gas that just precedes the element you are working with. 2

2

1

2

1

Ex: Boron is ls 2s 2p The noble gas preceding Boron is He, so the short way is [He]2s 2p . 2 2 6 2 4 2 4 Sulfur is ls 2s 2p 3s 3p Short way: [Ne]3s 3p Example: Titanium

More Practice Problems: Write electron configurations for each of the following atoms: l. boron 2. sulfur 3. vanadium Draw orbital diagrams for these: 4. sodium

5. phosphorus

Write shorthand electron configuration for the following: 6. Sr

7. Mo

Electron configurations for Ions-First, determine if the element will lose or gain electrons. Secondly, what number of electrons will be gained or lost? It is recommended that you write the e.c. for the atom and then determine what will happen. For cations (positive ions) – look at the element and decide how many electrons will be lost when it ionizes and keep that in mind when writing the E. C. The last number in the E. C. will now be LESS than what is written on your periodic table. Ex. Write the electron configuration for magnesium ion: [Ne]3s 2+ 2 2 6 (outer) electrons, so the e.c. for Mg is 1s 2s 2p Practice: l. #3

2

is for the atom. Mg is a metal and will lose its valence

2. #12 3. #19 For anions (negative ions) – look at the element and decide how many electrons that element will GAIN when it ionizes. The last number in the E. C. will be MORE than what is written on the periodic table. 2

2

4

2

4

Ex. Sulfide ion: Sulfur atom is 1s 2s 2p . Sulfur is a nonmetal with 6 valence electrons (2s and 2p ) and will gain 2 2 2 6 electrons: 1s 2s 2p is for the sulfide ion. Practice: 1. #17 2. #7 3. #16