chapter 11 modern atomic theory

Chapter 11

Modern Atomic Theory

Copyright © Houghton Mifflin Company. All rights reserved. 11 | 2

Rutherford’s Atom

• The concept of a nuclear atom (charged electrons moving around the nucleus) resulted from Ernest Rutherford’s experiments.

• Question left unanswered: how are electrons arranged and how do they move?


Rutherford’s Atom (cont.)


Electromagnetic Radiation

• Electromagnetic radiation is given off by atoms when they have been excited by any form of energy, as shown in flame tests.

• Flame tests:Sample: Color:LiClSrCl2

NaClKClCaCl2


Electromagnetic Waves

• Velocity = c = speed of light – 2.997925 x 108 m/s– All types of light energy travel at the same speed.

• Amplitude = A = measure of the intensity of the wave, i.e.“brightness”


Electromagnetic Waves (cont.)

• Wavelength = = distance between two consecutive peaks or troughs in a wave– Generally measured in nanometers (1 nm = 10-9 m)– Same distance for troughs

• Frequency = = the number of waves that pass a point in space in one second– Generally measured in Hertz (Hz), – 1 Hz = 1 wave/sec = 1 sec-1

• c = x


Types of Electromagnetic Radiation


Planck’s Discovery

• Showed that for certain applications light energy could be thought of as particles or photons


Emission of Energy by Atoms/Atomic Spectra

• Atoms that have gained extra energy release that energy in the form of light.


Atomic Spectra

• Line spectrum: very specific wavelengths of light that atoms give off or gain

• Each element has its own line spectrum, which can be used to identify that element.

• This is the basis of atomic absorption spectroscopy.


Atomic Spectra (cont.)

• Hydrogen atoms have several excited state energy levels.• Different colors are produced when the excited atoms return to the

ground state.• The line spectrum of hydrogen must be related to energy changes in the

atom.


Atomic Spectra (cont.)

• The atom is quantized, i.e. only certain energies are allowed.

Continuous levels Quantized levels


Bohr’s Model (Niel’s Bohr 1885-1911)

• Energy of the atom is quantized– An atom can only have certain specific energy

states called quantum levels or energy levels.– When an atom gains energy, an electron

“moves” to a higher quantum level.– When an atom loses energy, the electron

“moves” to a lower quantum level.– Lines in a spectrum correspond to the

difference in energy between the levels.


Bohr’s Model (cont.)

• Ground state: minimum energy of an atom

• The ground state of hydrogen corresponds to having its one electron in the n=1 level

• Excited states: energy levels higher than the ground state


Bohr’s Model (cont.)

• Distances between energy levels decrease as the energy increases– 1st energy level can hold 2 electrons, the 2nd

level 8 electrons, the 3rd 18 electrons, etc.


Problems with the Bohr Model

• Only explains hydrogen atom spectrum (and other 1-electron systems).

• Neglects interactions between electrons.

• Assumes circular or elliptical orbits for electrons (which is not true).


Wave Mechanical Model of the Atom

• Experiments later showed that electrons could be treated as waves: Louis De Broglie

• The quantum mechanical model treats electrons as waves and uses wave mathematics to calculate probability densities of finding the electron in a particular region in the atom.

– Schrödinger Wave Equation


Orbitals and Energy Levels

• Solutions to the wave equation give regions in space of high probability for finding the electron. These are called orbitals.

• Each principal energy level contains one or more sublevels. Sublevels are made up of orbitals.


Orbitals and Energy Levels (cont.)


Atomic Sublevels & Orbitals

• Each type of sublevel has a different shape each and energy.s orbital shape:

p orbital shapes:

• Each sublevel contains one or more orbitals.


Atomic Sublevels & Orbitals (cont.)


Pauli Exclusion Principle

• No orbital may have more than 2 electrons.• Electrons in the same orbital must have opposite

spins.• s sublevel holds 2 electrons (1 orbital)• p sublevel holds 6 electrons (3 orbitals)• d sublevel holds 10 electrons (5 orbitals)• f sublevel holds 14 electrons (7 orbitals)


Sublevels and Orbitals

n Sublevels : Types of Orbitals (and numbers)

1 1s(1)

2 2s(1) 2p(3)

3 3s(1) 3p(3) 3d(5)

4 4s(1) 4p(3) 4d(5) 4f(7)

For a multiple-electron atom, build-up the energy levels, filling each orbital in succession from lowest to highest.

Degenerate orbitals: orbitals with the same energye.g. Each p sublevel has 3 degenerate p orbitals


Orbital Filling

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d

7s 7pTherefore, the order of filling is:


Electron Configurations

• For a set of degenerate orbitals, fill each orbital half-way first before pairing the electrons.

• Electron configurations show how many electrons are in each sublevel of an atom – describes where the electrons are.

- The electron configuration for a ground state Li atom is:

- The electron configuration for a ground state N atom is:


Electron Configurations (cont)

What is the electron configuation of:

A. potassium

B. cobalt


Electron Configurations (cont.)

• Valence shell: highest energy level– Electrons in the valence shell are called valence

electrons.

– Core electrons: electrons not in the valence shell

– Often use symbol of previous noble gas in brackets to represent core electrons:

phosphorus: 1s22s22p6 3s23p3 =

core valence


Electron Configuration and the Periodic Table

• Elements in the same column on the periodic table have: – Similar chemical and physical properties – Similar valence shell electron configurations

• same numbers of valence electrons• same orbital types, but different energy levels

Be:

Mg:


s1

d1 d2 d3 d4 d5 d6 d7 d8 d9 d10

p1 p2 p3 p4 p5

p6

f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14

1234567

s2


Orbital diagrams

Orbital diagrams (or box diagrams) show orbitals as boxes grouped by sublevels with arrows representing the electrons.

Sulfur has an electron configuration of :

The orbital diagram for sulfur is:


Atomic Properties and the Periodic Table


Metallic Character: Metals

• Metals– Malleable & ductile

– Shiny, lustrous

– Conduct heat and electricity

– Form cations in solution

– Lose electrons in reactions - oxidized


Metallic Character: Metalloids

• Metalloids-Also known as semi-metals

-Show some metal and some nonmetal properties


Metallic Character: Nonmetals

• Nonmetals-Brittle in solid state

-Form anions and polyatomic anions

-Gain electrons in reactions - reduced


Metallic Character (cont.)

• The ease of losing an electron varies as follows:

Cs > Rb > K > Na > Li

• The same trend is seen in the Group 2 metals:


Trend in Ionization Energy

• Minimum energy needed to remove a valence electron from an atom in the gas phase:

M(g) → M+(g) + 1e-

• The lower the ionization energy, the easier it is to remove the electron.

– Metals in general have low ionization energies and nonmetals relatively high ones.

• Ionization energy decreases down the group.

– Valence electron is farther from the nucleus

• Ionization energy increases left to right across the period.


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