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Modern Atomic Theory and Chemical Bonding

Copyright Wayne Suits 2004

Rutherford’s Atom

• Rutherford showed that:– An atom is composed of a positive nucleus

surrounded by electrons (negative).

– The nucleus contains protons (positive) and neutrons (neutral).

– The nucleus is very small compared to the size of the entire atom.

• Questions left unanswered:– How are the electrons arranged around the nucleus

and what are their travel paths like?

Please recall Rutherford’s concept of the nuclear atom

The Rutherford atom – a tiny, dense nucleus that contains all of the positive charge and most of the mass of the atom surrounded by electrons

Electromagnetic Radiation (Light) as a Key to Understanding Electron Paths

• Early scientists discovered that Electromagnetic radiation (light) is given off by atoms of an element when they have been excited by some form of energy

• Furthermore, atoms of different elements give off different colors of light when they are excited.

When salts containing Li, Cu, and Na dissolved in methyl alcohol are set on fire, brilliant colors result.

LiLi

Cu

Na

Show Clip From Brown and LeMay

Spectral Analysis of Emitted Light from Excited Atoms

• When the emitted light from excited atoms was passed through a prism a curious spectrum of discrete lines of separate colors, separate energies was observed rather than a continuous spectrum of ROY G BIV.

• Furthermore, different elements show totally different line spectra.

• In fact, line spectra are used to identify the presence of different elements

Fig. 10-3, p. 265

Interpretation of Line Spectrum of Elements

• Atoms which have gained or absorbed extra energy from some excitation energy source (flame, electric discharge etc.) release that energy in the form of light

• The light atoms give off contain very specific wavelengths called a line spectrum– light given off = emission spectrum

• Each element has its own line spectrum which can be used to identify it

The Continuous Spectrum

Continuous Spectrum

Atomic Line SpectrumNa

Interpretation of Atomic Spectra

• The line spectrum must be related to energy transitions in the atom.– Absorption = atom gaining energy– Emission = atom releasing energy

• Since all samples of an element give the exact same pattern of lines, every atom of that element must have only certain, identical energy states

• The energy of an atom is quantized – limited to discrete values– If the atom could have all possible energies, then the

result would be a continuous spectrum instead of lines

Interpretation of Atomic Line Spectra in terms of Electron Paths

electrons may be thought of as traveling in concentric shells or energy levels about the nucleus.

the energy of the shells increase as one proceeds away from the nucleus.

When an atom absorbs energy, electrons are promoted from an inner, low energy, shell to an outer, higher energy shell.

Conversely, when an excited atom emits energy, electrons drop down from an excited outer higher energy shells to an inner lower energy shells

E1

E2E3

Note: E1 is less than E2 which is less than E3

n+

Atomic Absorbtions

Atomic Emissions

Electronic Transitions of the Nuclear Atom

Emitted Energy is in the form of light of specific wavelenths, specific energies, specific colored lines - the line spectrum

An excited lithium atom emitting a photon of red light to drop to a lower energy state.

(a) A sample of H atoms receives energy from an external source)

(b) The excited atoms (H) can release the excess energy by emitting photons.

When an excited H atom returns to a lower energy level, it emits a photon that

contains the energy released by the atom.

When excited hydrogen atoms return to their lowest energy state, the ground state, they emit

photons of certain energies, and thus certain colors.

Hydrogen atoms have several excited-state energy levels.

Each photon emitted by an excited hydrogen atom corresponds to a particular energy change in the hydrogen atom.

(a) Continuous energy levels. Any energy value is allowed. (b) Discrete (quantized) energy levels. Only certain energy states are allowed.

The difference between continuous and quantized energy levels can be illustrated by

comparing a flight of stairs with a ramp.

Show Clip from Brown and LeMay

Niels Hendrik David Bohr (1885-1962)

Source: Emilio Segre Visual Archives

The Bohr model of the hydrogen atom represented the electron as restricted to

certain circular orbits around the nucleus.

I.  Atomic Structure (Review)

Atoms are primarily composed of 3 sub atomic

particles. Sub Atomic Particle

Charge Mass(amu)

proton (p) +1 1

neutron(n) 0 1

electron(e-

)-1 0

B. An atom is neutral if

# e-’s = # p’s.

• If a neutral atom gains extra electron(s) then it becomes a negatively charged species called an anion.

• If a neutral atom loses electron(s) then it becomes a positively charged species called a cation.

An atom is completely characterized by two

numbers; the atomic #(Z) and the

atomic mass # (A).1) Atomic # (Z) - the # of

protons in the nucleus - responsible for identity of the element.

2) Mass # (A)- the total # of protons plus neutrons.

Representing Atoms of an Element

• An atom may be represented as its Symbol preceded by its subscripted atomic number, Z, and its superscripted atomic mass number, A.

SymbolZ

A

In the case of the element Carbon

C6

12

Arrangement of the subatomic particles within the atom

At the center is the nucleus which contains the protons and neutrons.

electrons may be thought of as traveling in concentric shells or energy levels about the nucleus.

the energy of the shells increase as one proceeds away from the nucleus.

There is a max. # of e-’s that can be accommodated

in each shell.Shell Max.

capacity

e-’s1 2

2 8

3 8 or 18

4 32

Shell diagram for neutral atom of Phosphorus (P)

15 p16 n

Further development of atomic model.

•Each shell is composed of 1 or more subshells.

•Each shell has as many subshells as its own number.

–1st shell has 1 subshell.–2nd shell has 2 subshells.–3rd shell has 3 subshells.– 4th shell has 4 subshells.

There are only four different kinds of subshells.These subshells are labeled, in order of increasing energy, by the letters s, p, d & f.Each subshell can accomodate a different # of e-’s

Energy

Increases

subshell e- capacity

s 2

p 6

d 10

f 14

Thus the total capacity of shell is distributed amongst its subshells.

8

Shell/subshell diagram for phosphorus

The ground state electron configuration for phosphorus:

15 p16 n

1s

1s2

2s

2s2

2p

2p6

3s

3s2

3p

3p3

Atomic subshells in order of increasing energy,

filling order. 4f _____ 4d _____ 4p _____ 3d _____ 4s _____ 3p _____ 3s _____ 2p _____ 2s _____ 1s _____

NOTE: Although the 4th shell is higher in energy than the 3rd shell, not all subshells of the 4th shell are higher in energy than all subshells of the 3rd shell. In fact, the highest subshell of the 3rd shell (3d) is higher in energy than the lowest subshell of the 4th shell (4s)

Further development of atomic model

Our latest model of the atom identifies electrons as dots traveling about the nucleus in concentric subshells. The truth is that we can never know the exact position of an electron at any point in time. In 1926, however, Erwin Schrödinger (of University of Zurich) developed a theory known as Quantum mechanics in which he worked out a mathematical expression to describe the motion of an electron in terms of its energy.

Further development of atomic model

These mathematical expressions are called wave equations since they are based upon the concept that e-’s show properties not only of particles but also of electromagnetic waves. These wave equations have a series of solutions called wave functions which allow us to predict the volume of space around a nucleus in which there is a high probability of finding a particular e-. This volume of space in which an electron is most likely to be found is called an orbital.

Now, to fully develop our theory of atomic structure we must

understand that the subshells (s, p, d, f)

of our earlier atomic model consist of orbitals that are not

all concentric in shape. Furthermore, any one orbital can

only accommodate 2 e-’s. Consequently, the number of orbitals that comprise a subshell can easily be

calculated by simply dividing the subshell capacity by 2.

Number of Orbitals in each Subshell

• Any s subshell has a capacity of 2 e-’s– The number of orbitals that comprise any s

subshell is 1.• Any p subshell has a capacity of 6 e-’s

– The number of orbitals that comprise any p subshell is 3.

• Any d subshell has a capacity of 10 e-’s– The number of orbitals that comprise any d

subshell is 5.• Any f subshell has a capacity of 14 e-’s

– The number of orbitals that comprise any f subshell is 7.

Orbitals (s p d + f)

• All orbitals of the same kind have the same 3 dimensional shape but different sizes. The size increases with the energy level. All s subshells consist of one s orbital that is spherically symmetrical about the nucleus. An s orbital can accommodate 2 e- This accounts for the 2e- capacity of the s subshell

s Orbitals

1s ORBITAL 2s ORBITAL

Each p subshell actually consists of a set of three p orbitals of equal energy;

px py pz. • Each of the three p orbitals is dumbbell shaped and all are oriented in space perpendicular to one another.

• The max. capacity of each p orbital is 2e-.

• This accounts for the total capacity of the p subshell as being 6 e-’s.

Each p subshell consists of a set of three p orbitals of equal energy, px py pz

Shown together the three p orbitals look like

this:

The d subshell actually consists of a set of five d orbitals of equal energy. Each d orbital can hold a maximum of 2e-. This accounts for the total capacity of the d subshell as being 10 e-’s. The d orbitals do not play as important a role in the chemistry that we will be discussing therefore their shapes and names need not

be memorized.

Electron Spins

• Electrons spin on their axis

• Physics tells us that any charged species that spins, generates a magnetic moment. That is to say, it acts like a tiny bar magnet with a North and a South Pole.

• Furthermore, the “Right Hand Rule” tells us that if we wrap the fingers of our right hand around the spinning species, in the direction of the spin, then our thumb will be pointing to the magnetic north.

N

S N

S

Represeanting Electrons

• Therefore, because of their magnetic moments, we generally represent electrons using a single barbed arrow. The tip of the arrow points to the magnetic north of the electron.

Atomic Orbitals in order of Increasing Energy

3d__ 3d__ 3d__ 3d__ 3d__

4s___

3px__ 3py__ 3pz__

3s ___

2px__ 2py__ 2pz__

2s___

1s___

EN

ER

GY

Ground - state electron configurations

•This refers to the lowest energy arrangement of e-’s in orbitals about the nucleus.

• To obtain this ground - state electron configuration electrons are assigned to the orbitals of the previous slide according to the three rules.

Rules for Filling Orbitals

• Always fill the lowest energy orbitals first.

• The two electrons that occupy any orbital must have opposite spins.

• When filling orbitals of equal energy (those of the p,d,or f subshells) put one electron in each orbital with their spins parallel until all are half filled, then go back and pair them.

Orbital Electron Configurations

• Write the orbital electron configuration for P

• Write the orbital electron configuration for O

1s2 2s2 2px2 3s2 3px

1 3py1 3pz

12py2 2pz

2

1s2 2s2 2px2 2py1 2pz1

When liquid oxygen is poured between the poles of a

magnet, it “sticks” until it boils away. Oxygen is magnetic because of its unpaired

electrons

Source: Donald Clegg

Using the periodic table to write electron configurations

The P.T. is arranged such that each horizontal row (period) represents the filling of orbitals in their proper order.

More information from the Periodic Table

•The term valence electron refers to the # of e-’s in the outermost energy level or shell of an atom.

For all main group elements the # of the column (family) of the

Periodic Table in which the symbol for the element occurs = the # of

valence electrons. Element Number of

Valence e- s

Na 1

B 3

Cl 7

Lewis Structures

These are shorthand techniques for emphasizing the outer shell or valence e-’s of an atom by representing an atom as its symbol surrounded by its valence e-’s, the e-’s in the atoms outermost shell. Note that the symbol of the element represents the nucleus plus all inner shell e-’s.

Write Lewis dot structures for carbon,

hydrogen, oxygen, nitrogen and chlorine.

carbon C

hydrogen H

oxygen O

nitrogen N

chlorine Cl

Why do atoms react together to form

compounds? Atoms react with one another to form compounds in an attempt to achieve the e- configuration of their nearest noble gas neighbor (family 8). The reason for this is that the e- configuration of the noble gases represents an extremely stable situation.

There are two ways in which atoms can bond together so

as to achieve the e- configuration of their

nearest noble gas neighbor. 1. They can loose or gain the

necessary e-’s and thereby become ions and ultimately form ionic bonds.

2. Two or more atoms can share e-’s and form covalent bonds.

Ionic Bonds

These are formed when ions anions/cations of opposite charge come together. Generally ionic compounds are formed between metals (left of step) and nonmetals (right of step).

Consider the formation of the ionic compound magnesium

bromide.

Magnesium (Mg ) could achieve the e- config. of Neon by loosing 2e- .

Mg +2 + 2e-

Mg

isoelectronic with Ne

Bromine could achieve the e- config. of krypton by gaining

one e-.

Consequently one magnesium combines with two bromine atoms to form MgBr2.

Br

+

1e-

Br Note

:

Kr

Note all atoms in MgBr2

are isoelectronic with their nearest noble gas

neighbor.

Mg+2 + Br-1 = MgBr2

Covalent Bond • A covalent bond results from the sharing of an electron pair between two atoms.– Whenever two atoms share a pair of e-’s, it is as if each member of the bonded pair of atoms has gained an extra electron.

– As atoms bond together to become isoelectronic with their nearest noble gas neighbors, covalent bonds generally occur when two or more nonmetallic elements (right of step) bond together because the nearest noble gas neighbors for these elements lies ahead of them. Consequently, they all need to gain electrons to become isoelectronic with their nearest noble gas neighbors.

Can become isoelectronic with Ne by

gaining 4e-

Lewis Structure

for covalent

molecule of CH4

Kekulé structure

for covalently bonded molecule

Can become isoelectronic with He by gaining 1e-

How many hydrogen atoms bond to one carbon atom?

C + H4 C H

H

H

H C H

H

H

H

Let’s Look at the Water Molecule

O + H2 O

H

HIn the water molecule each oxygen is isoelectronic with:

Neon

In the water molecule each hydrogen is isoelectronic with: Helium

Kekulé structure for water molecule.

O

H H

Now let’s build the Ammonia Molecule

N + H3 NH

H

H or N H

H

H

Ammonia is composed of Nitrogen and Hydrogen

The Covalent Bond and Electronegativity

• The sharing of an e- pair between two atoms may be equal .– If this is the case then the resulting covalent bond is a nonpolar covalent bond.

• If, on the other hand the sharing is unequal then a polar covalent bond results.

The reason for this variance in bond polarity is due to the fact that different elements have different tendencies to

attract to themselves extra electrons. In other words,

each element has a different electronegativity

Electronegativity

The tendency of an atom, when in combination with other atoms, to attract to itself the bonded (extra)

e-’s.

Electronegativity values increase from left to right across any horizontal row (period) of the P.T. and they decrease going down any vertical column (family) of the P.T.

Consequently the most electronegative elements are

N, O, F, Cl, Br

Electronegativity values for selected elements.

• If two atoms are covalently bonded and one has a high electronegativity and the other has a low electronegative then the electron pair comprising that bond is not shared equally but spends more of its time closer to the more electronegative atom. The immediate result of this unequal sharing is that the more electronegative atom gains a partial negative charge (-) while the less electronegative element gains a partial positive charge ( +). This type of bond is called a polar covalent bond.

The degree to which a covalent bond is polarized is indicated

by the electronegativity difference between the two bonded atoms. Refer to next slide for electronegativity

values of elements. • If the electronegativity difference is greater than .5 but less than 2.0 then the covalent bond is polar.

• If the electronegativity difference is less than .5 then the covalent bond is nonpolar.

Polar Covalent Bonds in H2O

O

H H + +

-

Electronegativity Difference Between Oxygen and Hydrogen is:

1.4

A molecule typical of those found in petroleum. The bonds are not

polar.

Electronegativity Difference Between Carbon and Hydrogen is:

0.4

Electronegativity values for selected elements.

Ionic Bond and Electronegativity

•Consideration of electronegativity can demonstrate that ionic bonds are nothing more than an extreme case of a polar covalent bond. In fact…

•if the electronegativity difference between two atoms is greater than 2.0, then any bond between these two atoms would be ionic.

Molecular Polarity

If a molecule contains polar bonds, and if those polar bonds are located such that the + charges are at one end of the molecule and the - charges are at the other end, then the molecule is a polar molecule. The measure of molecular polarity is a quantity called the dipole moment (D).

Like Dissolves Like

• Polar molecules dissolve in Polar Solvents

• Nonpolar molecules dissolve in nonpolar solvents

• Polar molecules do not dissolve in nonpolar solvent

• Nonpolar molecules do not dissolve in polar solvents

: An oil layer floating on water. The oil is nonpolar and the water is polar

Polar water molecules interact with the positive and negative ions of a salt. Ionicly bonded materials are

the extreme case of polar substances