modern atomic theory a.k.a. the electron...

MODERN ATOMIC THEORY

A.K.A.

THE ELECTRON CHAPTER W O R L D O F C H E M I S T RY C H A P T E R 1 1

WHERE DID WE LEAVE OFF?

Ernest Rutherford (1906)

Gold Leaf Experiment Results:

• Atom is mostly empty space.

• Center of atom is positively-charged and

concentrated.

Suggested that electrons traveled around the nucleus.

• Couldn’t explain why the negative electrons

weren’t attracted to positive nucleus, causing

the atom to collapse

MODERN ATOMIC THEORY

Understanding the nature of light and how it

transmits energy allowed us to understand the

structure of the atom.

In increasing energy, ROYGBV

ELECTROMAGNETIC SPECTRUM

ELECTROMAGNETIC

RADIATION

Light exhibits properties of waves as well as

particles. This is known as wave-particle duality

of light.

Most subatomic particles behave as PARTICLES and

obey the physics of waves.

ELECTROMAGNETIC

RADIATION

Wavelength

represented by

“lambda”

Waves have a frequency = the number of wave cycles per second (Hz or s-1)

ELECTROMAGNETIC

RADIATION

Long wavelength low frequency

Short wavelength high frequency

increasing

frequency

increasing

wavelength

ELECTROMAGNETIC

SPECTRUM

Quick Check: X-rays vs Radio waves

Which has the lower frequency?

Which has the shorter wavelength?

Which has more energy?

ELECTROMAGNETIC

SPECTRUM

Radio waves

X rays

X rays

EXCITED GASES

& ATOMIC STRUCTURE

Electricity on.

Excited electrons.

Gas glows.

EMISSIONS OF ENERGY

BY ATOMS

When all e– are in lowest possible energy state, an atom is

in the ground state.

If “right” amount of energy is absorbed by an e–, it can

“jump” to a higher energy level. This is an unstable,

momentary condition called the excited state.

When e– falls back to a lower-energy, more stable orbital (it

might be the orbital it started out in, but it might not), atom

releases the “right” amount of energy as light.

Any-old-value of energy to be absorbed or released is NOT

OK. This explains the lines of color in an emission spectrum.

EMISSIONS OF ENERGY

BY ATOMS

EVERY ELEMENT GIVES A UNIQUE

EMISSION SPECTRUM

ENERGY LEVELS

OF HYDROGEN

• When we add a lot of energy to a sample of H atoms

certain types of photons are produced.

• Photons are “bundles” of energy that make up light.

• We see only selected colors; this means that the

hydrogen atom must have certain discrete energy

levels.

ENERGY LEVELS

OF HYDROGEN

Energy levels are quantized, i.e. only certain values are

allowed. Excited hydrogen atoms always emit photons with

the same discrete colors (wavelengths).

Bohr’s greatest contribution to science

was in building a simple model of the

atom. It was based on an understanding

of the LINE EMISSION SPECTRA

of excited atoms.

• Problem is that the model only

works for Hydrogen.

Niels Bohr

(1885-1962)

ATOMIC LINE EMISSION SPECTRA

AND NIELS BOHR

ATOMIC SPECTRA AND BOHR

Bohr said classical view is wrong!

• e- can only exist in certain discrete orbits

• e- only possess certain amount of energy

called quanta

Bohr received Nobel prize for atomic structure in 1922.

One view of atomic structure in early 20th century was that an electron (e-) traveled about the nucleus in an orbit.

Planetary Model

EARLY MODEL OF ATOM

PROVEN TO BE

WRONG!!!!

B O Z E M A N E M I S S I O N A N D A B S O R P T I O N S P E C T R A ( 5 : 1 7 M I N )

CRASH COURSE CHEMISTRY:

THE ELECTRON

Schrodinger applied idea of e- behaving as a

wave to the problem of electrons in atoms.

He developed the WAVE EQUATION.

Solution gives set of math expressions

called WAVE FUNCTIONS.

Each describes an allowed energy state of

an e-

E. Schrodinger

1887-1961

QUANTUM OR WAVE MECHANICS

(CURRENT MODEL)

ELECTRONS OCCUPY

ORBITALS

What is an orbital?

• NOTHING like an orbit

• probability map (based on mathematical

calculations)

• the distance the electron will most likely be from

the nucleus

• does not tell us when the electron occupies a

certain point in space or how it moves

Electrons in atoms are

arranged as LEVELS

SUBLEVELS

ORBITALS

ARRANGEMENT OF

ELECTRONS IN ATOMS

• Each energy level has a number called the

PRINCIPAL ENERGY LEVEL, n

• Currently n can be 1 thru 7, because there are

7 periods on the periodic table.

ENERGY LEVELS

ENERGY LEVELS

n=2

n=3

n=4

n=1

n=6 n=7

n=5

TYPES OF SUBLEVELS

• Sublevel(s) make up energy levels.

• The most probable area to find these electrons

takes on a shape. They are named s, p, d, and f.

• Orbitals make up sublevels. They can hold up to

2 e- with opposite spins.

• An electron can spin in one of two directions; we represent as or

• Two electrons must have opposite spins to occupy the same orbital.

PAULI EXCLUSION PRINCIPLE

2

s 1 2 8

s 1 2

s 1 2

p 3 6 18

p 3 6 d 5 10

32

REL ATIVE S IZES OF THE SP HERICAL

1 S , 2 S , AND 3 S ORBITAL S OF

HYDROGEN.

The three p orbitals lie 90o apart in space.

P ORBITALS

THE SHAPES AND LABELS

OF THE FIVE 3D ORBITALS.

F ORBITALS

• f sublevel has 7

orbitals

LABEL THE SUBLEVELS

n=2

n=4

1s

4s

3s

2s

4p

4d

4f

3d

3p

2p

DIAGONAL RULE

• Must be able to write it for the test! This will be question #1 ! Without it, you will not get correct answers!

• The diagonal rule is a memory device that helps you remember the filling order of the orbitals from lowest energy to highest energy. This is called the Aufbau principle.

• Aufbau means building up or construction in German!

s

s 3p 3d

s 2p

s 4p 4d 4f

s 5p 5d 5f 5g?

s 6p 6d 6f 6g? 6h?

s 7p 7d 7f 7g? 7h? 7i?

1

2

3

4

5

6

7

By this point, we are past the current periodic table

so we can stop.

DIAGONAL RULE Steps:

1. Write the energy levels top to

bottom.

2. Write the orbitals in s, p, d, f

order. Write the same number of

orbitals as the energy level.

3. Draw diagonal lines from the top

right to the bottom left.

4. To get the correct order, follow

the arrows!

W H Y A R E D A N D F O R B I T A L S A LWAY S I N

L OW E R E N E RG Y L E V E L S ?

• d and f orbitals require LARGE amounts of energy

• It’s better (lower in energy) to skip a sublevel that

requires a large amount of energy (d and f orbitals)

for one in a higher level but lower energy.

This is the reason for the diagonal rule!

BE SURE TO FOLLOW THE ARROWS IN ORDER!

ELECTRON CONFIGURATIONS

… it is a list of all the electrons in an atom (or ion)

1. Locate element on periodic table to determine

# of electrons (atomic #)

2. Fill orbitals in proper order using diagonal rule.

3. Check total number of electrons.


2p4 Energy Level

Sublevel

Number of

electrons in the

sublevel

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc.

LET’S TRY IT!

Write the electron configuration for the following elements:

1H

3Li

7N

10Ne

1s1

1s2 2s1

1s2 2s2 2p3

1s2 2s2 2p6

LET’S TRY IT!

19K

30Zn

82Pb 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6

…5s2 4d10 5p6 6s2 4f14 5d10 6p2

1s2 2s2 2p6 3s2 3p6 4s2 3d10

1s2 2s2 2p6 3s2 3p6 4s1

SHORTHAND NOTATION

• A way of abbreviating long electron

configurations

• Since we are only concerned about the

outermost electrons, we can skip to places we

know are completely full (noble gases), and

then finish the configuration

• Step 1: Find the closest noble gas to the atom (or ion), WITHOUT GOING OVER the number of electrons in the atom (or ion). Write the noble gas in brackets [ ].

• Step 2: Find where to resume by finding the next energy level.

• Step 3: Resume the configuration until it’s finished.

SHORTHAND NOTATION

• Longhand is 1s2 2s2 2p6 3s2 3p5

• You can abbreviate the first 10 electrons with a noble gas, neon.

[Ne] replaces 1s2 2s2 2p6

• The next energy level after Neon is 3

• So you start at level 3 on the diagonal rule (all levels start with s) and finish the configuration by adding 7 more electrons to bring the total to 17

[Ne] 3s2 3p5

EX) CHLORINE

PRACTICE SHORTHAND

NOTATION

Write the shorthand notation for each of the following atoms:

N

K

Ca

I

Bi

[He] 2s2 2p3

[Ar] 4s1

[Ar] 4s2

[Kr] 5s2 4d10 5p5

[Xe] 6s2 4f14 5d10 6p3

TRY THESE!

Write the shorthand notation for:

Cu

W

Au

[Ar] 4s2 3d9

[Xe] 6s2 4f14 5d4

[Xe] 6s2 4f14 5d9


We need electron configurations so that we can

determine the number of electrons in the outermost

energy level (valence shell). These are called valence

electrons. The number of valence electrons determines

bonding in molecules and compounds.

VALENCE ELECTRONS

Electrons are divided between core and valence electrons (involved in bonding).

B 1s2 2s2 2p1

Core = [He] , valence = 2s2 2p1

Br [Ar] 3d10 4s2 4p5

Core = [Ar] 3d10 , valence = 4s2 4p5

OCTET RULE

The maximum number of valence electrons is

eight.

Atoms like to either empty or fill their

outermost level to achieve a full valence shell.

RULES OF THE GAME

Number of valence electrons for main group

atoms = Group number (for 1A – 8A groups)

• Ex) Mg is in group 2A and has 2 valence e-

• How many valence electrons do krypton

atom have?

IONS AND ELECTRONS

• Recall… negative ions have gained electrons, positive

ions have lost electrons.

• The electrons that are lost or gained should be

added/removed from the highest energy level

(not the highest orbital in energy!)

Skip this section….

Ex) Sn Atom: [Kr] 5s2 4d10 5p2

Sn+4 ion: [Kr] 4d10

Sn+2 ion: [Kr] 5s2 4d10

Note that the electrons came out of the highest energy level, not the

highest energy orbital!

IONS AND ELECTRONS

Atom lost 4e-

from 5s and 5p

Atom lost 2e-

from 5p


• Ex) Bromine

Atom: [Ar] 4s2 3d10 4p5

Br- ion: [Ar] 4s2 3d10 4p6

Note that the electrons went into the highest energy level, not the

highest energy orbital!

IONS AND ELECTRONS

Atom gained 1e-

into 4p


TRY SOME IONS!

• Write the shorthand

notation for these:

Br-

Ba+2

Al+3

• Write the longhand

notation for these:

F-

Li+

Mg+2

1s2 2s2 2p6

1s2

1s2 2s2 2p6

[Kr]

[Xe]

[Ne]


ORBITALS AND THE

PERIODIC TABLE

s block p block d block

f block

Ex) Consider EC of the halogens.

9 F: 1s2 2s2 2p5

17 Cl: 1s2 2s2 2p6 3s2 3p5

35 Br: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5

53 I: … 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p5

85 At: …………………………...…. ?

The last orbital to fill is the same for

elements in the same group but

different principal energy levels.

EXPLAIN THIS COMIC:

“WHO” IS THIS NEW KID?

ORBITAL DIAGRAMS

• Graphical representation of an electron

configuration

• One arrow represents one electron

• Shows spin and which orbital within a sublevel

• Same rules as before (Diagonal Rule)

ORBITAL DIAGRAMS

One additional rule: Hund’s Rule

• In orbitals of EQUAL ENERGY,

place one electron in each orbital

before making any pairs

• All single electrons must spin the same

way

LITHIUM

Atomic number = 3

1s22s1 3 total e-

CARBON

Atomic number = 6

1s2 2s2 2p2 6 total e-

Here we see for the first time

HUND’S RULE. When placing

electrons in a set of orbitals having

the same energy, we place them singly

as long as possible.

DRAW THESE ORBITAL

DIAGRAMS!

• Oxygen (O)

• Iron (Fe)

• Chromium (Cr)

• Mercury (Hg)

GENERAL PERIODIC TRENDS

As you go down a group:

Larger orbitals

Electrons held less tightly

As you go across a period:

Electrons held more tightly

Higher effective nuclear charge

TRENDS IN ATOMIC SIZE measured in a tomic radi i

As you go down a group… atoms get bigger. Why?

• electrons are added further from the nucleus and there is less attraction

• additional energy levels • shielding effect- each additional energy

level “shields” the electrons from being pulled in toward the nucleus.

As you go across a period…

atoms get smaller. Why?

• increased nuclear charge (positive charge from

protons) in same principal energy level

• each added electron feels a greater and greater +

charge because the protons are pulling in the

same direction, where the electrons are

scattered.

Small

TRENDS IN ATOMIC SIZE

TREND IN ATOMIC SIZE

WHICH IS BIGGER?

• Na or K ?

• Na or Mg ?

• Al or I ?

• CATIONS are SMALLER than the atoms

from which they come.

• The electron/proton attraction has gone UP

and so size DECREASES.

Li,152 pm 3e an 3p

+ Forming a

cation.

ION SIZES

Li+,78 pm 2e and 3p


ION SIZES

• ANIONS are LARGER than the atoms from which

they come.

• The electron/proton attraction has gone DOWN and

so size INCREASES.

• Trends in ion sizes are the same as atom sizes.

Forming

an anion.


TRENDS IN ION SIZES

Figure 8.13


WHICH IS BIGGER?

• Cl or Cl- ?

• K+ or K ?

• Ca or Ca+2 ?

• I- or Br- ?


This is called the FIRST ionization

energy because we removed only the

OUTERMOST electron

This is the SECOND IE.

IE = energy required to remove an electron from an

atom (in the gas phase).

IONIZATION ENERGY

Mg

• IE increases across a period because

the positive charge increases.

• IE decreases down a group because

size increases (Shielding Effect)

TRENDS IN

IONIZATION ENERGY

WHICH HAS A HIGHER 1 ST

IONIZATION ENERGY?

• Mg or Ca ?

• Al or S ?

• Cs or Ba ?

ELECTRONEGATIVITY

… is a measure of the ability of an atom in a

molecule to attract electrons to itself.

Concept proposed by

Linus Pauling

1901-1994

PERIODIC TRENDS:

ELECTRONEGATIVITY

• Electronegativity decreases down a group of elements

because atoms with more energy levels attract electrons

less (more shielding).

• Electronegativity increases across a period of elements

because more protons, while the energy levels are the

same, means atoms can better attract electrons.

WHICH IS MORE

ELECTRONEGATIVE?

• F or Cl ?

• Na or K ?

• Sn or I ?

modern atomic theory a.k.a. the electron...

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