Unit 4.1

Arrangement of electrons in an Atom

Rutherford’s atom

• Rutherford’s model of the atom was incomplete (doesn’t explain where the electrons are)

The Wave-like Electron

Louis deBroglie

The electron propagates through space as an energy wave. To understand the

atom, one must understand the behavior of electromagnetic waves.

c = νλ c = speed of light, a constant (3.00 x 108 m/s) ν = frequency, in units of hertz (hz, sec-1) λ = wavelength, in meters

Electromagnetic radiation(a form of energy) propagates through space as a wave moving at the speed of light.

E = hν E = Energy, in units of Joules (kg·m2/s2)

h = Planck’s constant (6.626 x 10-34 J·s) ν = frequency, in units of hertz (hz, sec-1)

The energy (E ) of electromagnetic radiation is directly proportional to the frequency (ν) of the radiation.

Long Wavelength

= Low Frequency

= Low ENERGY

Short Wavelength

= High

Frequency =

High ENERGY

Wavelength Table

The Electromagnetic Spectrum

Wave-Particle Duality JJ Thomson won the Nobel prize for describing the electron as a particle.

His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron.

The electron is a particle! The electron is

an energy wave!

Niels Bohr’s Atom

• Electrons orbit the nucleus in orbits, like a solar system.

Planetary Model

Electrons cannot exist between orbits (energy is quantized)

Bohr’s atom • Electrons closest to the

nucleus are lowest in energy.

• Ground state- electrons are

in the lowest energy level possible

• If energy is put into the

atom, the electrons will jump up in energy- move away from the nucleus (excited state).

Bohr’s atom • Excited electrons naturally

go back to ground state. In order to do this, energy must leave the atom. Because energy is quantized in an atom, the amount of energy that leaves is the difference in energy between orbits If this energy is in the visible light range, we will see certain colors (line emission spectrums)

This produces bands of light with definite wavelengths.

Electron transitions involve jumps of definite amounts of energy.

…produces a “bright line” spectrum

Spectroscopic analysis of the hydrogen spectrum…

Flame Tests

strontium sodium lithium potassium copper

Many elements give off characteristic light which can be used to help identify them.

Electron orbitals

Unit 4.2

The Bohr Model of the Atom

Neils Bohr

I pictured electrons orbiting the nucleus much like planets orbiting the sun. But I was wrong! They’re more like bees around a hive.

Quantum Mechanical Model of the Atom

Mathematical laws can identify the regions outside of the nucleus where electrons are most likely to be found. These laws are beyond the scope of this class…

Erwin Schrödinger 1887-1961 • Born in Vienna, Austria • University of Berlin professor • Nobel Prize 1933 • Based on waves of light and probability.

• Quantum Mechanical Model –

Heisenberg Uncertainty Principle

You can find out where the electron is, but not where it is going.

OR…

You can find out where the electron is going, but not where it is!

“One cannot simultaneously determine both the position and momentum of an electron.”

Werner Heisenberg

Energy levels (shells)- n

• n = 1, 2, 3,.. • n= 1 is lowest in energy and closest to the

nucleus • As n increases, the distance from the nucleus

increases

• ) ) ) ) ) n=1 2 3 4 5

Orbital shapes also referred to as sublevels or subshells are defined as the surface that contains 90% of the total electron probability. Shapes or sublevels are located inside a shell or energy level In the first shell there is one shape In the second shell there are two shapes In the third shell there are three shapes, etc.

Sublevels/subshells/shapes

The s shape is a spherical shape centered around the origin of the three axes in space. The s shape has one orbital. Each orbital can hold 2 electrons

s Orbital shape

There are three dumbbell-shaped p orbitals each is assigned to its own axis (x, y and z) in space. Each orbital can hold 2 electrons so the p shape can hold a total of 6 electrons

p orbital shape

Things get a bit more complicated with the five d orbitals that are found in the d sublevels. To remember the shapes, think of “double dumbells” …and a “dumbell with a donut”! Each orbital can hold 2 electrons so the d shape can hold a total of 10 electrons.

d orbital shapes

f orbital shape

There are 7 f orbitals that are found in the f sublevel.

Each orbital can hold 2 electrons so the f shape

can hold a total of 14 electrons.

Shape of f orbitals

Energy Levels, Orbitals, Electrons Energy Level (n)

Orbital type in the

energy level (types = n)

Number of Orbitals

(n2)

Number of Electrons

Number of electrons per Energy level

(2n2)

1 s 1 2 2 2 s

p 1 3

2 6

8

3 s p d

1 3 5

2 6 10

18

4 s p d f

1 3 5 7

2 6 10 14

32

Electron Spin Electron spin describes the behavior (direction of spin on internal axis) of an electron within a magnetic field. Possibilities for electron spin:

12

−12

+

Clockwise or counter clockwise

Electron configuration

Unit 4.3

Electron configuration

Tells where all the electrons are in an atom – Orbital notation (uses arrows to represent

electrons) – Electron configuration (uses superscripts to

represent electrons) – Noble gas configuration (uses a noble gas to

represent core electrons and superscripts for electrons)

Pauli Exclusion Principle

Two electrons occupying the same orbital must have opposite spins

Wolfgang Pauli

Aufbau principle

• Electrons want to have the lowest energy possible.

• First energy level fills up before 2nd , then 3rd, etc. • Within a shell s orbital fills up, then p orbitals, d

orbitals, and last f orbitals • 1s 2s 2p 3s 3p 4s 3d (refer to handout)

• 4s fills before 3d because it is actually lower in

energy

Hund’s rule

• Electrons want to be in their own orbital before they pair up (without breaking aufbau’s principle)

• All electrons in the same subshell must have the

same spin (be in their own orbital before pairing up)

_↓_ __ __ __ __ (orbital notation) 1s 2s 2p

Practice

• Write the orbital notation for oxygen • Oxygen has 8 electrons

__ __ __ __ __ 1s 2s 2p Electron configuration: 1s2 2s2 2p4

Periodic table and e- configuration

• The periodic table is set up to show the electron filling order (lowest to highest energy)

• We can use the periodic table to help write electron configurations and noble gas configurations

Periodic table

• Groups 1-2: s block (2 columns) • Groups 13-18: p block (6 columns) • Groups 3-12: d block (10 columns) • Lanthanides and actinides: f block (14 columns)

• Periods tell the highest shell that is filled

There are some exceptions, but you will not have to

worry about them.

Orbital filling table

Practice • Write the electron configuration for

– carbon: • 1s2 2s2 2p2

– Iron:

• 1s2 2s2 2p6 3s2 3p6 4s2 3d6

– Lead:

• 1s2 2s2 2p6 3s2 3p6 4s2 3d104p6 5s2 4d10 5p6 6s2 5d10 4f14 6p2

Electron configuration gets tedious for bigger atoms so we simplify

with noble gas configurations

Valence electrons

• Valence electrons: electrons in the highest shell that are involved in bonding

• Inner shell electrons (core electrons):

electrons not in the highest shell and don’t bond with other atoms

Valence electrons – carbon:

• 1s2 2s2 2p2

• valence electrons: 4 – Iron:

• 1s2 2s2 2p6 3s2 3p6 4s2 3d6

• valence electrons: 2

– Lead:

• 1s2 2s2 2p6 3s2 3p6 4s2 3d104p6 5s2 4d10 5p6 6s2 5d10 4f14 6p2

valence electrons: 4

Noble gas configuration • Noble gas replaces most of the core electrons • Find the noble gas before the element and follow the periodic table

to finish – carbon:

• 1s2 2s2 2p2

[He] 2s2 2p2 – Iron:

• 1s2 2s2 2p6 3s2 3p6 4s2 3d6

[Ar] 4s2 3d6 – Lead:

• 1s2 2s2 2p6 3s2 3p6 4s2 3d104p6 5s2 4d10 5p6 6s2 5d10 4f14 6p2

[Xe] 6s2 5d10 4f14 6p2

Element Configuration notation

Orbital notation Noble gas notation

Lithium 1s22s1 ____ ____ ____ ____ ____ 1s 2s 2p

[He]2s1

Beryllium 1s22s2 ____ ____ ____ ____ ____ 1s 2s 2p

[He]2s2

Boron 1s22s2p1 ____ ____ ____ ____ ____ 1s 2s 2p

[He]2s2p1

Carbon 1s22s2p2 ____ ____ ____ ____ ____ 1s 2s 2p

[He]2s2p2

Nitrogen 1s22s2p3 ____ ____ ____ ____ ____ 1s 2s 2p

[He]2s2p3

Oxygen 1s22s2p4 ____ ____ ____ ____ ____ 1s 2s 2p

[He]2s2p4

Fluorine 1s22s2p5 ____ ____ ____ ____ ____ 1s 2s 2p

[He]2s2p5

Neon 1s22s2p6 ____ ____ ____ ____ ____ 1s 2s 2p

[He]2s2p6

Periodic properties

Unit 4.4

Valence electrons

• Valence electrons: electrons in the highest shell that are involved in bonding

• Inner shell electrons (core electrons):

electrons not in the highest shell and don’t bond with other atoms

Valence electrons – carbon:

• 1s2 2s2 2p2

• valence electrons: 4 – Iron:

• 1s2 2s2 2p6 3s2 3p6 4s2 3d6

• valence electrons: 2

– Lead:

• 1s2 2s2 2p6 3s2 3p6 4s2 3d104p6 5s2 4d10 5p6 6s2 5d10 4f14 6p2

valence electrons: 4

Patterns in the periodic table

• S and P block are the main group elements

• Valence electrons – Group 1: 1 valence electron – Group 2: 2 valence electrons – Group 13: 3 valence electrons – Group 14: 4 valence electrons – Group 15: 5 valence electrons – Groups 16: 6 valence electrons – Group 17: 7 valence electrons – Group 18: 8 (valence shell is full)

Atomic Radii

• Size of an atom gets smaller as you go across a period

-electrons are added to the same energy level. Protons are being added to the nucleus pulling the electrons closer to the nucleus. More protons that pull the electrons in and all have the same valence shell

• Size of an atom gets bigger as you go down a group – A shell is added to the atom

practice

• Mg, Cl, Na, P Which is the largest atom? Answer: Na Ca, Be, Ba, Sr, Which has the largest radius? Answer: Ba

• Smaller atoms have an easier time picking up an electron and a harder time giving an electron up.

• Bigger atoms have an easier time giving an electron away, but a harder time picking an electron up.

Ions-an atom with a charge

• Metals (bigger atoms) tend to lose their electrons and become cations Cations- a positive ion When an atom loses electrons it become a cation Na has 11 protons and 11 electrons net charge = 0 Na+ has 11 protons and 10 elecrons net charge = +1

cations • A cation is smaller than its atom

– Na+ is smaller than Na (less electrons for protons to hold onto)

Metals will lose all their valence electrons -the inner shell electrons are being held onto to tightly to leave group 1: loses 1 electron to have a charge of +1 group 2: loses 2 electrons to have a charge of +2 group 13: loses 3 to have a charge of +3 Cs +

Al3+

Ba2+

anions

• nonmetals (smaller atoms) tend to gain electrons and become anions anions- a negative ion When an atom gain electrons it become an anion F has 9 protons and 9 electrons net charge = 0 F- has 9 protons and 10 elecrons net charge = -1

anions • An anion is bigger than its atom

F- is bigger than F (more electrons for protons to hold onto) nonmetals will gain electrons until the valence shell is full

(they want 8 electrons in the highest shell like the noble gases)

group 14: gains 4 electrons to have a charge of -4 group 15: gains 3 electrons to have a charge of -3 group 16: gains 2 electrons to have a charge of -2 group 17: gains 1 electron to have a charge of -1

F- C4- S2- N3-