Unit 2. Atoms are measured in picometers, 10 -12 meters Hydrogen atom, 32 pm radius Nucleus tiny compared to atom If the atom were a stadium, the nucleus.

Atomic Structure and Bonding

Unit 2

Major Subatomic Particles

•Atoms are measured in picometers, 10-12 meters Hydrogen atom, 32 pm radius

• Nucleus tiny compared to atom If the atom were a stadium, the nucleus would be a marble• Radius of the nucleus is on the order of 10-15 m

• Density within the atom is near 1014 g/cm3

Name Symbol Charge Relative Mass (amu)

Actual Mass (g)

Electron e- -1 1/1840 9.11x10-28

Proton p+ +1 1 1.67x10-24

Neutron no 0 1 1.67x10-24

Elemental Classification•Atomic Number (Z) = number of protons (p+) in the

nucleus Determines the type of atom

• Li atoms always have 3 protons in the nucleus, Hg always 80

• Mass Number (A) = number of protons + neutrons [Sum of p+ and nº]

Electrons have a negligible contribution to overall mass

• In a neutral atom there is the same number of electrons (e-) and protons (atomic number)

Nuclear Symbols•Every element is given a corresponding symbol

which is composed of 1 or 2 letters (first letter upper case, second lower), as well as the mass number and

atomic number

E A

Z

elemental symbol

mass number

atomic number

ATOMIC NUMBER AND MASS NUMBER

the number of protons in an atom

the number of protons and neutrons in an atomH

e2

4

Atomic Number

Mass Number

Number of electrons = Number of protons

in a neutral atom 5

•Find the number of protons number of neutrons number of electrons atomic number mass number

W184 74

F199 Br80

35

IonsCation is a positively charged particle.

Electrons have been removed from the element to form the + charge.

ex: Na has 11 e-, Na+ has 10 e-

Anion is a negatively charged particle. Electrons have been added to the atom to form the – charge.

ex: F has 9 e-, F- has 10 e-

Isotopes•Atoms of the same element can have different

numbers of neutrons and therefore have different mass numbers

• The atoms of the same element that differ in the number of neutrons are called isotopes of that

element

• When naming, write the mass number after the name of the element

H11Hydrogen-1

H21

Hydrogen-2

H31Hydrogen-3

Calculating AveragesAverage = (% as decimal) x (mass1) + (% as decimal)

x (mass2) + (% as decimal) x (mass3) + …

Problem:

Silver has two naturally occurring isotopes, 107Ag with a mass of 106.90509 u and abundance of 51.84 % ,and 109Ag with a mass of 108.90476 u and abundance of 48.16 % What is the average atomic mass?

Average = (0.5184)(106.90509 u) + (0.4816)(108.90476 u)

= 107.87 amu

• If not told otherwise, the mass of the isotope is the mass number in ‘u’

• The average atomic masses are not whole numbers because they are an average mass value

• Remember, the atomic masses are the decimal numbers on the periodic table

Average Atomic Masses

• Calculate the atomic mass of copper if copper has two isotopes 69.1% has a mass of 62.93 amu The rest (30.9%) has a mass of 64.93 amu

• Magnesium has three isotopes 78.99% magnesium 24 with a mass of 23.9850 amu 10.00% magnesium 25 with a mass of 24.9858 amu The rest magnesium 26 with a mass of 25.9826 amu What is the atomic mass of magnesium?

More Practice Calculating Averages

BohrProposed electrons (e-) orbit around the nucleus

in circular pathsSaid e- in a particular path have a fixed energy

(energy levels)e- can go from any energy level to another by

gaining or losing a specific amount of energy = a “quantum of energy”

When e- absorbs a quantum of energy, it goes from it’s ground state (where it’s normally found) to an excited state

The excited state is at a higher energy level

Bohr postulated that: Fixed energy related to the orbit Electrons cannot exist between orbits The higher the energy level, the further it is

away from the nucleus An atom with maximum number of

electrons in the outermost orbital energy level is stable (unreactive)

Think of Noble gases

Atomic Line Emission Spectra and Niels BohrAtomic Line Emission Spectra and Niels Bohr

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 HydrogenNiels Bohr

(1885-1962)

Spectrum of White Light

Spectrum of Excited Hydrogen Gas

Line Emission Spectra of Excited AtomsLine Emission Spectra of Excited AtomsExcited atoms emit light of only

certain wavelengthsThe wavelengths of emitted light

depend on the element.

Drawback to BohrBohr’s theory did

not explain or show the shape or the path traveled by the electrons.

His theory could only explain hydrogen and not the more complex atoms

Energy level populations (Science10)

Electrons found per energy level of the atom.

The first energy level holds 2 electronsThe second energy level holds 8 electronsThe third energy level holds 18 electrons

Examples for group 1Li 2.1 Na 2.8.1 K 2.8.8.1

The Quantum Mechanical ModelEnergy is quantized. It comes in chunks.A quanta is the amount of energy needed to

move from one energy level to another.Since the energy of an atom is never “in

between” there must be a quantum leap in energy.

Schrödinger derived an equation that described the energy and position of the electrons in an atom – an ORBITAL

Orbits (Bohr) vs Orbitals (Quantum Mechanics)

Bohr said electrons travel in an orbit – can predict exact location of electron at any point in time.

Schrodinger used mathematics (calculus) to find the region in space where an electron will be found 90% of the time - this region is called an orbital. There are 4 main types of orbitals – s, p, d, and f.

Modern View of the Atom The modern view of the atom

suggests that the atom is more like a cloud.

Atomic orbitals around the nucleus define the places where electrons are most likely to be found.

23

s orbitals

1 s orbital forevery energy level

1s 2s 3sSpherical shapedEach s orbital can hold 2 electronsCalled the 1s, 2s, 3s, etc.. orbitals

p orbitalsStart at the second energy level 3 different directions3 different shapesEach orbital can hold 2 electrons

The d sublevel contains 5 d orbitalsThe d sublevel starts in the 3rd energy

level 5 different shapes (orbitals)Each orbital can hold 2 electrons

The f sublevel has 7 f orbitalsThe f sublevel starts in the fourth energy levelThe f sublevel has seven different shapes

(orbitals)2 electrons per orbital

Electron ConfigurationWe use e- configuration as a shorthand

to show how e- are arranged around a nucleus

Example: Carbon is …

1s2 2s2 2p2

Electron ConfigurationsThe way electrons are arranged in atoms.Aufbau principle- electrons enter the

lowest energy first.This causes difficulties because of the

overlap of orbitals of different energies.Pauli Exclusion Principle- at most 2

electrons per orbital - different spinsHund’s Rule- When electrons occupy

orbitals of equal energy they don’t pair up until they have to .

Summary

Sublevel

# of shapes

(Orbitals)

Max number

of e-

Starts at

energy level

s 1 2 1

p 3 6 2

d 5 10 3

f 7 14 4

Electron Arrangement

1st Rule: The Aufbau Principlee- fill orbitals of the lowest energy firstWe can use the periodic table to help

us!

The Diagonal Rule

Example #1

Oxygen

1 s

2 s 2 p

1s2 2s2 2p4

Example #2

Magnesium

1 s

2 s 2 p

3 s

1s2 2s2 2p6

3s2

Example #3

Iron

1 s

2 s 2 p

3 s 3 p

4 s

3d

1s2 2s2 2p6

3s2 4s23d63p6

PracticeBoronArgonCalciumIodineSodiumZincLead

AbbreviationsWe can abbreviate electron

configurations using the Noble GasesEx: Sulfur

1s2 2s2 2p6 3s2 3p4

[Ne] 3s2 3p4

Ex: Lead 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6

6s2 4f14 5d10 6p2

[Xe] 6s2 4f14 5d10 6p2

2nd Rule: Pauli Exclusion Principle

Each orbital orientation can hold up to 2 e-

e- must have opposite spins (up/clockwise or down/counter clockwise)

Therefore:s has up to 2 e- (1 orientation)p has up to 6 e- (3 orientations)d has up to 10 e- (5 orientations)f has up to 14 e- (7 orientations)

We can use the 2nd rule to draw Orbital Diagrams

Example #1Oxygen: 1s2 2s2

2p4

1s 2s 2p

Example #2Magnesium: 1s2 2s2 2p6 3s2

1s 2s 2p

3s

Example #3Iron

1s2

2s2

2p6

3s2 4s23d63p6

3rd Rule: Hund’s Rulee- will not pair up until each orbital

orientation has 1 e- in itThe first e- in a pair will spin up, the

second will spin downExample: Oxygen is 1s2 2s2 2p4

1s 2s 2p

Orbital NotationOrbital Notation shows us visually the

arrangement and spin of electronsExample: Carbon is 1s2 2s2 2p2

1s 2s 2p

Energy Level DiagramsEnergy Level

Diagrams give us the same information as orbital diagrams, plus they show us the different energy levels of each orbital

Example: Carbon is 1s2 2s2 2p2 1s

2s

2p

Incr

easi

ng e

nerg

y

1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

Phosphorous, 15 e- to place

The first to electrons go into the 1s orbital

Notice the opposite spins

only 13 more

Incr

easi

ng e

nerg

y

1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

The next electrons go into the 2s orbital

only 11 more

Incr

easi

ng e

nerg

y

1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

• The next electrons go into the 2p orbital

• only 5 more

Incr

easi

ng e

nerg

y

1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

• The next electrons go into the 3s orbital

• only 3 more

Incr

easi

ng e

nerg

y

1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

Incr

easi

ng e

nerg

y

1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

• The last three electrons go into the 3p orbitals.

• They each go into separate shapes

• 3 unpaired electrons

• 1s22s22p63s23p3

Orbitals fill in order Lowest energy to higher energy.Adding electrons can change the energy of

the orbital.Half filled orbitals have a lower energy.Makes them more stable.Changes the filling order

Write these electron configurations

Titanium - 22 electrons

1s22s22p63s23p64s23d2

Vanadium - 23 electrons

1s22s22p63s23p64s23d3

Chromium - 24 electrons

1s22s22p63s23p64s23d4

Electronic Structure - Questions

Copy and complete the following table:

Atomic no.

Mass no.

No. of protons

No. of neutrons

No. of electrons

Electronic structure

Mg 12 1s2 2s2 2p6

3s2

Al3+ 27 10

S2- 16 16

Sc3+ 21 45

Ni2+ 30 26