All matter is composed of atoms.

Understanding the structure of atoms is critical to understanding the properties of matter

HISTORY OF THE ATOM

1808 John Dalton

suggested that all matter was made up of

tiny spheres that were able to bounce around

with perfect elasticity and called them

ATOMS

DALTONS ATOMIC THEORY

Particle Mass

(g) Charge

(Coulombs) Charge (units)

Electron (e-) 9.1 x 10

-28 -1.6 x 10

-19 -1

Proton (p) 1.67 x 10-24

+1.6 x 10-19

+1

Neutron (n) 1.67 x 10-24

0 0

mass p = mass n = 1840 x mass e-

Each proton and neutron has a mass of approximately 1 dalton.

HISTORY OF THE ATOM

1898 Joseph John Thompson

found that atoms could sometimes eject a far

smaller negative particle which he called an

ELECTRON

HISTORY OF THE ATOM

1910 Ernest Rutherford

oversaw Geiger and Marsden carrying out his

famous experiment.

They passed Helium nuclei (particle α )

through a piece of gold foil which was only a

few atoms thick.

They found that although most of them

passed through. About 1 in 10,000 hit

atomic radius ~ 100 pm = 1 x 10-10 m

nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m

Rutherford’s Model of the Atom

Atomic Structure

Atoms are composed of

-protons – positively charged particles

-neutrons – neutral particles

-electrons – negatively charged particles

Protons and neutrons are located in the nucleus. Electrons are found in orbitals surrounding the nucleus.

HELIUM ATOM

+ N

N

+

- -

proton

electron neutron

Shell

Atomic Structure

Every different atom has a characteristic number of protons in the nucleus.

atomic number = number of protons

Atoms with the same atomic number have the same chemical properties and belong to the same element.

ATOMIC NUMBER (Z) = number of protons in nucleus

MASS NUMBER (A) = number of protons + number of neutrons

= atomic number (Z) + number of

neutrons

ISOTOPS are atoms of the same element (X) with different numbers

of neutrons in the nucleus

X A

Z

H 1 1 H (D)

2 1 H (T)

3 1

U 235 92 U 238

92

Mass Number

Atomic Number Element Symbol

H 1 1 H (D)

2 1 H (T)

3 1

*An element can have several naturally occurring isotopes.

*These isotopes of a element behave in the same way.

*In calculating the relative atomic mass of an element with

isotopes, the relative mass and proportion or percentage of

each is taken into account.

Isotope Relative isotopic

mass

Relative abundance

(%)

Cl 34.969 75.80

Cl 36.966 24.20

Ar = (relative isotopic mass X1 % abundance) + relative isotopic mass X2 % abundance)

100

Ar (Cl) = (34.969 X 75.8) + ( 36.966 X 24.2)

100

Ar (Cl) = 2650.65 + 894.58

100

Ar (Cl) = 35.45 amu (atomic mass unit)

HISTORY OF THE ATOM

1913 Niels Bohr

studied under Rutherford at the Victoria

University in Manchester.

Bohr refined Rutherford's idea by adding

that the electrons were in orbits. Rather

like planets orbiting the sun. With each

orbit only able to contain a set number of

electrons.

MULTIELECTRON ATOMS

*The principal quantum number, n, describes the energy

level on which the orbital resides.

*Largest E difference is between E levels

*The values of n are integers > 0.

*1, 2, 3,...n.

Quantum Number

Azimuthal Quantum

Number, l

*defines shape of the subshell

*Allowed values of l are integers ranging from 0 to n − 1.

Value of l 0 1 2 3

Type of subshell s p d f

So each of these letters corresponds to a shape of orbital.

*Describes the three-dimensional orientation of the orbital.

*Values are integers ranging from -l to l:

−l ≤ ml ≤ l.

*Therefore, on any given energy level, there can be up to:

* 1 s (l=0) orbital (ml=0),

* 3 p (l=1) orbitals, (ml=-1,0,1)

* 5 d (l=2) orbitals, (ml=-2,-1,0,1,2)

* 7 f (l=3) orbitals, (ml=-3,-2,-1,0,1,2,3)

The energy of electron only depends on the value of n

shell = all orbitals with the same value of n subshell = all orbitals with the same value of n and l

an orbital is fully defined by three quantum numbers, n, l, and ml

Each shell of QN = n contains n subshells n = 1, one subshell n= 2, two subshells, etc Each subshell of QN = l, contains 2l + 1 orbitals l = 0, 2(0) + 1 = 1 l = 1, 2(1) + 1 = 3

*Value of l = 0.

*Spherical in shape.

*Radius of sphere

increases with

increasing value of n.

*Value of l = 1.

*Have two lobes with a nodal plane between

them.

Note: always 3 p orbitals for a given n

*Value of l is 2.

*2 nodal planes

*Four of the five

orbitals have 4

lobes; the other

resembles a p

orbital with a

doughnut around

the center.

Note: always 5 d orbitals for a given n.

*This leads to a fourth quantum number, the spin quantum number ms.

*The spin quantum number has only 2 values +1/2 and -1/2

*Describes magnetic field vector of electron

*

*No two electrons in the

same atom can have

exactly the same energy.

*For example, no two

electrons in the same

atom can have identical

sets of quantum

numbers.

*Name of each electron unique

*Name consists of four numbers:

*n,l,ml,ms

*Example:

*Mr. George Herbert Walker Bush

*We must learn to name our

electrons

*Unlike people, there is a lot in the

“name” of an electron.

*Distribution of all

electrons in an atom.

*Consist of

*Number denoting the

energy level.

*Letter denoting the type

of orbital.

*Superscript denoting the

number of electrons in

those orbitals.

*Each box represents

one orbital.

*Half-arrows represent

the electrons.

*The direction of the

arrow represents the

spin of the electron.

1s22s1

(of maximum multiplicity)

“For degenerate orbitals,

the lowest energy is

attained when the

number of electrons with

the same spin is

maximized.” NOT:

*We fill orbitals in increasing order of energy.

*Different blocks on the periodic table, then

correspond to different types of orbitals.

*Remember: The periodic table was arranged the way it was based on chemical properties.

*Totally empirical, until now. Based only on observation.

Short cut for writing electron configurations

Some irregularities

occur when there

are enough

electrons to half-fill

s and d orbitals on a

given row.

Atomic Structure

Neutral atoms have the same number of protons and electrons.

Ions are charged atoms.

-cations – have more protons than electrons and are positively charged

-anions – have more electrons than protons and are negatively charged

An ion is formed when an atom, or group of atoms, has a

net positive or negative charge (why?).

If a neutral atom looses one or more electrons

it becomes a cation.

If a neutral atom gains one or more electrons

it becomes an anion.

Na 11 protons

11 electrons Na+ 11 protons

10 electrons

Cl 17 protons

17 electrons Cl- 17 protons

18 electrons