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Atomic Structure &

The Periodic Table

The Greek Philosophers

• Democritus – believed that all matter is made up of tiny particles that could not be divided

• Aristotle -- thought that matter was made of only four elements

• Fire

• Earth

• Air

• Water

Dalton’s Atomic Theory

• All elements are composed of atoms

• All atoms of the same element have the same mass, and atoms of different elements have different masses*

• Compounds contain atoms of more than one element

• In a particular compound, atoms of different elements always combine in the same way

J. J. Thomson

• Provided the first evidence that atoms are made up of even smaller particles.

• “Plum Pudding” model – electrons were

evenly spread out among a positively

charged mass (chocolate chip ice

cream)

Thomson’s Experiment

• Hypothesis – The beam was a stream of charged particles that caused the air to glow.

• Experiment – Positive and negative plates were put on either side of the tube.

• Results – The beam was bent towards the positive plate.

• Conclusion – There are negative particles being given off of the atoms that are attracted to the positive plate. (Atoms are made of smaller pieces!)

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Ernest Rutherford’s Gold Foil Experiment

Hypothesis: that alpha particles would pass straight through a thin sheet of gold

Results

The alpha particles did not pass straight through, but instead were deflected in

various directions.

Conclusion

The positive charge of an atom IS NOT evenly spread out, but concentrated in

a very small area (nucleus).

•Writing Elemental Symbols

• First letter – ALWAYS UPPERCASE

• Second letter – always lowercase

• Hg

• NOT HG or HG

Subatomic Particles• Proton (p+)

• Positive charge

• Part of the nucleus

• Neutron (n)• Neutral (no charge)

• Part of the nucleus

• Approximately the same mass as a p+

• Electron (e-)• Negative charge

• Found orbiting around the nucleus

• MUCH smaller than p+ or n (1/1836)

Atomic Number

• The number of protons found as a part of the nucleus in an atom

• Unique for each element

• Cannot change without changing the

type of element

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•For neutral atoms

# of p+ = # of e-

Mass Number

• The total number of particles that make up the nucleus

• Mass number = p+ + n

• n = mass number – p+

• Atoms of the same element may have

varying mass numbers because the

number of neutrons can vary

Isotopes

• Atoms of the same element with

different mass numbers (numbers of neutrons)

Ways to write isotopes

• Element – mass number(Carbon – 12)

• Symbol – mass number

(C – 12)

• Mass number Symbol (12 C)

•Ways to write isotopes, cont.

C)( SymbolNumber MassNumberAtomic

126

Niels Bohr

• Focused on electrons

• Electrons move at constant speed in fixed spherical orbits around the

nucleus

• Proposed the existence of energy levels

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Energy Levels

• The possible energy an electron

can have

• Similar to steps

• Higher steps have more energy

• Going down a step means energy was released

Energy Levels, cont.

• Seven possible energy levels that

correspond to the rows on the periodic table

• Also called “shells”

Bohr’s Model of the Atom

• Each circle represents an energy

level

• 1st level: 2 electrons

• 2nd level: 8 electrons

• 3rd level: 18 electrons

• 4th level: 32 electrons

Erwin Schrödinger

• Developed mathematical model

to describe the motion of electrons

• Work leads to electron cloud

model

Orbital

• Region of space where an electron

is likely to be found

• Shapes are created by taking

many pictures of the electron’s

location over a period of time

Pauli Exclusion Principle

• Only two electrons with opposite

spins will occupy the same orbital

• A series of the same type of orbital

is also called a sub-level or sub-

shell.

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s-orbital

• Shaped like a sphere

• Only one type (b/c it can only point one direction!)

Electron Probablity Map – p orbital

p - orbital

• Has one shape that can point

three ways

• Total of 3 types of p-orbitals (px,

py, pz)

Electron Probability Maps

d orbitals

d - orbital

• Five total types

f - orbitals

• Seven possible types

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•How Orbitals Interact

http://micro.magnet.fsu.edu/electromag/java/atomicorbitals/index.html

Aufbau Principle

• Electrons fill orbitals that have the

lowest energy first

• Follow the DIAGONAL RULE!

Electron Configurations(More sophisticated than Bohr’s!)

• Know the total number of electrons

in the atom!

• Use the diagonal rule

•The Diagonal Rule

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d

7s 7p

Draw the electron configurations for

the following atoms”

Sodium (Na)

Sulphur (S)

Silver (Ag)

Tin (Sn)

Orbital Diagrams

Shows the placement of electrons in the orbitals…

Ex. Sodium

1s22s22p63s1

1s 2s 2p 3s

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Hund’s Rule

Electrons will fill each orbital on a sublevel before pairing with each other.

Ex. Carbon

Quantum Numbers

Principal Quantum Number (n)

n = energy level (1, 2, 3, …)

Azimuthal Quantum Number (ℓ)

ℓ = type of orbital (s=0, p=1, d=2, f=3)

Magnetic Quantum (which orbital)

m = ℓ

Electron Spin (first or second electron)

s = ½

What are the quantum numbers for the 3rd

and 6th electrons in neon?

N 1s22s22p6

3rd electron n=2, ℓ=0, m=0, s=+½

6th electron n=2, ℓ=1, m=-1, s=-½

1s 2s 2p

Organizing the Periodic Table - Revisited

Shorthand electron configurations

Use Noble Gas shortcuts ONLY!

Kr (36 e-)

1s22s22p63s23p64s23d104p6

Mo (42 e-)

1s22s22p63s23p64s23d104p65s24d4

[Kr] 5s24d4

John Newlands

• Arranged the 16 known

elements in order of increasing

atomic mass

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Dmitri Mendeleev

• Arranged the periodic table by

increasing atomic mass and stacking elements with similar

properties in the same column

• Was able to predict the properties of elements not yet discovered by

looking at the blank spaces in his table

Henry Moseley

• Fine-tuned the periodic table

by placing the elements in

increasing atomic number

(protons were not known during

Mendeleev’s time)

Other Periodic Arrangements •Atomic Informationfrom the Periodic Table

1

1.00794

H

Atomic Number

Atomic Mass

Element Symbol

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Atomic Mass

• The weighted average of all

possible isotopes for an element

• Atomic Mass =

(%A)(Mass of A) + (%B)(Mass of B)

Organizing the Periodic Table

• Metals are left of the stair-step

line

• Non-metals are to the right of

the stair-step line

• Metalloids (have properties of

both) are found adjacent to

stair-step line

Period

• A row in the periodic table

corresponding to the energy

level on which the electrons

exist

Group

• A column on the periodic table

in which the elements have

similar properties

• Similar properties are created

by similar electron

configurations

1A

2A 3A 4A 5A 6A 7A

8A• A Groups = Representative Elements

B groups = transition elements

These are called the inner transition elements, and they belong here

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Group 1A are the alkali metals (but NOT H)

Group 2A are the alkaline earth metals

H

• Group 8A are the noble gases

• Group 7A is called the halogens

Valence Electrons

• Electrons that are in the outermost

shell

• Involved in bonding with other

atoms

• Same amount as “A” group #

• Examples: H has 1, N has 5, P has 5

Octet Rule

• Atoms are most stable when they have full outer shells

• 8 electrons (s2, p6) for most

representative elements (s2 for H, He)

Ion

• an atom (or group of atoms) that has gained or lost electrons

• anion – negative charge – gained electron(s) – tend to be larger than parent atom

• cation – positive charge – lost electron(s) – tend to be smaller than parent atom

Monoatomic Ions

• Ions from single atoms

• Charge can be predicted by location on the periodic table

• Cation – positively charged ion – lost e-

• Anion - negatively charged ion –gained e-

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+1 +2

+3 -3 -2 -1

Atomic Size

• Measure the Atomic Radius - this is half the distance between the two nuclei of a diatomic molecule.

}Radius

Atomic Size - Group trends• As the atomic number increases

each atom has another energy level,

so the atoms get

bigger.

H

Li

Na

K

Rb

Atomic Size - Period Trends

• Going from left to right across a period,

the size gets smaller.

• Electrons are in the same energy level.

• But, there is more nuclear charge.

• Outermost electrons are pulled closer.

Na Mg Al Si P S Cl Ar

Ionization Energy• Ionization energy is the amount of energy required to completely remove an electron (from a gaseous atom).

• Removing one electron makes a 1+ ion.

IE - Group trends

• As you go down a group, the first IE decreases (less energy required, easier to remove electron) because the electron is further away from the attraction of the nucleus

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IE - Period trends

• All the atoms in the same period have the same energy level.

• But, increasing nuclear charge

• So IE generally increases from left to right (higher energy, more

difficulty to remove electron)

Trends in Electronegativity

• Electronegativity is the tendency for an

atom to attract electrons to itself when

it is chemically combined with another

element.

• They share the electron, but how

equally do they share it?

• The higher the EN the “stronger” an atom is at attracting electrons.

Electronegativity Group Trend

•The further down a group, the farther the electron is away from the nucleus, plus the more electrons an atom has.

•Thus, more willing to share.

•Low electronegativity.

Electronegativity Period Trend

• Metals let their electrons go easily

• Thus, low electronegativity

• Nonmetals want more electrons.

• Try to take them away from others

• High electronegativity.

Electron Affinity

• The amount of energy RELEASED when an electron is added to an atom

• the more negative, the more energy

that is released, the more likely an atom

will add an electron

Trends in three atomic properties

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Trends in metallic behavior