Preliminary Course

Atomic Structure 1 + 2

Colm Healy

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Outline

Part One• Matter

• Atomic Structure

• Atomic Model(s)

Part Two• Periodic Table

• Main Group Properties

• Ions

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ChemistryThe Study of Stuff

“Matter” – physical substances

with mass and volumeProperties depend on underlying

molecular or atomic structure

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The Basics

All Matter is made up of Atoms

Tiny (~1x10-10 m) particles

Ultimately the source of all

properties of matter

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The Basics

There are 118 different types of atoms – these are called the Elements.

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The Basics

Atoms (mostly) don’t like being on their own, so they form chemical bonds.

A discrete collection of atoms bound together is called a Molecule.

Single atomsComplicated molecules Very Complicated molecules

Atoms form molecules in a very predictable way, based on their elements

Molecules are discrete entities with discrete properties

Bond

Atom

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The Basics

A molecular formula is built up from the number of atoms of each element

contained in the molecule:Examples:

Water H2O

Peroxide H2O2

Carbon dioxide CO2

Carbon monoxide CO

Ammonia NH3

e.g. Water; H2O

Two atoms of

Hydrogen

(H)

One atom of

oxygen (O)

So two H atoms and one O atom in every molecule of water

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The Basics

If a substance contains more than one

element chemically bound together, this

is called a Compound

If a substance contains multiple

elements or compounds that are

not bound together, it is called a

MIXTURE.

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To Recap:Classifications of Matter

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In more detailScientific Models

Science as a whole tries to DESCRIBE

the world by constructing MODELS.

Models that do not cover everything are

not useless – they are helpful tools to

describe their own situation

As better or more general models are

developed (based on experimental

results) they replace the old models

We’re going to go through a few simple

atomic models (VERY briefly)….

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In more detail

The first atomic theory:

Atomic Theory

~500BC, Democritus

hypothesises about tiny

particles called atoms

No experimental evidence –

just an idea

Guesses they might have

different shapes – like a jigsaw

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In more detailAtomic Theory

A better theory (John Dalton, 1803):

• Matter is made of extremely small particles called atoms.

• Atoms of different elements differ in size, mass, and other

properties.

• Atoms cannot be subdivided, created, or destroyed.

• Atoms of different elements combine in simple whole-number

ratios to form chemical compounds.

• In chemical reactions, atoms are combined, separated, or

rearranged.

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Atomic TheoryExperimental Evidence

The Law of Multiple Proportions: Combinations of atoms always occur in whole number ratios:

Examples:

Water H2O

Peroxide H2O2

Carbon dioxide CO2

Carbon monoxide CO

Ammonia NH3

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Atomic Theory

Dalton’s atomic theory didn’t explain WHY atoms were combining into compounds:

Charged particles

Around 1900 people discover electrically

charged particles even smaller than atoms (the

electron and the nucleus):

JJ Tomson’s experiments (1897)

Electromagnetic forces hold atoms and

molecules together in compounds

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Electromagnetic particles

Thompson and Rutherford introduce ideas of the

electron and the nucleus:

Electrons are NEGATIVELY charged

Nucleus is POSITIVELY charged

Electrons orbit the nucleus – the

atom is mostly empty spaceRutherford’s experiments 1910

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Some Basic Physics

With (electrically) charged particles,

OPPOSITES ATTRACT

Coulomb’s Law

Atomic AttractionConversely, like charges repel

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Atomic Theory

Chadwick divides the nucleus

into Protons and Neutrons

The last piece of the puzzle

The nucleus is very small :

Atom approx. 10-10 m

Nucleus approx. 10-14 m

But most of the mass (>99%) is

in the nucleus

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Atomic TheoryParticles Recap

Name(Symbol)

Electron (e-)

Neutron (n0)

Proton (p+)

Nucleus

Charge

1+

0

1-

Mass (amu*)

1.00727

1.00866

0.00054858

Atoms are generally neutral

ie. Number of protons = number of electrons

*1 amu = 1.66053904 × 10-24 grams

Neutrons just add mass, they don’t affect chemical properties - ISOTOPES

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Atomic Structure

Now that we’ve developed a sense of the

atom, lets look at how electrons behave

Electrons are the outer shell of the atom –

they dictate how the atom interacts with

the rest of the world – i.e. they determine

how an atom behaves chemically.

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Electronic Models

In the early 1900’s, physicists begin to develop Quantum Mechanics:

Quantum Mechanics

A branch of physics describing how very small particles behave:

i.e. electrons, protons, and some very light atoms

Relatively complicated mathematics, so we’re not going to cover

the details here

De Broglie

Wave-Particle Duality

Schrodinger

Wave EquationEinstein Planck

QUANTISATION

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Electronic ModelsThe Bohr Model

Quantisation introduced the idea of energy levels

The idea is that electrons are only allowed have

certain amounts of energy – never an intermediate

amount

Bohr applied this to electrons orbiting around a

nucleus – the Bohr Model

The levels are labelled as n=1, n=2, n=3 etc….Bohr Model

An electron could hop from one to the other – but

could never be somewhere in between

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Bohr ModelExperimental Evidence

To move an electron up a level, you have to put in (an exact

amount of) energy

When an electron falls down a level, it emits (an exact amount

of) energy – usually as light

So you only get out specific colours (energy) back out:

Rainbow

(all colours):

Hydrogen emission

(specific colours):

The colour depends on the element – used in fireworks

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In more detail…Orbitals

In reality it’s slightly more complicated….

Each level (n=1, n=2) etc. can be divided up

into one or more sublevels called orbitals

The sub-levels are labelled s, p, d and f

(for historical reasons)

Note the order of the orbitals

(4s before 3d)

These are sometimes labelled by a lower

case L (l=0, l=1, l=2 and l=3)

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In more detail…Orbitals

The orbitals have specific shapes:

- s is a sphere

- p is a dumbbell

An electron in an s-orbital can be anywhere

inside that sphere

s:

p:

d:

f:

Each orbital can take two electrons

l = 0

l = 1

l = 2

l = 3

These shapes are very important, as they

define how an atom can bond

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In more detail…Orbitals

The orbitals fill up in a specific order (Aufbau Principle):

We can write the electronic congifuration for an atom (in its

ground state) as follows (e.g. for hydrogen):

1s1Energy Level

(n)

Sublevel

(s,p,d or f)

# of electrons

(1 or 2 for s-

oribitals)

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s:

p:

d:

f:

To RecapAtomic Structure 1:

We covered:

Atoms, elements and molecules

The structure of an atom : negative electrons orbiting a

positive nucleus (containing positive protons and neutral

neutrons)

The Bohr Model : quantisation and electronic energy levels

Orbitals : s, p, d and f orbitals, their shapes

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The Periodic Table

Lists all elements

(atom types)

Very useful tool

to determine

properties of

elements

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The Periodic Table

Columns are called

GROUPS. Elements

in a group all have

similar properties.

Rows are called

PERIODS.

Properties change

from period to

period in a

predictable way.

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The Periodic TableHow to read it:

Mass number (A)

Atomic number (Z).

Number of protons in the nucleus

Always a whole number

Number of protons = number of electrons

Mass/weight of an atom in amu

Need this to calculate masses etc.

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Isotopes

Atomic number (Z) = #protons

Mass number (A) ≈ #protons + #neutrons

Therefore #neutrons = A - Z

ISOTOPES have same Z but different A

(ie. Different numbers of neutrons)

A on periodic table is average of all

naturally occurring isotopes, so not

necessarily a whole number

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The periodic tableA tool for predicting properties

The periodic table is arranged

by increasing Z (atomic number)

IE. The different elements are

defined by the number of protons

Adding protons also adds

electrons

Properties “repeat” because

electrons “shells” (n=1, n=2) get

filled up

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The Periodic TableS-block, p-block and d-block

Because electrons fill their shells in a specific order (Aufbau principle), we can divide up the table

into certain “blocks” based on what type of orbital their outermost electrons are in – s, p, d or f.

S-block P-blockD-block

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Group 1 – Alkali Metals

• Soft, silvery metals

• Low melting points

• Very reactive with

water or air

• Produce H2 when

reacted with water

• Increasing reactivity

down the group

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Group 2 – Alkali Earth Metals

• Harder, silvery metals

• Produce H2 when

reacted with water

• Much less reactive

• Increasing reactivity

down the group

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Group 3-11 – The transition metals

• Various properties

• Quite reactive (good

catalysts)

• Very important in

chemistry and biology

• Examples include iron,

copper, gold, platinum,

mercury….

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Group 17 - Halogens

• Very reactive and toxic

• Used as disinfectants

• Form acids with

hydrogen (HCl, HBr)

• Decreasing reactivity

down the group

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Group 18 – The Noble Gases

• Colourless, odourless

gases

• Almost completely

inert

• Almost never react

• Electronically Stable

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The Noble GasesA special case

The noble gases are the most stable elements

(in terms of reactivity towards other elements)

This is to do with their electron configuration

(filled energy levels)

Because the (early) noble gases appeared

(roughly) every eight elements, this is

sometimes called the OCTET RULE

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The “octet rule”

In other words, atoms

want to obtain a “noble

gas configuration”

Want to LOSE ELECTRONS Want to GAIN ELECTRONS

Atoms will try and find

the shortest path to a

noble gas configuration

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Sodium

Neon (noble gas) has 10 electrons (Z=10). It is unreactive.

Sodium (group 1 metal) has 11 electrons (Z=11)

Therefore sodium desperately wants to lose an

electron to gain a noble gas configuration

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Chlorine

The nearest noble gas is argon (Ar), which has 18 electrons (Z=18).

Similarly, chlorine (group 17 halogen) has 17 electrons (Z=17)

Chlorine will therefore pull an electron from something to satisfy the octet rule

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Forming Salt

So if chlorine wants to gain an electron, and sodium wants to lose one, we can expect them to react together:

Sodium Chlorine Sodium Chloride

(table salt)

Salts like this are made up of IONS (permanently charged atoms)

The positive one (lost an electron) is called the CATION

The negative one (gained an electron) is called the ANION

Remember, opposites attract, so there is a force binding the ions together into a compound

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Cations + Anions

Want to LOSE ELECTRONS Want to GAIN ELECTRONS

Metals tend

to form

CATIONS

NON- Metals

tend to form

ANIONS

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Multiple ions

We can predict a lot about stable compounds from this

For example, group 2 metals need to lose two electrons to obtain noble gas configuration

Li and Na (Group 1)

Mg and Ca (Group 2)

Li+ Na+

Mg2+ Ca2+

So Calcium would have to react with TWO chlorines – Calcium forms CaCl2, not CaCl

(Notice charge = group)

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Multiple Ions

Similarly, with groups 15-17 (non-metals):

Nitrogen (Group 15)

Oxygen (Group 16)

N3-

O2-

Chlorine (Group 17) Cl-

(Notice charge = group - 18)

Oxygen wants to gain two electrons – it will form Na2O or CaO

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Working out compounds

This method is not 100% foolproof, as different types of bonding

come in to play (especially with carbon or transition metals), but

this predicts a huge amount of compounds

Examples:

Water H2O

Carbon dioxide CO2

Ammonia NH3

NaCl

KBr

HCl

Add up positive charges (number

of groups from the left) and

subtract negative charges

(number of groups from right)

until things are neutral

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To RecapAtomic Structure 2:

We covered:

The Periodic Table, Rows and Periods

Atomic numbers and Mass numbers

The Noble Gases and their electronic stability

A simple bonding case : Sodium chloride

Ions : positive cations and negative anions

Best of luck with first year!