Download Preliminary Course Atomic Structure 1 + 2

Preliminary Course
Atomic Structure 1 + 2
Colm Healy
Outline
Part One
• Matter
• Atomic Structure
• Atomic Model(s)
Part Two
• Periodic Table
• Main Group Properties
• Ions
2
Chemistry
The Study of Stuff
“Matter” – physical substances
Properties depend on underlying
with mass and volume
molecular or atomic structure
3
The Basics
All Matter is made up of Atoms
Tiny (~1x10-10 m) particles
Ultimately the source of all
properties of matter
4
The Basics
There are 118 different types of atoms – these are called the Elements.
5
The Basics
Atoms bind together to form Molecules.
Single atoms
Complicated molecules
Very Complicated molecules
Molecules are discrete entities with discrete properties
Atoms form molecules in a very predictable way, based on their elements
6
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 detail
Atomic Theory
How exactly does atomic structure dictate properties?
~500BC, Democritus
hypothesises atoms are
different shapes and sizes
Atoms can recombine in
different ways
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In more detail
Atomic Theory
John Dalton (1803) updates the theory:
• 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 Theory
Experimental 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
Charged particles
Dalton’s atomic theory didn’t explain WHY atoms were combining into
compounds
Around 1900 people discover
electrically charged particles
even smaller than atoms
JJ Tomson’s experiments (1897)
12
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 space
Rutherford’s experiments 1910
13
Some Basic Physics
With (electrically) charged particles,
OPPOSITES ATTRACT
Coulomb’s Law
Conversely, like charges repel
Atomic Attraction
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Atomic Theory
The last piece of the puzzle
Chadwick divides the nucleus
into Protons and Neutrons
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 Theory
Particles Recap
*1 amu = 1.66053904 × 10-24 grams
Name(Symbol)
Mass (amu*)
Charge
Proton (p+)
1.00727
1+
Neutron (n0)
1.00866
0
Electron (e-)
0.00054858
1-
Nucleus
Atoms are generally neutral
ie. Number of protons = number of electrons
There can be different numbers of neutrons – this gives rise to ISOTOPES
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Atomic Structure
There are more details
(to do with the location
of the electrons and the
amount of energy they
can have), but these
will be dealt with in
more detail later
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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 Table
How to read it:
Atomic number (Z).
Number of protons in the nucleus
Always a whole number
Number of protons = number of electrons
Mass number (A)
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 table
A 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 get filled up
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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 Gases
A special case
The noble gases are the most stable
elements
This is to do with their electron
configuration (more details later)
Because the 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”
Atoms will try and find
the shortest path to a
noble gas configuration
Want to LOSE ELECTRONS
Want to GAIN ELECTRONS
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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
30
Chlorine
Similarly, chlorine (group 17 halogen) has 17 electrons (Z=17)
The nearest noble gas is argon (Ar), which has 18 electrons (Z=18).
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
32
Cations + Anions
Metals tend
to form
CATIONS
Want to LOSE ELECTRONS
NON- Metals
tend to form
ANIONS
Want to GAIN ELECTRONS
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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)
Li+ Na+
(Notice charge = group)
Mg and Ca
(Group 2)
Mg2+ Ca2+
So Calcium would have to react with TWO clorines – Calcium forms CaCl2, not CaCl
34
Multiple Ions
Similarly, with groups 15-17 (non-metals):
Nitrogen
(Group 15)
N3-
Oxygen
(Group 16)
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
Add up positive charges (number
of groups from the left) and
subtract negative charges
(number of groups from right)
until things are neutral
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
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Best of luck with first year!