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Preliminary Course
Atomic Structure 1 + 2
Colm Healy
2
Outline
Part One• Matter
• Atomic Structure
• Atomic Model(s)
Part Two• Periodic Table
• Main Group Properties
• Ions
3
ChemistryThe Study of Stuff
“Matter” – physical substances
with mass and volumeProperties depend on underlying
molecular or atomic structure
4
The Basics
All Matter is made up of Atoms
Tiny (~1x10-10 m) particles
Ultimately the source of all
properties of matter
5
The Basics
There are 118 different types of atoms – these are called the Elements.
6
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
7
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
8
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.
9
To Recap:Classifications of Matter
10
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)….
11
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
12
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.
13
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
14
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
15
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
16
Some Basic Physics
With (electrically) charged particles,
OPPOSITES ATTRACT
Coulomb’s Law
Atomic AttractionConversely, like charges repel
17
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
18
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
19
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.
20
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
21
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
22
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
23
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)
24
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
25
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)
26
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
27
The Periodic Table
Lists all elements
(atom types)
Very useful tool
to determine
properties of
elements
28
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.
29
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.
30
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
31
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
32
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
33
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
34
Group 2 – Alkali Earth Metals
• Harder, silvery metals
• Produce H2 when
reacted with water
• Much less reactive
• Increasing reactivity
down the group
35
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….
36
Group 17 - Halogens
• Very reactive and toxic
• Used as disinfectants
• Form acids with
hydrogen (HCl, HBr)
• Decreasing reactivity
down the group
37
Group 18 – The Noble Gases
• Colourless, odourless
gases
• Almost completely
inert
• Almost never react
• Electronically Stable
38
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
39
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
40
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
41
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
42
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
43
Cations + Anions
Want to LOSE ELECTRONS Want to GAIN ELECTRONS
Metals tend
to form
CATIONS
NON- Metals
tend to form
ANIONS
44
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)
45
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
46
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
47
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!