Download 2015 VCE Chemistry Unit 1 -Miss Fitzsimmons

You’re Officially Welcomed
to VCE Chemistry 2015!
How we will start the year…
Unit 1, Area of Study 1
1.The historical development
of the Periodic table
2.The historical development
of The Atomic Structure
3.Mass numbers, isotopes and
electron configurations
4.Trends of the periodic table
1. The historical
development of the
Periodic table
The historical development
of the Periodic table
(you learnt this stuff in transition…
• The main guys…
here’s a recap!)
Priestley and Lavoisier: In 1774, Priestley was the first to isolate elemental
Oxygen, and Lavoisier discovered that Oxygen was the element that’s involved
in combustion reactions.
Davy: This clever chap used electrolysis to isolate the elements potassium and
sodium. He then discovered calcium, barium and strontium and was the first
to produce elemental aluminium.
Dalton: Around the early 1800’s, Dalton used maths to prove that:
All matter consisted of indivisible atoms
Atoms of a particular element are identical in mass and have identical
properties.
Atoms of each element have a unique mass.
Atoms are neither created or destroyed in reactions
Compounds are formed from the combination of two or more elements.
The historical development
of the Periodic table
(you learnt this stuff in transition…
here’s a recap!)
• The main guys… (continued)
Berzelius: Assigned symbols to stand for the name of the elements. He also
calculated accurately the relative atomic masses of many of the elements.
Mendeleev: The guy who created The
Periodic Table!
Dimitri Mendeleev…
Mendeleev first arranged the elements in order of atomic mass
and then examined the properties of the elements to discover
that they changed periodically.
He arranged the elements with similar properties into vertical
groups and elements in order of increasing atomic mass into
horizontal periods.
Using the periodic trends and the chemical properties of
elements, Mendeleev was able to accurately predict the
appearance and properties of elements that had not yet been
discovered.
Dimitri Mendeleev…
Dimitri Mendeleev…
Dimitri Mendeleev…
Dimitri Mendeleev…
Dimitri Mendeleev…
How we will start the year…
Unit 1, Area of Study 1
1.The historical development
of the Periodic table √
2.The historical development
of The Atomic Structure
3.Mass numbers, isotopes and
electron configurations
4.Trends of the periodic table
2. The historical
development of The Atomic
Structure
The historical development
of The Atomic Structure
(also covered in transition… here’s
the recap!) around 400BC
• Democritus:
Democritus was the first scientist to create a model of the atom. He was the first
one to discover that all matter is made up of invisible particles called atoms. He
created the name "atom" from the Greek word "atomos", which means uncuttable.
He also discovered that atoms are solid, insdestructable, and unique. Democritus
didn't know about the different elements, or about a nucleus or electrons, all he
knew was that everything is made of atoms.
Leaf
atom
The historical development
of The Atomic Structure
(also covered in transition… here’s
the recap!)
• Dalton:
1808
John Dalton was an English chemist that created the Atomic Theory of Matter, a
composition of previous findings by Democritus and his own findings. He
included in this theory that all matter is made of atoms, that atoms cannot be
created nor destroyed and also, atoms of different elements combine in whole
ratios to form chemical compunds. His theory would later contribute to an
advance in the atomic model.
Carbon
atom
The historical development
of The Atomic Structure
(also covered in transition… here’s
• Sir JJ Thompson: the recap!)
1897
J.J. Thomson was a very important scientist when it came to the atomic model.
Up until his time, all models of the atom looked like a big solild ball. J.J.
Thomson discovered the electron, which led him to create the "plum pudding"
atomic model. In this model, he thought that the atom was mostly positive, and
negative electrons wandered around the atom. The "plum pudding" model
influenced other scientists to make better atomic models.
The historical development
of The Atomic Structure
(also covered in transition… here’s
• Lord Rutherford: the recap!)
1911
Ernest Rutherford was another scientist that changed the atomic model. He found
that J.J. Thomson's model was incorrect, so he created a new one. He created the
nucleus, and said that instead of the positive matter being the whole atom, it was
just in the middle. He said the atom was mostly empty space and that the
electrons surrounded the positive nucleus. This model influenced one of his own
students to perfect the atomic model later on.
The historical development
of The Atomic Structure
(also covered in transition… here’s
the recap!) 1913
• Neils Bohr:
Niels Bohr was a Danish scientist that was a student of Rutherford. He decided
to make a new model based off of Rutherford's model, but changed the orbit of the
electron. Also, he created energy levels in the atom, where only a certain amount
of electrons could fit on one energy level of the atom. Bohr also used Planck's
ideas in order to create quantum mechanics, his new concept regarding energy.
This model is still used to this day.
The historical development
of The Atomic Structure
(also covered in transition… here’s
the recap!)
• Schrdinger/Plank/Heisenberg
1900-1927
Erwin Scrhodinger was an Austrian scientist that worked with the Quantum
model of the atom. He disagreed with Bohr's theory, so he created his own. He
thought that the only way to find the location and energy of an electron in an
atom was to calculate its probability of being a certain distance from the nucleus.
This equation influenced the Quantum mechanical model of the atom.
Max Planck developed his ‘Quantum theory’, which states that energy exists in
fixed amounts called quanta.
Werner Heisenberg was a German scientist that proposed the uncertainties of the
Quantum model. He said that you can't know the exact velocity and momentum
of the electron at the same time, which means you can't know the exact location of
the electron. This principle proves error in Bohr's model because of the
uncertainty of the location of an electron.
The historical development
of The Atomic Structure
(also covered in transition… here’s
the recap!)
• Sir James Chadwick:
1932
James Chadwick was an English scientist that discovered the neutron. Before this
discovery, Rutherford had concluded that the nucleus was made of positive
matter. It made sense that the atom was neutral because the negative electrons
and the positive protons cancelled out. But, Chadwick started to question why
there was a difference between the atomic mass and the number of protons.
Chadwick then found that the missing component was a neutral part: the
neutron.
The historical development
of The Atomic Structure
(also covered in transition… here’s the
recap!)
• The Current model:
You’re invited!!!
Write this in your diary!
Atomic Theory Party
You are invited to celebrate the superb discoveries and the contributions
made by many learned Legendary Scientists!
• When: Chemistry class
• Where: A8 – our knowledge inspiring laboratory
• What to bring: Any item of food that can be
shared that you feel is representative of any one of
the models of the Atomic Theory.
How we will start the year…
Unit 1, Area of Study 1
1.The historical development
of the Periodic table√
2.The historical development
of The Atomic Structure√
3.Mass numbers, isotopes and
electron configurations
4.Trends of the periodic table
3. Mass numbers, isotopes
and electron
configurations
3. Mass numbers, isotopes
and electron
configurations
Subatomic Particles
(the recap continues…)
Atoms are composed of three subatomic particles: protons, neutrons and
electrons. The two important properties of these particles are mass and
charge:
Relative Relative
Particle
mass
charge
proton
1
+1
neutron
1
0
electron
1/1840
-1
The mass of electrons is negligible when compared to the mass of protons and
neutrons, so their mass is not included when calculating the mass of the
atom.
3. Mass numbers, isotopes
and electron
configurations
The number of protons in an atom is known as the atomic number or
(the recap continues…)
proton number and is represented by the symbol Z.
The mass number of an atom is the number of protons plus the number of
neutrons, and is represented by the symbol A.
When an atom is represented
by its symbol, the mass
number, and sometimes the
atomic number, are shown.
mass
number (A)
atomic
number (Z)
3. Mass numbers, isotopes
and electron
configurations
Isotopes are atoms of the same element that contain different numbers of
neutrons.
(the recap continues…)
mass number
is different
atomic number
is the same
carbon-12
carbon-13
The reactivity of different isotopes of an element is identical because they
have the same number of electrons.
The different masses of the atoms means that physical properties of isotopes
are slightly different.
3. Mass numbers, isotopes
and electron
configurations
(the recap continues…)
electron configuration
In year 10, you learnt the model of the
atom states that a nucleus is surrounded
by shells of electrons. Each shell holds a
different maximum number of electrons:



1st shell : 2 electrons
2nd shell : 8 electrons
3rd shell : 8 electrons.
At VCE, this model is slightly different. Instead of electrons being arranged
in shells that are a different distance from the nucleus, they are arranged in
energy levels.
It was found that there are actually sub-levels to the shells you have previously
learnt about.
electron configuration
There are four sub-levels (subshells), labelled in order of
increasing energy: s, p, d and f.
Each holds a different number of
electrons.
Each principal
energy level
(shell) contains a
different number
of sub-levels.
sub-level max no. electrons
s
2
p
6
d
10
f
14
principal energy
level, n
1
2
3
4
sub-levels
max no.
electrons
1s
2
2s, 2p
8
3s, 3p, 3d
18
4s, 4p, 4d, 4f
32
electron configuration
electron configuration
electron configuration
As part of his work on electron configuration, Niels Bohr developed the
Aufbau principle, which states how electrons occupy sub-levels.
The Aufbau principle states that the
lowest energy sub-levels are occupied first.
This means the 1s sub-level is filled
first, followed by 2s, 2p, 3s and 3p.
However, the 4s sub-level is lower in
energy than the 3d, so this will fill
first.
electron configuration
electron configuration
electron configuration
Although the 3d sub-level is in a lower
principal energy level than the 4s sublevel, it is actually higher in energy.
This means that the 4s sub-level is filled
before the 3d sub-level.
Example: what is the electron structure of vanadium?
1. Count number of electrons in atom
23
2. Fill sub-levels, remembering
4s is filled before 3d
1s22s22p63s23p64s23d3
electron configuration
electron configuration
of ions
When an atom loses or gains electrons it produces a charged
particle called an ion.
An ion that has gained electrons forms a negative ion.
An ion that has lost electrons forms a positive ion.
That’s because
electrons are
negatively
charged!
electron configuration
ions
When writing the electronof
configuration
of ions, it is important to add or subtract
the appropriate number of electrons.
For non-transition metals, the sub-levels are
then filled as for atoms.
For negative ions
add electrons.
For positive ions
remove electrons.
Example: what is the electron structure of O2-?
1. Count number of electrons in atom
8
2. Add or remove electrons due to charge
8 + 2 = 10
3. Fill sub-levels as for uncharged atom
1s22s22p6
electron configurations
When transition metals form ions, it is the 4s electrons that are removed
before the 3d electrons.
Example: what is the electron structure of Ni2+?
1. Count number of electrons
in atom
28
2. Fill sub-levels, remembering
4s is filled before 3d
1s22s22p63s23p64s23d8
3. Count number of electrons
to be removed
2
4. Remove electrons starting
with 4s
1s22s22p63s23p63d8
electron configurations
Valence electrons
The electrons in the outermost shell of an atom are called the
valence electrons.
The valence electrons have the lowest ionisation energy (require
the lowest energy to remove) and are responsible for most
chemical reactions.
Chemists can usually predict the nature of an element by
knowing the number of valence electrons.
Eg: Sodium has 1 valence electron.
Eg: Chlorine has 7 valence electrons.
How we will start the year…
Unit 1, Area of Study 1
1.The historical development
of the Periodic table√
2.The historical development
of The Atomic Structure√
3.Mass numbers, isotopes and
electron configurations√
4.Trends of the periodic table
4. Trends of the periodic
table
Trends of the periodic
table
The electronic arrangement of an atom made it possible to see why element
properties are periodic.
The alkali metals in the first group of the Periodic Table all have outer
shell electron configurations of s1.
The halogens all have outer shell electron configurations of s2p5.
Mendeleev’s periodic law can be restated as:
Periodic variations in the chemical properties of elements arise from
periodic variations in the electronic structures their atoms.
Trends of the periodic table
So, How is it organised?
The genius of the periodic table is that it is organized like a big
grid. The elements are placed in specific places because of the
way they look and act. If you look at a grid, you know that
there are rows and columns. The periodic table has rows and
columns, too, and they each mean something different.
Trends of the periodic
table
Rows are called Periods
Even though they skip some
squares in between, all of the
rows go left to right. When you
look at a periodic table, each of
the rows is considered to be a
different period (Get it? Like
PERIODic table.)
Trends of the periodic table
Rows = periods
In the periodic table, elements have something in common if they are in
the same row.
All of the elements in a period have the same number of atomic orbitals
(shells).
Every element in the top row (the first period) has one shell for its
electrons. All of the elements in the second row (the second period) have
two shells for their electrons. It goes down the periodic table like that.
That’s why the top row only has Hydrogen (1s1) and Helium (1s2) – The
first shell can only hold two electrons.
Trends of the periodic table
And then, there are groups…
The periodic table has a special
name for its columns, too. When a
column goes from top to bottom,
it's called a group.
Trends of the periodic table
columns = groups
The elements in a group have the same number of electrons in their outer
orbital.
Every element in the first column (group one) has one electron in its
outer shell. Every element on the second column (group two) has two
electrons in the outer shell. As you keep counting the columns, you'll
know how many electrons are in the outer shell.
There are some exceptions to the order when you look at the transition
elements, but you get the general idea.
Trends of the periodic table
blocks of elements
The Periodic Table can be divided into four main blocks of elements.
The s-block contains the Group I and II elements. These elements have s1 or
s2 subshells as their highest energy subshells.
The p-block contains the Group III-VIII elements. These elements have
outer shell configurations of s2p1 to s2p6.
The elements in s- and p-block are sometimes referred to as main group
elements.
Trends of the periodic table
blocks of elements
The d-block contains elements known as the transition metals. It
contains elements in which the 5 orbitals of the d-subshell are being
progressively filled.
The f-block contains the lanthanides and the actinides.
The lanthanides contain elements where the 4f-subshells are being
filled.
The actinides contain elements where the 5f-subshell is being filled.
Trends of the periodic table
blocks of elements
Trends of the periodic table
blocks of elements
Trends of the periodic table
blocks of elements
Trends of the periodic table
blocks of elements
Trends of the periodic table
blocks of elements
Trends of the periodic
table
TREND
Atomic Radius
INCREASES Down
a Group
Atomic Radius
DECREASES across
a period
EXPLANATION
Electrons occupy most of the volume of
the atom.
The size of atoms decreases due to the
increasing positive charge of the
nucleus. This greater core charge pulls
the electrons closer causing the volume
of an atom to reduce.
Trends of the periodic
table
TREND
EXPLANATION
First Ionisation
Energy
DECREASES down
a group
Atoms become larger due to more shells
and the outer electrons are further from
the nucleus feeling less core charge. The
energy required to remove these outer
shell electron decreases.
First Ionisation
As the strength of attraction between the
Energy
outer electrons and the nucleus
INCREASES across increases, the energy required to remove
a period
the outermost electron increases as they
feel more core charge.
Trends of the periodic
table
TREND
EXPLANATION
Electronegativity
DECREASES Down
a Group
Electronegativity
INCREASES across
a period
As the outer electrons become more
distant, electrons feel less core charge so
have a weaker attraction to the nucleus.
The electron attracting ability of atoms
increases as the pull on the outer
electrons increases.
Trends of the periodic
table
Trends of the periodic table
patterns in groups
Even though elements in the same group have very similar
properties there are some major differences as you move down the
group.
Some of the following trends are noticeable as you move down a
Group:
•
•
•
•
•
•
The size of the atom increases
The first ionisation energy decreases
The atom becomes more metallic
Reducing strength of metals increases
Oxidising strength of non-metals decreases
Electronegativity decreases
Trends of the periodic table
patterns in periods
As you move across the period, not only does the number of outer
shell electrons increase but the number of protons increases which
increases the core charge within the nucleus.
Due to this reason, as you move across a period:
The size of the atom decreases
First ionisation energy increases
Elements become non-metallic
Reducing strength decreases
Oxidising strength increases
Electronegativity increases
Trends of the periodic
table
Trends of the periodic
table
The periodic table
Group 1
The periodic table
Group 2
The periodic table
Group 3
The periodic table
Group 4
The periodic table
Group 5
The periodic table
Group 6
The periodic table
Group 7
The periodic table
Group 8
The periodic table
transition metals
The periodic table
lanthanides
The periodic table
actinides
The periodic table
transactinides
How we will start the year…
Unit 1, Area of Study 1
1.The historical development
of the Periodic table√
2.The historical development
of The Atomic Structure√
3.Mass numbers, isotopes and
electron configurations√
4.Trends of the periodic
table√
5. Masses of particles
Masses of particles
The relative isotopic mass of an atom is its mass compared to the isotope
carbon 12 (12C – a carbon atom with 6 protons and 6 neutrons).
12C is said to have a mass of exactly 12 units. All other masses are
relative to this.
Relative atomic mass –Why does the periodic table have
decimal numbers for the mass numbers?
Most elements have more than one isotope. The relative atomic mass of the
element is the average mass of the isotopes taking into account the abundance
of each isotope.
Example: what is the Ar of boron?
In a sample of boron, 20% of the atoms are 10Br and 80%
are 11Br.
If there are 100 atoms, then 20 atoms would be 10Br and
80 atoms would be 11Br.
The relative atomic mass is calculated as follows:
Ar of Br = (20 × 10) + (80 × 11)
100
Ar of Br = 10.8
Masses of particles
relative atomic mass
To calculate the relative atomic mass use:
Ar
(relativeisotopicmass  %abundance)


100
Example 1:
Calculate the relative atomic mass of chlorine.
Masses of particles
Masses of particles
Example 2:
The relative atomic mass of copper is 63.54. If copper
has two isotopes, 63Cu with a relative isotopic mass of 62.95 and
65Cu with a relative isotopic mass of 64.95, calculate the
percentage abundance of the isotopes.
Masses of particles
relative molecular mass
The relative molecular mass (Mr) is the sum of all the relative
atomic masses (Ar) of each atom present in a molecule.
Eg: Oxygen:
Atomic Mass = 16 (From periodic table)
Molecular Mass = Add all the atomic masses together.
Mr (O2) = 2 X 16
= 32
We multiply by 2 because there a two
oxygen atoms in the molecule
Masses of particles
relative molecular mass
If you have an iPad or iPod you might want to get this app to make
calculating molecular mass quicker (there is also an android version)…
ChemCalc
How we will start the year…
Unit 1, Area of Study 1
1.The historical development
of the Periodic table√
2.The historical development
of The Atomic Structure√
3.Mass numbers, isotopes and
electron configurations√
4.Trends of the periodic
table√
6. The mole concept
The mole concept
The mole concept
Because every atom has a different atomic mass, it is difficult to
be consistent when comparing the weights of different substances.
Chemists find it easier to deal with the number of particles of
the substance.
The standard amount of particles is known as a mole
The mole concept
1 mole contains the same amount of particles contained in
exactly 12g of 12C.
The number of particles in 1 mole of any substance is
6.02  1023.
This amount is known as Avogadro's Constant (NA).
The mole concept
When calculating the amount of a substance we use the formula:
Where:
n is the amount of the substance in moles
N is the number of particles present
NA is Avogadro’s constant
The mole concept
molar mass
The Molar Mass (M) of a substance is the mass of 1 mole of
that substance.
Since 1 mole of 12C is exactly 12g, the molar mass of an element
is the same as its relative atomic mass.
Given that molar masses vary greatly between substances, masses
of one mole of different substances can also vary greatly.
The mole concept
molar mass
Figure 4.10 The equivalent of one mole of some common elements and
compounds.
The mole concept
counting by weighing
Using the molar mass we can calculate the amount of a
substance if we know the mass. This is done by using the
formula:
Where:
n is the amount in moles
m is the mass in grams
Mr is the Molar Mass
m
n
Mr
The mole concept
formulas of compounds
• Using molar masses we can calculate the percentage
composition of each of each element within a compound.
• This can be calculated by:
%composition 
molar _ mass _ of _ element _ in _ compound
100
total _ molar _ mass _ of _ compound
The mole concept
finding the composition of a
compound
Using chemical analysis we can find the simplest whole number
ratio of atoms in a compound.
This is called the empirical formula.
Empirical formulas are determined experimentally, usually by
finding the mass of each element in a given compound.
The mole concept
empirical formula
Calculating the empirical formula:
Measure the mass of each
element in the compound
Calculate the amount in mole of
each element in the compound
Calculate the simplest whole
number ratio of moles of each
element in the compound
Empirical formula of the
compound
The mole concept
empirical formula
Example 1
A compound of carbon and oxygen is found to contain
27.3% carbon and 72.7% oxygen. Calculate the
empirical formula of the compound.
Example 2
9.0g of a compound of carbon, hydrogen and oxygen is
found to contain 4.8g of oxygen and 3.6g of carbon.
Calculate the empirical formula of the compound.
The mole concept
molecular formula
The molecular formula of a compound gives the actual
number of atoms present in a molecule of a compound.
For example, the empirical formula of glucose is CH2O. Its
molecular formula is 6 times larger – C6H12O6.
A molecular formula can be deduced from the empirical
formula of a compound if the molar mass is known.
The mole concept
molecular formula
Example
A sample of a hydrocarbon was found to contain 7.2g of
carbon and 1.5g of hydrogen. The molar mass of the
compound is 58.
What is the molecular formula of the compound?
How we will start the year…
Unit 1, Area of Study 1
1.The historical development
of the Periodic table√
2.The historical development
of The Atomic Structure√
3.Mass numbers, isotopes and
electron configurations√
4.Trends of the periodic
table√
Shhh…