Unit 1/2 Chemistry
Miss Schram 2015
Elements
•
Elements on the periodic table are represented by symbols. Eg: C, N, Hg.
(Always capitals, then lower case)
•
The periodic table is a collection of information. On the periodic table, the
symbol has the atomic number (Z) at the top, and atomic mass (A) at the
bottom.
•
When we write symbols of elements in a text or such, the atomic number (Z) is
actually at the bottom, and the atomic mass (A) on the top.
•
The Atomic number (Z) represents the number of protons in the element.
•
The Atomic Mass (A) represent the number of protons and neutrons in the
element.
•
The Z number, determines what element the atom is. And each element has a
different symbol.
Formulas for elements
• Most elements exits as a single atom.
• The exceptions are the "magic 7" and phosphorus and Sulfur.
• Magic 7 are H, N, O, F, Cl, Br, I, who all exist as diatomic atoms
(X2). P4 and S8.
In class:
• Complete Lucarelli Set 5 (Q1, 2, 3)
Atomic Structure
Atoms are composed of 3 sub-atomic particles:
• Proton
• Neutron
• Electron
•
The nucleus is where protons and neutrons are contained. The nucleus
is quite small in size but contains nearly all of the atoms mass.
•
Electrons are found very far from the nucleus in the electron cloud.
Electrons move extremely fast, are very small and are spread out from
each other.
•
The nucleus has as overall positive charge and the electron cloud is
overall negative.
•
This current understanding of the atom is has evolved over time, and
your first assignment is to discover about the history of the atom's
evolution.
Atomic Structure
Sub
Atomic
Particle
Symbol
Charge
Relative
Mass
Location
Proton
p
+1
1
1800
Nucleus
Neutorn
n
0
1
1800
Nucleus
Electron
e
-1
1/1800
1
Electron
cloud
Atomic Structure
Sub
Atomic
Particle
Symbol
Charge
Relative
Mass
Location
Proton
p
+1
1
1800
Nucleus
Neutorn
n
0
1
1800
Nucleus
Electron
e
-1
1/1800
1
Electron
cloud
In class:
• Complete Lucarelli Set 2 (Q 6-16)
The Nucleus (pg 107)
• What happens when you hold two same ends of a magnet together.
• Electrostatic repulsion is what happens when two of the same
charges are near by
• So why doesn’t the protons in the nucleus, (multiple areas of dense
positive charge), repel the neculus and the atom apart?
The Nucleus
• There is another force, a strong nuclear force which balances out
the repulsion.
• This occurs between all particles (protons and neutrons) in the
nucleus, and is attractive.
• The strong nuclear force and electrostatic repulsion are balanced
and so the nucleus is stable and will not decay over time.
The Nucleus and the Electron Cloud
• The nucleus is over all positive and electrons are negative in
charge.
• electrostatic attraction is in action (2 opposite ends of a magnet)
• So in theory the electrons should be pulled into the nucleus as they
are lighter in mass?
The Nucleus and the Electron Cloud
• This doesn’t happen as the electrons are traveling at very fast
speeds around the nuclue ( like a bucket over youe head full of
water, if you spin it fast enough the water will not fall out)
• But if the electrons are traveling so fast, why don’t they fly off?
The Nucleus and the Electron Cloud
• The electrostatic attraction is what hold the atom together.
• Electrostatic attraction is the attractive force between two opposite
charges. Eg + and -.
• The positive nuclues attracts the electrons and they don’t fly off. The
forces are balanced again.
In Class:
•
Complete Question set 1.1 on pg 108.
Relative Atomic Mass (Ar) and Isotopes
• If A-Protons and Z= Protons + Neutrons, and the masses are all
considered an even 1, why are the masses of some elements, not
perfect whole numbers?
Relative Atomic Mass (Ar) and Isotopes
• Isoptopes are elements with a different number of neutrons in that
element.
• Carbon has 3 isotopes, C-12, C-13 and C-14.
• Determin the number of protons, neutrons and thus the atomic mass
for these each.
• Each has a different weight, how did we decide on the exact mass
for carbon?
Relative Atomic Mass and Isotopes
• This is because the mass is an average of the weight of that
element due to the different abundances.
• Eg carbon has 3 common isotopes of carbon-12 (98.9%), carbon-13
(1.1%) and carbon-14 (very minimal).
• The mass number is calculated via the calculation below.
• Mass number = [(abundance % x atomic mass) + (abundance %)] /
100
• http://www.kentchemistry.com/links/AtomicStructure/atomicmasscalc
.htm
Rea
• Some isotopes are stable as mentioned due to the attractive and
repulsive forces in the nucleus. Others are unstable and undergo
radioactive decay in an attempt to become stable.
• They have the same chemical properties (due to ?)
• Different physical properties (due to ?)
Relative Atomic Mass (Ar) and Isotopes
•
The mass number (calculated from the abundance of each element) is used
to determine the relative atomic mass (Ar).
•
Relative atomic mass of an atom, is the mass of an atom compared to the
mass of a carbon-12.
•
Carbon is considered 12 and every other atom is weighed against this.
•
Eg: 12 H atoms weigh the same as 1 C atom, there for 1 H atom must be
1/12 the mass of a C atom.
•
1 Mg = 2 C so Mg must be 2/1x mass of C. so 24.
•
Carbon was chosen as it is common and is relatively chemically stable.
In class:
•
Complete Question set 1.2 pg 110.
• Complete Lucarelli Set 2 Q6-16
The periodic table
• The periodic table was first developed by Dmitri Mendeleev (1869).
• Had rows and periods and elements were arranged by increasing
atomic number and grouped by trends.
• He left gaps for elements he believed to exist.
The Periodic Table
• The elements are listed in order of their atomic number
• Periods across
• Groups are down
• Groups are numbered and named.
– valence electrons thus similar properties
• Groups= number of valence electrons
• Period= number of energy levels/shells
Electrons
• the number of electrons = number of
protons (=atomic number)
• eg: element 47 has 47 protons so 47
electrons
• this is because we have 47 positive
charges so we require 47 negative
charges to balance out
• This only works for neutral or ground
state atoms (those with no charge)
Electron Shells
• Electron cloud- general area where
electrons found
– electron shells- discrete energy levels
where electrons are
Ions
• atoms with an unbalanced number of
protons and neutrons
• lost an electron= positive
– more neutrons than electrons
• gained an electron= negative
– more electrons than protons
• eg: element 67 forms charges of 3+.
• (67+) + (64-) = (+3)
Determining Charges
• atoms lose electrons to have as many electrons as the Nobel gases
(full shell)
• Group number
– tells us how many valence electrons the next shell down is full
– Group 1-13 positive
– group 14-17 negative
– Noble gases neutral
The Periodic Table Trends
Atomic Radius
• Down groups: increase
– Lower elements have more filled shells. Distance between
nucleus and outer levels increases, and so does the atomic
radius.
• Across periods L-> R: decrease
– Extra protons in nucleus so it is more positive. Negative
electrons are pulled in more strongly so they move in more.
Atomic Radius
Ionisation Energy:
Energy required to remove electrons from gaseous atom
• Down Groups: decrease
– Electrons in the shells are held further out, and so are less
strongly bound to the nucleus. Takes less energy as there is less
attraction.
• Across periods L->R: Increase
– Electrons are closer to the nucleus and so they are more
strongly bound to the nucleus, and it takes more energy to
remove the electrons.
Ionisation Energy
Electronegativity
energy required to attract and form bonds with electrons
• Down groups: Decrease
– Empty shells are further away from the nucleus and so any new
electrons are not as attracted to the nucleus.
• Across periods L->R : increase
– The tightly held in electron shells are closer and so there is a
stronger attraction between the nucleus and near by electrons of
other atoms.
Electronegativity
Electron affinity:
ability of atom in gas state to accept an electron and form negative ion,
full addition of a electron
• Down Groups: Decreases
• Across periods: L-> R: Increases
Reasons same as electronegativity
Electron affinity
Metallic character
• Down: Increasing
• Across: Decreasing
The Periodic Table Trends Summary
Lewis structures/ion formation:
• Lewis: shows valence electrons only
• Bohr: All electrons
• Electrons are placed around the symbol separate then pairs.
• Electrons are shown as dots or crosses (be consistent)
• Try the first 10
Lewis Structures/Ion Formation
• Electrons are lost, so they have configurations of the closest noble
gas. (octect rule)
– happens as electrons are very stable and elements want to be stable.
• Charge is determined by the number of valence electrons and
nature of the element (ionisation and electron affinity trends).
• Metals: positive and charge
• Non-metals: negative, square brackets, charge outside
Electron arrangement
• Shells/energy levels are 1-7. and each element has these
shells/energy levels .
• Each energy level has a specific amount of energy.
• Number 1 is the lowest energy and 7 highest.
• Electrons have to be in a energy level. They can not sit between
levels.
• Electrons can transition between levels by the gain and loss of
energy. (prac later)
Electron Arrangment
• Each energy level can only hold a specific number of electrons.
Formula we can use is : 2n2 n=energy level
– Or remember 2, 8, 18, 32
– Fill the lowest (1) to the highest (7) and the shell must be completely filled before
the next shell can start to be filled.
– Electrons always occupy the lowest energy levels first
Electron configuration
•
Each element has electrons arranged in a specific way that is unique to that
element.
•
–
like coordinates to the element
–
How many numbers= how many occupied energy level (eg period)
–
Last number= number of e in valence shell (group)
–
Add all electrons together= atomic number of neutral atom
H= 1
–
•
C= 2,4
–
•
Period 1, group 1, atomic number 1
Period 2, group 4, atomic number 6
If we work in reverse, we can write our own electron configurations 
In Class
• Complete the work sheet for the electron configurations for the first
18 elements.
Electron Configuration- Exceptions
• Potassium:
– Potassium is Group 1, Period 4.
– Must be 2,8,8,1
– The 3rd shell, can hold a max of 18 electrons, but potassium does not have the
configuration of 2,8,9 as expected.
– This is because this is not energetically favourable
• Elements in the transition group (21-30), once the 4th shell has 2
electrons, the 3rd shell continues to fill.
• Once the 3rd shell is completely full to 18, then the 4th shell
continues to fill.
In Class
• Finish off electron Configurations Sheet
Answers
• K= 2, 8, 8, 1
• Ca= 2, 8, 8, 2
(now full valence shell, shell 3 continues to fill up)
• Sc= 2, 8, 9, 2
• Ti= 2, 8, 10, 2
Electron Configurations
• Summary:
– Strictly speaking the maximum number of electrons in each shell is 2, 8, 18, 32.
– Electron filling order:
Electrons
in the
atom
1st
2nd
3rd
4th
Up to 20
2
8
8
2
18
8
2
18
8
21-38
39-56
• Complete hand out.
5th
6th
2
In Class:
• Lucarelli set 9 (q 1-12)
Elemental Spectra
•
Element's electrons are in different energy levels and all
element have these 7 shells.
• Each electrons shells have different spaces.
• This means each elements shells have different amounts of
energy.
Elemental Spectra
• Electrons can move between shells by the gain
or loss of energy.
• Gaining energy means going up a energy level.
– This excited state, only lasts a few seconds before electron
releases energy and moves down
• Loss of energy means going down a level.
Atomic Emission Spectroscopy
• This energy is released as light.
• White light we see is a spectrum or can be split into smaller
segments. (rainbows)
• The energy released by each element is of a specific wavelength.
Atomic Emission Spectroscopy
• Flame test (we see one colour as we have not split the light, it is an
accumulation)
• We can split the light (analysis of light is spectroscopy) emitted with
a spectroscope to find the "line emission spectrum" or the specific
energy wavelengths emitted by the element.
• This pattern is the same and specific for each atom of each element.
Like a fingerprint
• https://www.youtube.com/watch?v=L5jEePYQfCM (13 min)
In Class:
• Prepare for the experiment on Wednesday (Experiment 8 Flame)
Atomic Absorption Spectroscopy
•
Element must be known.
•
A lamp made with that element is turned on.
•
sample is then vaporised,
•
atoms absorb the light from the lamp. (opposite of the emission spectra)
•
Only the element will absorb the light
•
Light and it passes through a slit, and monochromators (selects one
wavelength (colour) for analysis)
•
Detector measures the intensity of the remaining light.
•
It's a measure of how much light has not been absorbed by the sample.
ATOMIC ABSORPTION SPECTROSCOPY
• Quantitative: Determines how much of element is present. Destroys
sample.
• Video AAS(9min) (3 min mark how it works)
Atomic Absorption Spectroscopy
•
Element must be known.
•
A lamp made with that element is turned on.
•
sample is then vaporised,
•
atoms absorb the light from the lamp. (opposite of the emission spectra)
•
Only the element will absorb the light
•
Light and it passes through a slit, and monochromators (selects one
wavelength (colour) for analysis)
•
Detector measures the intensity of the remaining light.
•
It's a measure of how much light has not been absorbed by the sample.
Atomic Absorption Spectroscopy
• A calibration curve must be made before hand.
• Exact known quantities for that element are sent through and the
amount of light absorbed is recorded.
Mass Spectrometry
• Based of the different masses of atoms in a
sample.
• Determines: what elements present, and or
what isotopes are present.
• focus on isotopic composition
• Must be carried out at low pressure
Mass Spectrometry
• Atomisation
• Bombardment with high energy electrons or
UV light
– removes electrons making all atoms charged
• Ions are accelerated through electric field
• Pass through magnetic field and are deflected
– Lighter ions deflect more
• Detector (amount of ions) and read out is
given (mass/charge (m/z) vs abundance)
Mass Spectrometry
Read: Nelson Pg 22-26
On isotope uses and
formation
Chapter 2: Classifying and separating
substances
Matter
• Atoms are arranged to form many different materials.
• these substances can be classified by either the substances
the material is made of and the way they are combined.
•
•
•
•
Mixture: 2 or more different substances combined
Pure Substances: made up of one type of particle
Elements: one type of atom
Compounds: pure substances with more than one type of
atom, in a fixed proportion.
Matter
• Heterogeneous: material does not have a uniform
composition throughout.
• Homogeneous: material has a uniform composition
throughout.
• Solution: consists of a solute dissolved in a solvent
• Solute: a dissolved material
• Solvent: the dissolving material
the solute in dissolved in the solvent to make a solution
Classification of Materials
Matter
Homogeneous
Pure
substances
Elements
Compounds
Solutions
Heterogenous
Mixtures
Properties of Materials
Physical- Determined
without changing comp
mp/bp, strength,
density, malleability,
ductilty, electrical,
thermal, solubility, state,
etc
Chemical- Reactions to
from new substances
decomposition by
heat, effect of light,
reactions with water,
acids, bases, oxygen, etc
Chemicals can undergo chemical and physical changes.
Discuss.
Signs of Chemicals Change
•
•
•
•
•
Temperature change
Precipitate
Insoluble solid disappears
Gas produced
Colour change
Separating Mixtures
• Mixtures: not chemically combined. Can be
separated by physical or mechanical means
•
•
•
•
•
•
Difference in particle size
Boiling points
Density
Solubility
Magnetism
Electrostatic attraction
Difference in particle size
Sieving
(Solids and solids, solids and
liquids)
• colander
• Material with holes
• Catches larger particles
• Smaller ones fall through
Filtration
(Solids and liquids, solids and
gas)
• Filter paper
• Residue is trapped
• Filtrate passes through
Boiling points
Vaporisation (large Distillation
Fractional
Distillation
bp difference)
(at least 50 difference)
(Solid and liquid)
• Solvent is
boiled or
evaporated off
• Solute
remains
behind
(Solid and liquid, (BP much closer)
liquid and liquid) (liquid and liquid)
• Vapor is cooled • Heated and
and collected
components
as distillate
with different
BP rise to
different
heights
Boiling points
Boiling points
Boiling points
Density
• (Liquid and liquid) Separating funnel is used to
separate two immiscible liquids.
• (Solid and Solid) gold panning
Solubility
(Solid and Solid)
Add solvent suitable to one solid.
Filter with filter paper and evaporate off solvent.
Magnetism
• Degree to which a substance is attracted to a
magnetic field
• Iron, nickel, cobalt (and their ores/some
alloys)
Electrostatic attraction
• Difference in electrical charge
• Mixture brought into a field, particles are
attracted and repelled, caught separately
• Zircon, rutile and monazite
In Class Work
Read 33-48
Bonding