Chapter 3: The Periodic Table

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Chapter 3 The Periodic Table
Week 2, Lesson 1
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Why is the Periodic Table Important?
Week 2

• Mendeleev developed the periodic table based on
  discoveries by Boyle, Lavoisier, Dalton and
  others.
• It organised gathered information based on
  chemical properties and behaviour.
• It also became a helpful tool for making
  predictions and guiding future developments.
• Butit did not adequately explain why substances
  had certain properties and behaved as they did.

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Continued

• The ideas from Chapter 1 and the understanding
  discussed in Chapter 2 has made the periodic
  table the most useful reference available.
• It means that chemists do not have to memorise
  facts about individual elements, rather they can
  just identify patterns.
• It also provides a framework in which to organise
  knowledge.

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Why are Element Properties Periodic?

• Mendeleev discovered sets of elements with
  similar chemical properties.
• He discovered that as he looked at the elements
  in increasing atomic mass they changed from
  metals, to non metals and back to metals again.
• Similar patterns existed for other properties.

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Mendeleevs Periodic Law

• The properties of elements vary periodically
  with their atomic weights.
• The number of protons in an atom is now known to
  be the fundamental difference between atoms of
  different elements, not atomic weight.
• Because of this we know see the periodic table
  organised based on ascending atomic number.

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Group 1 Elements

• Period 1 Elements are known as alkali metals.
• These elements all have similar properties and
  are relatively soft metals that react highly with
  water and oxygen.
• Have a look at their electron configurations.
• Li 1s22s1
• Na 1s22s22p63s1
• K 1s22s22p63s23p64s1
• Rb 1s22s22p63s23p63d104s24p65s1
• Cs 1s22s22p63s23p63d104s24p64d105s25p66s
• They all have similar outer-shell configurations
  with their valence electrons being in the s
  subshell.

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Group VII or 17 Elements

• Are known as halogens.
• Again these elements have similar properties in
  that they are all coloured and are also highly
  reactive.
• They have the following electron configurations
• F 1s22s22p5
• Cl 1s22s22p63s23p5
• Br 1s22s22p63s23p63d104s24p5
• I 1s22s22p63s23p63d104s24p64d105s25p5
• These elements all have the same valence electron
  configuration, being s2p5.

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Periodicity and Element Properties

• The arrangement of electrons in atoms is
  responsible for the periodicity of element
  properties.
• Properties such as melting temperature,
  electrical conductivity, formulas of compounds
  formed when two elements react, and many others
  depend on the way the electrons are arranged,
  particularly the outer shell configuration as
  these are the electrons that are involved in
  bonding

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Mendeleevs Law Restated

• Variations of the chemical properties of
  elements across a period and similarities down a
  group are all associated with electronic
  configurations of their atoms.

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Patterns in Electronic Structure

• Vertical columns (groups) contain elements with
  similar outer-shell electron configurations. They
  are numbered I-VIII or 1-18.
• Horizontal rows (periods) each period contains
  elements with electrons in the same outer shell.
  The number of the period is the same as the
  number of the outer shell.

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Blocks of Elements

• There are four main blocks of elements, with the
  elements in each block filling the same type of
  subshell.
• s-block contains elements in groups 1 and 2.
• p-block contains elements in groups 13-18
• Although helium has an electron configuration of
  1s2, it is usually located in group 18 because
  its atoms are unreactive.
• d-block contains the transition metals.
• f-block contains the lanthanides and actinides.
  Lanthanides are a set of 14 elements with the
  atomic numbers 58-71 and actinides have atomic
  numbers 90-103.

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Trends in Properties
Week 3, Lesson 1

• The periodic variation on the properties of
  elements reflects the periodic variation in their
  electron configurations.
• This is more clearly seen if you look at how
  properties change from left to right across a
  period.
• There are similarities between elements in a
  group but there are also significant differences.

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Trends in Properties cont

• Although elements in the same group have the same
  number of electrons in their outer shells, the
  atoms become larger as you move down and this
  also effects the properties of the elements.
• We can look at the properties of elements in two
  ways
• The properties that relate to individual atoms of
  the element, ie electron configuration.
• The properties that reflect the way atoms or
  groups of atoms interact with each other.

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Atomic Properties

• Radius, ionisation energy and electronegativity
  all depend on the strength of the attraction
  between the valence electron and the nucleus.
• This attraction will generally depend on
• The positive charge that attracts the valence
  electrons, or,
• The distance of these electrons from the nucleus.

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Electrons and the Periodic Table

• The electrons of atoms of elements in the same
  period are located in the same outer shell.
• As you move across a period the number of protons
  increases, meaning the numbers of electrons also
  increases.
• The outer shell electron of lithium is attracted
  towards the three protons. However it does not
  feel the full charge as the two inner electrons
  act as a shield. The outer electron experiences a
  core charge of 1.
• The outer shell electrons in fluorine experience
  a core charge of 7 as the two inner electrons
  shield the outer electrons from the 9 protons.

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Electrons and the Periodic Table cont

• The outer shell of a sodium atom has a core
  charge of 1 because the 1 outer electron is
  attracted to the 11 protons however the first two
  complete shells are shielding the attraction.
• The core charge is the same moving down a group
  but the number of electron shells increases. As
  such the valence electrons are held less
  strongly.
• The electrons of elements in the same period are
  in the same outer shell. As you move across the
  period the core charge increases, therefore
  increasing the attraction towards the nucleus.

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Trends in the Periodic Table
Trend Explanation
Atomic Radius increases down a group Electrons occupy most of the volume of an atom.
Atomic Radius decreases across a period The size of atom decreases due to the increasing positive charge of the nucleus attracting the electrons.
First Ionisation Energy decreases down a group The atoms get larger and as such valence electrons are further away from the nucleus. The energy needed to remove the outermost electron decreases.
First Ionisation Energy increases across a group The attraction to the nucleus is higher and as such more energy is required to remove the outer electrons.
Electronegativity decreases down a group As the outer electrons become further away from the nucleus, electrons are more weakly attracted to an atom.
Electronegativity increases across a period The electron-attracting ability of atoms increases as the pull on the outer electrons increases.
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Defintions

• Atomic Radius The size of the atom.
• First Ionisation Energy The amount of energy
  required to removed the first electron.
• Electronegativity The ability of an atom to
  attract an electron towards itself.

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Metallic and Non-Metallic Character

• Elements on the right hand side of the table are
  non-metals, while the other elements are metals.
• Moving from left to right across a period,
  elements become less metallic and exhibit more of
  the non-metal properties.
• There is also a variation in metallic character
  within groups. Looking at group 14, Carbon is a
  non-metal, where tin and lead are both metals.

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Metalloids

• Some metals like germanium, silicon, arsenic and
  tellurium which are located across periods
  display both metallic and non-metallic
  properties.
• These are called metalloids.

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Chemical Reactivity of Elements

• Reactivity of Metals
• The way in which metals react with water can give
  us an indication of their relative reactivity.
• Those elements in group one are more reactive
  than those in group 2.
• As you move down a group the reactivity of the
  metal increases.
• Generally Reactivity increases down and group
  and decreases across a period.

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Chemical Reactivity cont

• Reactivity of Non-Metals
• We can get an indication of the reactivity of
  non-metals through reacting them with aqueous
  solutions, for example potassium iodide.
• This shows that the reactivity decreases down a
  group. This is because the electronegativity, or
  the ability to attract electrons is higher at the
  top of a group.

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Noble Gases

• Noble gases are all unreactive .
• These elements have very low melting and boiling
  temperatures and all are gases at room
  temperature.
• Their lack of reactivity is due to the electron
  configuration. Each noble gas has an outer shell
  that is considered full or stable.
• As such they do not want to react with other
  elements and unbalance their stability.

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Compounds
Week 3, Lesson 2

• Both compounds and elements are considered to be
  pure substances.
• Pure Substance matter that always has the exact
  same composition.
• Compounds are formed when atoms of two or more
  elements chemically combine in fixed proportions.
• Each compound has its own set of properties and
  these can be quite different from the elements
  that it is made up of.

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Compounds cont

• If we look at salt as an example
• Salt, sodium chloride is made up of sodium and
  chlorine.
• On the next slide is a table highlighting the
  significant different characteristics of sodium,
  chlorine and sodium chloride.

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Sodium (Metal) Chlorine (Non Metal) Sodium Chloride
Melts at 98 degrees Melts at -101 degrees Melts at 801 degrees
Conducts electricity when solid Does not conduct electricity when solid Does not conduct electricity when solid
Conducts electricity when molten Does not conduct electricity when molten Conducts electricity when molten
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Molecule

• A molecule is two or more non metal atoms
  chemically combined, for example water.
• Both hydrogen and oxygen are both gases under
  normal conditions, yet water is a liquid.
• Most compounds that are produced from two
  chemically combined non-metals generally have
  similar chemical properties to water.