Chapter 3: The Periodic Table

Week 2, Lesson 1

• 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.

• But…it did not adequately explain why substances had certain properties and behaved as they did.

Week 2

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.

• 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.

Mendeleev’s 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.

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.

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.

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

Mendeleev’s 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.”

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.

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.

Trends in Properties

• 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.

Week 3, Lesson 1

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.

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.

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.

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.

Trends in the Periodic TableTrend 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.

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.

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.

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.

Chemical Reactivity of ElementsReactivity 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.

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.

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.

Compounds• 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.

Week 3, Lesson 2

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.

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

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.