Periodicity Flashcards
Define “Ionisation Energy” and its use
the minimum energy required to remove the most loosely bound or outermost electron from a gaseous atom in its ground state
- useful for determining the properties of the element
- indicator for the strength of the force operating between the nucleus and the electrons
What affects the magnitude of ionisation energy?
- the distance of the electrons from the nucleus
the size of the nuclear charge - the presence of inner shells (shielding)
What is the equipment used to remove an electron?
a discharge tube where the electrons always move away from the cathode.
- the electron from the cathode strikes a gaseous atom X and ionises it, and an electron is removed from X and a gaseous ion is formed
How many ionisation energies would an electron with n electrons have?
n amount
Why does it take more energy to remove an electron in each successive ionisation energy?
it is harder to remove an electron from a positive ion compared to a neutral atom because the electron has more electrostatic attraction and less electron repulsion.
What happens if there is a massive jump in ionisation energy?
- a cation is removing an electron from a different and stable shell
- successive elements where the noble electron configuration with a full shell requires a lot of ionisation energy
Which atoms would be located at the peaks (for ionisation energy)
noble or group 7 gases
Which atoms would be located at the troughs (for ionisation energy)
group 1 and group 2 metals
How does ionisation energy increase?
increase across a period
decrease down a group
Why would oxygen have a lower ionisation than nitrogen and sulfur less than phosphorous?
both nitrogen and phosphorous have half-filled p orbitals, which is a stable configuration.
Are the reactivities or metal consitent with the ionisation values>
No, because reactivities depend on how metals react with water and acids, as well as how the metals are packed in their lattices, and the energy required to disrupt the lattice.
What are the representative elements?
The elements of the main block, Group I - VIII
What are the two types of radioactive elements?
synthetic and natural
What is the significance of elements in the same group?
- they have similar chemical properties as they will lose or gain the same number of electrons
- this is so that they can attain a noble gas configuration
Define “electronegativity”?
- a measure of the electron attracting power of an atom in a molecule
- determined using the Pauling scale
Define “metallic character”
having the physical properties which are usually attributed to metals:
- ductility
- malleability
- hardness
- ability to conduct electricity and heat well
Define “non-metallic character”
having the physical properties which are attributed to non-metals:
- being brittle
- poor conductor of heat and electricity
Define “chemical reactivity”
the tendency to form compounds and the rate at which this occurs
Define “electron affinity”
a measure of the energy released when one electron is added to an atom of an element
Why is there a change from non-metallic to metallic properties as a group descends?
increase in shell number means that the valence electrons are located further from the nucleus and therefore more easily lost by atoms
Why is there a change in atomic radius across a period?
as the nuclear charge increases, so do the number of protons and therefore pulls the electron shells inward and decreases the atoms size
Descending a group, how does ionisation energy change and why?
decreases, as the outer electrons are further from the nucleus, therefore less electrostatic attention, less energy needed to lose an electron
Descending a group, how does atomic radius and why?
increases, new shell of electrons
Descending a group, how does electronegativity change and why?
decreases, because the atomic radius increases there is less attraction for electrons
Descending a group, how does reactivity change and why?
increases for metals
decreases for non-metals
- the most reactive metals are the least electronegative
- the most reactive non-metals are the most electronegative
Descending a group, how does valency change and why?
stays the same for Group I, II, VI, VII, because valence electrons are well shielded from the nucleus, they are lost more easily
Across a period, how does ionisation change change and why?
increases, as the electrons are smaller
Across a period, how does atomic radius change and why?
decreases, nuclear charge decreases therefore an the atom’s outer electrons more strongly
Across a period, how does electronegativity
change and why?
increases, smaller electrons have the highest electronegativity
Across a period, how does reactivity change and why?
decreases for metals
increases for non-metals
Across a period, how does density change and why?
increases, because the atomic radius decreases, whilst the atomic mass increases
What is electronegativity used for?
determining the bonding type of a compound
What type of bonding occurs between large differences in electronegativity?
ionic, metals tend to have low electronegativity and non-metals will have high electronegativity; the strongest ionic bonds occur between alkali metals and halogens
What type of bonding occurs between small differences in electronegativity?
covalent, which can be separated into different levels of polarity. molecules can range from very polar to non-polar.
What results in a non-polar substance
an electronegativity difference of zero
Define “polarity”
a measure of how evenly the bonding electrons is shared
non-polar: bonding electrons shared evenly
polar: bonding electrons shared very unevenly.
What is the diagonal relationship
one step down one explains the similar between one element and another. –> this can be used with electronegativity as an explanation
Define effective nuclear charge
the net nuclear charge after considering the lessening of nuclear charge from the closed inner shells
Cations are smaller than their parent atoms
Anions are larger than their parent atoms
Describe the conductivity of non-metals
they are negligible
How does an element conduct electricity?
must contain electrons that are free to move, otherwise known as mobile charge carriers
Why are metals good conductors?
they have metallic bonding which means positive metals ions are attracted to delocalised and mobile electrons
- This means they are free to move and carry charge
Why are metalloids not good conductors?
they exist in giant lattice structures
- which means few electrons have enough energy at room temperature to enter high energy levels
- this means there are only a few delocalised electrons
- but at higher temperatures electrons are promoted to the higher energy shells
- therefore there are more delocalised electrons to carry charge
Why are non-metals not good conductors?
they are held in strong covalent bonds
What are melting and boiling points depenent on?
strength of forces between particles
- the stronger the force the more energy needed to overcome
- therefore higher temperature
What happens when a substance melts?
attractive forces are broken or loosened, such that the particles can move freely around each other but are still close
What happens when a substance boils?
the remaining attractive forces are broken so the particles can move freely and far apart
What is the trend of melting points across a period?
metals - high
metalloids - very high
non metals - low
Why do metals have high boiling points?
they have metallic bonding where the positive metal ions are attracted to delocalised mobile electrons
Why do metalloids have high boiling points?
some exist in network covalent bonds, which is a giant lattice structures in different arrangements like tetrahedral.
- they for a giant molecule or macromolecule
- they have high melting points because the covalent bonds are strong.
