module 3 - the periodic table Flashcards
giving Mendeleev the credit he deserves
How did Mendeleev arrange the elements across a period?
In order of atomic mass!
How did Mendeleev arrange the elements in groups?
He put the elements in groups with similar properties
what does the number of the period tell you?
the highest energy shell number, in an element’s atoms
What is periodicity??
A repeating trend in properties of elements across each period
How is the modern periodic table arranged?
In order of increasing atomic number
What is the chemistry of each element determined by?
Its electron configuration
What does each period start with?
An electron in a new highest energy shell
What are the 5 properties that show periodicity across the periods?
- electron configuration
- metal to non-metal
- ionisation energies
- structure
- melting point
What is the trend in electron configuration for each period?
The s and p sub-shells are filled in the same way
What is the trend in electron configuration in each group?
Elements in each group have atoms with the same number of electrons in each sub-shell
What des the similarity in each group give to those elements in that group?
Similar chemistry/ chemical reactions
What name is given to group 1 elements?
Alkali metals
What name is given to group 2 elements?
Alkaline earth metals
Which group did Mendeleev miss out in the periodic table? why was he unaware of it?
Group 18 (noble gases). They’re very unreactive and no elements from that group had been discovered at the time.
What does ionisation energy measure?
How easily an atom loses electrons to form positive ions
What is the first ionisation energy?
The energy required to remove 1 electron from each atom in 1 mole of gaseous atoms of an element to form 1 mole of gaseous 1+ ions
Write the equation for the 1st ionisation energy of sodium.
Na (g) —> Na+ (g) + e -
what energy level is the first electron lost from? What’s its attraction to the nucleus like?
- highest energy level
- experiences the least attraction from the nucleus
what 3 factors affect the attraction between the nucleus and the outer electrons of an atom (and the ionisation energy)?
- atomic radius
- nuclear charge
- electron shielding
how does atomic radius affect nuclear attraction with the outermost electrons?
- greater distance between nucleus and outer electrons = less nuclear attraction
- atomic radius has a large effect
how does nuclear charge affect nuclear attraction with the outer electrons?
more protons in nucleus = greater attraction between nucleus and outer electrons
how does electron shielding affect nuclear attraction with the outer electrons?
- inner shell electrons repel outer shell electrons = shielding effect
- reduces attraction between nucleus and outer electrons
In helium, why is the 2nd ionisation energy greater?
- after the 1st electron is lost, the last electron is pulled closer to the nucleus
- the nuclear attraction on the remaining electron increases and more ionisation energy is needed to remove it
define the second ionisation energy?
the energy required to remove 1 electron from each atom in 1 mole of gaseous 1+ ions of an element to form 1 mole of gaseous 2+ ions
when looking at an equation that shows ionisation energy, how can you tell which ionisation energy it is?
it’s the same as the charge on the ion produced
what do successive ionisation energies provide evidence of?
of the different electron energy levels in an atom
on a graph showing successive ionisation energy, what goes on the x and y axis?
x axis = ionisation number
y axis= ionisation energy
what trend is usually shown on a graph showing the successive ionisation energies of an atom of an element?
increase in ionisation energy
what does the large increase between 2 successive ionisation energies show?
- electron with the higher successive ionisation energy must be removed from a different shell
- that’s closer to the nucleus and with less shielding
what do periodic trends in first ionisation energies provide important evidence of?
the existence of shells and sub-shells
when looking at a graph that shows the trend in 1st ionisation energy across periods, what 2 key patterns can be seen?
- a general increase in 1st ionisation energy ACROSS EACH PERIOD
- a sharp decrease in 1st ionisation energy between the end of one period and the start of the next period
What happens to 1st ionisation energy when you go down a group?
it decreases
1- Down a group, what are the 2 main factors for the decrease in 1st ionisation energy?
2- which outweighs the nuclear charge more?
1- increased atomic radius
1- more inner shells, so increased shielding
2- increased radius
what happens to the number of shells as atomic radius increases?
there’s more of the inner shells
how does the increased atomic radius and shielding make the 1st ionisation energy decrease down a group?
because the nuclear attraction on outer electrons decreases
what goes on the x and y axis of a graph showing the trend in 1st ionisation energy across periods?
x axis = atomic number
y axis = 1st ionisation energy
what is the most important factor across a period for the general increase in first ionisation energy?
- the increased nuclear charge
List 4 reasons why the 1st ionisation energy across a period increases?
across a period:
- nuclear charge increases (number of protons)
- all the atoms have the same number of shells = they all have similar shielding
- nuclear attraction increases
- atomic radius decreases as outer electrons are pulled closer
across period 2, which two places does the ionisation energy fall?
from Be - B and N-O
in period 2, why is there a drop between Be and B?
- from beryllium to boron, the 2p sub-shells begin to fill.
- the 2p sub-shell in boron has a higher energy than the 2s sub-shell in beryllium
- so, in boron, the 2p electron is easier to remove than a 2s electron in beryllium
- the 1st ionisation energy of boron is less than that of beryllium
in which places does the ionisation energy drop slightly in period 3 and why?
- from Mg to Al
- from P to S
- for the same reasons as the ones in period 2
what does the fall in 1st IE from nitrogen to oxygen mark?
- the start of electron pairing in the p-orbitals of the 2p sub-shell
why is there a fall in 1st IE from nitrogen to oxygen?
- in oxygen, the paired electrons in one of the 2p orbitals repel each other
- this makes it easier to remove an electron from an oxygen atom rather than a nitrogen atom
- so, the 1st IE of oxygen is less than of Nitrogen
explain why successive ionisation energies always increase?
- As each electron is removed, the outer shell is drawn closer to the nucleus.
- Nuclear attraction is greater and more energy is needed to remove the next electron
Explain the sharp drop in 1st ionisation energy between Ne and Na ?
- sharp drop because of the addition of a new shell with a resulting increase in distance and shielding
- this decreases the nuclear attraction on the outer electrons and decreases the IE
The first six ionisation energies of an element in period 3 are 787, 1577, 3232, 4356, 16091, 19805 kj mole^-1. identify the element and explain your reasoning.
- Silicon
- because there’s a large increase between 4th and 5th ionisation energy, so the 5th electron was removed from an inner shell
- so there’s 4 electrons in the outer shell, so it’s in group 4
Explain why Al has a lower 1st IE than Mg ?
the 3p sub-shell in Al has a higher energy level than the 3s sub-shell in Mg
- the 3p electron is easier to remove
Explain why Sulfur has a lower first ionisation energy than Phosphorus ?
- Phosphorus has 3 electrons in the 3p sub-shell with 1 electron in each 3p orbital.
- sulfur has 4 electrons in the 3p sub-shell, with 2 electrons paired in one 3p orbital and 1 in each of the others
- The paired electrons in sulfur repel each other, making it easier to remove 1 of them rather than an unpaired electron.
which elements are metalloids and why are they called that?
- elements near to the metal-non metal divide e.g. B, Si, Ge, As e.t.c.
- they can show in-between properties
at the divide between metal and non-metal, what is the trend down the groups?
- non-metal to metal
are germanium and silicon metals or non-metals?
semi-metals!!
when are all metals in a solid state, and what is the exception?
- at room temp
- except mercury
what property do al metals have?
they can conduct electricity
in a solid metal structure, what has each atom donated and where to?
- its negative outer shell electrons to a shared pool of electrons which are delocalised throughout the whole structure
in a metal structure, what is left behind when the atoms donate their electrons?
the cations (consisting of nucleus and inner electron shells of the metal atoms)
define metallic bonding?
the strong electrostatic attraction between cations and delocalise electrons
in a metallic structure, can the cations move?
no, they’re fixed in position, to maintain the structure of the metal
what word describes what the delocalised electrons can do?
- they’re mobile
- can move throughout the structure
in a metal structure, what are the metal atoms held together in and what by?
- a giant metallic lattice
- by metallic bonds
what are metals’ mp and bp like?
HIGH!!
in what states can metals conduct electricity?
- solid and liquid states
How can metals conduct electricity?
- when voltage applied across a metal
- delocalised electrons move through structure, carrying a charge
which metal has the highest melting point?
tungsten (w)
which group of metals all have lower melting points below 200 degrees?
group 1 metals
what does the meting point of metals depend on?
the strength of the metallic bonds holding together the atoms in the giant metallic lattice
what is the solubility of metals like?
metals don’t dissolve
what structure do non-metallic molecules form? what does this lead to?
a simple molecular lattice held together by weak intermolecular forces
- therefore, they have low mp and bp
what structure does c, si and b form?
what holds it together?
- giant covalent lattice
- strong covalent bonds
what’s the name of the structure for diamond?
tetrahedral
what are the mp and bp of giant covalent lattices like? why?
- high
- because high temp is necessary to give the large amount of energy that’s needed to break the strong covalent bonds
what’s the solubility of giant covalent lattices like? why?
- they’re insoluble in almost all solvent
- the covalent bonds in the lattice are to strong to be broken by interaction with solvents
can giant covalent lattices conduct electricity? what are the 2 exceptions?
no!
- graphene and graphite
why can’t diamond conduct electricity but graphene can?
- in diamond, all 4 outer shell electrons are involved in covalent bonding, so none are available for conducting electricity
- in graphene, a different structure is formed where 1 of the electrons is free to conduct electricity
in graphene, what does carbon form and what happens to the remaining free electrons?
- planar hexagonal layers
- the remaining electrons are released into pool of delocalised electrons between the layers
- so can conduct electricity
which 2 carbon structures have layers ?
whats the bond angles?
- graphene and graphite
- 120
what is graphene?
a single layer of graphite
in graphite, what are the layers of graphene bonded by?
- weak London forces
How is carbon fibre linked to the structures of graphite and graphene?
the structure in carbon fibre is the hexagonal arrangement as in graphene and graphite
what are the 8 molecules with simple molecular structure? what forces do they have between their molecules?
- N2, O2, F2, Ne,p4, s8, CL2, Ar
- weak London forces!
