Chapter 7 - Periodicity Flashcards
what are the positions of the elements in the periodic table linked to?
their physical and chemical properties
how does atomic number change going along the periodic table? (left to right)
the elements are arranged in order of increasing atomic number
what is a group?
a vertical column on the periodic table
what are the similarities between elements in a group?
- same number of outer shell electrons
- similar properties
what is a period?
a horizontal row on the periodic table
what are the similarities between elements in a period?
they all have the same highest energy electron shell number
what is periodicity?
a repeating trend in properties of the elements across a period
what is the trend in electron configuration across a period?
each period starts with an electron in a new highest energy shell:
- across period 2, the 2s sub-shell fills with 2 electrons then the 2p sub-shell fills with 6 electrons
- across period 3, the 3s sub-shell fills with 2 electrons then the 3p sub-shell fills with 6 electrons for each period, the s- and p- sub shells are filled in the same way. this is a periodic pattern
what is the trend in electron configuration down a group?
elements in each group have atoms with the same number of electrons in the outer sub-shell
what is a block in the periodic table?
the elements in the periodic table can be divided into blocks corresponding to their highest energy sub-shell
what are the blocks in the periodic table?
- s-block
- d-block
- p-block
- f-block
what is first ionisation energy?
the energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions e.g: Na (g) -> Na+(g) + e-
which electron is lost in the first ionisation energy?
the highest energy electron, which experiences the least attraction from the nucleus
what 3 factors affect ionisation energy?
- atomic radius
- nuclear charge
- electron shielding
how does atomic radius affect ionisation energy?
the greater the distance between the nucleus and the outer electrons, the less the nuclear attraction. less energy is needed to remove an electron if nuclear attraction is less.
how does nuclear charge affect ionisation energy?
the more protons there are in the nucleus of an atom, the greater the attraction between the nucleus and the outer electrons
how does electron shielding affect ionisation energy?
electrons are negatively charged and so inner-shell electrons repel outer shell electrons. this is called the shielding effect. the shielding effect reduces the attraction between the nucleus and the outer electrons
how do you determine how many ionisation energies an element has?
an element has as many ionisation energies as it has electrons
why does ionisation energy increase as you remove more electrons? (e.g. from 1st ionisation energy to 2nd)
after an electron is lost, the other electrons (in the outer shell) are pulled closer to the nucleus. the nuclear attraction on the remaining electrons increases and more ionisation energy is needed to remove them.
how do sucessive ionisation energies provide evidence for different energy levels in an atom?
- large differences between ionisation energies can be observed
- they suggest that the latter electron is being removed from a different shell with less shielding and stronger nuclear attraction
what can sucessive ionisation energies be used to predict?
- the number of electrons in the outer shell
- the group of the element
- the identity of an element
how can first ionisation energies across a period be used as evidence of the existence of shells and sub-shells?
- across each period, first ionisation energies generally increase
- this is because nuclear charge increases, but elements have the same radius and number of inner shells so shielding is the same. nuclear attraction is higher so more energy is required - between the end of one period and the start of the next, there is a sharp decrease in first ionisation energy
- this is because elements have a larger radius and more inner shells, so shielding is increased (this outweighs the increase in nuclear charge). nuclear attraction is lower so less energy is required
why does first ionisation energy fall in two places in period 2 and 3?
from Be to B:
- B has its highest energy electron in the 2p sub-shell
- the 2p sub-shell has higher energy than the 2s sub-shell
- so the 2p electron is easier to remove
From N to O:
- both have their highest energy electrons in the 2p sub-shell
- in oxygen, the paired electrons in one of the 2p orbitals repel one another
- so it is easier to remove an electron from an oxygen atom
what are semi-metals/metalloids?
elements near to the metal/non-metal divide
what is the state of metals at room temperature?
all metals except mercury are solids
what is the one constant property of all metals?
they can all conduct electricity
what is metallic bonding?
the strong electrostatic attraction between cations and delocalised electrons
- the atoms have donated its negative outer-shell electrons to a shared pool of electrons, which are delocalised throught the whole structure
- the cations are fixed in position, maintaining the structure and shape of the metal
- the delocalised electrons are mobile and are able to move throughout the structure
what are the properties of metals?
most metals have:
- strong metallic bonds
- high electrical conductivity
- high melting and boiling points
how can metals high electrical conductivity be explained?
- metals conduct electricity in solid and liquid states
- when a voltage is applied, the delocalised electrons can move through the structure, carrying charge
how can metals high melting and boiling points be explained?
- the melting point depends on the strength of the metallic bonds holding together the atoms in the giant metallic lattice
- for most metals, high temperatures are necessary to provide the large amount of energy needed to overcome the strong electrostatic attraction between the cations and the electrons
how can metals solubility be explained?
- metals do not dissolve
- any interactions between polar solvents and the charges in a metallic lattice results in a reaction, not dissolving
what is the structure of boron, carbon and silicon?
they form giant covalent lattices, with billions of atoms held together with strong covalent bonds
what structure do carbon and silicon have?
- they have 4 electrons in their outer shells, so form 4 covalent bonds to other atoms
- so they have a tetrahedral structure
what are the melting and boiling points of giant covalent lattices?
they have high melting and boiling points because covalent bonds are strong and require lots of energy to break
what is the solubility of giant covalent lattices?
they are insoluble in almost all solvents because the covalent bonds holding together the atoms in the lattice are too strong to be broken by interaction with solvents
what is the conductivity of giant covalent lattices?
giant covalent lattices are non-conductors of electricity, with graphene and graphite being exceptions
why do graphite and graphene conduct electricity?
graphite and graphene have structures in which one of the electrons is available for conductivity, so they are able to conduct electricity
what is graphene?
a single layer of graphite, composed of hexagonally arranged carbon atoms linked by strong covalent bonds it has the same conductivity as copper and is the thinnest and strongest material ever made
what is the trend of melting points across periods 2 and 3?
- melting point increases from group 1 to group 4 (14)
- sharp decrease in melting point from group 4 (14) to group 5 (15)
- melting points are comparatively low from group 5 (15) to group 8 (18)
why does melting point increase from group 1 to group 4 (14)?
- bonding changes from metallic bonding to giant covalent structure
- in metallic bonding:
charge on cations increases across a period, so electrostatic attraction is larger. giant metallic structures always have strong metallic bonds between cations and delocalised electrons - in giant covalent structures:
there are strong covalent bonds between atoms
why does melting point decrease sharply from group 4 (14) to group 5 (15)?
- structures change from giant covalent to simple molecular
- simple molecular structures only have weak London forces between molecules
- so have much lower boiling points
