Chapter 7 Flashcards
How did Mendeleev arrange elements?
1) in order of atomic mass
2) lining up elements in groups with similar properties
If the group properties did not fit, what did Mendeleev do with the elements?
Swapped them around and left gaps, assuming that the atomic mass measurements were incorrect and that some elements were yet to be discovered
What does the arrangement, pattern, and shape of the periodic table reveal?
trends among the elements
What are the positions of the elements in the periodic table linked to?
their physical and chemical properties
Why is the periodic table essential?
for predicting the properties of elements and their compounds
How are elements arranged?
in order of increasing atomic number
What is the same in each group of the periodic table?
Each element in a group has atoms with the same number of outer-shell electrons and similar properties
What does the number of the period give?
the number of the highest energy electron shell in an element’s atoms
What is periodicity?
A repeating trend in properties of the elements across each element
What is the most obvious periodicity in properties?
the trend from metals to non-metals
What properties of periodicity are there?
1) electron configuration
2) ionisation energy
3) structure
4) melting points
What is the chemistry of each element determined by?
its electron configuration, particularly the outer, highest energy electron shell
What does each period start with?
an electron in a new highest energy shell
What do the blocks in the periodic table correspond to?
their highest energy sub-shell
What are the 4 blocks in the periodic table?
2) s
2) p
3) f
4) d
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 on electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions
What are the factors involving ionisation energy?
1) atomic radius
2) nuclear charge
3) electron shielding
How are electrons held in their shells?
by attraction from the nucleus
Where will the first electron be lost from?
the highest energy level
How does atomic radius affect ionisation energy?
1) the greater the distance between the nucleus and outer electrons, the less the nuclear attraction
2) force of attraction decreases with increasing distance
How does nuclear charge affect ionisation energy?
The more protons in the nucleus, the greater the attraction between the nucleus and outer electrons
How does electron shielding affect ionisation energy?
inner-shell electrons repel outer-shell electron and reduces the attraction between nucleus and outer electrons
What is the shielding effect?
Electrons are negatively charged and so inner-shell electrons repel outer-shell electrons
How many ionisation energies does does an element have?
As many as there are electrons
What is the first ionisation energy of helium?
He(g) –> He⁺(g) + e⁻
What is the second ionisation energy of helium?
He⁺(g) –> He²⁺(g) + e⁻
Is the second ionisation energy of helium greater or smaller than the first ionisation energy?
greater
Why is the second ionisation energy of helium greater than the first ionisation energy?
1) after the 1st electron is lost, the single electron is pulled closer to the nucleus
2) nuclear attraction on the remaining electron increases
3) more ionisation energy will be needed to remove 2nd electron
What is the second ionisation energy?
The energy required to remove one electron from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions
What do successive ionisation energies provide evidence for?
the different electron energy levels in an atom
What do successive ionisation energies allow predictions to be made about?
1) number of electrons in the outer shell
2) group of the element in the periodic table
3) identity of an element
What do periodic trends in first ionisation energies provide evidence for?
the existence of shells and sub-shells
What are the patterns in the first ionisation energies for the first 20 elements in the periodic table?
1) a general increase in first ionisation energy across each period
2) a sharp decrease in first ionisation energy between the end of one period and the start of the next period
Do first ionisation energies increase or decrease down a group?
decrease
Why do first ionisation energies decrease down a group?
1) atomic radius increases
2) more inner shells so shielding increases
3) nuclear attraction on outer electrons decreases
Do first ionisation energies increase or decrease across a period?
increase
Why do first ionisation energies increase across a period?
1) nuclear charge increases
2) same shell: similar shielding
3) nuclear attraction increases
4) atomic radius decreases
What does the fall in first ionisation energy from beryllium to boron mark?
the start of filling the 2p sub-shell
Why is the first ionisation energy of boron less than beryllium?
1) 2p sub-shell in boron has a higher energy than 2s sub-shell in beryllium
2) 2p electron in boron is easier to remove than one of the 2s electrons in beryllium
What does the fall in first ionisation energy from nitrogen to oxygen mark?
the start of electron pairing in the p-orbitals of the 2p sub-shell
Why is the first ionisation energy of oxygen less than nitrogen?
1) highest energy electrons are in a 2p sub-shell
2) in oxygen, paired electrons in one of the 2p orbitals repel one another, making it easier to remove an electron from oxygen than nitrogen
What are the elements near to the metal/non-metal divide called?
semi-metals/metalloids
At room temperature, what state are all metals except mercury?
solid
What state is mercury at room temperature?
liquid
What is the one constant property of all metals?
ability to conduct electricity
What is metallic bonding?
the strong electrostatic attraction between cations (positive ions) and delocalised electrons
In a solid metal structure, what does each atom do?
donate its negative outer-shell electrons to a shared pool of electrons, which are delocalised throughout the whole structure
What is the structure of a metal?
giant metallic lattice
What are the properties of metals?
1) stong metallic bonds
2) high electrical conductivity
3) high melting and boiling points
How do electrons conduct electricity?
when a voltage is applied, delocalised electrons move through the structure, carrying charge
Can ionic compounds conduct electricity in the solid state?
no
What does the melting point of metals depend on?
the strength of the metallic bonds holding together the atoms in the giant metallic lattice
Why are high temperatures needed to melt metals?
to provide the large amount of energy needed to overcome the strong electrostatic attraction between the cations and electrons
Do metals dissolve?
no
What is the structure of many non-metallic elements?
simple molecular lattice
How are simple molecular lattice structures held together?
weak intermolecular forces
Do simple molecules have high or low melting and boiling points?
low
What is the structure of the non-metals boron, carbon, and silicon?
giant covalent lattice
How are giant covalent lattice structures held together?
strong covalent bonds
How many electrons do atoms of carbon and silicon have in their outer shells?
4
What is the shape of the structure carbon and silicon form when bonding to other carbon or silicon atoms?
tetrahedral
What are the bond angles in a tetrahedral structure?
109.5°
What are the properties of substances with a giant covalent lattice structure?
1) high melting and boiling points
2) insolubility
3) non-conductivity
Why do giant covalent lattices have high melting and boiling points?
1) covalent bonds are strong
2) high temperatures necessary to provide the large quantity of energy needed to break strong covalent bonds
Why are giant covalent lattices insoluble?
covalent bonds holding together atoms in the lattice are too strong to be broken by interaction with solvents
What giant covalent lattices can conduct electricity?
1) graphene
2) graphite
Why can giant covalent lattices not conduct electricity?
all four outer-shell electrons are involved in covalent bonding, none are available for conducting electricity
Why can graphene and graphite conduct electricity?
1) only three of four outer-shell electrons are used in covalent bonding
2) remaining electron is released into pool of delocalised electrons shared by all atoms in the structure
What is graphene?
a single layer of graphite
Does melting point increase/decrease from Group 14 to Group 15?
increase
Why is there a sharp decrease in melting point from Group 15 to Group 18
change from giant to simple molecular structures - weak forces to overcome)
