CHAPTER 7 - PERIODICITY Flashcards
How did Mendeleev arrange the periodic table
Order of atomic mass
Lined in groups with similar properties
Left gaps, assuming atomic mass was wrong or more elements were yet to be discovered
Predicted properties of the missing elements from group trends
What is a group of the periodic table (not an example)
Vertical columns of elements
Same number of outer shell electrons and similar properties
What are Periods in the periodic table
Horizontal rows
The number of the period gives the highest energy electron shell in the elements atom
What are blocks in the periodic table
Four distinct areas corresponding to their highest sub-shell
Mendeleev left gaps in his period table that were later filled when, eg. Scandium gallium and germanium were discovered. In addition, there is an entire group in the modern periodic table that Mendeleev omitted. State which group was missing and suggest why Mendeleev was unaware of it
Group 18 (0)
Contains unreactive noble gases, no elements from group 18 had been discovered at the time
What is ionisation energy
Measures how easily an atom loses electrons to form positive ions
What is first ionisation energy
The energy required to remove one electron from each atom in one mole of an element to form one mole of gaseous 1+ ions
eg. Na(g) –> Na+ + e-
What are the factors affecting 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
The force of attraction falls off sharply, so atomic radius has large effect (most important factor)
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 repulsion is called the shielding effect, which reduces the attraction between nucleus and outer shell electrons
Why do second ionisation energies (or successive) require more energy to separate electrons
After an electron is lost, the remaining electrons are pulled closer to the nucleus
Atomic radius decreases
Nuclear ATTRACTION (not charge) increases per electron
(pg 97)
What does a large increase in ionisation energy suggest
The electron must be removed from a different shell, closer to the nucleus with no/less shielding
What are the two key patterns of first ionisation energies
General increase across each period
Sharp decrease between end of one period and the start of the next
(pg 98)
Why does ionisation energy decrease down a group
Atomic radius increases
Inner shell shielding increases
Nuclear attraction to outer electrons decreases
Why does ionisation energy increase across a period
Nuclear charge increases
Similar shielding: same shell
Nuclear attraction increases
Atomic radius decreases
What is the Trends in the first ionisation energy cross period 2
3 rises, 2 falls
Rise from Lithium - Beryllium
Fall to Boron
Rise from Carbon to Nitrogen
Fall to oxygen
Rise from Fluorine to Neon
(Pg 99)
Why is there a decrease in first ionisation energy from beryllium to Boron
Filling of the 2p subshell in Boron, easier to remove than a 2s Subshell in Beryllium
(pg 100)
Why is there a decrease in first ionisation energy from Nitrogen to Oxygen
Pairing of electrons in 2p sub-shell in Oxygen
Electron pair-repulsion repels other other electrons, making it easier to remove than 2p in Nitrogen
(pg 100)
State the equations to represent the first two ionisation energies in sulfur
S (g) —> S+ (g) + e-
S+ (g) —> S2+ (g) + e-
What is second ionisation energy
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
Explain why successive ionisation energies 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 general trend in first ionisation energy from sodium to argon
Nuclear charge increases from sodium to argon and outer shell.
Electrons are in the same shell with similar shielding
atomic radius decreases
results in an increase in nuclear attraction on the outer electrons and an increase in first ionisation energy.
Explain the sharp drop in first ionisation energy between neon and sodium
Sharp drop reflects the addition of a new shell with a resulting increase in distance and shielding.
This decreases the nuclear attraction on the outer electrons,
decreasing the first ionisation energy.
Explain the trend in first ionisation energy is shown by helium, neon and Argon
Ionisation energy decreases from helium to neon to argon due to an increase in the number of Shells, so increasing atomic radius and shielding.
This causes a decrease in nuclear attraction on outer electrons
decreasing the first ionisation energy.
Explain why aluminium has a lower first ionisation energy than magnesium
The 3p sub shell in aluminium has a higher energy level than the 3s sub shell in magnesium.
The 3p electron is easier to remove.
Explain why sulphur has a lower first ionisation energy than phosphorus
Phosphorus has three electrons in three piece of shell one electron for each orbital.
So far has four electrons in the 3p sub shell, two electrons paired in one orbital and one electron pears in each of the other two.
The pair electrons in sulphur repel one another making it easier to remove one of these electrons than unpaired electron.
What is Metallic bonding
Bonding between 2 metals
Each atom has donated all of its outer shell of negative electrons to a shared sea of delocalised electrons
Positive Cations left behind consists of the nucleus and the inner electron shells the metal atoms
What forms as a result of metallic bonding and what allows it to form that
Giant Metallic lattice
Cations are fixed in position, maintaining structure and shape of the metal
Delocalised electrons are mobile and able to move throughout the structure
What are some properties of metals
Strong bonds - attraction between positive ions and delocalised electrons
High electrical conductivity - Electrons can carry the charge
High MP and BP
Insoluble - leads to a reaction rather than dissolving eg. Sodium and water
What do non-metallic elements exist as?
Simple covalently bonded-molecules, held together by weak intermolecular forces
What do non-metals like Boron and carbon form
Giant covalent structure
Billions of atoms held together by a network of strong covalent bonds instead of small molecules
Describe the structure of Carbon in diamond form
Giant covalent lattice
4 electrons in outer shell form 4 covalent bonds
all bonds 109.5 degrees due to electron pair repulsion
Tetrahedral structure around each carbon
What are properties of giant covalent lattice structures eg. diamond
High MP + BP due to bond strength
Insoluble - bonds are too strong to be broken by interaction with solvents
Cannot conduct electricity due to all four outer-shell electrons involved in covalent bonding, none available to conduct electricity
(exceptions Graphene and Graphite)
Why Can Graphene and Graphite
Form hexagonal structures
Using 3 out of 4 electrons
1 available to conduct electricity
Trigonal Planar - 120 degrees
What is graphene
A single layer of graphite - composed of hexagonally structured carbon atoms
(pg 104)
How is Graphite arranged
Stacks of parallel hexagonally arranged carbon atoms, held together by weak London forces
Why is there are sharp decrease between group 14 (4) and 15 (5)
change from giant covalent to simple molecular structures
(pg 105)
Explain what is meant by metallic bonding and why this type of bonding enables metals to conduct electricity
Strong electrostatic forces of attraction between cations and delocalised electrons
The delocalised electrons can move across a potential difference
Explain how the bonding in a simple molecular lattice differs from that of a simple covalent lattice
Simple molecular lattice has London forces between molecules
Giant covalent lattice has covalent bonds between atoms
Across period 4, the trend in properties is not exactly the same as across Periods 2 and 3. Suggest explanations for the following:
1) Germanium is a good conductor of electricity
2) Arsenic has a much higher melting point than nitrogen and phosphorus
1) the group 4 element, Germanium has a giant metallic lattice rather than a giant covalent lattice in carbon and silicon
2) The group 5 element, Arsenic has a giant covalent lattice whereas in period 2 and 3, the group 5 element has a simple molecular lattice
