Topic 3: Periodicity Flashcards
Define periodicity.
Periodicity of elements is a repeated pattern.
Name group 1 of the periodic table.
Alkali metals
Name group 2 of the periodic table.
Alkaline earth metals
Name group 15 of the periodic table.
Pnictogens
Name group 16 of the periodic table.
Chalcogens
Name group 17 of the periodic table.
Halogens
Name group 18 of the periodic table.
Noble gases
How many elements does the periodic table currently consists of?
118 elements
What is the period number?
Horizontal rows of elements in a periodic table are called periods and the periodic number is equal to the principal quantum number, n, of the highest occupied energy level in the elements of the period.
Describe metals.
Metals are:
1) good conductors of heat and electricity
2) malleable (capable of being hammered into thin sheets)
3) ductile (capable of being drawn into wires)
4) lustre (they are shiny)
What are amalgams?
Amalgam is a mercury alloy with another metal. Are formed through metallic bonding with the electrostatic attractive force of conduction electrons binding positively charged metal ions together into a crystal lattice structure.
Describe non-metals.
Non-metals are:
1) poor conductors of heat and electricity
2) typically gain electrons in chemical reactions (they are reduced)
Describe metalloids.
Metalloids have both metallic and non-metallic properties.
Some are semiconductors, such as silicon and germanium, due to their intermediate, highly temperature-dependent electrical conductivity
What is an atomic orbital?
Atomic orbital is a region of space where there is a high probability of finding an electron. Meaning that the position of an electron is not fixed.
How can you determine the radius of an atom?
In order to find the radius of an atom consider two non-metallic atoms chemically bonded together, I.e. a diatomic molecule and the distance between the two nuclei of the atoms, d, and the bonding atomic radius, R(b), is defined as:
R(b) = 1/2d
Sometimes it’s referred to as covalent radius.
For metals, d is d’ where the distance is between two atoms adjacent to each other in the crystal lattice of the metal.
What is the van der Waals’ radius?
Van der Waals’ radius is the non-bonding atomic radius
What is the effective nuclear charge, Z(eff)?
Effective nuclear charge is the net charge experienced by an electron. It’s the nuclear charge, Z, minus the charge, S, that is shielded or screened by the core electrons (all electrons except for valence electrons)
Describe the periodic trends in atomic radius.
Across the period from left to right, the atomic radii decrease. This is because of the increasing effective nuclear charge, which pulls the valence electrons closer to the nucleus, thus reducing the atomic radius.
Down the group from top to bottom, the atomic radii increase. In each new period, the outer-shell electrons enter a new energy level so are located further away from the nucleus. This has a greater effect than the nuclear charge, Z, because of the shielding by the core electrons.
Transition metals do not greatly change across the period, because the electron in the outermost energy level of the principal quantum number, n, remains almost constant across the period.
Describe the periodic trends in ionic radius.
The radii of cations and anions vary from the parent atom.
The radii of cations are smaller than those of the parent atoms, because there are more protons than electrons so the valence electrons are more strongly attracted to the nucleus.
The radii of anions are larger than those of the parent, because the additional electron in the anion results in stronger repulsion between the valence electrons.
Describe the periodic trends in ionisation energy.
Across the period from left to right the ionisation energy values increase, because:
1) As the effective nuclear charge increases from left to right, the valence are pulled closer to the nucleus, so the attraction between the electrons and the nucleus increases, thus making it more difficult to remove an electron from the atom
2) atomic radii decrease across the period, because the distance between the valence electrons and the nucleus decreases, it becomes more difficult to remove the electron.
Going down the group, the ionisation energy values decrease, because:
1) atomic radii increase down the group, making it easier to remove an electron from the atom
2) the shielding effect of the core electrons increases faster than the nuclear charge, weakening the attractive force between the nucleus and outer electrons in the atom.
Define electron affinity.
Electron affinity is the energy required to detach an electron from the singly charged negative ions in the gas phase.
X^(-) (g) —> X (g) + e
Another definition
It is the energy released when 1mol of electrons is attached to 1mol of neutral atoms or molecules in the gas phase.
X(g) + e —> X^(-) (g)
The process is exothermic. The more negative the E(ea) value, the greater is the attraction of the ion for the electron
Describe the periodic trends in electron affinity.
In general, across the period the electron affinity becomes more negative, with some exceptions.
Group 17, the halogens, are the most negative, this is expected, because on gaining an electron these elements attain the stable noble gas configuration.
However, exceptions occur in e.g. arsenic, you can see why when you look at the electron configuration: [Ar]3d(10).4s(2).4(px1, py1, pz1), so if an electron is added it will enter a 4p orbital that already contains one electron, causing repulsion.
Electron affinity values do not show the same clear tends down a group.
Define electronegativity.
Electronegativity, weird X, is the relative attraction that an atom has for the shared pair of electrons in a covalent bond.
Describe the periodic trends in electronegativity.
The trends across the period and down the group mirror those of ionisation energies for the same reasons.
I.e. across the period electronegativity increases, because the atomic radii decreases and effective nuclear charge increases.
Down the group electronegativity decreases, because the atomic radii increases and although the atomic number increases, its effect is shielded by the core electrons.
Describe the periodic trends in metallic and non-metallic character.
Metallic character deceased across the period and increases down the group.
Metals have low ionisation energy values, because they have a tendency to lose electrons during a chemical reaction, I.e. they tend to be oxidised.
Non-metals show highly negative electron affinities, because they have a tendency to gain electrons during a chemical reaction, I.e. tend to be reduced.
Define an oxide.
An oxide is formed from the combination of an element with oxygen.
Describe the trends in the properties of metal and non-metal oxides.
Metal oxides are basic, I.e. they react with water to form hydroxides.
E.g.
CaO (s) + H2O(l) —> Ca(OH)2 (aq)
Oxides of non-metals are acidic, ie they react with water to form acidic solutions.
Eg.
CO2 (g) + H2O (l) <==> H2CO3 (aq) carbonic acid
Silicon dioxide
Silicon dioxide does not dissolve in water, but is classified as an acidic oxide because it can react with sodium hydroxide to form sodium silicate, Na2SiO3 (aq) and water.
Aluminium oxide
Al2O3 is classified as an amphoteric oxide, meaning it can react both as an acid and as a base.
Define amphoteric.
Amphoteric is a chemical species that behaves both as an acid and as a base
Define amphiprotic.
Amphiprotic is a particular type of amphoteric species that are either proton (H^+) donors or proton acceptors. Eg self-ionising solvents like water; amino acids and proteins.
List group 1 elements.
Lithium, sodium, potassium, rubidium, caesium, and francium.
Describe the chemical properties within the Group 1, the alkali metals.
- Have one valence electron, therefore form the ion X^+ in ionic compounds by losing an electron, ie are oxidised
- Going down the group, the atomic radius increases and ionisation energy decreases.
- The reactions with water become more vigorous further down the group. They form metal hydroxide, which gives an alkaline solution (hydrogen gas is released as by-product)
Name two elements in the periodic table that exist as liquids.
Bromine and mercury
List the elements in group 17, the halogens.
Fluorine, chlorine, bromine, iodine and astatine (At)
Describe the chemical properties within the Group 17, the halogens.
1) Have 7 valence electrons, so they have a tendency to gain electrons to attain the noble gas configuration, ie they are reduced.
2) group 17 elements exists as diatomic molecules
3) fluorine and chlorine are gases, bromine is liquid and iodine and astatine are solids at room temperature and pressure.
4) form ionic compounds with metals to form ionic alkali metal halide salts
5) form covalent compounds with non-metals
6) highly reactive, though the reactivity decreases down the group, because the atomic radius increases down the group making it less easy to gain an electron
What happens in the the reaction between Cl2 and Br^(-)?
Displacement reaction where chlorine takes the place of bromide. You can observe a yellow/orange solution due to formation of Br2 (aq)
What happens in the the reaction between Cl2 and I^(-)?
Cl2 takes place of I^(-) and the solution turns dark red/brown due to formation of I2 (aq)
What happens in the the reaction between Br2 and I^(-)?
Br2 takes place of I^(-) and the solution turns dark red/brown due to formation of I2 (aq).
