3.2 Periodic trends Flashcards
Valence electrons
The outermost electrons of an atom.
Atomic radius
The distance from the nucleus to the outermost electron.
Ionic radius
The distance from the nucleus to the outermost electrons in an ion.
First ionization energy
The minimum energy required to remove one mole of electrons from one mole of gaseous atoms.
Electron affinity
The energy released when one mole of an electron is added to one mole gaseous atoms.
Electronegativity
A measure of the attraction an atom has for a shared pair of electrons in a covalent bond.
Periodicity
Repeating trends or patterns of physical and chemical properties in elements.
Atomic radius trend
- Increases down a group as the number of electron shells increases.
- Decreases across a period (electrons are added to the same energy level and nuclear charge increases resulting in the attraction between the nucleus and outer electrons increasing and atomic radius decreasing).
Ionic radius trend
- Ionic radius of positive ions (from Groups 1 to 14) decreases across a period (the number of protons in the nucleus increases but the number of electrons remain the same).
- Ionic radius of negative ions (from Groups 14 to 17) decrease across a period as the number of protons in the nucleus increases but the number of electrons remains the same
Ionization energy trend
- Increases across a period (The increase in nuclear charge across a period causes an increase in the attraction between the outer electrons and the nucleus makes the electrons more difficult to remove).
- Decreases down a group (The electron being removed is from the energy level furthest from the nucleus so it gets easier to remove valence electrons as atomic radius increases down a group)
There are regular discontinuities in the trend of increasing ionization energies along a period, due to the existence of sub-levels within the main energy levels.
Electron affinity trend
Generally, metals have a low EA and non-metals have a higher EA.
Electronegativity trend
Generally, electronegativity tends to increase across a period and decrease down a group.
Metals have low electronegativities because they lose electrons easily.
Non-metals have high electronegativities as they gain electrons to complete their outer shell.
Noble gases electronegativity
The noble gases are not assigned electronegativities as they do not readily form bonds with other elements.
Melting points trend
Melting points generally increase across a period until group 14, then they decrease (depends on the type of bonding (covalent, ionic or metallic), structure (ionic lattice, molecular covalent, giant covalent, or metallic structures), and strength of metallic bond ).
Metallic character
How easily an atom can lose electrons.
Displayed by metals, which are all on the left-hand side of the periodic table including alkali metals, alkalie earth metals, tarnsition metals the lanthanide and actinides, and the basic metals.
Metallic character trends
- Increases down a group;
- Decreases across a period.
Group 1 (The Alkali Metals) trends
(atomic/ionic radii, 1st ionization energy, electronegativity, melting points, reactivity)
- Atomic/Ionic radius increases down the group as there are more electron shells.
- First ionization energy decreases down the group as the valence electron is further from the nucleus so its easier to remove.
- Electronegativity decreases because of increased distance and shielding.
- Melting points decrease as atoms become larger and therefore metallic bonds becomes weaker.
- Reactivity increases down the group as the valence electron is easier to lose, due to shielding.
Group 7 (Halogens) trends
(atomic/ionic radii, 1st ionization energy, electronegativity, melting points, reactivity)
- Atomic/Ionic radius increases down the group as there are more electron shells.
- First ionization energy decreases down the group as the valence electron is further from the nucleus so its easier to remove.
- Electronegativity decreases because of increased distance and shielding.
- Melting points increase as Van der Waal forces becomes greater with more electrons.
- Reactivity decreases down group as with each consecutive element the outer shell gets further from the nucleus. So the attraction between the nucleus and electrons gets weaker, so an electron is less easily gained.
- The more reactive halogens displace the less reactive halogens from their compounds:
Period 3 metal properties
(sodium, magnesium and aluminum)
- Shiny solids;
- Excellent thermal and electrical conductors;
- Ductile and malleable;
- Reducing agents;
- Form cations;
- Their oxides and hydroxides behave like bases and neutralize acids.
Aliuminium oxide properties
- Insoluble in water;
- Amphoteric.
Period 3 non-metal properties
(phosphorus, sulfur, chlorine and argon)
- Can be solids, liquids or gas;
- Tend to be oxidizing agents;
- Form anions;
- Their oxides tend to be acidic and are neutralized by bases.
Period 3 oxides + water
All period three oxides will react with water to form either an acidic or alkali (basic) solution.
Period 3 general properties
Period 3 oxides trends
Oxides change from basic through amphoteric to acidic across a period.
Cations, atoms and anions size comparison
Cation < atom < anion
(Cations are smaller than their parent atoms, as the formation of positive ions involves the loss of the outer shell.
Anions are larger than their parent atoms, as their formation involves the addition of electrons into the outer shell. The increased electron repulsion between the electrons in the outer energy level increases the radius of the outer shell.)
Halides
Compounds formed during a reaction between Group 1 metals and halogens.
Vigorous reactions
The most vigorous reaction occurs between the elements which are furthest apart in the Periodic Table.
