Topic 3/13: Periodicity Flashcards
Define atomic radius.
The distance from the nucleus to the outermost electron
Define ionic radius
The distance from the nucleus to the outermost electrons in an ion
Define electron affinity
The energy released when one mole of an electron is added to one mole of gaseous atom
Define electronegativity
The power of an atom to attract a shared pair of electrons in a covalent bond
Describe the arrangement of the periodic table.
- list of chemical elements
- arranged in order of increasing atomic number
- sorted into groups and periods
Describe the trend of atomic radius.
down a group
- number of electron shells increases ➜ atomic radius increases
across a period
- electrons added to same main energy level ➜ nuclear charge increases ➜ atomic radius decreases
Describe the trend of ionization energy.
down a group
- valence electron removed from furthest energy level ➜ less tightly held ➜ IE decreases
across a period
* stronger nuclear charge ➜ increase in attraction between outer electrons and nucleus ➜ more tightly held ➜ IE increases
Describe trend in electron affinity.
- metals have low EA, non-metals have high EA
- the greater the distance between nucleus and outer energy level, the weaker the electrostatic attraction released
Describe the trend in electronegativity.
- metals have low EN as they lose electrons easily
- non-metals have high EN as they gain electrons to complete the outer shell
- EN increases across a period, decreases down a group
Describe the trends in melting points.
across a period
- stronger forces of attraction between the particles ➜ more energy needed to overcome forces ➜ increasing melting point until group 14
- depends on the bonding, structure, strength
Describe the trends in metallic character.
- Metallic character: How easily an atom can lose electrons
- Increases down a group, decreases across a period
Describe the trends down Group I (the alkali metals)
- Atomic/Ionic radius ⬆️ as there are more electron shells
- First IE ⬇️ as valence electron is further from the nucleus, easier to remove
- EN ⬇️ increased distance and shielding
- MP ⬇️ as atoms are larger, metallic bonds are weaker
- Reactivity ⬆️ as valence electrons are easier to lose due to shielding
Describe the trends down group 7 (halogen)
- Atomic/Ionic radius ⬆️ as there are more electron shells
- First IE ⬇️ as valence electron is further from the nucleus, easier to remove
- EN ⬇️ increased distance and shielding
- MP ⬆️ as Van dear Waal forces becomes greater with more electrons
- Reactivity ⬇️ as attraction between nucleus and outer electrons are weaker
Define transition metals.
Elements whose atoms have incomplete d-orbitals or can form positive ions with an incomplete d sub-level.
Is zinc considered a transition element? Explain.
Zinc is not a transition element as it does not have an incomplete d orbital.
Describe physical properties of transition metals.
- High electrical and thermal conductivity
- High MP
- High tensile strength
- Malleable
- Ductile
Describe chemical properties of transition metals.
- Variable oxidation states
- Formation of complex ions
- Colored compounds
- Catalytic Behavior
Define diamagnetic, paramagnetic, and ferromagnetic.
Diamagnetic
- No unpaired electrons, weakly repelled
Paramagnetic
- One or more unpaired d-orbital electrons
Ferromagnetic
- Only occurs in iron, cobalt, nickel
- Large numbers of unpaired electrons
Define ligands.
Molecules or ions with a lone pair of electrons that form coordinate covalent bonds with a central metal ion.
Define complex ions.
Formed when ligands dative covalently bond to a central metal ion by donating a pair of electrons
Outline factors that affect the color of a transition metal.
- Nature of transition element
- Identity of metal ion
- Identity of ligand
- Oxidation state
Describe the trends across period 3.
- sodium, magnesium, aluminum are metals (basic)
- aluminum oxide is amphoteric
- phosphorus and sulfur are non-metals (acidic)
- chlorine and argon are non-metals (no oxide)
