CHEM Periodicity Flashcards
Define first ionization energy.
Minimum energy required to remove 1 mole of electrons from 1 mole of gaseous atoms (to form one mole of univalent cations) in the gaseous state.
Define atomic radius.
Half the distance between nuclei of two bonded atoms.
Define electronegativity.
Tendency for an atom to attract electron density to it when covalently bonded.
Define electron affinity.
Energy change that occurs when 1 mole of atoms gain an electron to form 1 mole of 1- ions (univalent anions) in the gas phase.
State, and explain, the trends in FIRST IONIZATION ENERGY across a period and down a group.
Across a period: ⬆️ increasing ATTRACTIONS due to increasing nuclear change and # of VALENCE electrons.
Down a group: ⬇️ increasing DISTANCE between nucleus and VALENCE electrons.
State, and explain, the trends in ATOMIC RADIUS across a period and down a group.
Across a period: ⬇️ increasing ATTRACTIONS due to increasing nuclear change and # of valence electrons in the SAME ENERGY LEVEL.
Down a group: ⬆️ increasing DISTANCE between nucleus and VALENCE electrons.
State, and explain, the trends in ELECTRONEGATIVITY across a period and down a group.
Across a period: ⬆️ increasing ATTRACTIONS due to increasing nuclear change attracting the SHARED electrons.
Down a group: ⬇️ increasing DISTANCE between nucleus and SHARED electrons.
State, and explain, the trends in ELECTRON AFFINITY across a period and down a group.
Across a period: ⬆️ increasing ATTRACTIONS due to increasing nuclear change attracting the ADDED electrons.
Down a group: ⬇️ increasing DISTANCE between nucleus and ADDED electrons.
^ All periodic trends can be explained with which two causes?
1) Changes to nuclear charge and numbers of electrons (across a period).
2) Changes to distance between nuclear charge and electrons (down a group).
State, and explain, the relationship between a CATION’s radius and that of its parent atom.
Positive ions lose valence electrons / the remaining electrons are located in lower energy levels and are CLOSER to the nucleus. The ionic radius of CATIONS is ⬇️ than the corresponding atomic radius (increased proton-electron attractions).
State, and explain, the relationship between a ANION’s radius and that of its parent atom.
Negative ions gain valence electrons / the added electrons INCREASE electron-electron repulsion. Valence electrons are located in the SAME energy level as the atom, while nuclar charge is unchanged. The ionic radius of ANIONS is ⬆️ than the corresponding atomic radius.
^ Melting and boiling points can be explained using?
1) The structures formed
2) The type of bonding that exists in the structure
Explain how melting occurs in metals.
Melting overcomes metallic bonding.
Metallic bond strength depends on:
- charge on the metal ion
- number of delocalized (valence) electrons
- charge density (depends on ionic radius)
💡 melting points and metallic bonding strength ⬆️ as atomic radius increases
Explain how melting occurs in silicon.
Silicon has the highest melting point in Period 3 - it exists as a network covalent structure (crystal lattice). Melting must break COVALENT bonds.
Explain how melting occurs in non-metals.
Melting overcomes IMFs (i.e. LDF) only - lowest melting points in the period. Melting point depends on molecular size (larger molecules have ⬆️ melting point - more total # electrons present).
Explain how metals can conduct electricity.
Metals consist of positive metal ions surrounded by a sea of DELOCALIZED electrons, which are free to move through the lattice.
Explain the general trend in the melting points of metal oxides.
Melting overcomes IONIC bonding (strength depends on metal ion charge, ionic radius, #oxide ions). Melting point ⬆️ with increasing atomic number (exc. Al2O3).
Explain how metal oxides can conduct electricity.
In molten form or in aqueous solution, where IONS are able to move through the liquid freely creating electrical flux.
