Topic 3: Periodicity Flashcards
Effective nuclear charge
given by the atomic number and increases between succesive elements as protons are added to the nucleus/
- ability of an atom to attract increases with group number but remain the same down a group
Electron affinity
the energy change when 1 mole of electrons is added to one mole of gaseous atoms
X(g) + e- —> X-(g)
Ionization Energy
energy require to remove one mole of electrons from one mole of gaseous atoms in their ground state
- measure of attraction between nucleus and outer electrons
Electronegativity
ability of an atom/nucleus to attract electrons in a covalent bond/molecule.
- It is an artificial scale from 0.7 to 4.0, created by combining the ionisation energy and the electron affinity of the elements.
Alkali
bases soluble in water which from a basic solution (hydroxide ions)
Group
(down) elements with the same number of valence electrons
Period
(across) The period number (n) is the outer energy level that is occupied by electrons.
Atomic Radii
half the distance between the nucleus’s of 2 bonded atoms of the same element
-atomic radii decrease along a period as the nuclear charge increases and electrons are added
to the same outer shell. The attraction between the outer electrons and nucleus increases.
Ionic Radii
: Cations are smaller than their parent atoms, as the formation of positive ions involves the loss of
the outer shell.
Group 1
S; Alkali Metals; reactive metals
Group 17
P; Halogens; reactive group of non-metals
Group 2
S; alkaline earth metals
Group 18
P; Noble gases; stable non metals
groups 3-12
D; transition metals
transition metals
The transition metals are in the large section of d-block elements in the middle of the Periodic Table from Sc
to Zn, etc. Zn is not a transition metal because it does not form ions with incomplete d sub-levels.
metalloid
- has physical and chemical features of non metals and metals are are in the P block of the periodic table along the staircase
presence of inner electrons
reduces the attraction of nucleus to the outer electrons
nuclear charrge as period crosses left to right
one proton is added to nucleus; one electron added to calance energy shell= effective nuclear charge increases with nuclear charge as there is no change in number of inner electrons
nuclear charge down a group
- increase in nuclear charge= increase in number of inner electrons
- remains more or less the same down a group
IONIC radii trend
- positive ions are smaller than parent ions (Na+ smaller than Na)
- negative ions larger than parent ions (Cl- larger than Cl)
- ionic radii decrease from Groups 1-14 fo rpositive ions (increased attraction between nucelus and electrons pulls outer shell closer to nucleus)
- ionic radii decrease down Group 14-17 for negative ions (decrease in ionic radii due to icnrease in nuclear charge across period)
- ionic radii decrease down a group as number of electron energy level increases
two general trends of ionization energies
- ionization energies increase across a period
(due to increase in effective nuclear charge that causes an increase in attraction between outer electrons and nucleus; makes electrons harder to remove) - ionizzation energies decrease down a group
(electron removed is from the energy level furtherest form nucleus; effective nuclear charge about the same; increased distance between electrons and nucleus reduces attraction between them)
electron affinity trends
- g17 elements have incomplete outer energy levels; high effective nuclear charge; attract electrons the most
- g1 metals have low effective nuclear charge + attract electrons the least
electronegativity trend
- increases across period due to increase in nuclear charge (increase attraction between nucleus and bond electrons)
- decreases down a group (bonding electrons are furthest from nucleus; reduced attraction)
G1 Trend
- melting point decreases down group (attractions decrease with distance due to delocalized outer electron and positive ion attractive forces)
- reactivity increases down group (francium highest)
G17 Trends
- melting point increases down group (london forces increase with size)
- reactivity decreases down group (Fluorine highest)
melting points across period
- rise with maximum at G14, then fall at minimum to g18
Alkali metals physical properties
- good conductors of electricity and heat
- low densities
- grey shiny surfaces
Alkali metals chemical properties
- very reactive metals
- form ionic compounds with non metals
Noble Gases
- colorless gases
- monoatomic (exist as single atoms)
- very reactive
- stable octet (complete valence energy levels with 8 electrons)
Alkali reactions with water
produce hydrogen and a metal hydroxide
Halogen physical properties
- colored
- show gradual change from gases, to liquid, to solids
Halogen chemical propeties
- reactive non metals
- reactivity decreases down group
- form ionic compounds with metals
- form covalent compounds with other non metals
halogen reaction with group 1 metal
-form ioinic halides (halogen atom gains one electron from group 1 elements to form halid ion X-)
Displacement reaction
the more reactive halogen displaces the ions of the less reactive halogen form its compounds
period 3 oxide acidity trend
basic —> amphoteric—> acidic
bonding of period 3 oxides
giant ionic–> molecular covalent—> giant covalent
(METALLIC TO NON METALLIC CHARACTER)
- oxides become more inoicn down a group as electronegativity decreaes
-conductivity decreases across period
-oxidation state increases across period
oxides of metals
ionic and basic
oxides of non metals
covalent and acidic
amphoteric oxide
show both acidic and basic properties
e.g. Aluminum Oxie
Basic Oxide reaction with water
form alkaline solutions
form salt and a water
Acidic Oxide reaction with water
form acidic solutions
form acids
Ionic or covalent giant strutures
high melting point
Molecular covalent structures
low melting points
