Trends in the Periodic Table Flashcards
Periodic
certain intervals of atomic number there occur elements of very similar chemical and physical properties.
Period Table
arrangement of known elements organised in a way that highlights the periodic nature of their repeating properties
Arrangement
left to right: increasing atomic number
up and down: similar chemical characteristics
Periods
Horizontal rows
Groups
Vertical rows
Change from left to right
gradual change from metals to non-metals
Top to bottom down any group
chemical and physical properties tend to remain similar though increasing in metallic nature
Halogens
Group 17.
Quite reactive non-metal elements.
Produce ionic compounds
Noble Gases
Group 18.
Non-metal gases
Have low chemical reactivity
Alkali metals
Group 1.
Soft low melting points
react vigorously with water and acids to produce hydrogen gas
ionic
Alkali earth metals
Group 2.
Metals that react strongly with acids producing hydrogen gas.
react with water producing metal hydrogen and hydrogen gas
ionic
Metal
good conductors of electricity good conductors of heat malleable and ductile shiny solids at room temp (except mercury)
Non-metal
poor electrical and heat conductivity
gases at room temp
those that are solid are brittle
quantum mechanical model
explains repeating chemical behaviour by relating element’s chemical and physical properties to the number of electrons it has
electron configuration
shows shells (of an atom or ion) containing electrons and the number of electrons in each of those shells 2,8,8,2
Valance electrons
electrons in outer most shell
all elements in any given group have same #
Lewis structure
for an atom or ion shows one dot for each valence electron
Core electrons
electrons in shells below valance electrons
Electron configuration and ion formation
gain or lose electrons to meet octet rule
Atomic radius
A.R. of the elements increases down any group of the periodic table (due to higher shell number)
A.R. of the elements decreases across any period (due to increasing nuclear charge - increasing attraction of electrons)
Ionisation energy
a measure of hoe strongly an atom holds onto its electrons
affects elements tendency to form pos or neg ions
decreases with increasing atomic radius
First ionisation energy
minimum amount of energy needed to remove the single most loosely bound electron
factors for ionisation
atom’s nuclear charge:
greater nuclear charge - electrons are more strongly attached to nucleus
distance between nucleus and electron:
as A.R. increases, attraction between nucleus and electrons decrease.
shielding inner electrons
Electronegativity
the ability of an atom in a molecule to attract electrons to itself
NON-METALS: high (gains electrons to form ions)
METALS: low (loses electrons to form ions)
left and right: increases
vertically: decreases
Valance Electrons and Bonding Capacity
strong relationship between the element’s group number and its ionic bonding capacity
strong relationship between an element’s covalent bonding capacity and its group number
Metallic properties
low ionisation energy is essential if atoms want to form positive ions required for the metallic structure (unique physical properties)
only found on left or P.T. bc I.E are lower
I.E. decrease down group element’s metallic properties increase
Metallic - reducing agents
lose electrons and become oxidised when reacting with substances like acids, oxygen and water
associated with low ionisation energy
Non-metal properties
high hardness, brittleness, high melting and boiling points and semi conducting and non electrical conductivity.
Upper right of P.T
highest ionisation, highest electronagtivity, most valance electrons
diatomic molecules
composed of only two atoms, of either the same or different chemical elements
monatomic gases
one in which atoms are not bound to each other
Spectral Analysis
energy levels different for each element
line spectrum/emission spectrum - atomic fingerprint
energy absorbed/released is unique
