Topic 2 & 3 Flashcards
Atomic Radius Trend
Increases down a group
Number of electron shells increases thus increasing in size
Shielding: core electrons prevents valence electrons from the nucleus, thus experiencing less attractive force
decreases across a period
As the number of protons and electrons increases, the attraction between the nucleus and the
outer electrons increases, pulling electrons closer to the nucleus resulting in a smaller radius
Atomic Radius ____ down a group
increase
Ionic Radius
- Cations are smaller than their neutral atoms
- Anions are larger than their neutral atoms
Ionization energy (IE)
The energy required
to remove one electron from a gaseous
atom, measured by how tightly the outer
electrons are attracted by the nucleus.
First ionization energy
the IE to take away the
first electron from the valence shell
Trend of IE
Increases across a period
o 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
o The electron being removed is from the energy level furthest from the nucleus, removing valence
electrons further away becomes easier due to electron shielding
IE increase corresponds to…
Large jump
Small jump
Tiny jump
o a new fully-filled inner shell (large jump)
o a new fully-filled subshell (small jump)
o a half filled subshell (tiny jump)
Electron affinity (EA)
the change in energy when one mole of an electron is added to one mole gaseous
atoms.
Generally negative because the electron and the nucleus form an attraction. The
forming of an attraction results in the release of potential energy as heat, thus the enthalpy change is
negative.`
Trend of EA
_____ across a period
____ down a group
Increases (negatively) across a period
o This is as a result of increasing attraction from the nucleus due to the increasing number of
protons and the reduced distance from the nucleus to the valence electrons
Decreases down a group
o The greater the distance between the nucleus and the outer energy level, the weaker the
electrostatic attraction and the less energy is released when an electron is added to the atom
EA: Explain why elements in Group 17 have the most negative electron affinity.
Atoms are one electron away from fulfilling a full energy level
EA: Deduce the nature of metal ions and non-metal ions with reference to ionization energy and electron
affinity
Metals form cations therefore having a lower IE (easier to lose e-) while non-metals form anions,
requiring an addition of electron (electron affinity is more negative)
EA: Why is the second electron affinity for oxygen endothermic (energy is absorbed)?
The 2nd electron affinity corresponds to the addition of an electron to the O- ion. The O- ion has a
low proton to electron ratio and so the new electron is actually repelled by the electrons that are
already present, thus the 2nd electron affinity is positive as energy has to be put in to force the
electron onto the ion.
EA: Why is fluorine an exception of the vertical trend?
The electrons in the outermost shell of a fluorine atom are closer together. The electron gained also
feels a great amount of repulsion from the electrons originally in the outermost shell. Energy is
required to keep the gained electron in the shell, causing fluorine to have a smaller electron affinity
than chlorine.
Electronegativity
a measure of the attraction an atom has for a shared election
- attraction to pull e- toward itself
trend of electronegativity
increase across a period, decrease down a group
Explain the electronegativity trend across a period
Electronegativity increases across a period, this is as a result of increasing attraction from the nucleus due
to the increasing number of protons and the reduced distance from the nucleus to the valence electrons
(shielding electrons and proton/electron ratio are the same for elements in a period).
Explain the electronegativity trend down a group
Electronegativity decrease down a group due to the increasing distance and shielding electrons.
Metallic character
how easily an atom can lose electrons -
same as IE but differently interpreted
trend of metallic character
decrease across period
increase down a group
Relate and explain the metallic character trend across a period to IE and atomic radii
IE increases across a period due to having a higher
effective nuclear charge (proton/electron ratio) making it more difficult for atoms to lose electrons (less metallic)
Relate and explain the metallic character trend down a group to IE and atomic radii
how easily an atom can lose electrons to form cations, low IE would correspond to high
metallic character as electrons experience less attraction from the nucleus due to shielding, and more electron shells
Alkali metals
Configuration:
Reactivity ____ down the group:
Melting points ____ down the group
React readily with ___ to give an e-
Metal + water -
React vigorously w halogens to produce:
s1
increase
decrease
water
metal hydroxide + hydrogen gas
metal halide (ionic salt)
Halogens
configuration:
Reactivity ___ down the group
Melting points___ down the group
Process of halogen reacting with halides (ions) to remove e-
Halogens
p5
decrease; outer shell gets further, weaker attration to acquire e-
Increase; intermolecular forces between halogen molecules increase down the group. strong forces - more forces need to overcome.
Oxidation
Halogen - precipitation reaction occurs when ___ and ___ in aqueons solution combine to form an __
cation; anion; insoluble ionic solid
Oxides
Acidity - ___ across a period generally due to ___
Acidity - tend to attract
Amphoteric refers to a molecule that ___
increases; increase in electronegativity
e-
acts as both an acid or a base
Why does acidity increase with EN?
Acidity is the tendency to accept electrons. The higher the EN, the higher attraction there is, meaning that it would be easier to attract electrons towards itself (accept electrons.
Electronegativity of aluminum is 1.6. Explain how it is amphoteric.
Aluminium has an electronegativity of 1.6. Exactly at the boundary where oxides of elements start becoming acidic, allowing them to behave both ways.
∆EN between aluminium and oxygen is also minimum.
Describe and explain the general trend in melting point in the metals in periods 2 and 3
The melting point of the 3 first metals in periods 2 and 3 increases as the strength of the metallic bond
increases.
b) Explain the reason for the peak in melting point in C (carbon) and Si (silicon)
Giant covalent bonding
Atomic Radius ____ down a group
increase
Atomic Radius ____ down a group
increase
Relative atomic mass def
