Periodic Table and Trends Flashcards
Effective Nuclear Charge
Across - increases = adding protons, you are still in the same energy level so the inner electrons stay the same.
Down - stay the same = the inner shell electrons “shield” the outer electrons from the nucleus.
Electronegativity
Measure of the tendency of an atom in a molecule to attract a shared pair of electrons towards itself.
It is also a measure of the attraction between the nucleus and the outer shell bonding electrons.
An element with a high electronegativity has strong electron pulling power.
Across - increases = increasing effective nuclear charge
Down - decreases = outer shell electrons are farther away from the nucleus which decreases the attraction.
Atomic Radius
The “atomic radius” is not “how big the atom is” but is measured as half the distance between neighbouring nuclei.
Across - decreases = effective nuclear charge increases which increases the attraction between the nucleus and the outer shell electrons, making it smaller.
Down - increase = adding energy levels
Ionic Radius
Positive ions (cations) are smaller than their parent atoms because they have lost one or more electrons. Negative ions (anions) are larger than their parent atoms because they have gained one or more electrons. Increasing electron repulsions also cause the radius to increase.
Across groups 1-17 - decrease = increasing effective nuclear charge. The increased attraction between the nucleus and the electrons pulls the outer shell electrons closer to the nucleus.
However there is a jump in size when you move from positive to negative ions within a period.
Down - increases = increasing energy levels.
First Ionisation energy
The first ionisation energy of an element is the energy required to remove one electron from an atom in a gaseous phase.
A measure of the attraction between the nucleus and the outer shell electron.
Across - increases = increasing nuclear charge which makes it more difficult to remove the outer shell electron.
Down - decrease = distance from the nucleus and the shielding by the inner shell electrons.
Properties of metals
- Good conductors of heat and electricity
- Malleable – can be beaten into thin sheets
- Ductile – can be drawn into wires
- Shiny or have luster
- Tend to lose electrons in chemical reactions
Elements in Groups 1 to 13 lose electrons to adopt the arrangement of the nearest noble gas with a lower atomic number.
They are usually metals.
Properties of non-metals
- Poor conductors of heat and electricity
- Tend to gain electrons in chemical reactions
- Can be solid, liquid or gas
Elements in Groups 15 to 17 gain electrons to adopt the electron arrangement of the nearest noble gas to their right.
They are usually non-metals.
Properties of metalloids
Have both metallic and non-metallic properties
Reactivity of metals
Metallic character of an element is defined as the ease in which they lose electrons. Which is therefore related to ionisation energy.
So the more easily a metal loses electrons the more reactive they are as their valence electrons have a weaker attraction to the nucleus.
Across - decrease
Down - increase
Reactivity of non-metals
Atoms of non-metals gain electrons in chemical reactions to form a stable octet.
The more easily a non-metal gains electrons the more reactive it is (has a higher electron affinity).
Across - increase = increase effective nuclear charge
Down - decrease = harder to attract electrons with more shells
Electron affinity
The energy released when atom gains an electron is called electron affinity.
Oxides
An ionic bond forms between two atoms with a large difference in electronegativity (greater than 1.8).
Since oxygen is highly electronegative, it forms ionic bonds with all metals.
Across - ionic to covalent , basic to acidic
