Chemistry Video 6 Flashcards
Li, Na, K
All shiny, conduct heat and electricity. Reacts with oxygen atoms in 2:1 ratio
Ca, Sr, Ba
All soft and shiny, reacts with oxygen atoms in 1:1 ratio
Dmitri Mendeleev
Organized the final periodic table. His table made undiscovered elements
Periodic table element block
Symbol of element in centre, under symbol is atomic mass, atomic number in corner
Group 1
Alkali metals; shiny metals, conduct heat, conduct electricity, react 2:1 with oxygen
Group 2
Alkaline earth metals; metallic properties, react with oxygen in 1:! ration
Groups 3-12
Transition metals; most of them are easily oxidized.
Group 15
Pnictogens
Group 16
Chalcogens
Group 17
Halogens
Group 18
Inert Gases, noble gases, unreactive. Monoatomic.
Atomic radius/covalent radius
One half the distance between the nuclei of 2 identical atoms that are bonded to each other.
Increases down the group. Decreases going left to right on a period.
Bond length
Distance between the nuclei of 2 atoms forming a bond. Sum of the atomic radii for the 2 atoms that are participating in the bond.
Effective nuclear charge, (Zeff) and trend
Pull exerted by the nucleus on an electron. Core electrons shield valence electrons from the pull of the nucleus.
Zeff = number of protons - core electrons
Increases from left to right across a period, decreases going down a group.
Reasoning: Across a period, the numbers of protons are increasing with no increase in a shielding effect, which results in electrons being pulled closer to the nucleus due to a stronger attraction. Going down a group, more shielding causes the effective nuclear charge to decrease. As electrons get further away from the nucleus, the attractive force between protons and electrons naturally lessens.
Ionic radius
Cation = smaller ionic radius compared to non-ion version
Anion = larger ionic radius compared to non-ion version
When a neutral atom loses an electron and becomes a cation, its radius will decrease due to an increase in the effective nuclear charge allowing the protons to pull in the electrons closer to the nucleus.
Conversely, if an atom gains electrons and becomes an anion, its atomic radius will increase. This is due to a decrease in effective nuclear charge and causing a decrease in the pull from the protons in the nucleus.
Metals typically form cations resulting in their ionic radius to be less than their atomic radius and non-metals typically form anions resulting in their ionic radius to be greater than their atomic radius.
Isoelectronic species
Atoms/ions with the same electron configuration, but differing numbers of protons
the most positively charged atom will have the smallest radius because it has the most protons with the same number of electrons, and therefore the greatest attraction to the nucleus.
Ionization energy
Energy required to remove an electron (gas phase, ground state configuration). Higher ionization energy = more difficult to remove the electron. Successive ionization energies are greater than the previous. Always remove the outermost electron
Increases going left to right on a period. Decreases going down a group
He has the highest. Fr has the smallest
Exceptions: Orbitals are most stable when half full or completely full.
Reasoning: As you go from left to right across the periodic table, the number of electrons and protons increases. As the valence shell continues to fill, the electrons become harder to remove (require more energy) due to an increase in effective nuclear charge. Fluorine, for example, has a high ionization energy because its electrons are strongly attracted to the more positively charged atomic nucleus.
Note: Going down a group, the shielding effect is increasing. Outer electrons are not experiencing the full pull of the positive charge of protons due to increasing inner electrons shielding them. This results in lower ionization energy as you go down a group because electrons become easier to expel with increasing shells. Iodine, for example, has a lower ionization energy than fluorine because it contains more electron shells, which makes it easier to extract an electron from the atom’s valence shell.
Electron affinity
Energy associated with adding an electron (gas phase, ground state configuration); The amount of energy released when an electron is added to
an atom. Can be favourable (negative) or unfavourable (positive).
The harder it is to remove an electron or the higher the Zeff, the easier it is to add an electron, and thus a greater electron affinity. Going down a group, the attraction of an electron to the nucleus decreases due to shielding; Electron affinity, therefore, decreases.
Trend: Electron affinity increases going from left to right across a period and
decreases down a group.
F and Cl have large electron affinity.
Exceptions: noble gases have full shell. Half-stable and full-stable orbitals.
i.e. Silicon has a greater affinity for an electron than phosphorus because accepting an electron will give it half-filled subshell configuration.
There are successive electron affinities.
Electronegativity
How well an atom can attract electron density towards itself. The more strongly it can attract electrons, the greater its electronegativity.
Increases going left to right on a period. Decreases down a group. F has greatest and Fr has the lowest. Exclude noble gases from trend.
Inner transition metals
The two rows (which belong in periods 6 and 7) beneath the periodic table are known as the inner transition metals. Period 6 elements are the lanthanides and Period 7 elements are the actinides.
Diatomic atoms
hydrogen, nitrogen, fluorine, oxygen, iodine, chlorine, and bromine.
Have No Fear Of Ice Cold Beer
Metallic character
increases going from right to left across a period and increases going down a group.
Metals vs non metals
Metals:
- Malleable, lustrous
- Good conductors of electricity/heat
- Form basic oxides
- Lose electrons to form cations
- Usually solid at room temperature, with the exception of Hg (liquid)
- Generally, high melting and boiling points
Non-metals:
- Brittle, dull
- Poor conductors of electricity/heat
- Form acidic oxides
- Gain electrons to form anions
- Gas or solid at room temperature, with the exception of Br (liquid)
- Generally, low melting and boiling points
Atomic radius trend
Trend: The atomic radius decreases
from left to right across a period and
increases going down a group.
Reasoning: Across a period, the number
of protons in an atom increases. Increasing protons results in greater nuclear attraction between the protons and electrons, which results in shells being pulled closer to the nucleus (which equals a smaller radius). Going down a group, the number of electrons increases. Each additional energy level gets farther and farther away from the nucleus, which causes the atomic radius to increase.
Multiple ionization energies
The first ionization energy is the energy required to remove the outermost electron. Following removal of the first electron, elements can have second, third, fourth, etc. ionization energies. These values are always larger than the first ionization energy because subsequent electrons are more difficult to remove
Exceptions to multiple ionization energies concept
- The first is that alkaline earth metals have greater ionization energy than group 13 elements. This is because alkaline earth metals have completely filled orbitals and require more energy to remove an electron than group 13 elements
- The second exception is that group 15 elements have greater ionization energy than group 16 elements. Group 15 elements have half-filled orbitals, which is a more stable configuration than that of group 16 elements and thus require more energy to remove an electron
