Chapter 7: Trends in the Periodic Table (Definitions) Flashcards
Define atomic radius(covalent radius) of an atom.
4
- half the distance between the nuclei of two atoms
- of the same element
- that are joined together
- by a single covalent bond
Why does the values of the atomic radius increase down the groups in the Periodic Table?
(a) =3
(b) =4
(a) the additional electrons are going into a new energy level which is further from the nucleus
- since the outer electrons are becoming further away from the nucleus
- the atomic radius increases
(b) the screening effect of inner electrons
- the nuclear charge increases but it is lessened because of the screening effect
- the electrons in the inner energy level(s) help to screen the outer electrons from the positive charge of the nucleus
- thus, some of the attractive force of the positive charges in the nucleus is lessened by the inner energy levels of electrons
- therefore, even though the nuclear charge increases, the increase is counteracted by the shielding effect of the inner energy levels of electrons.
Why does the values of the atomic radius decrease across a period in the Periodic Table?
(a) =4
(b) =4
(a) Increase in effective nuclear energy:
- number of protons in the nucleus increases from left to right across a period
- there is greater attractive force on the outer electrons
- increased nuclear charge draws the energy levels closer to the nucleus
- this attractive force increases from left to right across the period.
(b) No increase in the screening effect
- the extra electron, moving left to right across a period, is being added into the same outer energy level
- therefore, there is no increase in the screening effect
- there is no additional energy level of electrons being added to counteract the increasing attractive force of the nucleus
- the size of the atom shrinks
Define first ionisation energy of an atom
3
- the minimum energy required to completely remove the most loosely bound electron
- from a neutral gaseous atom
- in its ground state.
What is the Equation for first ionisation energy(sodium)
Na(g)->Na+(g)+e-
Why does the values of the first ionisation energy decrease down a group in the Periodic Table
(a) =3
(b) =5
(a) Increasing atomic radius
- the atomic radius increases down any one group
- the outermost electrons are becoming further away from the attractive force of the nucleus
- therefore, it becomes easier to remove an electron from the outer energy level
(b) Screening effect of inner electrons
- the increase in nuclear charge is cancelled out
- by the corresponding screening effect of the
- intervening energy levels of electrons.
- the outermost electrons are shielded from the attractive force of the positively charged nucleus
- and so are easier to remove
Why does the values of the first ionisation energy increase across a period in the Periodic Table
(a) =5
(b) =4
(a) Increasing effective nuclear charge
- from left to right across a period the number of protons in the nucleus is increasing
- this means that the attraction between the nucleus and the outer electrons is steadily increasing
- without any corresponding increase in screening by electrons
- therefore, the electrons are being held more firmly
- and so it requires more energy to remove one of the electrons from the outermost energy level
(b) Decreasing atomic radius
- the atomic radius decreases from left to right across a period
- this means that an electron in the outermost energy level is becoming closer to the nucleus
- therefore, it is more difficult to remove the electron
- due to the increased attraction between it and the nucleus
Define second ionisation energy of an atom.
4
- the energy required to remove an electron
- from an ion
- with one positive charge
- in the gaseous state
What is the equation for second ionisation energy
X+(g)->X2+(g)+e-
Why does the values of electronegativity decrease down the groups in the Periodic Table
(a) =3
(b) =4
(a) Increasing atomic radius
- the atomic radius increases as you move down any group
- this means the outermost electrons are becoming further away from the attractive force of the nucleus
- therefore, there is a smaller attraction between the nucleus and the shared pair of electrons.
(b) Screening effect of inner electrons
- the increase in nuclear charge down a group is cancelled by the screening effect of the intervening energy levels of electrons
- this means that the outermost electrons are shielded from the attractive force of the positively charged nucleus
- since it is these outermost electrons that are involved in bonding
- the attraction of the nucleus for these electrons decreases going down the group.
Why does the values of electronegativity increase across the periods in the Periodic Table
(a) =3
(b) =3
(a) Increasing effective nuclear charge
- the number of protons in the nucleus increases from left to right across a period
- this means that the attraction between the nucleus and the outer electrons is steadily increasing
- therefore the electrons involved in bonding are being more strongly attracted to the nucleus
(b) decreasing atomic radius
- the atomic radius decreases from left to right across a period
- this means that the electrons in the outermost level are becoming closer to the nucleus
- therefore there is a greater attraction between the nucleus and these outer electrons