Chapter 7: Trends in the Periodic Table (Definitions) Flashcards

1
Q

Define atomic radius(covalent radius) of an atom.

4

A
  • half the distance between the nuclei of two atoms
  • of the same element
  • that are joined together
  • by a single covalent bond
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2
Q

Why does the values of the atomic radius increase down the groups in the Periodic Table?

(a) =3
(b) =4

A

(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.

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3
Q

Why does the values of the atomic radius decrease across a period in the Periodic Table?

(a) =4
(b) =4

A

(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

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4
Q

Define first ionisation energy of an atom

3

A
  • the minimum energy required to completely remove the most loosely bound electron
  • from a neutral gaseous atom
  • in its ground state.
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5
Q

What is the Equation for first ionisation energy(sodium)

A

Na(g)->Na+(g)+e-

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6
Q

Why does the values of the first ionisation energy decrease down a group in the Periodic Table

(a) =3
(b) =5

A

(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

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7
Q

Why does the values of the first ionisation energy increase across a period in the Periodic Table

(a) =5
(b) =4

A

(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

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8
Q

Define second ionisation energy of an atom.

4

A
  • the energy required to remove an electron
  • from an ion
  • with one positive charge
  • in the gaseous state
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9
Q

What is the equation for second ionisation energy

A

X+(g)->X2+(g)+e-

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10
Q

Why does the values of electronegativity decrease down the groups in the Periodic Table

(a) =3
(b) =4

A

(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.

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11
Q

Why does the values of electronegativity increase across the periods in the Periodic Table

(a) =3
(b) =3

A

(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

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