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UACE Chemistry > Inorganic Intro > Flashcards

Inorganic Intro Flashcards

1
Q

State Pauli’s exclusion principle

A

No two electrons in the same atom can have the same values for all the four quantum numbers

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2
Q

State Hund’s rule of maximum multiplicity

A

Electrons occupy degenerate (equal energy) orbitals singly with parallel spins before electron pairing occurs in any of the orbitals

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3
Q

State Aufbau’s principle

A

The orbitals are filled in order of increasing energy with the lowest filled first

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4
Q

Define atomic radius

A

This is half the internuclear distance between two atoms of the same element joined by a single covalent bond

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5
Q

Describe the factors that determine atomic radius of an element

A

Nuclear charge
- Increase in nuclear charge decreases atomic radius
- The higher the nuclear charge the more strongly electrons are attracted to the nucleus hence the smaller the atomic radius

Screening effect
- Increase in screening effect leads to increase in atomic radius
- The higher the screening effect, the lower the effective nuclear charge and the less strongly outer electrons attracted electrons are attracted to the nucleus; hence the larger the atomic radius

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6
Q

What is the trend of atomic radius across the period?

A

Atomic radius decreases across the period from left to right

  • Across the period from left to right, the nuclear charge progressively increases since a proton is added to the nucleus while screening effect remains constant since electrons are added to the same energy level
  • Effective nuclear charge increases and outer electrons become more strongly attracted and closer to the nucleus hence decrease in atomic radius
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7
Q

What is the trend of atomic radius down the group?

A

Atomic radius increases down the group

  • Down the group, both the nuclear charge (due to addition of protons) and screening effect (due to addition of shells of electrons) increase progressively.
  • However the increase in screening effect outweighs the increase in nuclear charge
  • Consequently, the effective nuclear charge decreases down the group, the outer electrons become less strongly attracted to the nucleus and atomic radius increases
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8
Q

Describe the trend of atomic radius among transition metals

A
  • Generally, along a given series of transition elements there is only a gradual decrease in atomic radius from left to right
  • Atomic radii of the first transition series decrease from Scandium to Chromium then remains almost constant till Nickel and increases from Copper to Zinc
  • There in increase in nuclear charge in the beginning of the series but ad electrons continue to be filled in the d-orbitals they screen the outer 4s electrons from influence of the nuclear charge
  • Increased nuclear charge and increased screening effect balance each other in the middle of the transition series and the atomic radius becomes constant
  • Towards the end of the series the repulsive interaction between electrons in the orbitals becomes very dominant. There is expansion in the electron cloud and atomic radius increases
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9
Q

Why is the radius of a cation smaller than that of the atom from which it is formed?

A
  • In a neutral atom, the number of protons in the nucleus is equal to that of electrons outside the nucleus
  • When a neutral atom loses one or more electrons in the outermost energy level, the electrons become fewer than protons
  • The remaining few electrons experience greater nuclear attraction making the resulting cation have a smaller radius than that of the atom from which it is formed.
  • For a given element, cationic radius decreases further with increase in number of electrons lost
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10
Q

Why is the radius of an anion larger than that of the atom from which it is formed?

A
  • When a neutral atom gains one or more electrons to form an anion, the number of electrons exceeds that of protons
  • The electron cloud expands and electrons experience less nuclear attraction making the resulting anion have a larger radius than that of the atom from which it is formed
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11
Q

Define ionization energy

A

The minimum energy required to remove one mole of electrons from one mole of free gaseous atoms.

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12
Q

What is first ionization energy?

A

The minimum energy required to remove one mole of electrons from one mole of free gaseous atoms to form one mole of (unipositively charged ions) ions with a single positive charge

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13
Q

Define second ionization energy

A

The minimum energy required to remove one mole of electrons from one mole of free gaseous unipositively charged ions to form one mole of ions with a double positive charge

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14
Q

State the factors affecting the magnitude of ionization energy

A
  1. Atomic radius
  2. Screening effect
  3. Nuclear Charge
  4. Electronic structure/ configuration
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15
Q

How does atomic radius affect first ionization energy of an element?

A
  • First ionization energy decreases with increase in atomic radius
  • Elements with big atomic radii> outermost electrons experience less nuclear attraction > little energy required to remove these electrons
  • Elements with small atomic radii> outermost electrons experience strong nuclear attraction > a lot of energy required to remove these electrons
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16
Q

How does nuclear charge affect first ionization energy of an element?

A
  • First ionization energy increases with increase in nuclear charge
  • High nuclear charge > outermost electrons are tightly held by the nucleus> high nuclear attraction = high ionization energy
  • Low nuclear charge > outermost electrons are less strongly held by the nucleus> low nuclear attraction = low ionization energy
17
Q

How does screening effect affect first ionization energy of an element?

A
  • First ionization energy decreases with increase in screening effect
  • High screening effect> more effectively outer electrons are shielded from nuclear attraction> lower ionization energy
  • Low screening effect> less effectively outer electrons are shielded from nuclear attraction> higher ionization energy
18
Q

How does electronic configuration affect first ionization energy of an element?

A

Completely filled or half filled orbitals tend to be very thermodynamically stable and require a lot of energy to remove an electron from them

19
Q

Describe the trend of first ionization energy in the periodic table down the group

A
  • First ionization energy generally decreases down the group
  • Down the group there is increase in both nuclear charge and screening effect. The progressive increase in screening effect outweighs the increase in nuclear charge causing a decrease in effective nuclear charge down the group
  • Decrease in effective nuclear charge implies that the outermost shell electrons become less attracted to the nucleus requiring less energy to be removed thus low first ionization energy
20
Q

Describe the trend of first ionization energy in the periodic table across the period

A
  • Generally, first ionization energy increases across the period from left to right
  • From one element to another, across the period from left to right, an electron is added to the same outer shell and a proton added to the nucleus as well
  • Screening effect of outer electrons by inner electrons remains almost constant while the nuclear charge progressively increases across the period
  • Effective nuclear charge of outermost electrons increases across the period leading to increase in first ionization energy
21
Q

Why is the first ionization energy of Magnesium higher than that of aluminum?

A
  • The electron in magnesium is obtained from a thermodynamically stable 3s sub-shell full of electrons which requires a lot of energy to be removed.
  • In aluminum, the 3p subshell from which the electron is removed contains only one electron which is effectively screened from nuclear attraction by inner electrons. This makes it thermodynamically unstable and little energy is required to remove the electron.
22
Q

Why is the first ionization energy of phosphorus higher than that of sulphur?

A
  • The electron in phosphorus is obtained from a thermodynamically stable half-full 3p sub-shell which requires a lot of energy to be removed.
  • In sulphur the 3p sub-shell is neither half full nor full of electrons therefore it is thermodynamically unstable and little energy is required to remove the first electron.
23
Q

Define electron affinity

A

The energy liberated when each atom in one mole of free gaseous atoms gains an electron

24
Q

Explain why the second electron affinity is endothermic first is exothermic

A

When a negatively charged ion is formed its negative charge repels any further incoming electrons to be added to it such that a lot of energy must be supplied for the next electron to be added.

25
Q

Describe the factors that affect electron affinity

A

Atomic radius
- The smaller the atomic radius the stronger the incoming electron is attracted to the nucleus and the greater the electron affinity
- The larger the atomic radius the less nuclear attraction on the incoming electron and the less the electron affinity

Nuclear charge
- The greater the nuclear charge the more strongly the incoming electron is attracted to the nucleus and the greater the electron affinity
- The smaller the nuclear charge the less the incoming electron is attracted to the nucleus and the smaller the electron affinity

Screening effect
- the greater the screening effect the less the incoming electron is attracted to the nucleus and the smaller the electron affinity
- The less the screening effect the more strongly the incoming electron is attracted to the nucleus thus the larger the electron affinity

26
Q

Describe the trend of electron affinity in the periodic table down the group

A

Down the group electron affinity generally decreases

  • Increase in both nuclear charge and screening effect > due to extra shells of electrons being added the screening effect outweighs the nuclear charge leading to decrease in the effective nuclear charge
  • Reduced ability of atoms to attract electrons hence reduction in electron affinity
27
Q

Define electronegativity

A

The measure of the ability of an atom or element to attract the bonding pair of electrons in a covalent bond towards itself

28
Q

What factors affect electronegativity?

A

Nuclear charge
- The greater the nuclear charge the more strongly the bonding electrons are attracted and the greater the electronegativity

Screening effect
- The less the screening effect the stronger the nuclear attraction and effective nuclear charge on the bonding electrons, and the higher the electronegativity
- The higher the screening effect, the less the nuclear attraction and effective nuclear charge on the bonding electrons and the lower the electronegativity

Atomic radius
- The smaller the atomic radius, the more strongly the bonding electrons are attracted by the nucleus as the electrons are close to the nucleus and the greater the electronegativity

29
Q

What is the trend of electronegativity in the periodic table down the group?

A

Electronegativity decreases down the group

  • Both nuclear charge and screening effect increase down the group, but the increase in screening effect due to extra shells outweighs the increase in nuclear charge, leading to decrease in effective nuclear charge
  • This means the atoms ability to attract electrons decreases down the group
30
Q

What is the trend of electronegativity in the periodic table across the period?

A

Electronegativity increases across the period

  • Across the period from left to right the nuclear charge increases from one elements to the next due to an extra proton added, but the screening effect remains constant as electrons are added to the same outer most shell.
  • Therefore, the effective nuclear charge increases and the atoms ability to attract electrons increases as well hence the increase in electronegativity.
31
Q

True or false
Binary compounds are ionic if the electronegativity difference between the two bonded atoms is > 1.8

A

True

32
Q

True or false
Binary compounds are covalent if the electronegativity difference between the two bonded atoms is <=1.8

A

True

33
Q

True or false
Binary compounds are non-polar if the electronegativity difference between two bonded atoms is 0

A

True

UACE Chemistry (20 decks)

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