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A-level Chemistry > 2.4 - Electrons, bonding and structure > Flashcards

2.4 - Electrons, bonding and structure Flashcards

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1
Q

Define the term ‘orbital’ ?

A

Orbital - A region of space about the nucleus holding up to 2 electrons with opposite spins

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2
Q

What are the different types of orbitals ?

A

s, p, d, f

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3
Q

What is the shape of s orbitals ?

A

Spherical

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4
Q

What is the shape of p orbitals ?

A

Dumbbell-shaped

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5
Q

What is the number of orbitals in the s, p, d and f sub-shells ?

A
  • s = one orbital ( two electrons )
  • p = three orbitals ( six electrons )
  • d = five orbitals ( ten electrons )
  • f = seven orbitals ( fourteen electrons )
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6
Q

Do you know how to write electron configuration of atoms and ions ?

A

Yes

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7
Q

What are the rules for how shells are occupied by electrons ?

A
  • The orbitals fills from the lowest energy levels
  • The 4s orbital fills up before the 3d orbital
  • The 4s orbital loses electrons first before the 3d orbital
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8
Q

Do you know how to write electron configuration using boxes ?

A
  • Electrons will prefer to occupy orbitals on their own unless that option is unavailable
  • Yes…
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9
Q

What are the exceptions to the rules of electron configuration ?

A
  • Chromium and Copper
  • The 4s orbital is occupied by one electron and is not completed as the 3d orbital is filled
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10
Q

Write the electron configuration for chromium and copper ?

A
  • Chromium = 1s² 2s² 2p⁶ 3s² 3p⁴ 4s¹3d⁵
  • Copper = 1s² 2s² 2p⁶ 3s² 3p⁴ 4s¹ 3d¹⁰
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11
Q

Define the term ‘ionic bond’ ?

A

Ionic bond - The strong electrostatic attraction between oppositely charged ions

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12
Q

What is the structure of ionic compounds ?

A
  • They have a Giant Ionic lattice structure
  • Each ion is surrounded by oppositely charged ions held together by strong ionic bonds
  • The regular cubic arrangement gives the crystal its cubic shape
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13
Q

Explain the structure of an ionic compound such as sodium chloride ( NaCl ) ?

A
  • Each positive sodium ion is surrounded by six negative chlorine ions forming a giant ionic lattice
  • The regular cubic arrangement gives the crystal its cubic shape
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14
Q

Draw a diagram representing the structure of an ionic compound ( eg. NaCl ) ?

A
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15
Q

What are the physical properties of ionic compounds ?

A
  • High melting/ boiling point
  • Only conducts electricity when molten/ solution
  • Brittle
  • Soluble in polar solvents ( eg. Water )
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16
Q

Explain why ionic compounds have a high melting/ boiling point ?

A

Strong ionic bonds/ electrostatic attraction between oppositely charged ions require a lot of energy to break resulting in a high melting/ boiling point

17
Q

Explain why ionic compounds can only conduct electricity when molten/ solution ?

A

Ions are mobile when an ionic compound is molten/ solution and therefore carry a charge. Ions are not able to move when solid.

18
Q

Explain why ionic compounds are brittle ?

A

Applying a force causes like charged ions to align and therefore to repel

19
Q

Explain why ionic compounds are soluble in polar solvents ?

A

Polar solvents such as water can often dissolve ionic compounds

20
Q

Do you know how to draw dot and cross diagrams for ionic compounds ?

A

Yes

21
Q

Define the term ‘covalent bond’ ?

A

Covalent bond - The strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms

22
Q

Define the term ‘dative covalent bond’ ?

A

Dative covalent bond - A covalent bond where both of the electrons come from the same species/ atoms

23
Q

Do you know how to draw dot and cross diagrams for covalent bonding ?

A

Yes

24
Q

What are double or triple covalent bonds ?

A
  • A double covalent bond is two shared pair of electrons
  • A triple covalent bond is three shared pair of electrons
25
Q

Do you know how to draw dot and cross diagrams to represent multiple and dative covalent bonding ?

A

Yes

26
Q

What is another name for a ‘dative covalent bond’ ?

A

A co-ordinated bond

27
Q

How are ions formed ?

A
  • Metal atoms lose electrons to form positive ions ( cations )
  • Non-metal atoms gains electrons to form negative ions ( anions )
28
Q

Define the term ‘Average bond enthalpy’ ?

A

Average bond enthalpy - The average bond enthalpy term is the average amount of energy needed to break a specific type of bond homolytically in gaseous molecules , measured over a wide variety of different compounds.

29
Q

What does average bond enthalpy mean in regards to covalent bond strength ?

A

The higher the average bond enthalpy, the greater the strength of the covalent bonds

30
Q

Define the term ‘electronegativity’ ?

A

Electronegativity - the ability of an atom to attract the electrons in a covalent bond

31
Q

What are the factors affecting electronegativity ?

A
  • Nuclear Charge
  • Atomic Radius
  • Shielding
32
Q

Explain how nuclear charge affects electronegativity ?

A
  • As the number of protons increases in the nucleus, the greater the nuclear charge
  • This leads to stronger attraction between the nucleus and bonding pairs of electrons leading to an increase in electronegativity
33
Q

Explain how atomic radius affects electronegativity ?

A
  • As the atomic radius decreases, the distance between the nucleus and bonding pairs of electrons decreases
  • This leads to stronger attraction between the nucleus and bonding pairs of electrons leading to an increase in electronegativity
34
Q

Explain how shielding affects electronegativity ?

A
  • As the number of electrons increases, the number of inner electrons shells and therefore shielding increases
  • This leads to weaker attraction between the nucleus and bonding pairs of electrons causing the electronegativity to decrease
35
Q

What is the trend in electronegativity down the group ?

A

As you go down the group, the electronegativity decreases

36
Q

What is the trend in electronegativity across the period ?

A

As you go across the group, the electronegativity increases

37
Q

Explain why electronegativity decreases down the group ?

A
  • The atomic radius increases due to an increased number of electrons shells
  • The shielding increases due to an increased number of electrons
38
Q

Explain why electronegativity increases across the period ?

A
  • The nuclear charge increases due to more protons
  • Atomic radius decreases due to more protons in the nucleus
  • The shielding stays the same as all elements are in the same period
39
Q

Explain how atomic radius decreases across a period ?

A
  • As you go across a period, the number of protons in the nucleus increases causing nuclear charge to increase
  • This means there is a greater attraction force between the nucleus and electrons
  • This means electrons within shells are held closer to the nucleus which decreases the atomic radius

A-level Chemistry (29 decks)

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