S2.2 The covalent model SL Part 1 Flashcards

1
Q

Octet rule

A

The tendency of an atom to achieve stability by ensuring its valence shell is full, either by gaining, losing or sharing electrons.

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2
Q

Covalent bond

A

The electrostatic force of attraction between two positive nuclei and a shared pair of electrons.

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3
Q

Lewis formula

A

This shows all valence electrons, whether these are pairs of electrons involved in covalent bonds or non-bonded pairs of electrons, known as non-bonding pairs. Also known as an electron dot or Lewis structure.

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4
Q

How do you deduce the Lewis formula for molecules and ions?

A

Calculate the total number of valence electrons, draw a skeletal structure with single bonds, add non-bonding pairs, and ensure all atoms have a full valence shell.

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5
Q

Which molecules are stable with fewer than an octet of electrons?

A

Molecules like beryllium dichloride (BeCl2) and boron trifluoride (BF3) are stable with incomplete octets by forming the maximum number of covalent bonds.

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6
Q

How to name a simple molecular compound from its chemical formula?

A

Use prefixes (mono-, di-, tri-, etc.) to denote the number of atoms, omitting the prefix for a single atom of the first element listed, unless it’s a single-element compound, which retains the element’s name.

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7
Q

How to calculate the total number of valence electrons for molecules and ions with double and triple bonds?

A

Sum the valence electrons of all atoms present. For ions, add an electron for each negative charge or subtract one for each positive charge.

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8
Q

What is the first step in drawing the skeletal structure for Lewis formulas?

A

Identify the central atom (usually the least electronegative), connect peripheral atoms with single bonds.

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9
Q

After connecting atoms with single bonds, what do you do with the remaining electrons?

A

Distribute them as non-bonding pairs (lone pairs) on the outer atoms to fulfill the octet rule.

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10
Q

How are double or triple bonds formed in the Lewis structure?

A

Use non-bonding pairs from outer atoms to form double or triple bonds with the central atom to complete octets.

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11
Q

Final step in deducing the Lewis structure?

A

Ensure the total number of electrons used matches the calculated number of valence electrons.

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12
Q

How does the number of bonds affect bond strength?

A

More bonds result in stronger bonds. Triple bonds are the strongest, followed by double, then single bonds.

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13
Q

Relationship between bond numbers and bond length?

A

Double and triple bonds are shorter than single bonds due to greater electrostatic attraction pulling nuclei closer.

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14
Q

Coordination bonds

A

A covalent bond where both of the electrons being shared in the bond have been donated by one atom.

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15
Q

How can coordination bonds in compounds be identified?

A

Coordination bonds occur when one atom donates both electrons for a bond, while the other atom accepts them without contributing electrons. They are a special type of covalent bond that differs in the source of electrons.

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16
Q

How do coordination bonds form?

A

Coordination bonds form between a donor species, with non-bonding electrons available for bonding, and an acceptor species that has space to accept electrons. This interaction facilitates stability and adherence to the octet rule.

17
Q

Why are coordination bonds important?

A

They are crucial for molecules that would not achieve a stable electronic configuration otherwise. Examples include carbon monoxide and complexes involving ammonia and boron trifluoride.

18
Q

How are coordination bonds represented?

A

They are depicted using Lewis formulas or an arrow notation that indicates electron donation from the donor to the acceptor.

19
Q

Electron domains

A

The region in which bonding and non-bonding pairs of electrons are most likely to be found. Non-bonding pairs, single bonds, double bonds and triple bonds each count as one electron domain.

20
Q

Electron domain geometry

A

The region in which bonding and non-bonding pairs of electrons are most likely to be found. Non-bonding pairs, single bonds, double bonds and triple bonds each count as one electron domain.

21
Q

Valence Shell Electron Pair Repulsion (VSEPR)

A

A theory used to explain and predict molecular geometry, based on number of electron domains (bonded and non-bonded electron pairs).

22
Q

How to predict electron domain and molecular geometry for up to four domains?

A

Count the electron domains around the central atom: linear (2 domains), trigonal planar (3 domains), tetrahedral (4 domains). Molecular geometry considers atom arrangement, including non-bonding pairs.

23
Q

How do non-bonding pairs and multiple bonds affect bond angles?

A

Non-bonding pairs are more repulsive, reducing bond angles. Multiple bonds count as one electron domain but affect molecule shape due to electron density.

24
Q

Electronegativity

A

A measure of how much an atomic nucleus attracts the shared electrons that are involved in a covalent bond.

25
Q

Pauling scale

A

The scale on which electronegativity is measured, ranging from 0.0 to 4.0 Pauling units.

26
Q

Non-polar bonds

A

A covalent bond formed between two atoms with equal electronegativity values which results in equal sharing of the electrons in their bond.

27
Q

Polar covalent bond

A

A bond which has a partially positive end and a partially negative end due to asymmetrical electron distribution caused by the difference in electronegativity of the two bonded species.

28
Q

How does molecular symmetry affect the net dipole moment?

A

Symmetry plays a crucial role; in asymmetrical molecules with polar bonds, these polarities contribute to a net dipole moment. Conversely, in symmetrical molecules, polarities may cancel out, resulting in no net dipole moment.

29
Q

What initial step is taken to deduce a molecule’s net dipole moment?

A

The first step involves identifying the presence of polar bonds within the molecule or ion to understand potential contributions to the net dipole moment.

30
Q

What effect does asymmetry have on a molecule’s dipole moment?

A

Asymmetry in a molecule with polar bonds typically leads to a net dipole moment, indicating distinct positive and negative ends within the molecule.