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bonding Flashcards

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1
Q

what is electronegativity

A
  • Electronegativity is the ability of an atom to attract a pair of electrons towards itself in a covalent bond
  • This phenomenon arises from the positive nucleus’s ability to attract the negatively charged electrons, in the outer shells, towards itself
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2
Q

what is the pauling scale

A

The Pauling scale is used to assign a value of electronegativity for each atom

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3
Q

which atom is the most electronegative

A
  • Fluorine is the most electronegative atom on the Periodic Table, with a value of 4.0 on the
  • It is best at attracting electron density towards itself when covalently bonded to another atom
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4
Q

how does nuclear charge affect electronegativity

A

an increased nuclear charge results in an increased electronegativity

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5
Q

why does nuclear charge increase electronegativity

A
  • Attraction exists between the positively charged protons in the nucleus and negatively charged electrons found in the energy levels of an atom
  • An increase in the number of protons leads to an increase in nuclear attraction for the electrons in the outer shells
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6
Q

how does atomic radius affect electronegativity

A

an increased atomic radius results in a decreased electronegativity

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7
Q

why does atomic radius decrease electronegativity

A

The atomic radius is the distance between the nucleus and electrons in the outermost shell
Electrons closer to the nucleus are more strongly attracted towards its positive nucleus
Those electrons further away from the nucleus are less strongly attracted towards the nucleus

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8
Q

how does shielding affect electronegativity

A

an increased number of inner shells and subshells will result in a decreased electronegativity

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9
Q

why does shielding decrease electronegativity

A

Filled energy levels can shield (mask) the effect of the nuclear charge causing the outer electrons to be less attracted to the nucleus
Therefore, the addition of extra shells and subshells in an atom will cause the outer electrons to experience less of the attractive force of the nucleus

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10
Q

what is the trend of electronegativity down a group

A

There is a decrease in electronegativity going down the group

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11
Q

why does electronegativity decrease down a group

A

The nuclear charge increases as more protons are being added to the nucleus
However, each element has an extra filled electron shell, which increases shielding
The addition of the extra shells increases the distance between the nucleus and the outer electrons resulting in larger atomic radii
Overall, there is decrease in attraction between the nucleus and outer bonding electrons

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12
Q

What is the trend of electronegativity across a period

A

Electronegativity increases across a Period

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13
Q

why does electronegativity increase across a period

A

The nuclear charge increases with the addition of protons to the nucleus
Shielding remains reasonably the same across the Period as no new shells are being added to the atoms
The nucleus has an increasingly strong attraction for the bonding pair of electrons of atoms across the Period of the Periodic Table
This results in smaller atomic radii

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14
Q

what kind of molecule is a non-polar molecule

A

Both atoms will have the same electronegativity value and have an equal attraction for the bonding pair of electrons leading to formation of a covalent bond
The equal distribution leads to a non-polar molecule

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15
Q

what kind of molecule is a polar covalent molecule (how are they formed)

A

When atoms of different electronegativities form a molecule, the shared electrons are not equally distributed in the bond
The more electronegative atom (the atom with the higher value on the Pauling scale) will draw the bonding pair of electrons towards itself
A molecule with partial charges forms as a result
The more electronegative atom will have a partial negative charge (delta negative, δ–)
The less electronegative atom will have a partial positive charge (delta positive, δ+)

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16
Q

what happens if there is a large difference in electronegativity of two atoms in a molecule

A

If there is a large difference in electronegativity of the two atoms in a molecule, the least electronegative atom’s electron will transfer to the other atom
This in turn leads to an ionic bond – one atom transfers its electron and the other gains that electron
The cation is a positively charged species which has lost (an) electron(s)
The anion is a negatively charged species which has gained (an) electron(s)

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17
Q

what value of difference in electronegativity leads to a covalent bond

A

<1.0

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18
Q

what value of difference in electronegativity leads to a polar covalent bond

A

1.0 - 2.0

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19
Q

what value of difference in electronegativity leads to an ionic bond

A

> 2.0

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20
Q

what is dative covalent bonding/coordinate bond

A

In simple covalent bonds the two atoms involved shares electrons
Some molecules have a lone pair of electrons that can be donated to form a bond with an electron-deficient atom
An electron-deficient atom is an atom that has an unfilled outer orbital
So both electrons are from the same atom

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21
Q

what is molecular orbital

A

When a covalent bond is formed, the atomic orbitals overlap to form a combined orbital containing two electrons

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22
Q

how does the overlap of the orbitals affect the strength of the bond

A

The greater the atomic orbital overlap, the stronger the bond

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23
Q

what are the two types of molecular orbitals

A
  • sigma

- pi

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24
Q

what are sigma bonds

A

Sigma (σ) bonds are formed by direct overlap of orbitals between the bonding atoms

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25
Q

what are pi bonds

A

Pi (π) bonds are formed by the sideways overlap of adjacent above and below the σ bond

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26
Q

where is the electron density in a sigma bond

A

The electron density in a σ bond is symmetrical about a line joining the nuclei of the atoms forming the bond
The pair of electrons is found between the nuclei of the two atoms

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27
Q

what bonds the atoms together in a sigma bond

A

The electrostatic attraction between the electrons and nuclei bonds the atoms to each other

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28
Q

how are pi bonds formed

A

Pi (π) bonds are formed from the sideways overlap of adjacent p orbitals

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29
Q

what is the overlap/structure of pi bond

A

The two lobes that make up the π bond lie above and below the plane of the σ bond
This maximises overlap of the p orbitals

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30
Q

how are pi bonds drawn

A

A single π bond is drawn as two electron clouds one arising from each lobe of the p orbitals
The two clouds of electrons in a π bond represent one bond containing two electrons

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31
Q

what is hybridisation

A

the mixing of atomic orbitals to form covalent bonds

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32
Q

how does hybridisation happen

A

The p atomic orbitals can also overlap end-on to form σ bonds
In order for them to do this, they first need to become modified in order to gain s orbital character
The orbitals are therefore slightly changed in shape to make one of the p orbital lobe bigger

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33
Q

what is an sp3 hybridisation

A

Mixing an s with three p orbitals is called sp3 hybridisation (each orbital has ¼ s character and ¾ p character)

34
Q

what is an sp2 hybridisation

A

Mixing an s with tw p orbital is called sp2 hybridisation

35
Q

what is an sp hybridised orbital

A

mixing an s with one p-type orbitals

36
Q

what is expanding the octet rule

A

In some instances, the central atom of a covalently bonded molecule can accommodate more or less than 8 electrons in its outer shell
Being able to accommodate more than 8 electrons in the outer shell is known as ‘expanding the octet rule’

37
Q

what is electron deficient

A

Accommodating less than 8 electrons in the outer shell means than the central atom is ‘electron deficient’

38
Q

what is an ionic bond

A

Ionic bonds tend to be formed between elements whose atoms need to “lose” electrons to gain the nearest noble gas electronic configuration and those which need to gain electrons.
The electrons are transfered from one atom to the other.

39
Q

what is the structure of ionic bonds

A

ions are held together in a crystal lattice by electrostatic attraction

40
Q

what is the structure of metallic bonds

A

a metallic solid consists of a lattice of positive ions with the outer electrons forming a sea of delocalised electrons.

41
Q

what is a metallic bond

A

electrostatic force of attraction between the delocalised electrons cloud and the metal cations constitute the metallic bond.

42
Q

what does the strength of metallic bonds depend on

A

the strength of metallic bond depends on the size of the cation and the charge on the cation.

  • size of the cation bigger, the strength of metallic bond smaller
  • charge of the cation bigger, the strength of the metallic bond bigger.
43
Q

how do metal atoms achieve stability/how does metallic bonding happen

A

Metal atoms achieve stability by “off-loading” electrons to attain the electronic structure of the nearest noble gas. These electrons join up to form a mobile cloud which prevents the newly-formed positive ions from flying apart due to repulsion between similar charges.

44
Q

what is a covalent bond

A

An electrostatic force of attraction between two non-metals to share a pair of electrons.
Usually formed if the difference in electronegativity between the atoms are small.

45
Q

why are atoms held together in covalent bonds

A

because their nuclei which

have an overall positive charge are attracted to the shared electrons

46
Q

when are covalent bonds formed (between which atoms)

A
  • between atoms of the same element
  • between atoms of different elements on the RHS of the table
  • when one of the elements is in the middle of the table
  • with head-of-the-group elements with high charge density
47
Q

what is orbital theory

A

Covalent bonds are formed when orbitals, each containing one electron, overlap. This forms a region in space where an electron pair can be found; new molecular orbitals are formed.

48
Q

which compounds are an exception to the octet rule

A

Compounds with an odd number of electrons eg: NO, NO2
Compounds with less than 8 electrons in the outer shell per atom eg: BeCl2, AlCl3
Compounds with more than 8 electrons in the outer shell per atom eg: PCl5, SF6, BrF5

49
Q

which elements are an exception to the octet rule

A

Only elements of Period 3 and beyond can expand their octet to accommodate more than 8 electrons.

50
Q

do period 2 elements follow the octet rule

A

Period 2 elements can only accommodate a maximum of 8 electrons in its outer shell thus cannot expand their octet.

51
Q

what is the electronegativity of covalent bonds (polar or nonpolar)

A

Single covalent bonds are formed by sharing a pair of electrons between two atoms
In diatomic molecules the electron density is shared equally between the two atoms
Eg. H2, O2 and Cl2
Both atoms will have the same electronegativity value and have an equal attraction for the bonding pair of electrons leading to formation of a covalent bond
The equal distribution leads to a non-polar molecule

52
Q

what is the electronegativity of ionic bonds (polar or nonpolar)

A

If there is a large difference in electronegativity of the two atoms in a molecule, the least electronegative atom’s electron will transfer to the other atom
This in turn leads to an ionic bond – one atom transfers its electron and the other gains that electron
The cation is a positively charged species which has lost (an) electron(s)
The anion is a negatively charged species which has gained (an) electron(s)

53
Q

how are covalent bonds formed

A

covalent bonds are formed when the orbitals of two neighbouring atoms overlap so that both nuclei attract the pairs of electrons between them

54
Q

what are the two types of covalent bonds

A

sigma and pi

55
Q

how are sigma bonds formed

A

Sigma (σ) bonds are formed from the end-on overlap of atomic orbitals
S orbitals overlap this way as well as p orbitals

56
Q

what are the characteristics of sigma bonds

A

The electron density in a σ bond is symmetrical about a line joining the nuclei of the atoms forming the bond
The pair of electrons is found between the nuclei of the two atoms
The electrostatic attraction between the electrons and nuclei bonds the atoms to each other

57
Q

how are pi bonds formed

A

Pi (π) bonds are formed from the sideways overlap of adjacent p orbitals
The two lobes that make up the π bond lie above and below the plane of the σ bond
This maximises overlap of the p orbitals
pi bonds can only arise upon the formation of a sigma bond

58
Q

which bonds does a double bond have

A

1 sigma 1 pi

59
Q

which bonds does a triple bond have

A

1 sigma and 2 pi

60
Q

what is hybridisation

A

The p atomic orbitals can also overlap end-on to form σ bonds
In order for them to do this, they first need to become modified in order to gain s orbital character
The orbitals are therefore slightly changed in shape to make one of the p orbital lobe bigger
This mixing of atomic orbitals to form covalent bonds is called hybridisation

61
Q

what kind of hybridisation is Mixing an s with three p orbitals

A

sp3 hybridisation (each orbital has ¼ s character and ¾ p character)

62
Q

what kind of hybridisation is Mixing an s with two p orbital

A

sp2 hybridisation

63
Q

what kind of hybridisation is mixing an s with one p-type orbitals

A

sp hybridised orbitals

64
Q

what is bond energy

A

the energy required to break one mole of a covalent bond between two atoms in the gaseous state
The larger the bond energy, the stronger the covalent bond is

65
Q

what is bond length

A

-the distance between the nuclei of the two atoms joined by a covalent bond (and hence depends on the size of the atoms)

66
Q

how does bond length and energy correlate

A
  • the shorter the bond length, the stronger is the bond

- the greater the bond energy, the stronger is the bond

67
Q

what dictates the shape of a covalent molecule

A

the electron pairs around the central atom of a molecule dictate the shape of the molecule

68
Q

what is the electron pair repulsion model

A

electron pairs in a molecule (both bond pairs and lone pairs) repel each other and move as far apart as possible

69
Q

what is the order of strength of repulsion in the electron pair repulsion model (biggest to smallest)

A

lone pair-lone pair repulsion > lone pair-bond pair repulsion> bond pair-bond pair repulsion

70
Q

what are bond pairs

A

electrons that are alone in an orbital and can bond with electrons from another atom

71
Q

what are lone pairs

A

electrons that are in a pair in an orbital, they are non-bonding which are under the control of only one atom

72
Q

what is VSEPR theory

A

-the repulsion between electron pairs is increased by the increase in electronegativity of the central atom

72
Q

what is VSEPR theory

A
  • the repulsion between electron pairs is increased by the increase in electronegativity of the central atom
  • the number of lone pairs and bond pairs around the central atom dictates the shape of the molecule
  • electron pairs stay as far apart as possible to overcome the repulsion
73
Q

what are the steps to find the shape of a molecule

A
  1. Find out the central atom (least amount in the formula)
  2. write the electronic configuration for the central atom
  3. check if the central atom can get excited
  4. count the number of electron pairs
74
Q

which atom cannot be a central atom

A
  • hydrogen because it can only form one bond with its limited number of electrons and central atoms need to be able to form multiple bonds
  • helium and lithium dont form covalent bonds
75
Q

what is the excited state

A

when an electron in a full orbital jumps to the next orbital with more empty spaces so it is more favourable for bonding

75
Q

what is the excited state

A

when an electron in a full orbital jumps to the next orbital with more empty spaces so it is more favourable for bonding

76
Q

what are the rules of bond shapes

A
  • same group/family means the molecules/bonds have the same shape because the central atom is the same
  • same formula does not mean same shape because the central atom isnt the same
77
Q

what is a coordinate/dative bond

A
  • if both the electrons in a covalent bond come from only one of the atom
  • once the bond is formed its identical to other covalent bonds
78
Q

what must dative/coordinate bonding atoms have

A
  • the donor group must have a lone pair of electrons in its outermost shell while the acceptor group must have vacant orbitals in its outer shell
  • this enables the acceptor to accomodate the lone pair of electrons from the donor

AS Chemistry (22 decks)

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