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Chemistry > 3 Chemical bonding > Flashcards

3 Chemical bonding Flashcards

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1
Q

Electronegativity

A

the ability of an atom to attract a pair of electrons towards itself in a covalent bond

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2
Q

Which element is the most electronegative ?

A

-Fluorine is the most electronegative with a value of 4.0 as it is the best at attracting electron density towards itself when covalently bonded to another atom

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3
Q

Factors influencing electronegativity

A
  • nuclear charge
  • atomic radius
  • shielding by inner shells and sub-shells
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4
Q

Nuclear charge

A
  • An increase in the number of protons leads to an increase in nuclear attraction for the electrons in the outer shells
  • Therefore, an increased nuclear charge results in an increased electronegativity
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5
Q

Atomic radius

A
  • 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
  • Therefore, an increased atomic radius results in a decreased electronegativity
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6
Q

Shielding by inner shells and sub-shells

A
  • Filled energy levels can shield the effect of the nuclear charge causing the outer electrons to be less attracted to the nucleus
  • Therefore, an addition of extra shells and subshells in an atom will cause the outer electrons to experience less of the attractive forces of the nucleus
  • An increased number of inner shells and subshells will result in a decreased electronegativity
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7
Q

Across a period electronegativity increases, because…

A
  • The nuclear charge increases with the addition of protons to the nucleus
  • Shielding remains reasonably the same as no new shells are being added to the atoms
  • The nucleus has an increasing strong attraction for the bonding pair of electrons of atoms across the period
  • Results in smaller atomic radii
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8
Q

Down a group electronegativity decreases, because…

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 the shielding
  • The addition of the extra shells increases the distance between the nucleus and the outer electrons resulting in larger atomic radii
  • Decrease in attraction between the nucleus and outer bonding electrons
  • increased shielding and increased atomic radius outweigh the affects of the increased nuclear charge
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9
Q

Using the Pauling electronegativity values to predict the formation of ionic bonds

A
  • When atoms of different electronegativities form a molecule, the shared electrons are not equally distributed in the bond
  • The more electronegative atom (atom with higher value) will draw the bonding pair of electrons towards itself and a molecule with partial charges forms
  • The more electronegative atom will have a partial negative charge (delta negative)
  • The less electronegative atom will have a partial positive charge (delta positive)
  • Leads to a polar covalent molecule
  • If there is a large difference in electronegativity of the two atoms in a molecule, the least electronegative atoms electrons will transfer to the other atom which leads to an ionic bond
  • The cation is positively charged and has lost electrons
  • The anion is negatively charged and has gained electrons
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10
Q

Using the Pauling electronegativity values to predict the formation of covalent bonds

A
  • Single covalent bonds are formed by the sharing of electrons between the two atoms
  • In diatomic molecules, the electron density is shared equally between the two atoms
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11
Q

Ionic bonding (transfer of electrons from a metallic element to a non-metallic element)

A
  • the electrostatic attraction between oppositely charged ions (positively charged cations and negatively charged anions)
  • This form of attraction is very strong and requires a lot of energy to overcome therefore causing high melting points in ionic compounds
  • Ions form a lattice structure which is an evenly distributed crystalline structure arranged in a regular repeating pattern (positive charged cancel out negative ones therefore final lattice is overall electrically neutral)

Metals-lose electrons from their valence shell forming positively charged cations

Non-metals- gain electrons forming negatively charged anions

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12
Q

Metallic bonding

A
  • the electrostatic attraction between the positive metal ions and the sea of delocalized electrons
  • Metal atoms are tightly packed together in lattice structures and the electrons in their outer shells are free to move throughout the structure (delocalized electrons)
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13
Q

Covalent bonding

A

the electrostatic attraction between the nuclei of two atoms and a shared pair of electrons

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14
Q

Expanding the octet rule (elements in period 3)

A

when the central atom of a covalently bonded molecule can accommodate more than 8 electrons in its outer shell

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15
Q

Electron deficient

A

accommodating less than 8 electrons in the outer shell

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16
Q

σ bonds

A
  • Sigma bonds are formed by direct overlap of orbitals between the bonding atoms
  • The electron density in the bond is symmetrical about a line joining the nuclei of the atoms forming the bond
  • The pair of electrons are found between the nuclei of the 2 atoms
  • The electrostatic attraction between the electrons and nuclei bonds the atoms to each other
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17
Q

π bond

A

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

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18
Q

Hydrogen

A
  • The hydrogen atom has only one s orbital
  • The s orbitals of the two hydrogen atoms will overlap to form a sigma bond.
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19
Q

Ethene

A
  • Each carbon atom uses three out of its four electrons to form sigma bonds.
  • Two sigma bonds are formed with the hydrogen atoms and one sigma bond is formed with the other carbon atom.
  • The 4th electron from each carbon atom occupies a p orbital which overlaps sideways with another p orbital on the other carbon atom to form a pi bond.
  • Means that the C-C is a double bond: one sigma and one pi.
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20
Q

Ethyne

A
  • Molecule contains a triple bond formed from two pi bonds (at right angles to each other) and one sigma bond.
  • Each carbon atom uses two of its four electrons to form sigma bonds.
  • One sigma bond is formed with the hydrogen atom and one sigma bond is formed with the other carbon atom.
  • Two electrons are used to form two pi bonds with the other carbon atom
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21
Q

Hydrogen cyanide

A
  • Molecule contains a triple bond.
  • One sigma bond is formed between the H and C atom (overlap of an sp C hybridized orbital with the 1s H orbital.
  • A second sigma bond if formed between the C and N atom (overlap of an sp C hydridised orbital with an sp orbital of N).
  • The remaining two sets of p orbitals of nitrogen and carbon will overlap to form two pi bonds at right angles to each other
22
Q

Nitrogen

A
  • Nitrogen too contains a triple bond.
  • The triple bond is formed from the overlap of the sp orbitals on each N to form a sigma bond and the overlap of two sets of p orbitals on the nitrogen atoms to form two pi bonds at right angles to each other.
23
Q

Hybridisation

A
  • mixing of atomic orbitals to form covalent bonds
  • p atomic orbitals can also overlap end-on to form sigma bonds, but for them to do this they need to first become modified in order to gain s orbital character
  • The orbitals are therefore slightly changed in shape to make one of the p orbital lobes bigger

1s orbital + 3p orbitals= sp3 hybridisation

1s orbital + 2p orbitals= sp2 hybridisation

1s orbital + 1p orbital= sp hybridised orbitals

24
Q

Bond energy

A
  • the energy required to break one mole of a particular covalent bond in the gaseous state
  • The larger the bond energy, the stronger the covalent bond is
25
Q

Bond length

A
  • the internuclear distance of two covalently bonded atoms
26
Q

Explanation of bond length

A
  • The greater the forces of attraction between electrons and nuclei, the more the atoms are pulled closer to each other
  • This decreases the bond length of a molecule and increases the strength of the covalent bond
  • Triple bonds are the shortest and strongest covalent bonds due to the large electron density between the nuclei of the 2 atoms
  • This increases forces of attraction between the electrons and nuclei of the atoms and as a result the atoms are pulled closer together causing a shorter bond length
  • The increased forces of attraction also means that the covalent bond is strong
27
Q

The reactivity of a covalent bond is greatly influenced by:

A
  • The bond polarity
  • The bond strength
  • The bond type (sigma/pi)
28
Q

How to use bond energy values and the concept of bond length to compare the reactivity of covalent molecules

A

long bond length +smallest bond energy=most reactive as it takes the least amount of energy to break apart and visa versa with short bond length

29
Q

Valence shell electron pair repulsion theory (VSEPR)

A

predicts the shape and bond angles of molecules

30
Q

VSERP rules

A
  • Valence shell electrons are those electrons that are found in the outer shell
  • Electron pairs repel each other as they have similar charges
  • Lone pair electrons repel each other more than bonded pairs
  • Repulsion between multiple and single bonds is treated the same as for repulsion between single bonds
  • Repulsion between pairs of double bonds are greater
  • The most stable shape is adopted to minimize the repulsion forces
31
Q

Different types of electron pairs have different repulsive forces

A
  • Lone pairs of electrons have a more concentrated electron charge cloud than bonding pairs of electrons
  • The cloud charges are wider and closer to the central atoms nucleus
  • Order of repulsion is therefore: lone pair-lone pair > lone pair-bond pair > bond pair-bond pair
32
Q

Hydrogen bonding

A
  • Strongest form of intermolecular bonding
  • A type of permanent dipole-permanent dipole bonding
  • For hydrogen bonding to take place a species which has a O or N (very electronegative) atom with an available lone pair of electrons and a species with an –OH or –NH group is needed (angle between the –OH/-NH and the hydrogen bond is 180 degrees)
  • When hydrogen is covalently bonded to an electronegative atom (O/N) the bond becomes very highly polarised
  • The H becomes so delta positive that it can form a bond with the lone pair of an O/N atom in another molecule
33
Q

What are the properties of water (ice and water)

A
  • Its relatively high melting and boiling points
  • Its relatively high surface tension
  • The density of the solid ice compared with the liquid water
34
Q

High melting and boiling points

A
  • Caused by the strong intermolecular forces of hydrogen bonding between the molecules
  • In ice and water, the molecules are tightly held together by hydrogen bonds
  • A lot of energy is therefore required to break the water molecules apart and melt/boil them
35
Q

High surface tension

A
  • Surface tension is the ability of a liquid surface to resist any external forces
  • The water molecules at the surface of liquid are bonded to other water molecules through hydrogen bonds and these molecules pull downwards on the surface molecules causing the surface ones to become compressed and more tightly together at the surface
  • Increases surface tension
36
Q

Density (why ice has a lower density than liquid water)

A
  • Solids are denser than liquids as the particles in solids are more closely packed together than in their liquid state
  • In ice though the water molecules are packed in a 3D hydrogen-bonded network in a rigid lattice and each oxygen is surrounded by hydrogen atoms
  • This way of packing the molecules in a solid and the relatively long bond lengths of the hydrogen bonds means that the water molecules are slightly further apart than in the liquid form
  • Therefore, ice has a lower density than liquid water
37
Q

Non-polar

A

when 2 atoms in a covalent bond have the same electronegativity

38
Q

Polar

A
  • when 2 atoms in a covalent bond have different electronegativities
  • As a result of this the negative charge center and positive charge centre do not coincide with each other (electron distribution is asymmetric)
  • The less electronegative atom gets a partial charge of δ+ (delta positive)
  • The more electronegative atom gets a partial charge of δ- (delta negative)
  • The greater the difference in electronegativities the more polar the bond becomes
39
Q

Dipole moment

A
  • a measure of how polar a bond is
  • the direction of the dipole moment is show by an arrow which points to the partially negatively charged end of the dipole
40
Q

When assigning polarity to molecules you need to consider?

A
  • Polarity of each bond
  • How the bonds are arranged in the molecule
41
Q

Why some molecules have polar bonds but are overall not polar

A

some molecules have polar bonds but are overall not polar, because the polar bonds in the molecule are arranged in such a way that the individual dipole moments cancel each other out

42
Q

van der Waals forces and dipoles

A

the intermolecular forces between molecular entities other than those due to bond formation (use as a generic term to describe all intermolecular forces)

43
Q

Intermolecular forces

A

forces between a molecule

44
Q

Intramolecular forces

A

forces within a molecule

45
Q

2 types of van der Waals forces:

A
  • Instantaneous dipole-induced dipole (id-id) forces
  • Permanent dipole-permanent dipole (pd-pd)
46
Q

Instantaneous dipole-induced dipole (id-id) forces

A
  • Exist between all molecules/atoms
  • The electron charge cloud in non-polar molecules/atoms are constantly moving and during this, the electron charge cloud can be more on one side of the atom than the other and this causes a temporary dipole to arise
  • This temporary dipole can induce a dipole on neighboring molecules and when this happens the δ+ end of the dipole in one molecule and the δ- end of the dipole in a neighboring molecule are attracted towards each other and because the electron clouds are always moving the dipoles are only temporary
47
Q

Id-id forces increase with:

A
  • increasing number of electrons (and atomic number) in the molecule
  • increasing the places where the molecules come close together
48
Q

Permanent dipole-permanent dipole (pd-pd)

A
  • Polar molecules have permanent dipoles
  • The molecules with always have a negatively and positively charged end
  • Forces between 2 molecules that have permanent dipoles are called permanent dipole-permanent dipole forces
  • The δ+ end of the dipole in one molecule and the δ- end of the dipole in a neighboring molecule are attracted towards each other
  • For smaller molecules with the same number of electrons, pd-pd forces are stronger than id-id
49
Q

Hydrogen bonding as a special case of permanent dipole- permanent dipole

A
  • Hydrogen bonding is an intermolecular force between molecules with an –OH/-NH group and molecules with an N/O atom
  • Hydrogen bonds are stronger than pd-pd forces
  • The hydrogen is bonded to an O/N which is so electronegative that almost all the electron density from the covalent bond is drawn towards the O/N atom
  • This leaves the H with a large delta positive and the O/N with a large delta negative charge resulting in the formation of a permanent dipole in the molecule
  • A delta positive H in one molecule is electrostatically attracted to the delta negative O/N in a neighboring molecule
50
Q

What types of bonding are stronger than intermolecular forces?

A

ionic, covalent, and metallic bonding

51
Q

Order of forces

A

hydrogen, pd-pd, id-id

Chemistry (17 decks)

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