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Chemistry : Physical Chemistry AS > bonding > Flashcards

bonding Flashcards

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1
Q

define ionic bonding

A

-electrostatic attraction between oppositely charged ions in an ionic lattice

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2
Q

what is the formula of sulfate

A

SO4 2-

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3
Q

formula of hydroxide

A

OH-

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4
Q

formula for nitrate

A

NO3 -

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5
Q

formula for carbonate

A

CO3 2-

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6
Q

formula for ammonium

A

NH4 +

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7
Q

describe the physical properties of ionic bonding

A

1)- High MP : strong electrostatic forces of attraction between oppositely charged ions that hold the ionic lattice together
lots of energy required to break the bonds

2)-dissolve in water (soluble) : water is polar so they can attract the positive and negative ions and break up the structure

3)-conduct electricity in molten and in aqueous solution: the ions are free to move and so can carry charge

4)- cannot conduct in solid as the ions are not free to move as they are fixed in tight spaces in the giant ionic lattice

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8
Q

define covalent bonding

A

shared pair of electrons

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9
Q

define dative bonding/coordinate bonding

A

when the shared pair of electrons came in the covalent bond come from one atom.

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10
Q

what are the two types of covalent bonding?

A

-giant covalent/macromolecular structure
-simple molecular structure

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11
Q

what molecules are examples of macromolecular structure?

A

Diamond and Graphite

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12
Q

-describe the property of Diamond

A

1)giant lattice structure:

2)High Melting Point: carbon atoms is covalently bonded to four other carbon atoms so strong rigid lattice

3) Doesn’t conduct electricity : no free delocalised electrons to carry the current through the structure.

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13
Q

-describe the property of Graphite

A

1)each carbon atom is covalently bonded to 3 other carbon atoms

2)High Melting Point+ BP : strong covalent bonds

3)Soft and Slippery: weak intermolecular forces ( Van Der Waals ) so can slide past each other.

4)Good conductor of electricity: free delocalised electrons to carry current through the structure.

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14
Q

-what molecules are example of simple molecular

A

-ammonia, iodine , water, methane

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15
Q

-describe the properties of simple molecular

A
  • held together by covalent bonds with weak intermolecular forces( Van der waals )
    -Low MP and BP ; Van Der Waals forces are weak so not a lot of energy is required to overcome this bond
  • Poor conductors : structure contains no charge
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16
Q

-define metallic bonding

A

-consists of a lattice of positive metal ion attracted to a sea of delocalised electrons

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17
Q

-describe the properties of metallic bonding

A
  • good conductors of electricity: ‘sea’ of delocalised electron is able to move and carry a flow of charge
  • malleable : layers of positive ions are able to slide past over each other
    delocalised electrons prevent fragmentation as they can move around the the lattice

-high MP: electrostatic forces of attraction between the positive ions and delocalised electrons are very strong and so lots of energy required to overcome this bond

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18
Q

-what are the factors that affect the strength of metallic bonding

A

-numbers of delocalised electrons: more delocalised electron the more stronger the bond because it has higher charge density

-charge of the ion- the greater the charge, the stronger the attraction between the sea of delocalised electrons and the positively charged ion

-size of the ion: the smaller the ion the stronger the bond

19
Q

-define polar covalent bond

A

-forms when the elements in the bond have different electronegativities.

-has an unequal distribution of
electrons in the bond and produces a charge separation, (dipole) δ+ δ- ends.

20
Q

-describe shape linear

A

2 bonding pairs
0 lone pairs
180 degrees

21
Q
  • describe shape trigonal planar
A

3 bonding pairs
0 lone pairs
120 degrees

22
Q

-describe shape trigonal bipyramidal

A

5 bond pairs
0 lone pairs
90 degrees + 120

23
Q

-why does a molecule adopt a particular shape?

A

the electrons position themselves as far away as possible to minimise repulsion

24
Q

-describe shape trigonal pyrimadal

A

3 bond pairs
1 lone pairs
107.5 degrees

25
Q

-describe shape bent/ non linear

A

2 bond pairs
2 lone pairs
104.5

26
Q

-define electronegativity

A

relative tendancy of an atom to attract a pair of electrons in a covalent bond

27
Q
  • describe shape octahedral
A

6 bond pairs
0 lone pairs
90 degrees

28
Q

-describe shape tetrahedral

A

4 bonding pairs
0 lone pairs
109.5 degrees

29
Q

-describe shape Tshape

A

3 bond pairs
3 lone pairs
90 and 180 degrees

30
Q

-describe shape seesaw shape

A

four bonding pairs
one lone pair
90 and 120 degrees

31
Q

-what is electronegativity measured on

A

-pauling scale

32
Q

-what are the factors that affect electronegativity

A

-number of protons
-distance from the nucleus
-shielding

33
Q

-how does atomic radius affect electronegativity

A

-Electrons closer to the nucleus are more strongly attracted towards its positive nucleus

-electrons further away from the nucleus are less strongly attracted towards the nucleus

-therefore, an increased atomic radius results in a decreased electronegativity

34
Q

-why does electronegativity decrease down a group

A

Increased Atomic Radius: As you move down a group, the number of electron shells increases
-less attraction between the bonding electron and nucleus

Shielding Effect:more electron shells present, increases shielding: less attraction between nucleus and bonding electrons

-nuclear charge: even though the number of protons increase, increased distance and shielding outweighs this increase in charge

35
Q

-how does shielding affect electronegativity

A

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
E.g
-Sodium (Period 3, Group 1) has higher electronegativity than caesium (Period 6, Group 1) as it has fewer shells and therefore the outer electrons experience less shielding than in caesium

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

36
Q

-why does electronegativity increase across a period?

A

-nuclear charge increases- more protons so more attraction

-atomic size decreases: nucleus attracts the bonding electrons more strongly

37
Q

-how does the number of protons/ nuclear charge affect electronegativity?

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

38
Q

-Van der waals

A

(induced dipole dipole ) forces exist between atoms and molecules

39
Q

explain why some molecules with polar bonds do not have a permanent dipole?

A

contain polar bonds but it is symmetrical so the polar bonds cancel each other out

E.G : CCl4

40
Q

how are van der waals formed

A

present in all molecules

The distribution of the electron cloud changes constantly
* This produces a temporary dipole that induces temporary dipoles in surrounding molecules
* The positive part of the dipole attracts the negative side of each surrounding dipoles
* The negative part of the dipole attracts the positive side of each surrounding dipoles

41
Q

How does increasing the size of a molecule affect the strength of van der Waals forces?

A
  • It increases the strength of the forces
42
Q

Why does increasing the size of a molecule increase the strength of van der Waals forces?

A
  • Larger molecules have more electrons
  • The size of the temporary dipole increases and creates a stronger attraction
43
Q

What do permanent dipole-dipole forces act between?

A
  • Polar molecules
44
Q

How are permanent dipole-dipole forces formed?

A
  • Each polar molecule has a partial positive charge and partial negative charge
  • The partial positive charge of one molecule attracts the partial negative charge on the other
  • The partial negative charge of one molecule attracts the partial negative charge of the other
  • The molecules rotate and the attracted partial charges align

Chemistry : Physical Chemistry AS (10 decks)

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