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Forces Of Atrraction Flashcards

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1
Q

Why are atoms not generally found in their free states

A

Because they have lower potential energies when in combination with other rather than in isoloation

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2
Q

What is a chemical Bond formed

A

When the two atom’s attractive forces exceed the repulsive forces

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3
Q

What are intramolecular forces of attraction

A

attraction that act within particles
comprised of different atoms to hold the atoms tightly together

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4
Q

What is intermolecular forces of attraction

A

forces of attraction act between
particles to keep them loosely near to each other as solids,
liquids, or gases

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5
Q

What are types of Intramolecular forces

A

Ionic, Covalent (Simple molecular, Giant molecular), Metallic,

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6
Q

What are types of intermolecular forces

A

Van der Waals, Hydrogen Bonding, Permanent dipole

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7
Q

What physical properties of matter with ionic bonding (solid, m.p/b.p, electrical conductivity, solubility)

A

Solid at room temperature
high melting and boiling point
conductors in molten or aqueous state
soluble in polar solvent

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8
Q

What physical properties of matter with Giant covalent bonding (solid, m.p/b.p, electrical conductivity, solubility)

A

Solid at room temperature
high melting and boiling point
non conductors except graphite
insoluble

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9
Q

What physical properties of matter with simple covalent bonding (solid, m.p/b.p, electrical conductivity, solubility)

A

Solid liquid or gas at room temperature
low melting and boiling point
non-conductors
usually soluble in polar or non polar solvent

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10
Q

What physical properties of matter with metallic bonding (solid, m.p/b.p, electrical conductivity, solubility)

A

Solid at room temperature
high boiling and melting points
good electrical conductivity
Dissolves in other metals to form alloys

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11
Q

The strength of metallic bonding increases with

A

increasing positive charge on the ions

decreasing size of the metal ions

increasing number of delocalised electrons

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12
Q

Explain Metallic bonding

A

In metallic bonding, metal atoms come together and donate
their valence electrons to form a “sea” of mobile electrons
that is communally shared among the resulting cations.

The cations organise themselves into a lattice structure (aka a
crystal lattice – an ordered 3D arrangement containing
repeating basic units) and the delocalised electrons move
freely throughout the lattice.

*The positive charges are held together by their strong electrostatic
attraction to the delocalised electrons. This strong electrostatic
attraction acts in all directions.

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13
Q

What are the types of structure for metals

A

Hexagonal close packed (hcp)
= 2 repeating layer arrangements ABABAB

face-centred cubic close packing (ccp or fcc)
= 3 repeating arrangements ABCABC…

Simple cubic (sc)

Body-centered cubic (bcc)

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14
Q

What type of metals have HCP structure

A

elemental metals including Be, Mg, Ti, Zr,

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15
Q

What type of metals have ccp (fcc) structure

A

Al, Ni, Cu, Ag

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16
Q

What type of metals have SC structure

A

Polonium (Po)

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17
Q

What type of metals have bcc structure

A

Fe, V, Cr, Mo, W

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18
Q

How to measure metallic radius

A

The distance between the nuclei of two metal atoms in a solid metallic lattice can be measured accurately by X-ray diffraction studies of
metal crystals. Half of this distance is called the metallic radius.

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19
Q

The density (mass per unit volume) of a metal depends on

A

– the atomic mass
– the size of the atom (i.e. the metallic radius)
– the type of packing

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20
Q

Properties of metals are

A

high density solids with high melting and boiling points. - It
takes a lot of energy to break the large number of strong forces of
attraction between the ions and the delocalised electrons

Good conductors of heat and of electricity - the delocalised
electrons are free to move when a charge is applied

Shiny- reflecting light of all wavelengths

Malleable and ductile - when stress is applied the cations can slide
over each other and so metals can be beaten into shapes and
drawn into wires

Insoluble in water (although some react with water) - the force of
attraction between the ions and the delocalised electrons is too
strong to be broken up and replaced by bonds to water molecules. Some metals react with water because they lose electrons

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21
Q

Why do metals tend to give up electrons more easily than non-metals

A

lower ionisation energies

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22
Q

Why do non metals tend to
accept electrons more easily than metals

A

have more exothermic electron affinities

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23
Q

What affects the rate at which ions become stable

A

the total of the ionisation energies or electron affinities involved to get them to noble gas configuration

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24
Q

What is ionic bonding

A

the transfer of electrons from a metal atom to a non-metal atom until the outer shells of the resulting ions have configurations similar to those of a noble gas

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25
Q

Why are ionic compounds said to have 3d crystal structures

A

each positive ion is surrounded by several negative ions, and each negative ion is surrounded by several positive ions in repeating units

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26
Q

Whats a formula ratio

A

the ratio in which ions are present in the crystal structure

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27
Q

How do we calculate ionic radius

A

from interionic distance measurements from the crystal structures of ionic compounds using the O2- ion as a reference.

the interionic distance is the distance between the centres of the cation and the anion. It is equal to the sum of the ionic radius of the cation (the cationic radius) and the ionic radius of the anion (the
anionic radius)

the radii of any cation can be determined by:
rcation = interionic distance – rO2-

the unknown radius of
another anion bonded to that cation can be found from:
ranion = interionic distance - rcation

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28
Q

What is lattice enthalpy

A

The lattice enthalpy is defined as the energy absorbed when 1 mole
of an ionic compound is formed from its gaseous ions. It always has
a negative value since this process is exothermic

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29
Q

What is lattice enthalpy used for

A

Lattice energies can be used as a measure of the relative stabilities
of ionic substances. The more exothermic the lattice energy (i.e. the more heat given off when the ionic bond is formed), the more
stable the resulting ionic compound

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30
Q

Lattice energy becomes more exothermic with:

A

– increasing ionic charge
– decreasing interionic distances or ionic radii since small distances
means that ions can more tightly pack together in the crystal.

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31
Q

Properties of Ionic Compounds

A

Strong, rigid, crystalline solids -owing to the regular arrangement of
ions held together by powerful electrostatic forces

High melting points, boiling points, heats of fusion, and heats of
vaporisation - it takes a lot of energy to separate the strongly-attracted ions

Hard and brittle - like charges repel each other when a force is applied

Conduct electricity when molten but NOT when solid because when
molten, ions have enough freedom to move, whereas movement is
restricted in the crystal lattice in the solid state

Soluble in water and other polar solvents. Individual ions need to be
stabilized by polar solvent molecules in order to effectively break apart the crystal lattice in solution

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32
Q

What is covalent bonding in terms of orbital

A

In terms of orbitals, a covalent bond forms when two atomic orbitals overlap.

The joined orbital is called a molecular orbital

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33
Q

What are sigma bonds

A

Sigma bonds ( bonds) or sigma molecular orbitals are formed by the overlap of atomic orbitals along a line drawn between the two nuclei

Sigma bonds give the maximum possible electron density
between two nuclei; therefore these are very strong bond

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34
Q

What is the electron density of sigma bond

A

symmetrical about a line joining the two nuclei

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35
Q

How are Pi bonds formed

A

Pi bonds ( bonds) or pi molecular orbitals are formed by the sideways overlap of two p atomic orbitals

A pi bond is weaker than a sigma bond; so compounds with pi bonds are reactive

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36
Q

How are electron density of Pi bonds distributed

A

The electron density of the bond formed is distributed above
and below the plane of the sigma bond

The electron density is zero along the axis of the nuclei

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37
Q

What is a double bond formed by

A

one sigma bond and one pi bond

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38
Q

What is a triple bond formed by

A

one sigma bond and two pi bonds

39
Q

What is bond order

A

the number of electron pairs being shared by two covalently bonded atoms

40
Q

What is bond length

A

the distance between the nuclei
of two atoms that are bonded together

Shorter bond lengths mean that the shared electrons are closer to the nuclei resulting in stronger bonds

41
Q

What is covalent radius

A

the distance between the nucleus of a covalently bonded atom and the
outermost shell.

42
Q

How do we calculate covalent radius

A

half the distance between the nuclei of two identical atoms bonded covalently via a single bond

43
Q

What is Electronegativity

A

Electronegativity is a measure of the ability of an atom to attract
an electron
It is derived by combining data on ionisation energies and electron
affinities for the atom

44
Q

How does electronegativity increase across the periodic table

A

Electronegativity increases from left to right across a period and
decreases down a group in the periodic table (i.e. it generally
follows the same trend as ionisation energy

Fluorine is the most electronegative element

45
Q

When is a bond said to be non-Polar

A

If two atoms forming a covalent bond have electronegativity values
that are the same in magnitude then the shared electrons will be
symmetrically distributed

46
Q

When is a bond said to be polar

A

If the two atoms forming the covalent bond have different
electronegativities, then the atom with the greater electronegativity
will pull the shared electrons closer to itself, causing the electrons to
be unsymmetrically distributed between the atoms

47
Q

Where is the partial negative and positive charge in a polar bond

A

the bond has a partial negative charge (–) at the more electronegative end, and a partial positive charge (+) at the less electronegative end

48
Q

Where does the dipole point to and from

A

from less electronegative atom to more electronegative atom

49
Q

What is bond enthalpy (bond dissociation enthalpy)

A

the energy required to disssociate 1 mole of molecules into atoms

50
Q

Factors affecting bond energy/bond strength include:

A

– Bond length – shorter bonds are stronger
– Bond order – multiple bonds are more difficult to break
– Types of bonds – sigma bonds are stronger than pi bonds
– Bond polarity – higher polarity leads to stronger bonds

51
Q

molecules that are unable to complete the octet of electrons
when they form covalent bonds are called

A

electron deficient

52
Q

Whats a coordinate bond or dative covalent bond

A

one atom provides both the electrons for the covalent bond

53
Q

For coordinate bonding to occur, we need:

A

– one atom with a lone pair of electrons
– a second atom with an unfilled orbital

54
Q

In the displayed formula, a coordinate bond can be indicated by

A

an arrow pointing away from the atom donating the electrons

55
Q

Molecules are said to have an expanded octet of electrons when…

A

number of electrons in their outer shell ismore than 8

example sulphur hexafluoride SF6

56
Q

What does VSEPR Theory stand for and what does it do

A

Valence Shell Electron Pair Repulsion

predicts the shapes and bond angles of simple covalent
molecules

57
Q

What does this theory suggest

A

the electron pairs around an atom
repel each other and arrange themselves in space so as to
minimize the repulsive forces

58
Q

What determines the shape of molecules

A

The shapes of molecules are determined primarily by the
numbers of bonding pairs and lone pairs of electrons around a particular atom called the central atom

59
Q

Interactions between electrons in two different bonds are called

A

bonding pair-bonding pair repulsions

60
Q

Interactions between lone pairs and bonding pairs are called

A

lone pair-bonding pair repulsions

61
Q

Interactions between two different lone pairs are

A

lone pair-lone pair repulsions

62
Q

The repulsion between electron pairs is inversely proportional to the ____ between them

A

distance

the shorter the distance the greater the repulsion

63
Q

the strength of repulsions increases in the order

A

Bond pair-bond pair < lone pair-bond pair < lone pair-lone pair

the strength of the repulsion depends on the proximity of
each pair to the central atom. Lone pairs are closer to the central atom
than bonding pairs since there is no other nearby positive nucleus pulling the electrons away

64
Q

2 groups attached without lone pairs

A

Two groups

Equal repulsion

Bond angles both 180

Shape is linear

65
Q

2 groups attached with lone pairs

A

Two groups

Two Lone pairs

Greatest repulsion between lone pairs

Less repulsion between bond pairs

Bond angles 104.5

Shape is non linear, V shaped, bent

66
Q

3 groups attached without lone pairs

A

Three bond pairs

Equal Repulsion

All angles 120

Shape is trigonal planar

67
Q

3 groups attached with lone pairs

A

On lone pair

Three bond pairs

Greater repulsion between lone and bonding pairs

Bond angles 107

Shape pyramidal

68
Q

4 groups attached without lone pairs

A

Four bond pairs

Equal Replusion

all bond angles 109.5

Shape is tretrahedral

69
Q

5 groups attached without lone pairs

A

Equal Repulsion

Five bond pairs

Three angles form the trigonal plane with bond angles for 120

Two remaining bond angles are 90

Shape: trigonal bipyramidal

70
Q

6 groups attached without lone pairs

A

Six bond pairs

Equal repulsion

Bond angles all 90

Shape octahedral

71
Q

What are the two questions that VSEPR theory does not address

A

How are the valence (or outer shell) electrons used to form
covalent bonds if the electrons are paired in various subshells?

How do we account for the shapes of the resulting molecules
using molecular orbital theory?

72
Q

What is Hybridization

A

Hybridization is the mixing of different atomic orbitals to form
new orbitals of equal energy

73
Q

What happens in sp3 hybridization

A

one s orbital and three p orbitals combine to form four equal in energy sp3 hybrid orbitals (sppp = sp3)

After hybridization, the four resulting sp3 hybrid orbitals are
oriented in a tetrahedral arrangement. When these four
orbitals form bonds to other atoms, the resulting bonds take
on a tetrahedral shape in the molecule. (109.5)

74
Q

What happens in sp2 hybridization

A

In sp2 hybridization, one s orbital and two p orbitals combine
to form three equal in energy sp2 hybrid orbitals (spp = sp2).
One p orbital remains unhybridised

After hybridization, the three resulting sp2 hybrid orbitals are
oriented in a trigonal planar arrangement. When these three
orbitalsform bonds to other atoms, the resulting bonds take
on a trigonal planar shape in the molecule. (120)

The remaining unhybridised p orbital is perpendicular to the
plane containing the sp2 hybrid orbitals and is used for 
bonding to create a double bond

75
Q

What happens in sp hybridization

A

In sp hybridization, one s orbital and one p orbitals combine
to form three equal in energy sp hybrid orbitals. Two p orbitalsremain unchanged.

After hybridization, the two resulting sp hybrid orbitals are
oriented in a linear arrangement. When these two orbitals
form bonds to other atoms, the resulting bonds take on a
linear shape in the molecule (180)

The remaining unhybridised p orbitals are perpendicular to
the plane containing the sp hybrid orbitals and are used for 
bonding to create a triple bond

76
Q

What a conjugated bond system

A

a system of alternating single and double bonds

Example A special type of sp2 hybridization occurs in the ring structure of benzene (C6H6)

77
Q

Covalently bonded substances can form the following crystal lattices:

A

– simple molecular lattices with weak intermolecular forces
between the separate molecules

– giant molecular structures with strong covalent bonds throughout
the entire lattice made up of different atoms

– giant atomic structures with individual atoms of one type
covalently bonded in a giant structure (e.g. diamond, graphite)

78
Q

Whats the differences between Simple Molecular Lattices and Giant Molecular Structures

A

In simple molecular lattices, weak intermolecular forces operate
between the separate molecules

In giant molecular structures, strong covalent bonds occur throughout the entire lattice

79
Q

What are the properties of Simple Molecular Substances

A

Since simple molecular solids have weak van der Waals forces
between the molecules; small amounts of heat will provide
sufficient energy to separate the molecules. So, molecular
substances have low melting and boiling points and are often
found as liquids or gases at room temperature.

Molecular solids have no free electrons and thus are poor
conductors of electricity

Molecular solids tend to be soft and brittle

Most molecular solids are only slightly soluble in water

80
Q

What are the properties of Giant Covalent Substances

A

Giant covalent structures have very high melting and boiling points
since it requires a lot of energy to break the strong covalent bonds
throughout

Giant molecular and atomic solids with no free electrons are poor
conductors of electricity (e.g. SiO2, diamond). Exception: Graphite is
a good conductor.

Because of the strong covalent bonds throughout, giant covalent
structures other than graphite are hard solids

Giant covalent structures are insoluble in water

81
Q

Why is graphite a conductor of electricty

A

Graphite contains unbonded or delocalised electrons that can move throughout the p orbitals above and below the plane.

82
Q

What are Intermolecular forces

A

the forces of attraction that operate between separate molecules holding them close to one another

83
Q

Which are stronger intermolecular or intramolecular

A

Intermolecular forces are weak when compared with the
intramolecular forces responsible for ionic, covalent, and metallic
bonding

84
Q

What are the three intermolecular forces

A

permanent dipole-dipole attractions, van der Waals forces, and hydrogen bonding

85
Q

The comparative strength of intermolecular forces varies as
follows:

A

van der Waals < permanent dipole – dipole < hydrogen bonding

86
Q

What is a non polar molecule

A

In a non-polar molecule, the electrons are on average evenly
distributed throughout the molecule

87
Q

Whats a polar molecule

A

In a polar molecule, the electrons on average spend more time at one
end of the molecule than the other. One end of the molecule is slightly
negatively charged (-) while the other end is slightly positively
charged (+)

88
Q

What are Permanent dipole – dipole forces

A

exist between polar molecules.
They are weak attractive forces between the + of the dipole of one
molecule and the - of the dipole of a neighbouring molecule

89
Q

What is an instantaneous dipole

A

An instantaneous dipole is a temporary dipole that occurs when a molecule’s electrons are distributed unevenly around the nucleus, creating a dipole of uneven charges. This happens due to the constant motion of the electrons.

90
Q

van der Waals forces determine

A

the distances between atoms in
liquids and solids

91
Q

The strength of the van der Waals Forces increases with

A

the # of electrons and molar mass of the atom or molecule

the shape of the molecule:
elongated molecules are more easily polarised than compact,
symmetrical molecules and therefore have stronger van der Waals forces

the size of the molecule:
small molecules have weak van der Waals forces between them

Larger molecules have more points of contact between them and thus
more opportunities for induced dipoles to form leading to stronger vander Waals forces

the degree of branching:
linear molecules have stronger van der Waals forces between them
than branched molecules. Branching reduces the number of points of contact between the two molecules leading to weaker van der Waals forces overall

92
Q

What are van der Waals forces

A

temporary instantaneous induced
dipole – dipole forces that exist in all atoms and molecules
including noble gases

93
Q

What is a hydrogen bond

A

The force of attraction between an electron-deficient hydrogen
bonded to fluorine, oxygen or nitrogen and a lone pair on a
neighbouring fluorine, oxygen or nitrogen is called a hydrogen bond

94
Q
A

Chemistry (20 decks)

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