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Chemistry S3 > 3 chem bond incomplete* > Flashcards

3 chem bond incomplete* Flashcards

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1
Q

structure and bonding of an ionic compound

A

Ionic compounds in the solid state have a giant ionic lattice structure held together by *strong electrostatic forces of attraction** between oppositely charged ions in an ionic compound.

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1
Q

what comes together to form ionic bonds and how to tell which is the anion and cation

A

Electrons are transferred from metal atoms to non-metal atoms to form ions that come together in a solid ionic compound.
● The atom that transferred the electron(s) gets positive charge(s) and becomes a cation.
● The atom that gained the electron(s) get negative charge(s) and becomes an anion.

hence metals become cations and non-metals become anions

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2
Q

Sodium forms the oxide Na2O. Explain why the oxide of sodium has this formula.

A

Each Na atom has one valence electron. To attain noble gas electronic configuration, each Na atom loses its valence electron to form a Na+ cation.

Each O atom has six valence electrons. To attain noble gas electronic configuration, each O atom gains two electrons from two Na atoms to form an O2– anion.

In the resultant giant ionic lattice formed, the ratio of Na+ : O2– is 2:1 such that the lattice is electrically neutral overall. Therefore, the formula is Na2O.

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3
Q

physical properties of ionic compounds

A

▪ High melting and boiling points
▪ Soluble in water
▪ Conducts electricity in molten and aqueous state
▪ Hard, rigid and brittle

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4
Q

Why do ionic compounds have high melting and boiling points?

A

Ionic compounds have a giant ionic lattice. In melting, large amount of energy is supplied to overcome the strong electrostatic forces of attraction between the oppositely charged ions to break the giant ionic lattice to form free ions.

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5
Q

how is the strength of ionic bonding determined by the charge of the ions, and their radii.

A

The higher the charge and the smaller the radii of the ions(ionic radii), the stronger the ionic bonding.

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6
Q

Between NaCl and MgO, which compound would be expected to have the higher melting point? Explain.

A

▪ Both NaCl and MgO have a giant ionic lattice.
▪ However, Mg2+ and O2─ have higher charge and smaller ionic radius than Na+ and Cl- (higher charge means more electromagnetic attraction = closer to nucleus)
▪ Hence the strength of the ionic bonds in MgO is larger than that of NaCl.
▪ More energy is needed to overcome the stronger electrostatic forces of attraction between the oppositely charged ions in MgO.
▪ MgO will have a higher melting point.

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7
Q

What happens when an ionic salt dissolves in water?

A

When an ionic salt dissolves, each ion on the crystal’s surface attracts the oppositely charged poles of the polar water(solvent) molecules and the ions become hydrated (or solvated if solvent is another polar solvent). This hydration process releases energy. The ionic crystal structure breaks down and the solid dissolves.

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8
Q

Why is it that ionic compounds do not conduct electricity in the solid state but conducts in the molten and in aqueous solutions?

A

Ionic compounds do not conduct electricity in the solid state as the ions are fixed in positions and are not free to move. In molten and in aqueous solutions, the ions are mobile. The ions are **free to move and flow towards the oppositely charged electrodes and act as mobile charge carriers, carrying a current.

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9
Q

Why are ionic solids Hard, rigid and yet brittle

A

In an ionic lattice, oppositely charged ions are held in fixed positions throughout the crystal lattice by strong ionic bonding. Moving the ions out of position requires large amounts of energy to break these bonds. Ionic lattices are therefore quite hard and rigid.
However, if enough pressure (e.g. by cutting or knocking) is applied, ions of like charge are brought next to each other. Repulsion between ions of like charges will cause the lattice to shatter apart. Ionic lattices are therefore said to be brittle.

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10
Q

structure and bonding of a metallic compound

A

Metals in the solid state have a giant metallic lattice structure, held together by metallic bonding. Metallic bonding is the electrostatic force of attraction between a lattice of metal cations and the sea of delocalised electrons in the metallic lattice.

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11
Q

what affects Strength of metallic bond s

A

Number of delocalised valence electrons in the structure.
–>The larger the number of valence electrons contributed per atom, the greater the number of delocalised electrons and the stronger the metallic bonding.

Charge and radius of the metal cation.
–>The higher the charge and the smaller the radius of the metal cation, the greater electrostatic force of attraction between the cations and the sea of delocalised electrons.

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12
Q

Physical properties of metals

A

▪ High melting and boiling points
▪ Conducts electricity
▪ Good conductor of heat
▪ Malleable and ductile
▪ Shiny surface
▪ Hard

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13
Q

y doe metals have a High melting and boiling points

A

_Most metals are solids with moderately high melting points and much higher boiling points. Metals have a giant metallic lattice structure which consists of a lattice of metal cations with a sea of delocalised electrons and the electrostatic forces of attraction between the cations and the delocalised electrons constitute the metallic bond. _
metallic bonding spam

Melting and boiling points of metals are relatively high because a large amount of energy is required to overcome the relatively strong electrostatic forces of attraction between the cations and delocalised
electrons.

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14
Q

Why are metals good electrical conductors?

A

Metals are good electrical conductors because the delocalised electrons which act as mobile charge carriers are free to move about in the solid structure.
When a piece of metal is attached to a battery, electrons flow from the negative terminal into the metal and replace electrons flowing from the metal into the positive terminal.

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15
Q

malleable meaning

A

can be beaten into a sheet

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16
Q

ductile meaning

A

can be pulled into wires

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17
Q

Reasons for malleability and ductility:

A

▪ In the giant metallic structure, the metal ions are held in an orderly array.
▪ When an external force (e.g. hammering) is applied to a piece of metal, the layers of metal ions in the solid structure slide past each other and end up in new positions.
▪ The overall shape changes (the metal deforms) but the metal does not break because the sea of delocalized electrons prevents repulsions among the cations as they move past one another.
▪ The metallic bond strength remains the same.( the metallic bonds are able to maintain their integrity, enabling the metal to undergo plastic deformation without fracturing. )
▪ Metallic bonding is termed non-directional bonding. (Because metallic bonds are not limited to specific orientations, metals can be easily deformed in any direction without weakening the overall structure.)

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18
Q

what are alloys

A

Alloys are mixtures of metals involving incorporation of small quantities of other element(s) into the pure
metal.

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19
Q

why are alloys hard

A

Alloying makes metals harder. This is because atoms of the added metal have a different size. This will disrupt the orderly arrangement of the main metal atoms that can no longer slide over each other easily when a force is applied.

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20
Q

what are covalent bonds

A

Covalent bonding is the electrostatic forces of attraction between the positively charged nucleus of both the bonded atoms and their shared pair of electrons.

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21
Q

Formation of Covalent Bonds

A

The covalent bond is formed by an overlap of valence atomic orbitals. The resultant electron cloud is called a molecular orbital.

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22
Q

how are sigma bonds formed

A

Sigma bonds are formed by the ‘head-on’ overlap of two atomic orbitals.

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23
Q

how many sigma bonds can tajke place between 2 atoms

A

There can only be one sigma bond between two atoms as there is no other way for another head-on overlap of the atoms’ orbitals to take place.

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24
Q

why is the only s orbital that can bond be from h

A

only valence electrons can bond and only h has the s orbital as the valence atomic orbital

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25
Q

p atomic orbital look like

A

dumbell

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26
Q

what are pi ponds

A

π bonds are formed by the ‘side-on’ overlap of two atomic orbitals.

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27
Q

is a pi bond or a sigma bond stronger

A

Side-on overlap is less efficient than head-on overlap as the overlap is poorer; hence a π bond is weaker than a σ bond and is more easily broken.

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28
Q

what is the only element that has to draw a circle instead of a dumbell when chemically bonding

A

hydrogen, as it is the only element with an s orbital as a valence shell

29
Q

which bond must exist first so the other can
sigma or pi

A

For a π bond to exist, a σ bond must first be present.

30
Q

how many pi bonds can be between 2 atoms

A

There may be one or two π bonds between two atoms depending on the number of atomic orbitals involved in the bonding. There can be a maximum of only 2 π bonds between two atoms.
(cuz got 3 p orbis, x y z, and x go to sigma bond so got 2 left that can become pi)

31
Q

what is the bond order

A

The number of covalent bonds formed between 2 atoms is called the bond order.
A bond order of 3 indicates a triple bond.
Multiple bonds are stronger and shorter than single bonds.

32
Q

whats the max number of covalent bonds 2 atoms can have

A

Depending on the type and number of atomic orbitals involved, it is possible to have more than 1 covalent bond formed between 2 atoms. the max is 4 as there can be atoms with 4 valence e

33
Q

names of 1 bond 2 bond and 3 bond

A

single bond, double bond and triple bond

34
Q

what is a bond pair

A

In a covalent compound, an electron pair that is shared between the bonded atoms is called a bond pair.

35
Q

what is lewis structure drawing method

A

use line to represent one pair of bond pairs (a coavlent bond)
still draw the other non bonded electrons on the atom

36
Q

What is electronegativity

A

It’s how much a molecule wants the electrons in a covalent bond

37
Q

general physcial properties of simple molecular substances

A

low melting point and low boiling point
cannot conduct electricity
insoluble in water
soluble in organic solvents

38
Q

Y do simple molecular substances have the term simple in front

A

“simple molecular substances” refers to substances made up of molecules containing only a few atoms bonded together

its also to contrast to other names like “ giant covalent structures”

39
Q

what is the bond length

A

The bond length of a covalent bond is the distance between the nuclei of the 2 bonded atoms when their atomic orbitals overlap to form molecular orbitals.
The bond length is essentially the sum of the covalent radii of the two atoms involved in the bond

40
Q

whats bond energy

A

The strength of a covalent bond is measured by the bond energy of that covalent bond. The bond energy of a covalent bond is the average energy required to break a particular bond in one mole of gaseous substance.

41
Q

relationship between size, length and strength of bonding in molecules

A

The smaller the atom, the shorter and the stronger the bond, and the larger the bond energy.

42
Q

The bond energy of the C─C single bond is 350 kJ mol–1 while the bond energy of a C=C double bond is 610 kJ mol–1. Why is the bond energy of the double bond less than double that of the single bond?

A
  • A C─C single bond is made up of a σ bond while a C=C double bond is made up of a σ and a π bond.
  • A π bond is weaker than a σ bond (due to lower degree of orbital overlap) and thus the energy of a C=C double bond is not twice that of a C─C single bond.
43
Q

what is bond polarity

A

When two atom have the same electronegativity, the bonding electrons are shared equally and the covalent formed between the two atoms is non-polar .
When two atom have different electronegativity, there is unequal sharing of bonding electrons and the covalent formed between the two atoms is polar .

44
Q

What’s partial positive/negative charge

A

The more eN atom has the greater tendency to attract the bonding electrons to itself –> partial negative charge (labeled δ–) due the build up of electron density around it.
The other less eN atom in the bond then acquires a partial positive charge (labeled δ+).

45
Q

how to calculate electronegativity

A

Electronegativity increases across a Period and decreases down a Group.
- bottom left is Fr least electronegative, top right is F most

46
Q
A
47
Q

whats a dipole moment

A

ONLY IN POLAR BONDS. +–> dipole moment arrow
# arrow points from the less eN atom to the more eN atom.
The greater the difference in electronegativity, the greater the bond polarity and the greater its dipole moment.

48
Q

is a polar bond or non-polar bond stronger

A

For a covalent bond of the same bond order and similar bond length, a polar bond is stronger than a non-polar bond.

49
Q

whats a non polar molecule

A

DOESNT MEAN theres NO polar bonds, its js that the polar bonds CANCEL out each other by going in opposing directions
no overall dipole / overall dipole moment is zero.

50
Q

what are intramolecular forces

A

covalent, ionic and metallic bonds
INSIDE THE MOLECULE

51
Q

what are intermolecular forces

A

OUTSIDE THE MOLECULESS, attactign them tgt
The attraction that exist between molecules in simple molecular elements or compounds are collectively referred to as intermolecular forces.
intermolecular forces are electrostatic forces of attraction which arise from the interactions of the net dipole of simple covalent molecules.

52
Q

what are the 3 intermolecular forces

A

(a) Dispersion forces (London forces or 3instantaneous dipole-induced dipole interaction)
(b) Permanent dipole-permanent dipole interactions
(c) Hydrogen bonds

53
Q

what are the 2 van der Waals’ forces

A

(a) Dispersion forces (London forces or instantaneous dipole-induced dipole interaction)
(b) Permanent dipole-permanent dipole interactions

54
Q

what are rewquirements for hydrogen bonds

A

● One molecule must contain a H atom directly bonded to a highly electronegative atom (nitrogen, oxygen or fluorine).
● The other molecule must contain an atom with a lone pair of electron on the highly electronegative atom (nitrogen, oxygen or fluorine).

55
Q

how do hydrogen bonds work

A

N-H O-H F-H bonds r super polar so all the electron density is withdrawn from H and the H has a poorly shielded proton aka significant δ+ charge.
it can experience a particularly strong attraction from a lone pair of electrons on an adjacent molecule,

56
Q

how to draw hydrogen bonds

A
  • rmbr to put the δ+- charges
  • draw the lone pairs for the N O F
  • hydrogen bonding is labbleed with like the zebra crossing |||||||||||||||| that
57
Q

whats the strongest intermolecular force

A

hydrogen bonding

58
Q

Factors affecting the strength of hydrogen bonding

A
  • Bond polarity, more polar = stronger
  • extensiveness, Extent means how many hydrogen bonds can form between molecules. eg. water can form 4 per molecule, 2 pairs of lone on its oxygen + 2 hydrogen
59
Q

Main factor affecting the strength of dispersion forces

A
  • The number of electrons per molecule. The greater the number of electrons in the molecules or atoms, the greater the electron cloud size, the greater the polarizability of the electron clouds, hence the stronger the dispersion forces.
  • shape, bigger surface area allow for more points of contact (more polarisable), resulting in stronger dispersion forces.
59
Q

effects of hydrogen bonding

A

stronger hydrogen bonds =
- Relatively high melting and boiling points
- Water has a high surface tension, high surface tension. hbonds hold it tgt
- Structure and density of ice, Each water molecule is thus surrounded by four water molecules in a roughly tetrahedral
arrangement.

59
Q

what are dispersion forces

A

Non-polar molecules are attracted to one another by instantaneous dipole – induced dipole interactions, also known as dispersion forces. However, dispersion forces exist in all molecules and can contribute significantly to the total intermolecular forces between polar molecules which have a large non-polar group in their structure.

59
Q

whats an instantaneous dipole

A

when a non polar molecule has an uneven distribution in the elctron cloud. like in its zone te elctrons js float around right but sometimes they both float to one side, then the neighbouring molecule will get an induced dipole cuz its own electroin gets repelled, so an instantaneous dipole – induced dipole interactions,
weak and very short-lived

60
Q

Factors affecting boiling points of non-polar molecules

A

The stronger the dispersion forces, the higher the boiling points.
NOTEEEE: During melting or boiling, only the weak intermolecular forces are broken/overcome, the strong covalent bonds are not broken.

60
Q

y do SIMPLE COVALENT MOLECULES have low Melting and boiling points

A

This is because less energy is needed to overcome the weaker intermolecular attractions between the discrete molecules

60
Q

y are SIMPLE COVALENT MOLECULES unable to conduct electricity

A

it has no no mobile charge carriers.

61
Q

Compound + Boiling point/°C
Methane(CH4):−161 Ethane(C2H6):−89 Propane(C3H8):−42
All three compounds have _________structure.
As the number of carbon atom increases from CH4 to C3H8, number of __________________ increases. Hence the size of the electron cloud _____________.
Strength of _____________ between the molecules increases which requires _______________ to overcome. Hence boiling point increases from CH4 to C3H6.

A

simple molecular
electrons
increases
dispersion forces
more energy

62
Q

attraction exists between noble gas atoms?

A

dispersion forces

63
Q

noble gas 1 2 3, the boiling points r increasing
Explain the difference in boiling points between these noble gases.

A

As the number of electrons per atom increases from 1 to 2 to 3, the size of the electron cloud of the atoms increases. Hence dispersion forces between the atoms becomes stronger and requires more energy to overcome.

63
Q
A
64
Q

ionic vs covalent

A

ionic:
-ionic bonding
-giant ionic lattice structure,
-stronger electrostatic forces of attraction between the oppositely charged ? n ? ions
covalent:
-covalent bonding
-simple molecular structure.
-weaker dispersion forces between ? molecules.

Chemistry S3 (10 decks)

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