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Chemistry Unit 1 > 4 Bonding > Flashcards

4 Bonding Flashcards

0
Q

Is the formation of a lattice and exothermic or endothermic process ?

A

Exothermic

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1
Q

Is the formation of ions in the gaseous state from elements in their standard states and endothermic or exothermic process ?

A

Endothermic

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2
Q

What is the definition of the lattice energy of an ionic crystal ?

A

The copy of formation of one mole of an ionic compound from gaseous irons under standard conditions

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3
Q

What two things affect the lattice energy of a compound ?

A

The charge and size of ions

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4
Q

Explain why the lattice energy for smaller ions is larger (more negative) than for large ions

A

In smaller ions the attractive force of the positive nucleus holds the outer electrons more tightly because they are closer to the nucleus

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5
Q

Explain why the lattice Energy for magnesium fluoride is larger ( More negative) than sodium fluoride

A

Magnesium has a 2+ charge so attracts fluoride more strongly

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6
Q

What is coulombs law ?

A

Force of attraction= kq1q2/r(2)

Q1 q2 charge of ions
R= distance between ions

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7
Q

What does coulombs law tell you about the relationship between the charge on the irons and the attractive force between them ?

A

Increasing the charge on either ion increases the attractive force between two oppositely charged ions in the lattice

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8
Q

What does coulombs law tell us about the relationship between the size of the ions and the attractive force between them ?

A

Decreasing the size of one or both ions decreases the distance between them and so increases the attractive forces

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9
Q

What Borne-Haber cycles be used for?

A

Predictably relative stabilities of different compounds and why particular ionic compounds exist

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10
Q

What two methods can be used to calculate lattice energy ?

A

Born-Haber cycles

Coulomb energies

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11
Q

What leads to these distortion of ionic bonds?

A

The attraction of the positive cation for the outer electrons of the negative anion

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12
Q

He polarising power depends on charge density. What does charge density depend on ?

A

It’s ionic radius and charge

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13
Q

Are smaller cations less polarising or more polarising than larger ones ?

A

More

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14
Q

Are cations with a larger charge more or less polarising them cations with a small charge?

A

More

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15
Q

What does polarising power mean?

A

The ability of a cation to distort the anion

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16
Q

What does polarisability mean?

A

How easily anion is distorted by a cation

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17
Q

What does the polarisability depend on ?

A

It’s ionic radius which affects how tightly they are held

Larger anions are more easily polarised

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18
Q

Give three properties that can be different between the elements that make up a compound and the compound

A

Appearance and state
Reaction with water
Electrical conductivity

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19
Q

Why does NaCl have such different properties to sodium and chlorine ?

A

Sodium chloride is an ionic compound
Sodium is a lattice of metal ions
Chlorine is a gas made of small molecules

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20
Q

How and why does atomic radius changes you go across the period?

A

Reduces across a period because there are no new shells but more positive charge so electrons are attracted more

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21
Q

How and why does the atomic radius change you go down the group ?

A

Increases down a group because the outer electrons enter new energy levels whilst the charge is being shielded by other electrons

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22
Q

Will K or K+ be larger? Why?

A

The radius of the positive ion is smaller than the atom as electrons are more strongly attracted to the nucleus

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23
Q

Will Br or Br- be larger? Why?

A

Negative ions are largely because additional negative charge means that all electrons abound less tightly to the nucleus

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24
How are electron density maps produced?
Passing x-rays tomorrow perfect crystal which is scattered all diffracted by the electrons in the atoms or ions in the structure, producing a diffraction pattern
25
What do electron density maps tell us?
The arrangement of atoms or ions in the crystal
26
What does the arrangement of ions in an ionic lattice depend on?
The relative sizes of the different ions present
27
The coordination number of the ions in NaCl is six. What does this mean ?
Each ion six nearest neighbours
28
Why is the structure of carsium chloride different to NaCl?
The caesium ion is larger than the sodium iontherefore more chloride ions can fit around it making it have a larger coordination number
29
Why do ionic compounds have very high melting and boiling temperatures?
They contain large numbers of ionic bonds between the oppositely charged ions which need to be broken all we can to melt an ionic solid
30
Why don't all elements form ionic compounds?
The energy released in the formation of the lattice of ions would be insufficient to overcome the energy required to form the ions in the first place
31
Where do you find elements that form covalent bonds in the periodic table?
Middle as the loss or gain of three or four electrons would require a great deal of energy
32
What happens to the electron density as two hydrogen atoms approach one another?
Electron density shifts as the electrons are attracted to the nuclei the most probable place to find the electrons is between the nuclei
33
What is a covalent bond in terms of electron density and nuclei?
A balance between the attractive force pulling the nuclei together due to the electron density between the nuclei and the repulsive force of the two positively charged nuclei pushing each other apart
34
What is bond length ?
The distance between the atoms when the attractive and repulsive forces are balanced
35
What is bond enthalpy?
The energy released as the two atoms come together to form the bond
36
Between which types of elements do covalent compounds form ?
Nonmetals
37
What is the double bond ?
When two pairs of electrons are shared
38
What do dot and cross diagrams tell you about the shapes of molecules?
How the atoms share electrons
39
What is a dative covalent bond?
When both of the electrons that make up a covalent bond come from the same atom
40
What is a common feature of compounds that readily form dative covalent bonds?
Unshared electron pairs
41
True or false dative bonds are longer than regular covalent bonds
False
42
True or false dative bonds are the same strength as the equivalent covalent bond
True
43
What do most covalent compounds exist as?
Discrete molecules
44
Give two properties of atomic crystals | What do these properties tell us about the strength of covalent bonds?
Very hard High melting temperatures Very strong
45
Why do covalent molecules have a fixed shape?
Each area of electronegativity repels the others
46
Does the shape of covalent molecules change when they change state ?
No, providing the evidence for the nature of the covalent bonds
47
What two properties does the shape of the covalent molecule determined about the compound?
Look up
48
Why is BeCl(2) linear?
To minimise the repulsing between the pairs of electrons around beryllium they must be on opposite sides of the beryllium
49
Why isn't water linear?
It has four clouds of electronegativity (the lone pairs of electrons on the oxygen atom make the covalent bonds at an angle to each other)
50
What is the definition for covalent bonding ?
Consists of a shared pair of electrons with one electron being supplied from each electron Atoms are held together because new cure attracted to shared electrons
51
When do covalent bonds form ?
Between atoms of the same element Between atoms on opposite sides of the periodic table When one of the elements is in the middle of the table Head of the group elements with high ionisation energies
52
Why do you covalent bonds form?
To get the nearest noble gas electron configuration Something I achieved and octet as they haven't got enough electrons Others share only son Atoms of elements in the third period onwards can exceed that octet
53
What properties do simple molecules have?
DON'T CONDUCT because no mobile ions or electrons LOW BOILING POINT- weak intermolecular forces in between molecules MORE SOLUBLE IN ORGANIC COMPOUNDS than in water; some are hydrolysed
54
What are the properties of giant covalent molecules?
VERY HIGH MELTING POINT- large number of strong covalent bonds which need breaking DON'T CONDUCT ELECTRICITY (other than graphite)- no mobile ions or electrons
55
Why does graphite conduct electricity?
Each atom only has 3 bonds | Allowing last electron to move freely between layers so conducts electricity
56
What are van dee Waals' forces?
Intermolecular forces
57
What is the theory of dative bonding?
Dative covalent bonds differ only in formation Both electrons of the shared pair of provided by one species Donor species will have lone pairs in their outer shells Exceptors species will be short of their octet (electron deficient)
58
What is meant by a Lewis base?
A lone pair donor
59
What is meant by Lewis acid ?
A lone pair acceptor
60
How is a dative bond represented in a dot and cross diagram?
By an arrow
61
Describe metallic bonding
The strong attack attraction between positive metal ions and the Sea of delocalised electrons that surrounds them
62
Explain why metals conduct electricity
Delocalised electrons are free to move through the lattice under the influence of an electric field Electrons are repelled by the negative electrode and attracted to positive
63
Why do metals have a high thermoconductivity?
Delocalised electrons in areas of high temperature have Hugh kinetic energy so randomly move to the cooler areas of the metal, transferring energy to the other electrons in the process
64
Why do metals have high melting and boiling temperatures?
Very strong force of attraction between positive atoms are negative delocalised electrons which requires lots of energy to separate
65
Explain why metals malleable and ductile
The positive metal nuclei can move within the sea of electrons and wherever they move they are still surrounded by negative electrons
66
What can't the model of positive nuclei a sea of delocalised electrons explain?
Wide range of metal melting temperatures
67
What happens during the formation of metallic bonds ?
Outershell electrons become delocalised
68
What does the strength of the metallic bonds depend on?
Number of electrons donated | The size of the metal atom/ion
69
How does the density of the electron cloud affect the melting point of metals ?
More delocalised electron results in a higher electron density making melting and boiling points higher
70
How does the ionic size/charge the fact Dave melting point of a metallic bond?
Smaller ions require more energy to melt

Chemistry Unit 1 (7 decks)

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