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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
Q

How are electron density maps produced?

A

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

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25
Q

What do electron density maps tell us?

A

The arrangement of atoms or ions in the crystal

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26
Q

What does the arrangement of ions in an ionic lattice depend on?

A

The relative sizes of the different ions present

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27
Q

The coordination number of the ions in NaCl is six. What does this mean ?

A

Each ion six nearest neighbours

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28
Q

Why is the structure of carsium chloride different to NaCl?

A

The caesium ion is larger than the sodium iontherefore more chloride ions can fit around it making it have a larger coordination number

29
Q

Why do ionic compounds have very high melting and boiling temperatures?

A

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
Q

Why don’t all elements form ionic compounds?

A

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
Q

Where do you find elements that form covalent bonds in the periodic table?

A

Middle as the loss or gain of three or four electrons would require a great deal of energy

32
Q

What happens to the electron density as two hydrogen atoms approach one another?

A

Electron density shifts as the electrons are attracted to the nuclei the most probable place to find the electrons is between the nuclei

33
Q

What is a covalent bond in terms of electron density and nuclei?

A

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
Q

What is bond length ?

A

The distance between the atoms when the attractive and repulsive forces are balanced

35
Q

What is bond enthalpy?

A

The energy released as the two atoms come together to form the bond

36
Q

Between which types of elements do covalent compounds form ?

A

Nonmetals

37
Q

What is the double bond ?

A

When two pairs of electrons are shared

38
Q

What do dot and cross diagrams tell you about the shapes of molecules?

A

How the atoms share electrons

39
Q

What is a dative covalent bond?

A

When both of the electrons that make up a covalent bond come from the same atom

40
Q

What is a common feature of compounds that readily form dative covalent bonds?

A

Unshared electron pairs

41
Q

True or false dative bonds are longer than regular covalent bonds

A

False

42
Q

True or false dative bonds are the same strength as the equivalent covalent bond

A

True

43
Q

What do most covalent compounds exist as?

A

Discrete molecules

44
Q

Give two properties of atomic crystals

What do these properties tell us about the strength of covalent bonds?

A

Very hard
High melting temperatures
Very strong

45
Q

Why do covalent molecules have a fixed shape?

A

Each area of electronegativity repels the others

46
Q

Does the shape of covalent molecules change when they change state ?

A

No, providing the evidence for the nature of the covalent bonds

47
Q

What two properties does the shape of the covalent molecule determined about the compound?

A

Look up

48
Q

Why is BeCl(2) linear?

A

To minimise the repulsing between the pairs of electrons around beryllium they must be on opposite sides of the beryllium

49
Q

Why isn’t water linear?

A

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
Q

What is the definition for covalent bonding ?

A

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
Q

When do covalent bonds form ?

A

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
Q

Why do you covalent bonds form?

A

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
Q

What properties do simple molecules have?

A

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
Q

What are the properties of giant covalent molecules?

A

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
Q

Why does graphite conduct electricity?

A

Each atom only has 3 bonds

Allowing last electron to move freely between layers so conducts electricity

56
Q

What are van dee Waals’ forces?

A

Intermolecular forces

57
Q

What is the theory of dative bonding?

A

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
Q

What is meant by a Lewis base?

A

A lone pair donor

59
Q

What is meant by Lewis acid ?

A

A lone pair acceptor

60
Q

How is a dative bond represented in a dot and cross diagram?

A

By an arrow

61
Q

Describe metallic bonding

A

The strong attack attraction between positive metal ions and the Sea of delocalised electrons that surrounds them

62
Q

Explain why metals conduct electricity

A

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
Q

Why do metals have a high thermoconductivity?

A

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
Q

Why do metals have high melting and boiling temperatures?

A

Very strong force of attraction between positive atoms are negative delocalised electrons which requires lots of energy to separate

65
Q

Explain why metals malleable and ductile

A

The positive metal nuclei can move within the sea of electrons and wherever they move they are still surrounded by negative electrons

66
Q

What can’t the model of positive nuclei a sea of delocalised electrons explain?

A

Wide range of metal melting temperatures

67
Q

What happens during the formation of metallic bonds ?

A

Outershell electrons become delocalised

68
Q

What does the strength of the metallic bonds depend on?

A

Number of electrons donated

The size of the metal atom/ion

69
Q

How does the density of the electron cloud affect the melting point of metals ?

A

More delocalised electron results in a higher electron density making melting and boiling points higher

70
Q

How does the ionic size/charge the fact Dave melting point of a metallic bond?

A

Smaller ions require more energy to melt

Chemistry Unit 1 (7 decks)

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