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Chemistry > electron bonding and structure > Flashcards

electron bonding and structure Flashcards

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1
Q

What is first ionisation energy

A

The energy required to remove one electron from each atone from 1mole of gaseous atoms to form 1mole of gaseous +1 ions

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2
Q

Successive ionisation energies.

A

The measure of energy required to remove each electron in turn

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3
Q

Second ionisation energy?

A

Energy required to remove one electron from each atom in one mole of gaseous atoms to form 1mole of +2 ions

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4
Q

What is plasma?

A

The mixture between positive ions and negative electrons

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5
Q

What does a large increase in successive ionisation energies show?

A

That the second electron has been removed from a different shell, closer to the nucleus and with less shielding than inner electrons

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6
Q

What is the filling order for shells?

A

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d

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7
Q

What is a semi metal?

A

Displays properties of metal and non metal

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8
Q

What is periodicity?

A

The trend in properties repeated a cross a period

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9
Q

Describe the periodic trend in terms of electron configuration across the periodic table?

A

For each successive element there is one more electron in the outer shell

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10
Q

What is an example of a metalliod and why?

A

Silicon because it is shiny like a metal but it brittle and it conducts electricity but very poorly

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11
Q

What is the distinct change between group 4 and group 5 in both physical and molecular structure?

A

Giant structures to simple molecular structures

Strong forces to weak forces

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12
Q

Why is group 2 called the alkaline earth metals

A

All have hydroxides that are alkaline

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13
Q

What are group 2 general properties

A

High melting and boiling points
Low densities
FORM COLOURLESS COMPOUNDS

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14
Q

Why do first ionisation energies decrease down a group ?

Example?

A
  • offsetting the nuclear charge is a greater distance from the nucleus and more screeningThe only factor left is the extra distance between the outer electron and the nucleus in
  • sodium and lithium
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15
Q

In both Periods 2 and 3, although there are some fluctuations, the main trend is for first ionisation energy to increase across both periods. Explain why that is.

A

In Period 2, the outer electrons are all in 2-level orbitals, which are all at much the same sort of distance from the nucleus. They are all screened by the 1s electrons. As you go across the period, the number of protons in the nucleus increases, increasing the attraction on the 2-level electron which will be removed. In Period 3, the same thing is true, except that we are now talking about 3- level electrons, screened by 1s, 2s and 2p electrons. Because attractions increase, ionisation energy will also increase.

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16
Q

Using magnesium and beryllium as an example, explain why first ionisation energies in Period 3 are smaller than the corresponding ones in Period 2.

A

Magnesium’s electronic structure is 1s2 2s2 2p6 3s2 ; beryllium’s is 1s2 2s2.The magnesium’s first ionisation energy is less than beryllium’s because the electron being removed is in the 3-level rather than the 2-level, and so is further from the attraction of the nucleus.

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17
Q

Explain why the elements from scandium to copper have similar first ionisation energies.

A

The first electron to be lost from all of the d-block elements comes from the 4s orbital. The 3d orbitals lie slightly closer to the nucleus. As you add an extra proton to the nucleus, you also add an extra electron to the 3d orbitals. That more or less screens the effect of the extra proton.

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18
Q

What does a tetrahedral shape look like?

A

1 up

3 bottom

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19
Q

What six elements exist naturally as single unbounded atoms

A

The noble gases

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20
Q

What makes the noble gases so unreactive

A
  • very stable outer shell of 8 electrons or 2 electrons

- two in s orbital and two each in 3p orbitals

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21
Q

What is the octet rule?

A

Unpaired electrons pair up by sharing or transferring electrons to form chemical bond
Results in each atom having the same electron config of a noble gas
Tendency to acquire this is octet rule

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22
Q

What is a compound

A

A substance formed from 2 or more chemically bonded elements in a fixed ratio usually shown by a chemical formula

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23
Q

What happens in ionic bonding

A

Electrons transferred from metal to non metal to form oppositely charged ions that attract

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24
Q

What happens in covalent bonding

A

Electrons shared between 2 non metals

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25
Q

What is an ionic bond

A

The electrostatic attraction between oppositely charged ions

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26
Q

What is the structure of a metallic bond and describe it

A

A giant ionic lattice: each ion surrounded by negative electrons forming a giant ionic lattice when in solid state

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27
Q

Why do ionic compounds have high melting and boiling points

A

Solids at room temp
Large amount of energy needed to break STRONG electrostatic forces that that hold oppositely charged ions together in lattice

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28
Q

How does charge affect melting and boiling points of ionic compounds

A

Greater the charge the stronger the electrostatic force between the ions and a higher amount if energy needed to break up the lattice durin melting

29
Q

How does an ionic lattice dissolve in polar solvents such as water and use sodium chloride as an example

A

Polar water molecules break down the lattice by surrounding each ion to form a solution

  • water molecules attract Na+ and Cl- ions
  • lattice breaks down as being dissolved.
  • Na+ attracts - charges on O atoms of water mol
  • Cl- attracts + charges on h atoms of water mol
30
Q

What is a dative covalent bond

A

A shared pair of electrons which has been provided by one of the bonding atoms only

31
Q

What happens when an acid is added to water?

A

Water molecules form Oxoniumions

32
Q

What can break the Octet rule?

A

There may not be enough electrons to reach an octet
and
more than four electrons may pair up bonding

33
Q

How many covalent bonds of atoms of nonmetals in group 5 form

A

3 or 5

34
Q

Group six

A

2 4 or 6

35
Q

How many can group 7 form

A

1 3 5 or 7

36
Q

What is the bond angle of the trigonal planar

How many electron pairs around the central atom

A

120

3

37
Q

What is the bond angle in a tetrahedral

How many electron pairs around central atom

A

109.5

4

38
Q

What is the bond angle in octahedral and how many electron pairs around the central atom

A

90

6

39
Q

What is a water molecule and what is its bond angle

A

Non linear

104.5

40
Q

Why does a hydrogen atom have a 100% covalent bond

A

Each hydrogen atom has an equal share of the pair of electrons in the bond

41
Q

What is electronegativity

A

A measure of the attraction of a bonded atoms for a pair if electrons in a covalent bond

42
Q

What is a permanent dipole

A

A small charge difference across the bond that results from a difference in the electro negativity of the bonded atoms

43
Q

What does a polar covalent bond have

A

A permanent dipole

44
Q

What is apolar molecule

A

It has an overall dipole when you take into account any dipoles across the bond

45
Q

What happens to electronegativity across a periodic table

A

Electronegativity increases towards the top right of the periodic table.
fluorine has the most electronegative atoms

46
Q

What is an ionic bond in relation to electronegativity

A

When there is a large difference in electronegativity one atom will more effectively have captured both the bonding electrons

47
Q

What is a permanent dipole dipole force

A

A weak attractive force between permanent dipoles in neighbouring polar molecules

48
Q

What are Vanderwall forces

A

Attractive forces between induced dipole’s in neighbouring molecules

49
Q

How are the atoms in a molecule bonded together

A

By strong covalent bonds

50
Q

What are the 3 common types of intermolecular forces

A

Hydrogen bonds
Permanent dipole dipole forces
Vanderwall’s forces

51
Q

How are Vanderwall forces caused

A

By the movement of electrons in shells this
movement causes the distribution of charge within the electron shells To be unbalanced at any moment
there will be instantaneous dipole across the molecule

52
Q

What is a hydrogen bond

A

An electron deficient hydrogen atom on one molecule and a lone pair of electrons on a highly electronegative atom on a different molecule

53
Q

Why is ice less dense than water

A

Ice is an open lattice with hydrogen bonds holding the water molecules apart
when ice melts the rigid hydrogen bonds collapse allowing the water molecules to move closer together

54
Q

Why do most metals have high melting and boiling points

A

There is a strong attraction between the positive ions and delocalised electrons which means higher temperatures are needed to break the metallic bonds and dislodge the ions from Rigid positions within the lattice

55
Q

What does ductile mean?

A

Can be drawn out and stretched. Ductility permits metals to be drawn into wires.

56
Q

What does malleability mean?

A

Can be hammered into shape

57
Q

How is metallic bonding different from covalent bonding?

A

Covalent bonding consists of localised electrons which Are positioned between the two atoms bonded together, these electrons are shared only between the two atoms forming a covalent bond. In metallic bonding electrons are delocalised and are shared by all the positive ions in the structure

58
Q

Within a shell state the number of S, T, D and F orbitals

A

1
3
5
7

59
Q

What is a giant ionic lattice

A

Three-dimensional structure of oppositely charged ions held together by strong ionic bonds

60
Q

What are the two types of compounds with covalent bonds

A

simple molecular lattice

giant covalent lattice

61
Q

Define a simple molecular lattice

A

3-d structure of molecules bonded by weak vanderwaals forces

62
Q

What is the structure of a solid simple molecular lattice

A

strong covalent bonds between atoms

weak forces between molecules

63
Q

why are simple molecular structures soluble in non polar solvents

A

vanderwaals forces form between the simple molecular structure and the non polar solvent which weaken the lattice structure

64
Q

define giant covalent lattice

A

3d structure of atoms bonded by strong covalent bonds

65
Q

what is the structure of diamond

A

tetrahedral held together by stong covalent bonds through lattice

66
Q

why is diamond so hard

A

because the tetrahedral shape allows forces to be spread throughout the lattice

67
Q

what is the structure of graphite

A

strong hexagonal layered structure but with weak vanderwaals between layers

68
Q

what are simple molecular structures made up from

A

o2. h2o, h2, n2.Ne

Chemistry (6 decks)

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