Module 2 - [ch5,6] (atoms) Module 3 - [ch7] (periodicity)

Question 1

Why does first ionisation energy decrease going down a group [3]

Answer 1

atomic radius increases,
increased distance between outermost electron and nucleus

more electron shells
more electron shielding

less nuclear attraction

Question 2

How are shells filled with electrons?

Answer 2

orbitals are filled in order of energy (lowest energy shells are filled first)
electrons are assigned individually before being paired

Question 3

what is an atomic orbital

Answer 3

a region around the nucleus holding max two electrons with opposite spin
(cloud of negative charge)
it’s impossible to know where exactly the electron is

Question 4

ammonium ion (charge)

Answer 4

NH4+

Question 5

what is the shape and bond angle of an NH4+ ion

Answer 5

tetrahedral , 109.5

Question 6

Discuss the conductivity of ionic compounds [2]

Answer 6

when solid ions are in fixed positions and are not free to carry charge
when molten or in aqueous solution ions are mobile and can carry charge

Question 7

define first ionisation energy [3]

Answer 7

energy required to remove one electron [1] from each atom in 1 mole [1] of gaseous atoms [1] to form 1 mole of gaseous 1+ ions

Question 8

why is second ionisation energy generally greater than the first

Answer 8

since an electron has been removed the remaining electrons experience less repulsion and are pulled closer to the nucleus
experience more nuclear attraction
harder to remove

proton : electron ratio is greater

Question 9

define isotopes [1]

Answer 9

atoms of an element with the same number of protons but different number of neutrons [1]

Question 10

what would be evidence for shells in an atom?

Answer 10

jump in first ionisation energy
suggests the electron is much harder to remove since is much closer to the nucleus

Question 11

Explain why the first ionisation energies show a general increase across
Period 2 [2]

Answer 11

increasing atomic number ( number of protons)
more nuclear attraction
similar shielding across period since same shell

Question 12

Explain why the first ionisation energy of B is less than that of Be. [2]

Answer 12

In B, an electron is being removed from a higher energy level
An s electron lost in Be and a p electron lost in B

Question 13

What determines which block an element is in?

Answer 13

the highest energy level shell being filled

Question 14

what is meant by ionic bonding [1]

Answer 14

electrostatic force of attraction between oppositely charged ions [1]

Question 15

predict the bond angle of an F2O molecule, explain your choice [3]

Answer 15

2bp , 2lp
bond angle 104.5
bent (non-linear)
lone pairs repel more than bonded pairs

Question 16

Describe and explain two anomalous properies of water which results from
hydrogen bonding [4]

Answer 16

when solid, water is less dense than when liquid
ice has an open lattice

relatively high mp and bp
hydrogen bonds are strong, require lots of energy to overcome

Question 17

what is metallic bonding

Answer 17

giant metallic lattice
electrostatic attraction between fixed positive ions and delocalised electrons

Question 18

describe the structure and bonding shown by cl, how does this explain it’s difference in melting point compared to Mg ?

Answer 18

simple molecular lattice
Cl has london forces between molecules
london forces weaker than metallic bonds hence a lower melting point than Mg

Question 19

Explain, in terms of bonding and structure, the properties of graphite [6]

Answer 19

graphite is maleable [1] , has a very high melting point [1] and a good conductor [1]
maleable - layers can slide over one another, weak london forces between layers [1]
high mp - strong covalent bonds require energy to overcome (giant covalent lattice) [1]
good conductor - delocalised electrons free to carry charge [1]

Question 20

shape and bond angle of an ammonia molecule

Answer 20

NH3
3bp 1lp
107
pyramidal
lone pairs repel more than bonded pairs

Question 21

diagram of hydrogen bonding?

Answer 21

H+ - O:~ —— H+ - O: ~

Question 22

Suggest why H2S has a much lower boiling point than H2O

Answer 22

no hydrogen bonding,
weaker IM forces, less energy required to overcome forces

Question 23

why is chlorine is a gas at room temperature but carbon does not boil until well over 4500 °C.
Explain this difference, in terms of bonding and structure.

Answer 23

to boil Cl2 weak london forces have to be broken
to boil carbon covalent bonds need to be broken
covalent bonds are much stronger than london forces, more energy required

Question 24

what is meant by hydrogen bonding

Answer 24

electrostatic attraction between a hydrogen atom in one polar molecule (e.g. water) and a small electronegative atom (O,N,F) in another molecule

Question 25

define electronegativity

Answer 25

ability of an atom to attract the pair of electrons in a covalent bond

Question 26

explain a polar bond in terms of electronegativity

Answer 26

one element is more electronegative than the other attract the pair of electrons better, imbalance of charge

Question 27

????? why does boiling point increase from Na to Al

Answer 27

From Na → Al, no of delocalised electrons increases charge on positive ion increases/ ionic size decreases/ charge density increases attraction between + ions and electrons increases/ metallic bonding gets stronger

Question 28

Name the shape and bond angle of an NCl3 molecule why does it have this shape and bond angle

Answer 28

pyramidal 107 electron pairs repel to get as far apart as possible lone pairs repel more than bonded pairs 3bp, 1lp surround central N atom

Question 29

base + acid --> salt + water why does electrical conductivity of solution decrease as acid is added ?

Answer 29

as acid is added, reactions take place using up ions which were carrying charge

Question 30

The O–H bonds in water and the N–H bonds in ammonia have dipoles why do they have dipoles? [1]

Answer 30

one of the elements is more electronegative than the other [1] more electronegative one attracts the pair of electrons closer

Question 31

Describe and explain the density of ice compared with water [2]

Answer 31

ice is less dense than water hydrogen bonds cause an open lattice structure in ice [2]

Question 32

What is the difference between a covalent bond and a dative covalent bond? [1]

Answer 32

dative covalent, bonded pair comes from same atom [1]

Question 33

why is chlorine is a stronger oxidising agent than iodine.

Answer 33

Cl has fewer shells electron will be more strongly attracted

Question 34

Explain why a water molecule has a different shape from a carbon dioxide molecule. [2]

Answer 34

electron pairs repel to get as far apart as possible water - non-linear 104.5 co2 - 180 linear Oxygen in water surrounded by 4 areas of electron density/2 bonds and 2 lone pairs Carbon in CO2 surrounded by 2 regions of electron density/2 double bonds

Question 35

Explain why water has polar molecules but carbon dioxide has non-polar molecules [2]

Answer 35

co2 is symmetrical whereas h20 isn't in co2 dipoles cancel out

Question 36

silver

Answer 36

Ag+

Question 37

zinc

Answer 37

Zn2+

Question 38

aluminium

Answer 38

Al3+

Question 39

ammonium

Answer 39

NH4+

Question 40

carbonate

Answer 40

C03 2-

Question 41

sulfate

Answer 41

SO4 2-

Question 42

hydroxide

Answer 42

OH-

Question 43

nitrate

Answer 43

NO3-

Question 44

how does solubility work

Answer 44

ionic lattice must be broken down water molecules must surround ions and attract them solubility decreases as ionic charge increases since attraction may be too strong to overcome

Question 45

properties of ionic compounds

Answer 45

high mp and bp soluble in polar solvents conductors in molten or aq state

Question 46

what is a covalent bond really

Answer 46

orbital overlap

Question 47

what does it mean to say a covalent bond is localised and an ionic bond is not

Answer 47

in an ionic bond the ion attracts in all directions in a covalent bond the attraction acts solely between the shared pair of electrons and the nuclei

Question 48

define average bond enthalpy

Answer 48

measurement of covalent bond strength measure of average energy needed to break the bond

Question 49

what is electron-pair repulsion theory

Answer 49

the electron pairs surrounding a central atom determine the shape of the molecule electron pairs will repel each other to get as far apart as possible to minimise repulsion lone pairs repel each other more strongly than bonded pairs

Question 50

4bp

Answer 50

tetrahedral 109.5

Question 51

3bp

Answer 51

trigonal planar 120

Question 52

2bp

Answer 52

linear 180

Question 53

1 more lone pair does what to the bond angle

Answer 53

decreases it by 2.5

Question 54

3bp 1lp

Answer 54

trigonal pyramidal 107

Question 55

2bp 2lp

Answer 55

non linear (bent) 104.5

Question 56

6bp

Answer 56

octahedral 90 SF6

Question 57

how is electronegativity measured

Answer 57

pauling electronegativity scale increasing from bottom left to top right

Question 58

rough values of pauling electronegativity

Answer 58

0 - covalent <= 1.8 - polar covalent > 1.8 - ionic

Question 59

how is the electron pair shared in a non-polar bond

Answer 59

equally between the bonded atoms both have the same / similar electronegativity

Question 60

what does the strength of london forces increase with

Answer 60

number of electrons

Question 61

which of the IM forces are the strongest and weakest

Answer 61

hydrogen bonding london forces

Question 62

what is soluble in what?

Answer 62

polar subtances in polar solvents non-polar substances in non-polar solvents

Question 63

3 factors affecting electronegativity

Answer 63

atomic radius nuclear charge (number of protons) shielding

Question 64

why is the 4s shell filled first? what are the exceptions to this and why?

Answer 64

lower energy than 3d chromium and copper half/full 3d subshell more stable

Question 65

properties of simple molecular substances

Answer 65

low bp weak intermolecular forces strong covalent bonds (intramolecular) not affected when boiled relatively small molecules doesn't conduct

Question 66

define covalent bonding

Answer 66

electrostatic attraction between nuclei and a shared pair of electrons

Question 67

why are ionic compounds soluble in polar solvents

Answer 67

polar solvents such a water have a polar bond the delta positive and delta negative charges are able to attract the charged ions

Question 68

what is a lone pair

Answer 68

a pair of electrons on the outer shell not involved in bonding

Question 69

what does expansion of the octet mean

Answer 69

when a bonded atom has more than 8 atoms in its outer shell

Question 70

what does expansion of the octet mean

Answer 70

when a bonded atom has more than 8 atoms in its outer shell

Question 71

describe the bonding in simple molecular substances

Answer 71

atoms within the same molecule are bonded together by strong covalent bonds different molecules are held together by weak intermolecular forces

Question 72

what are three properties of giant covalent structures

Answer 72

high melting and boiling points non-conductors except graphite insoluble in both polar and non-polar solvents

Question 73

what does it mean to say a bond is non-polar

Answer 73

the electrons in the bond are evenly distributed

Question 74

what is meant by intermolecular force

Answer 74

attractive force between neighbouring molecules

Question 75

how are the elements arranged in the periodic table

Answer 75

increasing atomic number

Question 76

what is meant by periodicity

Answer 76

repeating trends in physical and chemical properties

Question 77

Why does first ionisation energy decrease from group2 to group3

Answer 77

in group3, outermost electrons are in p-orbitals whereas in group2 they are in s-orbitals so the electron is more easily removed

Question 78

why does first ionisation energy decrease from group5 to group6

Answer 78

in group5 the 5 electrons in the p-orbitals are single whereas in group6 the outermost electrons are spin paired experiences repulsion and more easily removed

Question 79

does first ionisation energy increase or decrease between the end of one period and the start of another why?

Answer 79

decrease increased atomic radius more shells, more electron shielding

Question 80

properties of a giant metallic lattice

Answer 80

high melting and boiling point good conductivity malleable ductile

Question 81

what is a malleable metal

Answer 81

the metal can be shaped into different forms

Question 82

what is the solubility of giant covalent structures and why

Answer 82

insoluble in almost all solvents covalent bonds too strong to be broken by interaction with solvents