3.1 Periodicity

Question 1

periodicity

Answer 1

the repeating trends in physical and chemical properties

Question 2

what change happens across each period?

Answer 2

elements change from metals to non-metals

Question 3

how can electron configuration be written in short?

Answer 3

the noble gas before the element is used to abbreviate it
e.g. Li is 1s2 2s1
can also be written as Li is [He] 2s1

Question 4

1st IE

Answer 4

the energy required to remove 1 electron from each atom in 1 mole of the gaseous element to form 1 mole of gaseous 1+ ions

Question 5

1st IE og Magnesium equation example

Answer 5

Mg (g) -> Mg+ (g) + e-

Question 6

what factors affect ionisation energy?

Answer 6

-atomic radius
-nuclear charge
-electron shielding

Question 7

explain the changes in 1st IE across period 3

Answer 7

1st IE increases across period 3 because there is:
- an increase in nuclear charge
- a decrease in atomic radius
- the same electron shielding
– which means more energy is needed to remove the 1st electron
1st IE dips at Al because its outer electron is in a 3p orbital which is a higher energy level than Mg’s 3s orbital, therefore less energy is needed to remove an electron
1st IE dips at S because one 3p orbital contains 2 electrons -> repulsion between paired electrons -> less energy needed to remove one

Question 8

why does 1st IE decrease between group 2 to 3?

Answer 8

because in group 3 the outermost electrons are in p orbitals whereas in group 2 they are in s orbital, so the electrons are easier to be removed

Question 9

why does 1st IE decrease between group 5 to 6?

Answer 9

because in group 5 the outermost electrons in p orbital are single electrons and in group 6 the outermost electrons are spin paired with some repulsion, therefore the electrons are slightly easier to remove

Question 10

does 1st IE increase or decrease between the end of one period and the start of the next? why?

Answer 10

it decreases because:
-atomic radius increases
-electron shielding increases

Question 11

does 1st IE increase or decrease down a group? why?

Answer 11

it decreases because:
-shielding increases -> weaker attraction
-atomic radius increases -> distance between outer electrons and nucleus increases -> weaker attraction
-increase in number of protons is outweighed by an increase in atomic radius and shielding

Question 12

properties of giant metallic lattices

Answer 12

-high melting and boiling point
-good electrical conductors
-malleable
-ductile

Question 13

describe the structure, forces and bonding in every element across period 2

Answer 13

Li & Be
-giant metallic; strong attraction between positive ions and delocalised electrons; metallic bonding
B & C
-giant covalent; strong forces between atoms; covalent bonding
N2, O2, F2, Ne
-simple molecular; weak IMF between molecules; covalent bonding within molecules and IMF between molecules

Question 14

describe the structure, forces and bonding in every element across period 3

Answer 14

Na, Mg, Al
-giant metallic; strong attraction between positive ions and delocalised electrons; metallic bonding
Si
-giant covalent; strong forces between atoms; covalent bonding
P4, S8, Cl2, Ar
-simple molecular; weak IMF between molecules; covalent bonding within molecules and IMF between molecules

Question 16

explain why first IE show a general increase across period 3?

Answer 16

-atomic radius decreases
-nuclear attraction
-electron shielding stays the same
(no. of protons increases so greater nuclear charge)

Question 17

explain why the general increase in first IE across period 3 is not followed for Mg to Al

Answer 17

-due to orbitals
-an Mg outer electron is removed from 3s orbital but Al outer electron is removed from 3p orbital
-Al (3p) has higher energy level than Mg (3s) so Al electron is easier to lose

Question 18

explain how first IE decrease down groups

Answer 18

-nuclear attraction decreases
-atomic radius increases
-increased shielding

Question 19

state and explain the trend in atomic radius from Li to F

Answer 19

trend-nuclear charge increases

explanation
-greater nuclear attraction to outer electrons
-outer electrons experience the same shielding
-atomic radius decreases and so nuclear charge increases

Question 20

explain the difference in MP of P and Cl

Answer 20

-P has more electrons (high Mr) which means it has greater london forces
-more energy is needed to overcome the intermolecular forces

Question 21

describe the bonding in Mg and Si

Answer 21

-Mg is held together by metallic bonding - electrostatic attraction between delocalised electrons and positive metal ions
-Si is held together by covalent bonds - electrostatic attraction between shared pair of electrons and nuclei of the bonded atoms

Question 22

along a period…

Answer 22

1st IE increases due to:
-decreasing atomic radius
-decreasing electron shielding
-and greater forces of attraction

Question 23

down a group…

Answer 23

1st IE decreases due to:
-increasing atomic radius
-increasing electrons shielding
-which reduces effect of electrostatic attraction

Question 24

successive ionisation energies…

Answer 24

increase, because atomic radius decreases and there is greater attraction between outer electrons and nucleus

Question 25

high ionisation energy

Answer 25

means there is a strong attraction between the electron and nucleus so more energy is needed to overcome attraction and remove the electron

Question 26

G2 elements react with water to produce:

Answer 26

metal hydroxide + hydrogen

Question 27

G2 metals react with oxygen to produce:

Answer 27

a solid white oxide

Question 28

G2 metals react with dilute acid to produce:

Answer 28

salt + hydrogen

Question 29

G2 oxides and hydroxides are:

Answer 29

bases
-and most of are soluble in water so they are alkalis

Question 30

calcium hydroxide is used to:

Answer 30

neutral is acidic soils in agriculture

Question 31

magnesium hydroxide is used in milk of magnesia to?

Answer 31

-neutralise excess acid, producing salt + water

Question 32

calcium carbonate is used in?

Answer 32

building materials like limestone

Question 33

halogens get less reactive as you go down the group because:

Answer 33

-atomic radius increases
-increased electron shielding
-makes it harder for larger atoms to attract the electron needed to form an ion (despite increased charge on nucleus) so larger atoms are less reactive
-they also become less oxidising

Question 34

why are halogen atoms oxidising agents?

Answer 34

they react by gaining an electron in their outer shell to form 1- ions

this means they are reduce, and therefore they oxidise another substance - so they are oxidising agents

Question 35

test for halide ions

Answer 35

1. add 2 drops of dilute nitric acid to remove ions that might interfere with the test
2. then add silver nitrate solution

ionic eq:
Ag+ (aq) + X- (aq) -> AgX (s)
where X = Cl, Br or I

-colour of precipitate identifies the halide
Cl-white
Br-cream
I-yellow
then to be extra sure we can add ammonia solution
Cl-dissolves in dilute NH3(aq)
Br-dissolves in conc. NH3(aq)
I-insoluble in conc. NH3(aq)

Question 36

disproportion reaction

Answer 36

when a halogen is simultaneously oxidised and reduced

Question 37

bleach equation

Answer 37

Cl2(g) + 2NaOH(aq) -> NaClO(aq) + H2O(l)
(sodium chlorate(I) solution=bleach)
Cl goes from 0 to +1 and gets reduced
Cl also goes from 0 to -1 and gets oxidised

Question 38

adv and dis of using chlorine to sterilise water:

Answer 38

adv-kills bacteria
dis-Cl gas is very toxic
alternatives include ozone or UV light