1.6 The periodic table

Question 1

Define periodicity

Answer 1

the trend in properties from left to right is repeated across each period, predictions can be made about properties

Question 2

Why is magnesium in the s block?

Answer 2

b/c:
- 1s2 2s2 2p6 3s2 - the last subshell of the electronic configuration is its assigned block, the highest energy orbital occupied by electrons

Question 3

What properties of elements change as you move across a period?

Answer 3

• trends in electronegativity
• atomic no. increases
• metals -> non-metals
• nature of bonding changes
• increase in number of protons and electrons

Question 4

Define IE

Answer 4

The energy require to remove one mole of electrons from one mole of gaseous atoms to form one mole of ions with a single positive charge

Question 5

Why does IE decrease down a group?

Answer 5

shielding of outer electrons from inner electrons increases so ENC decreases

Question 6

Why does IE increase across a period?

Answer 6

bc ANC is increasing so greater attraction for electrons
- also little variation in atomic radius and electrons are being removed from the same shell so some degree of shielding

Question 7

Why is there a decrease in ionisation energy between grp 2 and 3?

Answer 7

bc grp 3 elements outer electron is in a new subshell of slightly higher energy level and partly shielded by s electrons so is easier to loose

Question 8

Why is there a general increase in a successive IE graph?

Answer 8

as removing electrons from positive ion

Question 9

What do the big jumps in a successive IE show?

Answer 9

change in shell

Question 10

What does the m.p depend on?

Answer 10

the strength of the forces holding the particles together in a solid

Question 11

why doe metallic structures have a high m.p?

Answer 11

bc strong electrostatic attraction between ‘sea’ of delocalised electrons

Question 12

In a m.p graph, what is the type of bonding in the first three elements in a period? What does this mean for melting points?

Answer 12

• metallic
• increasing bc metallic bonding gets stronger, metal ions have a greatercharge and theres an increased no. of delocalised electrons

Question 13

In a m.p graph, what is the type of element with the highest m.p.
What does this mean for its melting point?

Answer 13

• giant covalent
• each atom is covalently bonded to 4 other atoms (like diamond), high energy needed to break these bonds

Question 14

In a m.p graph, what is the type of bonding present in the next three elements in a period as the m.p decreases?
What does this mean for their m.p?

Answer 14

• simple covalent
• decreasing as intermolecular forces are weak and easy to break + shielding is increasing so outer electrons are easier to remove

Question 15

In a m.p graph, what is the type of bonding presents in the last element
What does this mean for its m.p?

Answer 15

• atomic
• lowest m.p and b.p bc exists as a separate atom held together by very weak temporary dipole forces

Question 16

Why is there a decrease in IE between group 5 and 6?

Answer 16

• the electron is removed from an orbital containing a pair of electrons, the repulsion between these electrons makes the electrons easier to remove.

Question 17

Define electronegativity

Answer 17

the ability of an element to attract a pair of covalently bonded electrons

Question 18

Why does electronegativity increase across a period?

Answer 18

As you move from left to right across the periodic table, atoms have a greater nuclear charge and a smaller covalent radius. This allows the nucleus to attract the bonding electrons more strongly.

Question 19

Why does electronegativity decrease down a group?

Answer 19

atomic size increases as additional electron shells are added, resulting in a larger distance between the nucleus and outer electrons.

Question 20

What do we see in a reaction between each of the group 1 metal elements and water? Why?

Answer 20

Li+ - floats on water, gently effervesces, doesn’t melt
Na+ - melts into a ball, dashes around surface
K+ -> melts into a ball, bursts into a lilac flame
bc atomic radius increases
the outer electron gets further from the nucleus
the attraction between the nucleus and outer electron gets weaker – so the electron is more easily lost.

Question 21

What is the general equation for group 1 metals and water, using M for metals?

Answer 21

2M + 2H20 -> 2MOH + H2

Question 22

What do we see in a reaction between the group 1 metals and oxygen? Why?

Answer 22

Li+ - shiny when cut, slowly goes dull
Na+ - shiny when cut, quickly goes dull
K+ - shiny when cut, rapidly goes dull
Bc atomic radius increases so outer electron gets further from nucleus, reducing the attraction between the nucleus and outer electron so the electron is more easily lost

Question 23

What is the general equation for group 1 metals and oxygen, using M for metals?

Answer 23

4M + O2 -> 2M2O

Question 24

What are the flame tests for the group 1 metals?

Answer 24

Li+ - red
Na+ - yellow-orange
K+ - lilac

Question 25

What is the reaction between group 2 metals and water?

Answer 25

Mg2+ - no reaction w/cold water, reacts w steam, bright light formed Ca2+ - lots of effervescing, steady stream of bubbles, liquid goes cloudy as a white ppt of Ca(OH)2 forms Ba2+ - greater effervescence and solution is clearer since more soluble

Question 26

What is the general equation between a group 2 metal and water? (using M as a substitute for the metal)

Answer 26

M + 2H2O -> M(OH)2 + H2

Question 27

What are all s-block metal oxides and what do they neutralise to form what? Give an example

Answer 27

- are strong bases and neutralise acids to form a salt and water - MgO + 2HCl -> MgCl2 + H2O

Question 28

Since hydroxides are soluble, what are they?

Answer 28

alkalis

Question 29

What do group 1 oxides and barium oxide react with to form a soluble hydroxide? Give an example.

Answer 29

- water - Na2O + H2O -> 2NaOH

Question 30

Why are saturated solutions of group 2 hydroxides (exception of Ba) very weakly basic?

Answer 30

the conc of hydroxide is very low

Question 31

Why can't all group 2 oxides be used for the same purposes?

Answer 31

different basic strength

Question 32

What is the flame test for the group 2 metals?

Answer 32

Mg2+ - no colour Ca2+ - brick red Ba2+ - apple green

Question 33

What are all nitrates?

Answer 33

soluble

Question 34

What are all carbonates?

Answer 34

insoluble

Question 35

Do hydroxides become more soluble or insoluble down a group? therefore...

Answer 35

more soluble Mg(OH)2 is insoluble Ba(OH)2 is soluble

Question 36

Do solutes become more soluble or insoluble down a group? therefore...

Answer 36

insoluble MgSO4 is soluble BaSO4 is insoluble

Question 37

What is the general ionic equation between a group 2 metal and a hydroxide?

Answer 37

M2+(aq) + 2OH-(aq) -> M(OH)2(s)

Question 38

What is the general ionic equation between a group 2 metal and sulphuric acid?

Answer 38

M2+(aq) + SO4^2-(l) -> MSO4(s)

Question 39

What is the general ionic equation between a group 2 metal and a carbonate?

Answer 39

M2+(aq) + CO3^2-(aq) -> MCO3(s)

Question 40

Why does thermal stability increase down the group 2 carbonates and hydroxides?

Answer 40

it becomes more difficult to break down/decompose

Question 41

What is the general equation between group 2 hydroxides decomposing on heating to the oxide and steam? Do all group 2 hydroxides do this?

Answer 41

M(OH)2(s) -> MO(s) + H2O (g) - yes

Question 42

What is the general equation between group 2 carbonates decomposing on heating to the oxide and carbon dioxide? Do all group 2 carbonates do this? What colour and states are all these carbonates? And the oxides? How do we test this?

Answer 42

MCO3 (s) -> MO(s) + CO2 - yes - all of these carbonates are white solids - all oxides produced are also white solids - can test this in a lab by heating and seeing how long it takes for products/ppt to form

Question 43

What are the salts halogens form called?

Answer 43

halides

Question 44

What do halogens exist as?

Answer 44

diatomic molecules containing a single covalent bond

Question 45

What is the colour of Cl, Br and I at room temp and what are their states?

Answer 45

Cl- green gas Br- red/brown liquid I- grey solid

Question 46

What happens down the halogen group (m.p)? Why?

Answer 46

- mp increases as elements atomic number increases so more temporary dipoles

Question 47

What are the halogens reactions with iron wool? And with sodium?

Answer 47

Fluorine w/iron wool= cold iron wool bursts into flames - w/sodium = reacts violently at room temp Chlorine w/iron wool= hot iron wool glows brightly - w/sodium = reacts vigorously with hot Na Bromine w/iron wool = hot iron wool glows (not as brightly) - w/sodium = reacts less vigorously with hot Na Iodine w/iron wool = hot iron wool shows faint red glow - w/sodium = reacts slowly with hot Na

Question 48

Why does reactivity decrease down group?

Answer 48

bc more shells so increase in shielding so harder to gain electron so oxidising power decreases

Question 49

What is the general equation for a halogen and iron wool? (with H representative of halogen)

Answer 49

2Fe + 3H2 -> 2FeH3

Question 50

What is the general equation for a halogen and sodium? (with H representative of halogen)

Answer 50

2Na + H2 -> 2NaH

Question 51

What is the general electronic configuration of halogens? which block is this?

Answer 51

ns2 np5 block p as last subshell is p

Question 52

Define a displacement reaction.

Answer 52

a more reactive element displaces a less reactive element from a solution of its salt

Question 53

What is the ionic displacement reaction between sodium bromide and chlorine? What is reduced and what is oxidised?

Answer 53

2Br + Cl2 -> 2Cl +Br2 -1 0 -1 0 Chlorine is reduced Bromine is oxidised

Question 54

Fluorine can displace all other halogens from their compounds, explain that in terms of electronegativity and atomic structure.

Answer 54

most electronegative element so pulls electrons towards itself, smaller radius so greater ENC