3.1 Periodicity Flashcards

1
Q

periodicity

A

the repeating trends in physical and chemical properties

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2
Q

what change happens across each period?

A

elements change from metals to non-metals

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3
Q

how can electron configuration be written in short?

A

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

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4
Q

1st IE

A

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

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5
Q

1st IE og Magnesium equation example

A

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

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6
Q

what factors affect ionisation energy?

A

-atomic radius
-nuclear charge
-electron shielding

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7
Q

explain the changes in 1st IE across period 3

A

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

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8
Q

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

A

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

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9
Q

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

A

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

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10
Q

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

A

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

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11
Q

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

A

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

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12
Q

properties of giant metallic lattices

A

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

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13
Q

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

A

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

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14
Q

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

A

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

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15
Q
A
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16
Q

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

A

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

17
Q

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

A

-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

18
Q

explain how first IE decrease down groups

A

-nuclear attraction decreases
-atomic radius increases
-increased shielding

19
Q

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

A

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

20
Q

explain the difference in MP of P and Cl

A

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

21
Q

describe the bonding in Mg and Si

A

-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

22
Q

along a period…

A

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

23
Q

down a group…

A

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

24
Q

successive ionisation energies…

A

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

25
high ionisation energy
means there is a strong attraction between the electron and nucleus so more energy is needed to overcome attraction and remove the electron
26
G2 elements react with water to produce:
metal hydroxide + hydrogen
27
G2 metals react with oxygen to produce:
a solid white oxide
28
G2 metals react with dilute acid to produce:
salt + hydrogen
29
G2 oxides and hydroxides are:
bases -and most of are soluble in water so they are alkalis
30
calcium hydroxide is used to:
neutral is acidic soils in agriculture
31
magnesium hydroxide is used in milk of magnesia to?
-neutralise excess acid, producing salt + water
32
calcium carbonate is used in?
building materials like limestone
33
halogens get less reactive as you go down the group because:
-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
34
why are halogen atoms oxidising agents?
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
35
test for halide ions
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)
36
disproportion reaction
when a halogen is simultaneously oxidised and reduced
37
bleach equation
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
38
adv and dis of using chlorine to sterilise water:
adv-kills bacteria dis-Cl gas is very toxic alternatives include ozone or UV light