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IB Chemistry > Topic 3.2: Periodic Trends > Flashcards

Topic 3.2: Periodic Trends Flashcards

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1
Q

Definition of atomic radius

A

a) Distance from the nucleus to the outermost electron

b) Half the internuclear distance in a molecule

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2
Q

Definition of ionic radius

A

Distance from the nucleus to the outermost electron in an ion

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3
Q

Definition of first ionization energy

A

Energy required to remove one mole of electrons from one mole of gaseous atoms

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4
Q

Definition of electron affinity

A

Energy change when one mole of electrons is added to one mole of gaseous atoms

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5
Q

Definition of electronegativityor

A

A measure of the attraction an atom has for a shared pair of electrons in a covalent bond

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6
Q

Definition of periodicity

A

Repeating trends of physical and chemical properties in elements

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7
Q

Why do elements in the same group react the same way?

A

Since reactions are determined by the valence electrons and elements in the same group posess the same number of valence electrons

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8
Q

Definition of nuclear charge

A

Attraction exerted by the nucleus on electrons

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9
Q

Trend of nuclear charge between successive elements

A

Nuclear charge is given by the atomic number and increases by one between successive elements

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10
Q

Why do outer electrons do not experience the full attraction of the nuclear charge?

A

Since they are shielded from the nucleus and repelled by inner electrons

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11
Q

What is the effective nuclear charge?

A

Nuclear charge experienced by an atom’s valence electrons

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12
Q

Trends in the effective nuclear charge across a period and down a group

A

It increases with the group number but remains approximately the same down a group

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13
Q

Explanation of the trends in atomic radius down a group

A

It increases
a) Number of electron shells increase and the shielding effect increases, counteracting any effects due to nuclear charge

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14
Q

Explanation of the trends in atomic radius across a period

A

It decreases
a) Electrons are added to the same main energy level and the effective nuclear charge increases due to no significant change in shielding.

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15
Q

Explanation of the difference between the atomic radius of cations and atoms

A

Cations are smaller than their parent atoms.

a) The positive nucleus remains the same with the same attractive force, but now pulling on fewer electrons
b) It involves the loss of outer shell

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16
Q

Explanation of the difference between the atomic radius of anions and atoms

A

Anions are larger that their parent atoms
a) The slightly increased electron repulsion between the electrons in the outer main energy level causes the electrons to move further apart and increases the radius of the outer shell

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17
Q

Explanation of the trends of cations’ ionic radius across a period

A

It decreases across a period

a) The number of protons in the nucleus increases but the number of electrons remain the same
b) The increased attraction between the nucleus and the electrons pulls the outer shell closer to the nucleus

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18
Q

Explanation of the trends of anions’ ionic radius across a period

A

It decreases across a period

a) The number of protons in the nucleus increases but the number of electrons remains the same
b) The increased attraction between the nucleus and the electrons pulls the outer shell closer to the nucleus

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19
Q

Explanation of the trends of ionization energy across a period

A

Ionization energy generally increase across a period

a) The increase in nuclear charge increases the attraction between outer electrons in the same energy level
b) Nucleus attracts the electrons more strongly and is more difficult to remove them

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20
Q

Explanation of the trends of ionization energy down a group

A

a) The electron being removed is further from the nucleus and less attracted by the nucleus
b) It gets easier to remove valence e- as atomic radius increases down a group

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21
Q

What are the regular discontinuities in the trends ionization energy across a period? (2)

A

a) Lower ionization energy in elements from Group 13 than 2

b) Lower ionization energy in elements from Group 16 than 15

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22
Q

Explanation of the lower ionization energy in elements from Group 13 than 2

A

a) The electron removed when the Group 13 elements are ionized is a p electron.
b) The electron removed when the Group 2 elements are ionized is an s electron.
c) Electrons in p orbitals are of higher energy and further away from the nucleus than s electrons, so they are easier to remove than electrons in an s orbital.

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23
Q

Explanation of the lower ionization energy in elements from Group 16 than 15

A

a) Group 15 elements have the configuration s2p1p1p1
b) Group 16 elements have the configuration s2p2p1p1
c) For Group 16 elements, the electron is removed from a doubly occupied 2p orbital. An electron in a doubly occupied orbital is repelled by its partner more strongly, so it is easier to remove than an electron in a half-filled orbital.

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24
Q

Explanation of trends across first, second, and third electron affinity

A

a) First electron affinity is exothermic since added electrons are attracted to the positively charged nucleus.
b) Second and third electron affinities are endothermic as the added electron is repelled by the negatively charged ion, so energy needs to be available for this to occur.

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25
Q

Why do metals have a low EA and non-metals have a higher EA?

A

a) The greater the distance between the nucleus and the outer energy level, the weaker the electrostatic attraction and the less energy is released when an electron is added to the atom

26
Q

Electron affinity of Group 1 and 17

A

a) Group 1 have the lowest effective nuclear charge and attract the extra electron the least
b) Group 17 have high effective nuclear charge and incomplete outer energy level, attracting the electrons the most.

27
Q

Exception in the trend of electron affinity of halogens

A

Going from F to Cl the electron affinity becomes more exothermic because the decrease in electron-electron repulsion outweighs the fact that there is less attraction between electron and the nucleus

28
Q

Why do electron affinities reach a maximum in Group 2?

A

a) Group 2 elements have an electron configuration s2
b) The added electron must be placed into a p orbital which is further from the nucleus
c) Electron experiences reduced electrostatic attraction due to shielding from electrons in s orbital

29
Q

Why do electron affinities reach a maximum in Group 15?

A

a) Group 15 elements have the configuration s2p1p1p1
b) Added electron must occupy a p orbital that is already singly occupied
c) Attraction between the electron and atom is less than expected as there is increased inter-electron repulsion

30
Q

Electronegativity in metals and non-metals

A

a) Metals have low electronegativities

b) Non-metals have high electronegativities

31
Q

Explanation in trends of electronegativity across a period

A

It increases across a period
a) Owing to the increase in nuclear charge with no change in shielding, resulting in an increased attraction between the nucleus and the bond electrons.

32
Q

Explanation in trends of electronegativity down a group

A

It decreases down a group

a) The bonding electrons are furthest from the nucleus and so there is reduced attraction

33
Q

Explanation of trend in melting point down Group 1

A

It decreases

a) They have metallic structures which are held together by attractive forces between delocalized outer electrons and the positively charged ions.
b) This attraction decreases with distance from the nucleus

34
Q

Explanation of trend in melting point down Group 17

A

It increases

a) Elements have molecular structures which are held together by dispersion forces
b) Dispersion forces increase with the number electrons, increasing the strength between molecules

35
Q

Description of trend in melting point across period 3

A

It increases reaching a maximum at G14 and falls reacing a minimum at G18

36
Q

Explanation of trend in melting point across Na, Mg, and Al (Metallic bond)

A

a) The nuclei of the atoms are getting more positively charged.
b) The number of delocalized electrons is getting more negatively charged.
c) Delocalized electrons are more attracted and getting nearer to the nuclei.
d) The strength of metallic bond increases, requiring more energy to be broken

37
Q

Explanation of high melting point of Si (Giant covalent structure)

A

Si atoms are attracted to each other by strong covalent bonds, which requires a very large amount of energy to be broken.

38
Q

Explanation of trend in melting point across P4, S8, and Cl2 (Van der Waals’ forces)

A

The strength of the van der Waals’ forces decreases as the size of the molecule decreases, reducing the energy required to disrupt these forces between molecules

39
Q

Definition of metallic character

A

How easily an atom can lose electrons

40
Q

Description of trends in metallic character across a period and down a group

A

a) It decreases across a period

b) It increases down a group

41
Q

Physical and chemical properties of Alkali Metals

A

a) Good conductors of heat and electricity
b) Low densities (soft)
c) Low melting points
d) Very reactive metals
e) Form ionic compounds with non-metals

42
Q

Reaction of alkali metals with water

A

2X(s) + H2O(l) => 2XOH(aq) +H2(g)

43
Q

Reaction of alkali metals with oxygen

A

4X(s) + O2(g) => 2X2O(s)

44
Q

Physical and chemical properties of Halogens

A

a) Colored (Yellow, Green, Orange, Purple)
b) Gradual change from gases to solids
c) Very reactive non-metals
d) Diatomic molecules
e) Form ionic compounds with metals and covalent compounds with other non-metals

45
Q

Reaction of halogens with Group 1 metals

A

2Na(s) + Cl2(g) => 2NaCl(s)

46
Q

Displacement reactions between halogens and halides

A

2KBr(aq) + Cl2(aq) => 2KCl(aq) + Br2(aq)

47
Q

Precipitation of halides

A

AgNO3(aq) + X-(aq) => AgX(s)

48
Q

Physical and chemical properties of Noble Gases

A

a) Colorless
b) Monoatomic
c) Very unreactive

49
Q

Properties of Na, Mg and Al

a) Structure
b) Formula of oxide
c) Acid-base character of oxide
d) Reaction with water

A

a) Giant metallic
b) Na2O / MgO / Al2O3
c) Basic | Amphoteric
d) Soluble | Insoluble

50
Q

Properties of Si

a) Structure
b) Formula of oxide
c) Acid-base character of oxide
d) Reaction with water

A

a) Giant covalent
b) SiO2
c) Acidic
d) Insoluble

51
Q

Properties of P, S, and Cl

a) Structure
b) Formula of oxide
c) Acid-base character of oxide
d) Reaction with water

A

a) Simple covalent
b) P4O10, P4O6 / SO3, SO2 / Cl2O7, Cl2O
c) Acidic
d) Soluble

52
Q

Reaction of basic oxides with water

A

Na2O(s) + H2O(l) => 2NaOH(aq)

MgO(s) + H2O(l) => Mg(OH)2(aq)

53
Q

Reaction of basic oxides with acids

A

Na2O(s) + 2HCl(l) => 2NaCl(aq) + H2O(l)

54
Q

Reaction of phosphorus oxide with water

A

P4O10(s) + 6H2O(l) => 4H3PO4(aq)

P4O6(s) + 6H2O(l) => 4H3PO3(aq)

55
Q

Reaction of sulfur oxide with water

A

SO3(l) + H2O(l) => H2SO4(aq)

SO2(g) + H2O(l) => H2SO3(aq)

56
Q

Reaction of chlorine with water

A

Cl2O7(l) + H2O(l) => 2HClO4(aq)

Cl2O(l) + H2O(l) => 2HClO (aq)

57
Q

Reaction of silicon with alkalis (No reaction with water)

A

SiO2(s) + 2NaOH => NaSiO3(aq) + H2O (l)

58
Q

Reaction of aluminum with acids (No reaction with water)

A

Al2O3(s) + 6HCl => 2AlCl3 (aq) + 3H2O(l)

59
Q

Reaction of aluminum with acids (No reaction with water)

A

Al2O3(s) + H2O(l) + NaOH(aq) => 2NaAl(OH)4(aq)

60
Q

Reactions of NO

A

a) No reaction with water

b) Oxidized in the air to form NO2

61
Q

Reactions of NO2

A

3NO2(g) + H2O(l) => 2HNO3(aq) + NO(aq)

2NO2(g) + H2O (l) => HNO2(aq) + HNO3(aq)

62
Q

Acidic rain explanation with chemical formulas

A

Acidic oxides dissolve in rainwater to form acidic rain (SO3 / NO2)
SO3(g) + H2O(l) => H2SO4(aq)
CaCO3(s) + H2SO4(aq) => CaSO4(aq) + H2CO3(aq)
H2CO3(AQ) => CO2(g) + H2O(l)

IB Chemistry (39 decks)

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