Topic 3: Periodicity Flashcards
Periodicity
The repeating pattern of physical and chemical properties shown by different periods in the periodic table
Atomic Radius
- Distance from the nucleus to the outermost electron
Group Trend of Atomic Radius
- Additional electrons are occupying additional energy levels
- Electron experience increases shielding and are further away from the nucleus
- As a result, there is a decrease in attractive forces of nucleus with the outer electrons
Periodic Trend of Atomic Radius
- Additional electrons are occupying the same energy level. As a result shielding is constant.
- However, additional protons are also being added to the nucleus. As a result, the outer electrons experience an increase in effective nuclear charge (ENC)
- As a result, there is an increase in attractive forces of nucleus with the outer electrons and the electrons are pulled closer to it.
Bonding Atomic Radius
Half of the distance between the centre of the nuclei of 2 covalently-bonded atoms
Why is atomic radius (RnB) > covalent radius (Rb)?
- The covalent bond is formed by the overlapping of atomic orbitals
- The overlapping region becomes common ground
Ionic Radius (IR)
Distance from the nucleus to the outermost electron in an ion
Cation vs Neutral Counterpart
- cation is smaller than the neutral counterpart due to the loss of electrons from the valence shell
- The outermost occupied energy level is one less (n-1) than in the neutral atom (n)
Anion vs Neutral Counterpart
- Anion is larger than the neutral counterpart due to the gain of electrons in the valence shell
- Increases electron repulsion in the outer energy level
Group Trend of Ionic Radius
Ionic Radius Increases
- Following the formation, valence electrons are still occupying an additional energy level relative to the one above them in the group
Period Trend of Ionic Radius
Ionic Radius Decreases
- Outer electrons occupy the same energy levels and protons are being added to the nucleus
- Therefore, shielding is constant and the nucleus increases the effective nuclear charge (ENC)
- Results in electrons more strongly attracted to the nucleus and pulled closer to it
First Ionization Energy (IE)
IS the energy required to remove the first electron from a neutral gaseous atom in its ground state.
X(g) + energy → X^-1 (g) + e^-
Group Trend of First Ionization Energy (IE)
First Ionization Energy Decreases
- Atomic radius is increasing, outer electrons are further away from the nucleus and sheilding effect increase greater than the nuclear charge, so electrons hold a stongly attracted to the nuclues
- Requires
Periodic trends of melting points
Across a period (→):
- m.pt depends on the element structure and the type of attractive forces holding atoms together
- metals: have metallic bonding, so ↑ no of e-s = ↑ m.pt
- metalloids: macromolecular covalent structures with strong bonds = ↑ m.pt
- non-metals: simple molecular structures with weak intermolecular forces = ↓ m.pt
- noble gases: monatomic molecules = ↓ m.pt
Down a group (↓):
- Gp 1: ↑ no. of shells = ↑ shielding effect = ↑ radius = ↓ m.pt
- Gp. 7: ↑ attractive forces between diatomic molecules = ↑ m.pt
First ionization energy
Energy required to remove one mole of electrons from one mole of gaseous atoms
Periodic trends of ionization energy
- across a period (→): ↑ no. of valence e-s = ↑ no. of protons = ↑ ionisation energy
- values don’t increase regularly across a period
- due to new sub-levels being filled
- and existence of paired electrons (paired e-s have greater repulsion between them so are easier to remove)
- down a group (↓): ↑ no. of shells = ↑ distance between nucleus and electron = ↓ ionisation energy
Discrepancies in periodic trends of ionization energy
Drop in value between:
- Be and B
- Mg and Al
- N and O
- P and S
- Be and B, and Mg and Al, due to existence of a new subshell in B and Al
- N and O, and P and S, due to the fact that electrons in O and S’s outer orbitals are paired
Electron affinity
Energy change when one mole of electrons is added to one mole of gaseous atoms
- ↑ tendency = ↑ Eea
Factors affecting electron affinity
- nuclear charge: ↑ nuclear charge = ↑ Eea
- size of atom: ↑ atomic radius = ↓ Eea
- e. config.: stabler configs = ↓ Eea
Periodic trends of electron affinity
- across a period (→): ↓ radius and ↑ nuclear charge = ↑ Eea
- down a group (↓): ↑ radius and ↑ nuclear charge = but effect of ↑ radius is greater = so ↓ Eea
Electronegativity
A measure of the ability of atoms to attract a shared pair of electrons toward itself in a covalent bond
Periodic trends of eleectronegativity
Polar covalent bond
Type of oxides
Group 1 trends
Reaction between alkali metals and water
Alkali metals: difference in reaction to water
Reaction between alkali metals and halogens
Group VII trends
Characteristics of halogens
Transition metals
Properties of transition metals
Blocks in which the periodic table is divided according to the last orbital of an element
a) s-block
b) p-block
c) d-block
d) f-block
s-block
Group 1, Group 2 and He
p-block
Group 13 - 18 (excluding He)
d-block
Groups 3 - 12 (including La and Ac)
f-block
Elements 58 - 71 and Elements 90 - 103
How are elements arranged in the periodic table?
According to increasing atomic number
Relationship between elements in the same group based on their electronic configuration
Group number is the same as the number of valence electrons (outer electrons)
Relationship between elements in the same period based on their electronic configuration
Period number is the same as the number of occupied main energy levels in the atom
Characteristics of metal
- Good conductors of heat and electricity
- Malleable (capable of being hammered into thin sheets)
- Ductile (capable of being drawn into wires)
- Lustre (shiny)
- Lose electrons in chemical reactions
Characteristics of non-metals
- Poor conductors of heat and electricity
- Gain electrons in chemical reactions
Characteristics of metalloids
-Both metallic and non-metallic properties
- Semiconductors
What elements conform metalloids?
B / Si / Ge / As / Sb / Te
Name of Group 1 elements
Alkali metals
Name of Group 2 elements
Name of Group 17 elements
Name of Group 18 elements
What is the name of groups 3-12?
What is the name of elements 57 - 71?
What is the name of elements 89 - 103?
Why do elements in the same group react the same way?
Since reactions are determined by the valence electrons and elements in the same group posess the same number of valence electrons
Nuclear charge
Attraction exerted by the nucleus on electrons
Trend of nuclear charge between successive elements
Nuclear charge is given by the atomic number and increases by one between successive elements
Why do outer electrons do not experience the full attraction of the nuclear charge?
Since they are shielded from the nucleus and repelled by inner electrons
Effective nuclear charge (ENC)
Nuclear charge experienced by an atom’s valence electrons
Trends in the effective nuclear charge across a period and down a group
It increases with the group number but remains approximately the same down a group
Explanation of the trends in atomic radius down a group
It increases
a) Number of electron shells increase and the shielding effect increases, counteracting any effects due to nuclear charge
Explanation of the trends in atomic radius across a period
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.
Explanation of the difference between the atomic radius of cations and atoms
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
Explanation of the difference between the atomic radius of anions and atoms
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
Explanation of the trends of cations’ ionic radius across a period
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
Explanation of the trends of anions’ ionic radius across a period
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
Explanation of the trends of ionization energy across a period
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
Explanation of the trends of ionization energy down a group
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
What are the regular discontinuities in the trends ionization energy across a period?
a) Lower ionization energy in elements from Group 13 than 2
b) Lower ionization energy in elements from Group 16 than 15
Explanation of the lower ionization energy in elements from Group 13 than 2
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.
Explanation of the lower ionization energy in elements from Group 16 than 15
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.
Explanation of trends across first, second, and third electron affinity
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.
Why do metals have a low EA and non-metals have a higher EA?
Electron affinity of Group 1 and 17
Exception in the trend of electron affinity of halogens
Why do electron affinities reach a maximum in Group 2?
Why do electron affinities reach a maximum in Group 15?
Electronegativity in metals and non-metals
Explanation in trends of electronegativity across a period
Explanation in trends of electronegativity down a group
Explanation of trend in melting point down Group 1
Explanation of trend in melting point down Group 17
Description of trend in melting point across period 3
Explanation of trend in melting point across Na, Mg, and Al (Metallic bond)
Explanation of high melting point of Si (Giant covalent structure)
Explanation of trend in melting point across P4, S8, and Cl2 (Van der Waals’ forces)
Metallic character
Description of trends in metallic character across a period and down a group
Physical and chemical properties of Alkali Metals
Reaction of alkali metals with water
Reaction of alkali metals with oxygen
Physical and chemical properties of Halogens
Physical and chemical properties of Halogens
Displacement reactions between halogens and halides
Precipitation of halides
Physical and chemical properties of Noble Gases
