Define isotope.
Isotopes are atoms of the same element with the same number of protons but different number of neutrons.
Why do isotopes have similar chemical properties but different physical properties?
- They have the same number and arrangement of electrons and most chemical reactions only involve the transfer of electrons between atoms.
- Also, they have different number of neutrons and hence different masses, hence they have different physical properties.
What is the formula for angle of deflection?
angle of deflection = |charge of particle/mass of particle|
Why are electrons deflected to a greater extent?
Even though electrons and protons have the same magnitude of charge, electrons are lighter and hence deflected to a greater extent.
Describe s orbitals.
S orbitals are spherical and there is only one s orbital in any subshell.
Describe p orbitals.
P orbitals are dumbbell shaped and they have two lobes of equal size. There are a total of three p orbitals in any subshell. All p orbitals are mutually perpendicular to one another. They all have the same energy (ie they are degenerate)
Putting all 3 of them together we get a sphere, hence a p orbital is 1/3 the volume of a sphere.
Describe d orbitals.
There are a total of 5 d orbitals in any subshell. They have different orientations in space and have the same energy.
Putting all 5 of them together we get a sphere, hence a d orbital is 1/5 the volume of a sphere.
State the two elements with anomalous electronic configurations.
- Chromium; Cr: [Ar]3d5 4s1
- Copper; Cu: [Ar]3d10 4s1
The 3d5 and 3d10 are energetically preferred configurations due to their symmetrical 3d electron cloud around the nucleus.
State the factors that affect the elctrostatic forces of attraction between the valence electrons and the nucleus.
- Charge of nucleus (nuclear charge): When there is more protons in the nucleus, larger positive charge of nucleus. hence stronger electrostatic forces of attraction between the nucleus and the valence electrons
- Shielding/Screening effect: It refers to the inter-electronic repulsion between the valence electrons and inner shell electrons. Fewer electrons in the inner shells which leads to a weaker shielding effect hence stronger electrostatic forces of attraction.
- Distance of electron from nucleus: smaller distance between nucleus and valence electron, stronger electrostatic forces of attraction (determined by principal quantum number)
What is effective nuclear charge?
The effective nuclear charge is the strength of the net electrostatic forces of attraction felt by the valence electrons after taking account the shielding effect of the inner shell electrons.
Effective nuclear charge = nuclear charge - shielding effect
Define the first ionisation energy.
The first ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of singly charged gaseous cations.
Explain why there is a general increase in first IE across the period.
Number of protons in the nucleus of the atom increases, thus nuclear charge increases. There is the same number of inner shell electrons hence shielding effect is approximately constant. Thus, there is an increase in effective nuclear charge and in electrostatic forces of attraction. Hence, more energy needed to remoce the valence electron and an increase in first IE across period.
Why is there a decrease in first IE from Group 2 to Group 13?
The np electron is of a higher energy level than the ns electron, hence less energy is required to remove the valence np electron from a Group 13 element.
Why is there a decrease in first IE from Group 15 to 16?
There is inter-electronic repulsion between the paired electrons in the same orbital for a Group 16 element, hence less energy to remove the np electron from Group 16.
Explain why there is a decrease in IE down the group.
Both nuclear charge and shielding effect increase. However, valence electrons are located in a shell with a larger principal quantum number and are further away from the nucleus. There is a decrease in electrostatic forces of attraction between the nucleus and the valence electron. Less energy is needed to remove the valence electron and hence a decrease in first IE down the group.
Why is there an increase in each successive IE?
Shielding effect decreases as electrons are removed, nuclear charge remains the same hence effective nuclear charge increases. Resulting in increase in electrostatic forces of attraction between the nucleus and the remaining electrons. More energy is needed for subsequent removal of electrons and hence successive IE of an atom always increases.
What does a large increase in IE represent?
It represents a change in principal quantum shell. Electron is removed from an inner principal quantum shell which is closer to the nucleus and hence experiences stronger electrostatic forces of attraction.
Define atomic radius.
Atomic radius is half the distance between the centres of two adjacent atoms found in the structure of the element.
Define metallic radius.
Metallic radius is half the distance between the centres of two adjacent atoms in the metal.
Define covalent radius.
The covalent radius of an atom is half the distance between the centres of two adjacent atoms that are covalently bonded.
Define van der Waals radius.
van der Waals radius of an atom is half the distance between the centres of two adjacent atoms which are not chemically bonded.
Define ionic radius.
Ionic radius is the average distance between the nuclei of two adjacent ions of the same element.
-
Atomic Structure22
-
Chemical Bonding32
-
Gaseous State9
-
Periodicity26
-
Mole Concept and Stoichiometry5
-
Energetics38
-
Reaction Kinetics36
-
Chemical Equilibria28
-
Acid Base Equilibria33
-
Solubility Equilibria11
-
Intro to Organic Chemistry25
-
Alkanes18
-
Alkenes21
-
Arenes15
-
Halogen Derivatives28
-
Hydroxy Compounds29
-
Carbonyl Compounds23
-
Carboxylic Acids and Derivatives24
-
NItrogen Compounds31
-
Electrochemistry18
-
Introduction to Transition Elements32
