Atomic structure

Flashcards on Ionization Energy, Electronic Configuration, Isotopes, Atomic Structure and Radius

What is ionization energy?

The amount of energy required to remove one mole of electrons from one mole of gaseous atoms of an element to form one mole of gaseous ions.

Under what standard conditions are ionization energies measured?

298 K and 101 kPa.

What are the units of ionization energy?

Kilojoules per mole (kJ mol-1).

Give an example of a first ionization energy equation (gaseous calcium).

Ca (g) → Ca+ (g) + e- IE₁ = +590 kJ mol-1

What is the second ionization energy (IE2) of an element?

The energy required to remove one mole of electrons from one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions.

Which group of metals have a relatively low ionization energy?

The group I metals.

What four factors affect the size of the first ionization energy?

Attractive forces between the nucleus and electrons, distance of outer electrons from the nucleus, shielding effect of inner electrons, and spin-pair repulsion.

How does the ionization energy change across a period?

The nuclear charge increases, the atomic radius decreases, and shielding remains constant.

Why is there is a rapid decrease in ionization energy between the last element in one period and the first element in the next period?

Increased distance between the nucleus and outer electrons due to adding a new shell, and increased shielding by inner electrons because of the added shell.

Explain the slight decrease in IE₁ between nitrogen and oxygen, including their electron configurations.

Nitrogen's electron configuration, which is 1s² 2s² 2px¹2py¹2pz. Oxygen's electron configuration, which is 1s² 2s² 2px²2py¹2pz. The slight decrease in IE₁ between nitrogen and oxygen is due to spin-pair repulsion in the 2px orbital of oxygen.

How does the ionization energy change down a group?

The distance between the nucleus and outer electron increases, and shielding by inner shell electrons increases.

Why do successive ionization energies of an element increase?

Removing an electron from a positive ion is more difficult than from a neutral atom. As more electrons are removed, the attractive forces increase due to decreasing shielding and an increase in the proton to electron ratio.

Name four factors that affect the magnitude of the ionization energy:

Positive nuclear charge, Shielding, Atomic/ionic radius, Spin-pair repulsion.

How are electrons arranged in an atom?

Electrons are arranged around the nucleus in principal energy levels or principal quantum shells.

What letters designate the subshells within principal quantum shells?

s, p, d, and sometimes f.

How many orbitals are in the s subshell, and how many electrons can it hold?

1 orbital (2 electrons).

How many orbitals are in the p subshell, and how many electrons can it hold?

3 orbitals (6 electrons).

How many orbitals are in the d subshell, and how many electrons can it hold?

5 orbitals (10 electrons).

How many orbitals are in the f subshell, and how many electrons can it hold?

7 orbitals (14 electrons).

What is the ground state electronic configuration?

The most stable electronic configuration of an atom, achieved by filling the sub-shells with the lowest energy first.

How do electrons fill orbitals within the same subshell?

Electrons will occupy separate orbitals in the same sub-shell first, and with the same spin, to minimize repulsion.

What is a free radical?

A species with one or more unpaired electrons.

How are ions formed, and how does this affect the electron configuration?

Negative ions are formed by adding electrons to the outer sub-shell, while positive ions are formed by removing electrons from the outer sub-shell.

Into what four main blocks is the Periodic Table split, depending on electronic configuration?

s, p, d, and f blocks.

What are isotopes?

Atoms of the same element that contain the same number of protons and electrons but a different number of neutrons.

What properties are the same, and what properties are different, for isotopes of the same element?

Same chemical symbol, same atomic number, different mass number.

How do the chemical and physical properties of isotopes compare?

Isotopes have similar chemical properties but different physical properties.

What are the three subatomic particles that make up an atom?

Protons, neutrons, and electrons.

How do electrons behave in an electric field?

Electrons are easily deflected away from the negative plate and towards the positive plate.

How do protons behave in an electric field?

Protons are deflected away from the positive plate and towards the negative plate.

How do neutrons behave in an electric field?

Neutrons are not deflected at all.

How to calculate the number of neutrons in the nucleus of an atom

Number of Neutrons = Mass number - Atomic number

What determines which element an atom or ion belongs to?

The atomic number (or proton number).

How is atomic radius defined?

Half the distance between the two nuclei of two covalently bonded atoms of the same type.

What are the general trends in atomic radii across the Periodic Table?

Atomic radii decrease as you move across a Period, atomic radii increase moving down a Group.

What are the trends in ionic radii in relation to charge?

Ionic radii increase with increasing negative charge; Ionic radii decrease with increasing positive charge