Chpt 3. PART 2

When going from top to bottom of a group, the n value (orbitals) of the valence electrons increases (has more orbitals) and hence the valence electrons are more likely to be further away from the nucleus (since they are in outer orbitals) and hence the atomic radius increases. [why atomic size increases down a group]

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Reason why transition metals lose electrons first from (n-1)d orbital rather than (ns) s orbital

1) The difference in energy between (ns) s orbital and (n-1)d orbital is very small. 2) As (n-1)d orbitals fill with electrons with increasing atomic #, the effective nuclear charge (Zeff) felt by their electrons increases more than Zeff of ns electrons. (this is because d electrons in a transition metal atom are shielded by less than ns electrons and hence ns electrons have more energy and ionize first)

d orbital (l=2) hold ____ electrons

10 electrons (5 orbitals)

f orbital (l=3) holds _____ electrons

14 electrons (7 orbitals)

s orbitals (l=0) hold ___ electrons

2 electrons (1 orbital)

p orbitals (l=1) hold ____ electrons

6 electrons (3 orbitals) [Note: each orbital can only contain two electrons]

Orbital Diagram

A depiction of the arrangement of electrons in an atom or ion. Boxes = orbitals. Single-headed arrows = electrons [shows 2p electrons are unpaired (in separate orbitals/boxes) By convention, the spin arrows of single electrons in separate orbitals (boxes) are drawn pointed up]

Electronegativity

Ability of an atom to hold electrons tightly

Elements in the same group have the same chemical properties (and same number of valence electrons).

Atoms DECREASE in size from left to right across a period and INCREASE in size from top to bottom down a group (atoms with a higher atomic # are hence not always larger than atoms with a smaller atomic number)

Anions of main group elements are much ______ than their parent atoms

Bigger

Deviation in valence electron filling patterns:

Chromium (Cr, Z = 24); [Ar]3d^5 4s^1 (3d orbital fills before 4s orbital). Same occurs with Copper (Cu, Z = 29).

Aufbau principle

Concept that electrons are placed in lowest-energy orbitals available first. (Fill lowest n-value, then in order of lowest l value (shape): s < p < d < f)

Isoelectric

Describes atoms or ions that have identical electron configurations.

In general EA becomes more negative with increasing atomic number across a row, but there are many exceptions.

Exception: Group 18 (Noble gasses) Because they already have stable atom configurations and adding an electron would actually require an input of energy.

For transition metals (groups/columns 3-12) [part of d block] The principle quantum number (n), starting in row/period 4, is always one less than the row number.

For F-block elements, the principle quantum number (n) is always 2 less than the row/period number.

Ionization energy trend

Harder to steal an electron away from upper right corner because it takes more energy to do so (has a higher ionization energy). (electrons in lower left corner are easy to steal away and hence have lose ionization energy)

An element in p-block forms a monatomic anion by gaining enough electrons to fill its valence p-shell orbitals.

Hence Na+ ion and F- ion have the same electron configuration as an atom of Ne. (shows F gaining an electron)

Electron affinity

How much an atom wants to gain an electron

Condensed electron configuration

Includes symbol of nearest noble gas, then the following occupied orbitals and valence electrons. Ex: [Ar]4s^2 3p^5

Effective nuclear charge (Zeff) defined (amount of (+) charge experienced by an electron)

Increases across a period (row) because the number of protons increases and the number of core (shielding) electrons stays the same. With a larger Zeff, valence electrons are more strongly attracted to the nucleus and atomic radius becomes smaller [why atomic size decreases across a row]

Ionization Energy (IE): energy required to remove an electron from an atom (further away an electron is from the nucleus, the easier it will be to pull it away)

Increases from left to right across a period (row) and decreases from top to bottom down a group. (an electron that is further away from the nucleus will be easier (take less nrg) to remove and hence will have a smaller ionization energy.

Pauli exclusion principle

No electrons in the same atom can have ID values for all 4 of their quantum numbers. [ex: 2p: n =2 and l =1] [ex: 4d: n=4 and l =2]

Transition metals cations: Nickel (Ni) atom losing 2 electrons:

Notice Nickel loses electrons from 4s orbital rather than 3d orbital.

Degenerate orbitals

Orbitals with the same energy

Hund's rule

The principle that the lowest-energy electron configuration of an atom has the maximum number of unpaired electrons, all of which have the same spin, in degenerate orbitals.

Expect to see radii of elements in the same group to increase with increasing atomic number (in column). Expect to see in a period (row) atomic radii to decrease as atomic number increases across the row (period).

Two interactions to keep in mind: 1) Increasing effective nuclear charge (Zeff). 2) Increasing repulsion between valence electrons (as atomic # increases so do the # valence electrons. More electrons generate electron-electron repulsions (like charges repel) which tends to increase the size of the atom.

For n ≥ 4 (rows), d orbitals fill before f orbitals (because electrons fill lowest energy orbitals first and f orbitals have a higher energy level than d orbitals)

[Ar] 3d^1 4s^2 (Displays order of increasing principle number (n) not the order valence electrons are filled in) (Since 4s fills first, but 3d has a higher quantum number so it is listed first)

Exercise: Write the condensed electron configuration of cobalt (Co, Z = 27)

[Ar] 3d^7 4s^2 (Note: 3d orbital is written first since it is in order of increasing n number, not order in which electrons are filled first = doesn't apply for normal elements: ex: Carbon = 1s^2 2s^2 2p^2 going right across)

Exercise: Write the condensed electron configuration for silver (Ag, Z = 47)

[Kr] 4d^10 5s^1 (a filled set of d orbitals is more stable than a partially filled set, which is the same for Cu Copper, the element above Ag on the table. Therefore, silver (Ag) has 10 electrons in its 4d orbital and only one in its 5s orbital)

Exercise: Using only the periodic table, arrange each set of particles by size from largest to smallest: a) O, P, S. b) Na+, Na, K

a) P > S > O. b) K > Na > Na+

l = 2

d subshell

l = 3

f subshell

Period (row) numbers (for the main group of elements) =

highest n-value

m sub l =

orbital orientation [values -l to +l]

l = 1

p subshell

l = 0

s subshell

l =

shape (of orbital) values [0 to (n-1)]

n =

size and energy (of orbital)

m sub s =

spin/charge [+1/2 or -1/2]

For the main group of elements (group up/down: 1-2, 13-18) the group number =

the number of valence electrons (in the highest orbital).

n-value (same as row/period number for main group) =

total number of orbitals in that energy level. [Ex: n=2 has 2^2 = 4 orbitals total in that energy level: one 2s orbital and three 2p orbitals]

For ions of main group elements: (losing an electron example)

Shows an atom of sodium forms a Na+ ion by losing its single 3s electron

Cations of main group elements are much ______ than their parent atoms.

Smaller

Effective nuclear charge (Zeff)

The attraction towards the nucleus experienced by an electron in an atom; the positive charge on the nucleus is reduced by the extent to which other electrons in the atom shield the electron from the nucleus.

Electron configuration

The distribution of electrons among the orbitals of an atom or ion. ex: 1s^1 (first one = principle quantum number (n); s = subshell; superscripted 1 = shows the number of electrons (1) in the 1s orbital)

Core electrons

The electrons in the filled, inner shells in an atom or ion; they are not involved in chemical reactions.

Electron Affinities (EA) textbook def:

The energy change that occurs when one mole of electrons combines with one mole of atoms or ions in the gas phase.

Ex: Cl^- + e^- ---> Cl^- EA = -349kJ/mol

The negative sign tells us that the energy is lost, or released, when chlorine forms chloride atoms. The more negative the value, the more energy is released when one mole of atoms combined with one mole of electrons to form one mole of anion with a 1- charge.

Valence electrons

The outermost occupied shell of an atom. (elements in the same group (column) have the same # of valence electrons)