Chemistry Ch. 4

anions

nonmetals form - charged ions

multi electron atom

not degenerate, energy depends on the values of l

main group elements

properties are largely predictable based on position on periodic table

transition elements (transition metals)

properties are less predictable based simply on their position in the periodic table

irregularities in transition elements

4s and 3d are very close in energy, therefore causing...

d sublevel

5 orbitals, 10 electrons

Trends involving electron affinity

-most groups don't exhibit any definite trend. -among group 1A metals, however, electron affinity becomes more positive as we move DOWN THE COLUMN (adding electron becomes less exothermic) -electron affinity becomes more negative (adding electron becomes more exothermic) as we move to the right across a period (row) in the periodic table.

why do transition metals have unexpected filling behavior?

-ns and (n-1) orbitals are very close in energy and, depending on exact configuration, can vary in relative ordering -as (n-1)d orbitals begin to fill in the first transition series, the increasing nuclear charge stabilizes the (n-1)d orbitals relativ to the ns orbitals/ This happens because the (n-1)d orbitals are not the outermost (or highest n) orbitals and are therefore nt effectively sheilded from the increasing nuclear charge by the ns orbitals.

s sublevel

1 orbital, 2 electrons

p sublevel

3 orbitals, 6 electrons

f sublevel

7 orbitals, 14 electrons

Exceptions to trends in first ionization energy

Boron is smaller than beryllium. Aluminum is smaller than gallium This is because of the change going from the s block to the p block. Oxygen is smaller than nitrogen. S and Se follows this as well. This is because of the repulsion between electrons when they occupy the same orbital.

ionization energy on periodic table

INCREASES as we across a period (move to the right) DECREASES as we move down a column

Why does metallic character decrease across a row ad increase down a group?

Ionization energy increases increases and electron affinity becomes more negative, therefore elements on the left side of the periodic table are more likely to lose electrons than elements in the right side of the periodic table (which are more likely to gain them)

Why would the second ionization energy of Na be larger than the second ionization energy of Mg?

Mg has 2 valence electrons, so the amount of energy it takes to remove the second electron is less than it is for Na. Na has only one valence electron. Removing a second electron from Na would be removing a core electron, which involves a tremendous amount of energy. thus resulting in the second ionization energy to be higher.

Metallic character trend

Moving RIGHT ACROSS A ROW, metallic character DECREASES

react to attain a noble gas configuration

elements without noble gas configuration...

Electron affinity (EA)

The energy change associated with the gaining of an electron by the atom in the gaseous state. Usually negative because an atom or ion usually releases energy when it gains an electron. Ex: Cl (g) + 1e- --> Cl- (g) EA = -349 kJ/mol

electron configuration of anions

add the number of electrons indicated by the magnitude of the charge of the anion Ex: F: 1s2 2s2 2p5 vs F-: 1s2 2s2 2p6

Diamagnetic

an atom or ion in which all electrons are paired is not attracted to an external magnetic field-it is instead slightly repelled Ex: zinc and zinc2+

paramagnetic

an atom or ion that contains unpaired electrons is attracted to an external magnetic field Ex: Ag

relationship between anions and their neutral atoms

are much larger than their corresponding neutral atoms Ex: Cl: [Ne] 3s2 3p5 vs Cl- [Ne] 3s2 3p6 the extra electron increases the repulsions among the outermost electrons, therefore making Cl- larger

relationship between cations and their atoms

are much smaller than their corresponding neutral atoms because it has lost electrons, making it smaller Ex: Na: [Ne]3s1 vs Na+: [Ne]

ions

atoms (or groups of atoms) that have lost or gained electrons

atomic radius

average bonding radii determined from the measurements on a large number of elements and compounds half the distance between nuclei

effective nuclear charge

average or net charge shielded by other electrons Zeff = Z (actual nuclear charge) - S (charge screened by other electron)

isoelectronic series

can see trends in ionic size by examining the radii of this it is ions with the same number of electrons

penetration

describes how one atomic orbital can overlap spatially with another, thus penetrating into a region that is close to the nucleus (and therefore less sheilded from nuclear charge) it gets a lower potential energy because it experiences a greater nuclear charge

coulomb's law

describes the attractions and repulsions between charged particles high pe causes like charges repel one another lowering pe causes unlike charges attract one another magnitude of interaction between charged particles increases as the charges of the particles increases. electron with 1- charge is more attracted to nucleus of 2+ than 1+

atomic radius trend

down a column, increases right across a period or row, decreases due to valence electrons

family/group

each column within the main group regions of the periodic table has similar properties

core electrons

efficiently shield electrons in the outermost principle energy level from nuclear charge, but the outermost electrons do not efficiently shield one another from nuclear charge

halogens

electron configuration: ns^2np^5 1 electron short of noble gas configuration 1- charge

aufbau principle

electron configurations can be built-up by filling orbitals from low to high energy

core electrons

electrons in complete principal energy levels and complete d and f sublevels do not form bonds nor react heavily

why does the ionization energy increase as we move right across a row (or period)?

electrons in the outermost principal energy level generally experience a greater effective nuclear charge (Zeff)

ionization energy

energy required to remov an electron from an atom or ion in the gaseous state always positive because it is endothermis and removing an electron always takes energy

second ionization energy

energy required to remove the second electron Ex: sodium Na+ --> Na2+ (g) + 1e- IE2 = 4560 kJ/mol it is the energy required to remove one electron from Na+ NOT to remove two electrons from Na

the row number of a main group element

equal the highest principal quantum number of that element

quantum mechanical theory

explains the electronic structure of atoms this in turn determines the properties of these atoms

transition metals

form various ions with different charges

how to deduce electron configuration of main group monotomic ion

from the electron configuration of the neutral atom and the charge of the ion

metals

good conductors of heat and electricity malleability ductility lose electrons in chemical changes --transition metals too but not to obtain noble gas configuration

alkali metals

group 1A elements may react violently to lose ns^1 to form ions with 1+ charge

alkaline earth metals

group 2A elements ns^2 electron configuration loses two electrons to form ions with 2+ charge

noble gases

group 8A most unreactive elements in the entire periodic table 8 valence electrons

Ag Cu Cr Mo

have half filled or filled d orbitals (you take one from the s orbital the fill up the d more)

stay about constant across each row

instead of decreasing in size, the radii of transition metals...

ground state

lowest energy state electron configurations are normally given in this -- atom with all electrons in the lowest energy orbitals possible

charge is the group number - 8

main group elements that form anions with predictable charge

charge is the group number

main group elements that form cations with predictable charge

metaloids

many are semiconductors (because of intermediate and highly temperature dependent electrical conductivity)

nonmetals

may be solid, liquid, or gas at room temperature poor conductors of heat and electricity gain electrons in chemical reactions

cations

metals form + charged ions

valence electrons

most important in chemical bonding those in the outermost principle energy level

why does ionization energy decrease down a column?

n increases down a column, and these electrons are farther away from the positively charged nucleus, and are therefore held less tightly, as we move down a column. This results in a lower ionization energy as we move a column.

pauli exclusion principle

no two electrons in an atom can have the same four quantum numbers each orbital can have a maximum of only 2 electrons with opposing spins arrow represents electron spin

periodic property

property predictable based on elements position within the periodic table

electron configurations for transition metal cations

remove electrons in the highest n-value orbitals first, even if this does not correspond to the reverse order of filling Ex: V: [Ar] 4s2 3d3 vs V2+: [Ar] 4s0 3d3 (same as [Ar] 3d3)

sheilding

repulsion of one electron by other electrons as shielding that electron from the full effects of the nucleus inner electrons shield outer electron from full nuclear charge

electron configurations

shows the particular orbitals that electrons occupy for that atom

orbital diagram

similar to electron configuration but symbolizes the electron as an arrow and the orbital as a box

Electron configurations for cations

subtracting the number of electrons indicated by the magnitude of the charge Ex: Li: 1s2 2s1 vs Li+: 1s2

the number of valence electrons they contain

the chemical properties of elements are largely determined by...

first ionization energy

the energy required to remove the first electron Ex: sodium Na (g) --> Na+ (g) + 1e- IE1= 496kJ/mol

greater nuclear charge in atoms or ions with the same number of electrons

the smaller the atoms or ion because if there is a greater number of protons, there is more of a pull from the protons to the electrons, causing the atom or ion to be smaller

1) Coulomb's Law 2) sheilding 3) penetration

what causes sublevels to split

the overall potential energy of the electrons that occupy that level is particularly low

when a quantum number is completely full...

hund's rule

when filling degenerate orbitals, electrons fill them singly first, with parallel spins

degenerate

when orbitals have the all have the same energy EX: H

periodic law

when the elements are arranged in order of increasing mass, certain sets of properties recur periodically

same number of valence electrons

why elements in a column have similar chemical properties