Unit 5 Flashcards
emission of energy by atoms
when atoms receive energy, they become excited. they can release the energy by emitting a light. the emitted energy is carried away by a photon
the energy of the photon corresponds exactly to the energy change of the emitting atom
high energy photons correspond to short wavelength light. low energy photons correspond to long wavelength light
the photons of red light have less energy than the photons of blue light cuz red light has a longer wavelength than blue light
energy levels of hydrogen
when we study the photons of visible light emitted, we only see certain colors
only certain types of photons r produced
because only certain photons r emitted, only certain energy changes r occurring
so, hydrogen atoms must have certain discrete energy lveles
we say the energy levels of H r quantized, that is, only certain values are allowed
energy levels of all atoms r quantized
hydrogen orbitals
the probability map is called an orbital. the orbital shown in the pic is called the 1s orbital & describes the ground(lowest) state of energy for H
size of orbital is defined by a sphere that contains 90% of the total electron probability
principal energy levels
discrete energy levels
designated by whole #’s symboized by n; n can equal 1, 2, 3, 4,… level 1 corresponds to n = 1, etc.
energy of teh level increases as value of n increases
describe size & shape. the s orbital is spherical. level 1 is smaller than level 2, which is smaller than level 3.
each principal energy level contains 1+ types of orbitals, called sublevels
sublevels
the # of sublevels present in a given principal energy level equals n.
Ex: level 1 contains one sublevel (1s); level 2 contains 2 sublevels (2 types of orbitals), the 2s orbital and the three 2p orbitals; and so on.
These are summarized in the pic. the # of each type of orbital is shown in parentheses.
n value always used to label orbitals of a given principal energy level & followed by letter tht indicates the type (shape) of the orbital
wave mechanical model
orbital
can be empty or can contain 1/2 electrons, but never more than 2
if 2 electrons occupy same orbital, must have opp spins
shape of an orbital doesn’t indicate the details of electron movement. merely indicates the prob distribution for an electron residing in that oribtal
spin
each electron appears to spin like a top on its axis
can only spin in 1 direction. we represent spin with up & down arrows.
Pauli Exclusion Principle
an atomic orbital can hold a max of 2 electrons & those 2 electrons must have opp spins
electron configuration
principal energy level followed by sublevel; # of electrons in the orbital placed as superscript
Ex: 1s1
orbital diagram
box diagram
valence electrons
the electrons in the outermost (highest) principal energy level of an atom
these r the electrons involved in bonding of atoms to each other
the atoms of elements in the same group have the same # of electrons in a given type of orbital, except that the orbitals are in diff princ energy levels. (except He, which is 1s)
elements w/ same valence electron arrangement show very similar chem behavior
orbital filling
in a principal energy level that has d orbitals, the s orbital from the next level fills before the d orbitals in the current level. That is, the (n + 1)s orbitals always fill before the nd orbitals.
Ex: the 5s orbitals fill for rubidium & strontium before the 4d orbitals fill
lanthanide series
after lanthanum, which has configuration [Xe]6s25d1
a group of 14 elements
corresponds to the fillling of the seven 4f orbitals
actinide series
after actinium, [Rn]7s26d1
14 elements
corresponds to filling of seven 5f orbitals
bond
a force that holds 2+ atoms together & makes them function as a unit
In water, the fundamental unit is the H-O-H molecule, which is held together by the two O-H bonds
ionic compounds
formed when an atom that loses an electron relatively easily reacts w/ an atoms that accepts an electron
when metal reacts w/ non-metal
resulting bonds = ionic bonds
electrons transferred
covalent bond
when 2 similar atoms form a bond, the electrons r equally attracted to the nuclei of the 2 atoms
electrons shared by nuclei
Ex: diatomic hydrogen H-H
polar covalent bonds
b/w the extremes
atoms r not so diff that electrons r transferred, but diff enough tht unequal sharing of the electrons results
electronegativity
the unequal sharing of electrons b/w 2 atoms is described by this property
the relative ability of an atom in a molecule to attract shared electrons to itself
the higher the electronegativity value, the closer the shared electrons tend to be to that atom when it forms a bond
F has highest electronegativity & so always polar
increasing electronegativity as goes right and up
polarity
depends on the diff b/w the electronegativity values of the atoms forming teh bonds
if similar electronegativites, the electrons r shared almost equally & bond shows little polarity
if very diff electronegativies, very polar bond is formed
stable electron configurations
representative (main-group) metals form ions by losing enough electrons to attain the configuration of the previous noble gas tht occurs before the metal
nonmetals form ions by gaining enough electrons to attain the configuration of the next noble gas
when a non-metal & a Group 1, 2, or 3 metal react to form a binary ionic bond, the ions form so that the non-metal completes the valence-electron config of the next noble gas & the metal empties the valence orbitals to achieve the config of the prev. noble gas
when 2 non-metals react to form a covalent bond, share electrons in a way that completes the valence-electron configuration of both atoms
Lewis structures
bonding involves just the valence electrons. this structure is a representaiton of a molecule tht shows how the valence electrons r arranged among the atoms in the molecule
H - duet rule
He - doesn’t form bonds cuz valence orbital already filled
2nd row nonmentals C thru F - octet rule
Ne - doesn’t form bonds cuz already has an octet of valence electrons
- Obtain sum of the valence electrons from all of the atoms.
- use 1 pair of electrons to form a bond b/w each pair of bound atoms. Use line.
- arrange the remaining electrons to satisfy the duet rule for hydrogen & the octet rule for each 2nd-row element (may need to guess & check w/ double bonds, etc.)
duet rule
H forms stable molecules where it shares 2 electrons
octet rule
2nd-row nonmetals C through F form stable molecules when r surrounded by enough electrons to fill the valence orbitals - that is, the one 2s and the three 2p orbitals.
8 electrons required to fill these orbitals
configuration tips
noble gas valences from top to bottom:
1s2
2s22p6
3s23p6
4s24p6
5s25p6
6s26p6
