(l) - shape of atomic orbitals (sometimes called a subshell) l=0,1,2,3...(highest number it can be= n-1)
Magnetic quantum number
(ml) - orientation of the orbital in space relative to the other orbitals in the atom. ex. l=0 then ml=0... l=1 then ml can = -1,0,1, etc. max # of ml= 2l+1
# of nodes= n-1 or l
Shell shaped area where electron can't be found
dumbbell shaped (can have different orientation on x y and axis)
clover shaped: can have different orientation on x y and z axis
Pauli exclusion principle
cannot have more than on electron with the same set of 4 numbers, therefore each shell can only hold two electrons with opposite spin.
atoms with more than one electron
Electron correlation problem
since the electron pathways are unknown, the electron repulsions cannot be calculated exactly
A 2s electron penetrates to the nucleus more than on in the 2p orbital
This causes an electron in a 2s orbital to be attracted to the nucleus more strongly than an electron in a 2p orbital.
Thus, the 2s orbital is lower in energy than the 2p orbitals in a polyelectronic atom.
As protons are added to the nucleus to build up the elements, electrons are similarly added to hydrogen-like orbitals.
The lowest energy configuration for an atom is the one having the max number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate (same energy) orbitals.
Energy required to remove an electron from a gaseous atom or ion.
Ionization energy increases left to right and as you go up the periodic table.
(why? the electrons being removed are farther from the nucleus)
Energy change associated with the addition of an electron to a gaseous atom.
Electron Affinity increases from left to right and going up the periodic table.
Atomic radius decreases going left to right and increases going down a group.
Group 1A name
Group 2A name
Alkaline earth metals
Middle elements name
Group 7A name
Group 8A name
Most chemically reactive of the metals
atomic mass unit
mass of an atom compared to 12C
a mole= 6.022x10^23
bonds formed between atoms by sharing electrons
Bonds formed due to force of attraction between oppositely charged ions.
atom or group of atoms that has a net positive or negative charge.
positive ion; lost electron(s)
negative ion; gained electron(s)
Composed of two elements
ionic and covalent compounds included
1. cation is named first and anion second
2. anion is named by taking the root of the element name and adding -ide.
3. type (II)- metals form more than one type of positive charge; charge on metal ion must be specified. (transition metal cations usually require a Roman numeral.)
multiple elements... just need to be memorized.
Hydrogen sulfate (or bisulfate)
NCS(-) or SCN(-)
Hydrogen Carbonate (or bicarbonate)
ClO(-) or OCl(-)
Binary Covalent Compounds
Formed between two nonmetals
1. first element is named first
2. second element is named as an anion
3. prefixes are used to denote the numbers of atoms present.
4. prefis mono- never used for naming first element.
Hydrogen appears first in the formula
Molecule with one or more H+ ions attached to an anion
Naming Acids without oxygen
acid is name with the prefix hydro- and the suffix ic
Naming Acids with oxygen
suffix -ic is added to the root name if the anion name ends in -ate
suffix -ous is added to the root name if the anion name ends in -ite
Forces that hold groups of atoms together and make them function as a unit.
Polar covalent Bond
unequal sharing of electrons between atoms in a molecule.
Results in a charge separation in the bond
ability of an atom in a molecule to attract shared electrons to itself.
electronegativity increases left to right and up the periodic table.
values range from 0.7 to 4.0
Difference in electronegativity between atoms
0-0.4 : Covalent bond
0.4-2.0 : Polar Covalent
Property of a molecule whose charge distribution can be represented by a center of positive charge and a center of negative charge.
Arrow represents dipole moment (points to the negative charge center)
Electron configurations in Stable Compounds
two nonmetals react to form covalent bond and share electrons so that it completes the valence electron configuration of both atoms
nonmetal and representative-group metal react to form binary ionic compound, form so that valence electron config. of the nonmetal achieves the electron configuration of the next noble gas atom. valence orbitals of the metal are emptied.
series of ions/atoms containing the same number of electrons.
To break bonds, energy must be added to the system (endothermic)
To form bonds, energy is released (exothermic)
the change in energy that takes place when separated gaseous ions are packed together to form an ionic solid.
Lattice Energy= k(Q1Q2/r)
Q1 and Q2= charges on the ions
R=shortest distance between the centers of the cations and anions
Ionic compound characteristic
any compound that conducts and electric current when melted will be classified as ionic.
Localized Electron Model
Electron pairs are assumed to be localized on a particular atom or in the space between two atoms:
Lone pairs- electrons localized on an atom
Bonding pairs- pairs of electrons found in the space between the atoms.
Shows how valence electrons are arranged among atoms in a molecule.
Hydrogen forms stable molecules where it shares two electrons
Elements form stable molecules when surrounded by eight electrons
Single Covalent Bond
covalent bond in which two atoms share one pair of electrons
Double covalent bond
covalent bond in which two atoms share two pairs of electrons
Triple Covalent Bond
covalent bond in which two atoms share three pairs of electrons
Steps for writing Lewis Structures
1. sum the valence electron from all atoms
2. use pair of electrons to form a bond between each pair of bound atoms.
3. Arrange the remaining electrons to satisfy the octet rule (or duet for hydrogen)
Boron exception to octet rule
tends to form compounds in which the boron atom has fewer than eight electrons around it.
Exceeding octet rule
sometimes necessary to exceed octet rule for several of the third-row(or higher) elements, place the extra electrons on the central atom.
More than one valid Lewis structure can be written for a particular molecule.
used to evaluate nonequivalent lewis structures
atoms try to achieve formal charges as close to zero as possible.
Formal Charge= (# valence e on free atom) - (# valence e assigned to the atom in the molecule)
lone pairs- count both of the electrons
bonds- count one of the electrons
valence Shell Electron-Pair Repulsion: the structure around a given atom is determined principally by minimizing electron pair repulsions.
Electronic Geometry arrangements
# of electron pairs Arrangement of pairs
3 Trigonal Planar
5 Trigonal bipyramidal
mixing of the native atomic orbitals to form special orbitals for bonding. examples... sp, sp2, sp3, dsp3, etc
electron pair is shared in an area centered on a line running between the atoms
forms double and triple bonds by sharing electron pair(s) in the space above and below the sigma bond. Uses the unhybridized p orbitals.
Using the localized Electron Model
1. Draw the Lewis structure(s)
2. Determine the arrangement of electron pairs using the VSEPR model.
3. Specify the hybrid orbitals needed to accommodate the electron pairs.