atomic trends and bonding/ IMF Flashcards
what is all matter composed of?
atoms
what is a molecule?
when two or more atoms are combined
what is a compound
when two or more molecules are combined
proton charge and weight
at center of atom, weigh 1 amu and have 1+ charge
1.6E-19 eV
electron weight and charge
in electron clouds, -1 or -1.6E-19
neutron weight
1 amu (same as proton_
atomic number (z)
identifies element, number of protons
mass number
P + N
isotopes
same atomic number but differnt number of neutrons
atomic mass
N+P
automic weight
average of all isotopic automic masses
polyatomic ions
contain multiple atoms with different ionic forms
oxyanions
polyatomics that contain oxygen and have suffix ite or ate
bohr model
outdated but explains how electrons exsist in clouds outside the nucleus in shells
ground state
lowest energy level of an electron n=1,2,3… and can jump to higher energy levels
photons
discrete amounts of energy that are emitted as electrons jump to excited states
energy of electromagnetic ratiation equation
E= hf where h is planks constant and f is the frequency of light (which is the speed of light divided by wave length)
E= h*c/ wavelength
rydberg formula
used to determine the energy held by an electron in a certain energy level
change in Energy = R (1/nf^2 - 1/ni^2)
r is constant
Heisenberg principle of uncertainty
cannot know the momentum and exact position of an electron at the same time
principle quantum number
electrons exist in orbitals
denoted as “n” and can be 1 or greater, higher n = higher energy and farther from the nucleus
- associated with rows of the periodic table
orbital
area of space that electrons are likely to be in, can only hold two electrons at a time
pauli exclusion principle
no two electrons in an atom can have the exact same values for all 4 quantum numbers
azimuthal/angular momentum quantum number
describes the subshell
s,p,d,f
magnetic quantum number
describes the spacial orientation of the orbital within its subshell ranging from + and - L. L=0 for s, L=1 for p, L=2 for d L= 3
spin quantum number
describes the spin orientation of the electron either +1/2 or -1/2. all electrons are paired with opposite spin
electron configuration
configuration of electrons in subshells , follow aufbau principle and lower energy levels fill first
hunds rule
electrons prefer to be in orbitals by themselves and will fill each orbital with 1 electron before sharing
where is the s block
the first 2 groups
where is the p block
groups 13-18
where is the d block
transition metals
where is the f block
lanthanide and actinide series at the bottom of the periodic table
valence electrons
outermost electrons that are involved in chemical bonds
which is a lower energy? 3d or 4s
4s so it fills before 3d (according to aufbau principle)
true or false: half filled and full filled orbitals are more stable than ones with 3/4/ or 9 electrons
true! the p and d will steal from s orbital
alkali metals
group 1: high conductivity, luster, duct-ability (drawn into thin wire) , malleability, reactivity, reactive
alkaline earth metal
group 2: harder and less reactive than alkali metals, good conductors
transition metals
groups 3-12. luster, hard, conduct electricity
nonmetals
not shiny, poor electricity conductors
chalcogens
group 16, feature O and S
halogens
group 17, highly reactive
noble gases
group 18, very stable. inert gases. low boiling points
effective nuclear charge
the attraction between the positively charged center of an atom and the electrons surrounding it. decreases down a group and increases to the right. increases with increased number of protons
atomic radius trend
increases down a group and to the left
ionic radius
cations are smaller than anions
ionization energy
the amount of energy required for an atom to give up an electron. increases across a period and up a group
electron affinity
the amount of energy released when an electron is added to an atom. increases up a group and to the right.
electronegativity
tendency of an atom to attract electrons, directly in correlation with effective nuclear charge
how many molecules are in one mole?
6.02E23 molecules
octet rule
each atom will act in a way that allows it to obtain 8 electrons in its valence shell for stability.
is bond formation typically endothermic or exothermic?
it is stable and releases energy so it is exothermic
incomplete octet
special case for hydrogen where it is a filled valence shell with a single bond H-H
others: Boron - 6
he- 2 li-2 be - 4
expanded octet
atom has more than 8 valence electrons. in case of PF5 there are 10
what is unique about octets in the third period?
they readily form expanded octets
intramolecular forces (definition and types)
forces that work inside a molecule to keep it together
ionic and covalent bonds
ionic bonds
form between a metal and a nonmetal, the two species have a large difference in ELECTRONEGATIVITY. form cations and anions
pauling scale
assigns electronegativity values.
fluorine is the most eneg @ 4.0 and the alkali metals have an eneg of around 1.0
** ionic bonds form when the difference in electronegativity between two atoms is 1.7 or greater
covalent bonding
sharing of electrons among nonmetals
- can be equal or unequal sharing
- the difference in electronegativity between the two atoms is small
nonpolar covalent bonds
when electronegativity between atoms sharing electrons are similar or the same and they share electrons equally
- <.5
polar covalent bonds
occur when difference in electronegativity is between .5 and 1.7 and electrons are shared unequally, with the more electronegative atom recieving more electrons and causing dipole moments
dipole moments
unequally sharing of electrons causes partial positive charges on less electronegative atoms pointing toward the more electronegative atom with a partial negative
coordinate covalent bond
when shared electrons are derived from a lone pair of one atom
bond order
number of bonds between two atoms
single, double, triple
sigma bond
forms due to end to end overlap of two atomic orbitals
a single bond forms 1 sigma bond
pi bond
forms when atomic orbitals point above and below the atoms making them parallel.
a double bond contains a pi and a sigma
bond length
inversely related to bond order. single bonds are the longest and triple bonds are the shortest
bond energy
energy required to break a covalent bond. the higher the bond energy, the stronger the bond
- increases with bond order
metallic bonding
an intramolecular force less common than ionic and covalent. occurs when metal atoms are joined together and electrons become delocalized in a sea of electrons that freely move through a solid and allows for electrical conduction and heat
intermolecular forces
attractive forces between molecules that are notably WEAKER than intramolecular forces and can be broken with heat . bonds driven by charge
types: london, dipole, hydrogen, ion-dipole
- stronger polarity = stronger intermolecular forces
- determine behavior and melt/boiling points
london-dispersion forces
weakest intermolecular force caused by random dipole moments, fluctuations in polar and nonpolar bonds
- larger compounds are more likely to experience london dispersion forces
dipole-dipole interactions
intermolecular force that is stronger than london dispersion. dipoles are stable (polar bonds only)
hydrogen bonding
between a hydrogen bound to F, O, N and a lone pair on another F, O, or N
ion-diple forces
strongest intermolecular force. ions are most polar and occur when two molecules have charge
ex: peptide hormone signalling with a basic aminoacid
lewis structure
establish relationships between atoms and their valence electrons
- valence electrons = group number (except transition metals)
- atom with lowest electronegativity goes in the middle
resonance structures
compounds with lone pairs that can move bonds to different positions
- can only shift lone pairs and double bonds, never single bonds
- delocalization of electrons increases stability of structure
formal charge
disparity between the number of valence electrons an atom has and the electrons it actually has.
formal charge = VE - bonds - lone pair electrons
orbital hybridization
mixing atomic orbitals to form hybrid orbitals
what does a triple bond consist of
2 pi and 1 sigma bond
two regions of hybridization (triple bond) yields what?
sp
three regions of hybridization (double bond) yields what?
sp2
4 regions of hybridization (single bond) yields what?
sp3
VSPER theory
uses lewis structures to determine shapes of molecules due to electron repulsion
- use number of lone pairs and bonded atoms to predict bond angle and shape
electronic geometry
takes into account lone pairs and bonded atoms to predict shape
molecular geometry
only takes into account boned atoms when predicting shape
tetrahedral molecular shape
4 electron rich regions and 4 bonded atoms
- CH4
- bond angle is 109.5
trigonal pyrimidal molecular shape
4 regions of electron rich. 3 bonded atoms and 1 lone pair
- bond angle is 107
nh3
bent molecular shape
2 bonds and 2 lone pairs, bond angle is 104
ex: h2o
linear molecular shape
2 bonds no lone pairs, bond angle is 180
trigonal planar molecular shape
3 bonds no lone pairs, bond angle 120
ex bf3
