Test Four – Chemistry Flashcard

bonding lowers the
potential energy between positive and negative
Atomic Size –
related to the distance between atoms in a sample of the element
Transition Metals:
the trend in size changes slightly since electron/electron repulsions counteract the decrease in size as we go across a period
Ionization Energy –
The energy required to remove an electron from an atom in the gas phase
As atomic radius decreases, the ionization energy _____
increases
Electron Affinity –
The energy change for a process in which an electron is acquired by the atom in the gas phase or “how bad an atom wants an electron.”
Isoelectronic Ions:
Ions that have the same number of electrons (but different number of protons). N-3, O-2, F-, Na+,and Mg+2
Chemical Bond:
When a chemical reaction occurs between two atoms, their valence electrons are reorganized so that a net attractive force occurs between atoms.
Ionic bond:
forms when one of more valence electrons is transferred from one atom to another. Produces a positive and negative ion. The “bond” is the attraction between the ions
Covalent bond: forms by
the sharing of valence electrons between atoms.
Metal with nonmetal bonding –
electron transfer and ionic bonding
Covalent bonding –
occurs most commonly between non mental atoms
Metal with metal bonding –
electron pooling and metallic bonding
Metallic bonding –
metal atoms sharing valence electrons but not by covalent bonding
Octet rule –
when atoms bind, they lose, gain, or share electrons to attain a filled outer level of eight electrons or two for H and Li
The lattice energy
the enthalpy change that occurs when 1 mol of ionic solid separates into gaseous ions
Ionic solids exist only because
the lattice energy exceeds the energy required for the electron transfer
Formation of a covalent bond always results in
a greater electron density between the nuclei
Bond order –
the number of electron pairs being shared by a given pair of atoms
Single bond –
the most common band and consists of one bonding pair of electrons
Double bond –
consists of two bonding pairs of electrons, four electrons shared between two atoms so the bond order is 2
Bond energy –
the energy needed to overcome this attraction and is defined as the standard enthalpy change for breaking the bond in 1 mol of gaseous molecules
Bond length –
the distance between the nuclei of the two bonded atoms
Bond order to directly related to _____and inversely related to ______
bond energy; bond length
Most covalent substances have low electrical conductivity because
their electrons are localized and ions are absent
Electronegativity –
the relative ability of a bonded atom to attract shared electrons
An important use for electronegativity is
determining an atom’s oxidation number
When atoms of different electro negativities form a bond, the bonding pair is shared ____
unequally
Polar covalent bond –
unequal distribution of electron density
Electronegativity difference –
the difference between the electronegativity values of the bonded atoms
Electronegativity difference is directly related to a bond’s _____
polarity
As the ?EN decreases, the bond becomes more ____
covalent
Bond Dissociation Enthalpy:
the enthalpy change for breaking a bond in a molecule with the reactants and products in the gas phase
Polar Bond:
When the bond between two atoms has a positive and negative end or pole.
Hydrogen atoms from ____ bonds
one
Carbon atoms from _____ bonds
four
Nitrogen atoms form _____ bonds
three
Oxygen atoms form ______ bonds
four
Fluorine is always a ______
surrounding atom
Resonance structures –
have the same relative placement of atoms but different locations of bonding and lone pairs
Resonance hybrid –
average of the resonance forms
Partial bonding in resonance hybrid, often leads to ______
fractional bond orders
Formal charge –
change an atom would have if the bonding electrons were shared equally
Electron deficient –
have fewer than eight electrons around the central nucleus
Free radicals –
species that contain a lone (unpaired) electron, which makes them paramagnetic and extremely reactive
Expanded valence shells occur only with _______ because they have d orbital available
nonmetals form period 3 or higher
VESPR Theory –
to minimize repulsions, each group of valence electrons around a central atom is located as far as possible from the others
Molecular shape –
three dimensional arrangement of nuclei joined by the bonding groups
The electron group arrangement is defined by
the bonding and nonbonding electron group
The Molecular shape is defined by
the relative positions of the nuclei, which are connected by the bonding groups only
Bond angle –
angle formed by the nuclei of the two surrounding atoms with the nucleus of the central atom at the vortex
Linear shape –
AX2 108o
Trigonal Planar–
AX3 120o
Tetrahedral –
AX4 109.5o
Trigonal Bipyramidal –
AX5 90o & 120o
Octahedral –
AX6 90o
Lewis electron dot symbols: developed by
Gilbert Newton Lewis
Bond pair:
The e- involved in the covalent bond.
Lone pair:
The e- not involved in bonding also called nonbonding e-.
Octet Rule:
The tendency of molecules and polyatomic ions to have structures in which eight e- surround each atom
Valence shell electron-pair repulsion is a method for
predicting the shapes of covalent molecules and ions.
Electron-pair geometry:
the geometry taken up by ALL the valence e- pairs around a central atom
Molecular geometry describes
the arrangement in space of the central atom and the atoms directly attached to it.
Valence Bond Theory created by
Linus Pauling
Bonding between the two atoms occurs when
the e- clouds on the two atoms interpenetrate or overlap
Orbital overlap increases the probability of
finding bonding e- in this region of space.
The idea that bonds are formed by overlap of atomic orbitals is the basis for
valence bond theory.
The covalent bond that arises from the overlap of the two s orbitals (1 from each H) is called
a sigma bond.
Liquids and solids resist
compression
Dipole-Dipole Attraction:
when one polar molecule encounters another, the positive end of one is attracted to the negative of the other, and via versa.
Hydrogen Bond:
The strong attraction between an electronegative atom with a lone pair and the hydrogen atom of the N–H, O–H, or F–H bond.
Dipole/Induced Dipole Forces:
polar molecules can induce a dipole in a molecules that do not have a permanent dipole.
As the water molecule approaches the O2 molecule a dipole is induced, this is referred to as _____
polarization.
The degree to which the e- cloud distorts is called ______
polarizability.
The larger the molar mass the greater the _____
polarizability of the molecule.
Dipole/Induced Dipole Forces are weaker than _____
electrostatic or dipole/dipole interactions.
Induced Dipole/Induced Dipole Forces are often referred to as _______
London dispersion forces
London forces arise between
all molecules both polar and non-polar
London dispersion forces are the only intermolecular forces that
allow non-polar molecules to interact
Vaporization:
or evaporation is the process in which a substance in the liquid phase becomes a gas
The standard molar enthalpy of vaporization, ?vapHo (kJ/mol):
the energy required to vaporize a sample.
Condensation:
when a molecule loses efficient energy to reenter the liquid phase. (exothermic)
Equilibrium Vapor Pressure:
the pressure exerted by the vapor in equilibrium with the liquid phase.
The tendency of its molecules to escape from the liquid phase and enter the vapor phase, referred to as _______
volatility.
Critical Point:
when a specific temp and pressure are reached, the interface between the liquid and the vapor disappears.
Critical Temperature:
the temperature at which this phenomenon is observed, and the corresponding pressure, critical pressure
Viscosity:
the resistance of liquids to flow
Surface Tension:
the energy required to break through the surface or to disrupt a liquid drop and spread the material out as a film.

Get instant access to
all materials

Become a Member