Terms in this set (324)
[blank] are microscopic particles that represent the smallest unit of matter with the properties of an element
[blank] are the fundamental building blocks of ordinary matter.-Free atoms are rare in nature.
are combinations of atoms bound together in specific geometric arrangements.
are those that are independent of the amount of substance -Example: density
are those at are dependent on the amount of substance-Examples: volume, mass
has no uncertainty, and thus do not limit the number of significant figures in any calculations. They originate from three sources:-Accurate counting of discrete objects-Defined quantities -Integral numbers that are part of an equation
refers to how close the measured value is to the actual value.
refers to how close a series of measurements are to one another or how reproducible they are.
John Dalton and the Atomic Theory
Each element is composed of tiny, indestructible particles called atoms.
2.All atoms of a given element have the same mass and other properties that distinguish them from the atoms of other elements.
3.Atoms combine in simple, whole-number ratios to form compounds.
4.Atoms of one element cannot change into atoms of another element. In a chemical reaction, atoms change they way they are bound together with other atoms to form a new substance.
derived from a single atoms
a group of atoms with an overall charge
developed the periodic law:
-When the elements are arranged in order of increasing mass, certain sets of properties recur periodically.
Alkaline earth metals
average atomic mass
for an element can be found on the periodic table along with the atomic number.
-Also called atomic mass, atomic weight, or average atomic weight
Ionic bonding occurs when ions assemble into an extended array called a [lattice] and are held together by the attraction between oppositely charged ions.
In [blank] bonds, bonding atoms share electrons in pairs.-The shared electrons have lower potential energy than they would in the isolated atoms because they interact with the nuclei of both atoms.
-[blank] bonded atoms are called molecules.
-Describes compounds in terms of their atomic substituents
-The complete inventory of atoms
-The lowest whole number ratio of atoms in a compound
Exist in nature with single atoms at their basic units (Examples: He, Ne, Ar)
-Do NOT normally exist with single atoms at their basic units
-Two or more atoms of the element bond together, and they exist as molecules (O2, P4, S8)
-Diatomic elements: H2, O2, F2, Br2, I2, N2, Cl2
The basic unit of an ionic compound is the [formula] (the smallest electrically neutral collection of ions).
are ions composed of two or more atoms that are covalently bonded. The charge is a property of the whole ion, not just of one of the atoms in it.
Pressure 1 and 2 need same unit but volume has to be in kelvins.
Conversion unit for Psi and atm
14.7 psi = 1 atm
How many liters are in a milliliter
1 L/1000 mL
Conversion unit for kPa and torr
101.3 kpa = 760 torr
Conversion unit for Celsius to kelvin
K = °C + 273.15
Conversion unit for kelvin to Celsius
Gay Lussac's Law
is one way to express how much of an element is in a given compound.
can be used to determine the empirical formula of a compound.
With the addition of the molar mass, you can also determine the molecular formula of the compound.
is a good way to obtain empirical formulas for unknown compounds, especially those containing carbon and hydrogen
is a term used to describe quantitative relationships in chemistry.•Any chemistry question that asks "How much?" of a particular substance consumed or formed in a chemical reaction is a [blank] problem.•All [blank] problems must start with a balanced chemical equation.
Many factors determine the amount of desired product actually produced in a reaction.
-Temperature of the reaction
-The possibility of side reactions
-Further reaction of the initial product to produce new product
The ideal amount of product that a reaction can make mathematically.
The amount the reaction produces in the laboratory.
The ratio of actual yield to theoretical yield.
Component present in greater amounts (usually much greater amounts)•Water is the [blank] in aqueous solutions.
•Component present in lesser amounts; the minor component
•In all solutions, solute dissolves in solvent.
-The measure of the amount of solute dissolved in the solvent
-The larger the value the more solute is dissolved per unit amount of solvent.
A unit of concentration in moles of solute per liter of solution.
-The process of lowering the concentration of a solution by adding more solvent.
-The number of moles of solute does not change during a [blank].
Substances that dissolve in water to produce a conductive solution.
Dissociate completely (soluble ionic compounds, strong acids, and strong bases)
Dissociate only a little bit (weak acids and weak bases)
Do not dissociate (covalent molecules)
•The process of a substance dissolving in solvent.•The molecules of the solute are surrounded by molecules of solvent in a process called solvation.
The molecules of the solute are surrounded by molecules of solvent in a process called
When ionic compounds dissolve, they
dissociate or break apart into their constituent ions.•The associated reaction is called a dissociation reaction.
Ionic compounds (strong electrolytes)NaCl(s) --> Na+(aq) + Cl-(aq)
Uneven distribution of electrons within the water molecule cause the O side to have a partial negative charge and the H side to have a partial positive charge.
When covalent compounds dissolve, the molecules themselves stay [blank].
Covalent compounds (nonelectrolytes)C12H22O11(s ---> C12H22O11(aq)Figure
A [blank] reaction is one in which the mixing of two solutions results in the formation of a solid, or precipitate
Aqueous chemical reactions can be written as a [blank]
The complete formula for each compound is shown.
Complete ionic equation
Emphasize dissociation of reactants and products by writing all strong electrolytes in their dissociated forms.
REMEMBER: Only aqueous ionic substances dissociate. Gases, liquids, solids, and covalent aqueous species NEVER dissociate.
Net ionic equations
Simplify the total ionic equation to only the components that undergo direct change in the reaction by omitting the spectator ions or species that appear exactly the same on both sides of the chemical equation.
-Force exerted per unit area•A low density of gas particles results in low pressure.•A high density of gas particles results in high pressure.
Conversion unit for atm and torr
1 atm = 760 torr
Conversion unit for mmHg and kPa
760 mmHg = 101.3 kPa
Conversion unit for moles to liter
1 mol/22.4 liter
Change in moles conversion
∆n = nF - nI
Ideal Gas Law
PV = nRT
Ideal gas constant: 0.08206 Lxatm/molxK
Conversion unit for atm and kPa
1 atm = 101.3 kPa
Combined Gas Law
P₁V₁/n₁T₁ = P₂V₂/n₂T₂
P₁V₁/T₁ = P₂V₂/T₂
1) Write balanced equation
2) Note: You can use coefficients to change from liter to liter like with mole to mole
3) You might need molar mass
4) You can use Pv = nRt to find moles for a conversion
5) To find limiting reactants use n = Pv/RT then divide by coefficient then use PV=nRt to find volume.
Units of STP
T = 0° = 273K
P = 1 atm
1 mol gas = 22.4 L
Molarity conversion units
M = mol/L
You can insert molarity for mol
Molar mass formula of a gas
M = mRT/PV
M = molar mass
m = grams
R = ideal gas constant
T = temperature
P = pressure
V = volume
Molar mass = g/mol
Density: d = m/v
Gas density formula
d = PM/RT
P = pressure (atm)
R = ideal gas constant
T = temperature (K)
M = molar mass (g/mol)
You can also use: 1 mol gas = 22.4 L
Density relationship with pressure formula
d₂/d₁ = P₂/P₁
Density relationship with temperature formula
d₂/d₁ = T₁/T₂
Density relationship with molar mass formula
d₂/d₁ = M₂/M₁ (Use molar masses of elements and current densities)
Total pressure formula
Partial pressures of elements equal total pressure
Xₙ = nA/nT or XA = PA/PT
nA = moles of element
nT = total moles
PA = partial pressure
PT = total pressure
Partial Pressure of Two elements
Pₙ = XₙPT
n = specified element
PT = total pressire
Xₙ = mole fraction
Mole fraction of an element
Xₙ = Pₙ/PT
n = element
Xₙ = mole fraction
Pₙ = partial pressure
PT = total pressure
Partial pressure of an element
Pₙ = (Xₙ)(PT)
Mole fractions of elements
X₁ + X₂ + X₃ =1
All mole fractions sum to 1
Partial pressure of elements
Pₙ = NₙRT/v or PT = nTRT/V
You might have to convert to moles then plug into, PₙV = nRT to get partial pressure
Collecting Gas over water
To find grams of a substance
1) Take into consideration of pressure from water vapor
2) Find partial pressure of said element
3) Use PV=nRT to find moles
4) Use molar mass to get to grams
Collecting Gas over Water
To find grams of an element consumed in a reaction
1) Write balanced equation
2) Partial pressure with given information
3) Find moles with Pv=nRT
4) Convert from moles of one element to another and then convert to grams
Collecting Gas over Water
To find mass percent with gases
1) Write balanced equation
2) Find partial pressure of select element
3) Use PV=nRT to find moles of element
4) Convert to grams of desired element
5) Divide grams of element by total grams then multiply by 100%
Charles's law, Boyle's law, and Avagadro's law were (based on observation).
The pressure due to any individual component in a gas mixture is the [blank] (Pn) of that component.
We can calculate the [blank] from the ideal gas law by assuming that each gas component acts independently.
A model which provides connections between the observed macroscopic properties of gases, the gas law equation, and the behavior of gas molecules on a microscopic scale.
1)The size of a particle is negligibly small.
2) The average kinetic energy of a particle is proportional to the temperature in kelvins.
3)The collision of one particle with another (or with the walls of the container) is completely elastic.
The average kinetic energy
of particles in a gas is proportional to the absolute temperature of the gas and does not depend on the identity of the gas.
Energy is related to mass: KE = (1/2)mv2.
Increasing temperature increases the kinetic energy and thus the (average) speed of the molecules.
The average kinetic energy is mentioned.
Speeds are not all the same.
Speeds are best described using a distribution function.
A Maxwell-Boltzmann distribution
of speeds can be used to predict the fraction of molecules at any given speed.
Higher temperature = greater speed (for the same gas)
Most probable speed is dependent on temperature and molar mass. If all of the gases are at the same temperature, then the molecule with greater molar mass has a slower speed.
Standard temperature and pressure (STP):
A specific set of temperature and pressure conditions. For a gas:
0 ̊C (273.15 K)
Standard molar volume:
-The volume occupied by one mole of ANY gas at STP. Can be used to determine the density of a gas.-Because the density and molar mass of a gas are directly proportional, we can derive a formula to relate density and molar mass.18
Mean free path
Molecules travel only short distances before colliding with other molecules and changing direction.
•The average distance a molecule travels between collisions.
•Increases with decreasing pressure
-The flow of energy between two objects as the result of differences in temperature.
-Heat is a process.
-Denoted as q.
-Another way that energy can be transferred.
The part of the universe under consideration
Whatever separates the system from the surroundings
Everything in the universe except the system.
root mean square speed
The (sometimes called root mean square velocity) can be calculated:
R = 8.3145 J/mol*K
T is in Kelvin
M is molar mass in kg/mol
Units are in m/s
System + surroundings
The SI unit of energy is the
1 joule (J) =
1 Lxatm =
1 calorie =
1 Calorie (nutritional calorie) =
Internal energy (E)
The sum of the kinetic and potential energies of all the particles that compose the system
First law of thermodynamics:
The total energy of the universe is constant.
Energy can be neither created nor destroyed, but it CAN be transformed.
Heat transfer stops when the samples of matter reach the same temperature, a condition called [blank]
Heat always flows from matter at higher temperature to matter at lower temperature.
pressure-volume work (PV work).
Work encompasses more than mechanical motion of macroscopic objects.
The most common type of work we will encounter in chemical processes is
Heat flow is measured in a .
Consists of a chemical equation together with a statement of the corresponding enthalpy change or heat of reaction (also called enthalpy of reaction)
mols of reactants and products.
The stoichiometric coefficients in thermochemical equations must be interpreted as
Heat of formation (∆Hfo)
Enthalpy change for the reaction in which ONE MOLE of the substance (in a specified state) is formed from its elements in their most stable forms or standard states.
Because the definition specifies only ONE mole of product, fractional coefficients are allowed for reactants.
The [blank]of the atom explains the strange behavior of electrons. It describes electrons as they exist within atoms and how the electrons determine the physical and chemical properties of elements.
A more accurate term for what we usually refer to simply as light.
There are many other light energies that are outside the bounds that our eyes can detect.
-The process by which gas molecules spread out in response to a concentration gradient
-Influenced by root mean square speed
-Heavier molecules diffuse more slowly than lighter ones
-The process by which a gas escapes from a container into a vacuum through a small hole
-Related to diffusion and root mean square speed, so heavier molecules effuse more slowly
-The study of the energetic consequences of chemistry
The full range of light energies including:
Gamma rays, X-rays, ultraviolet radiation, visible light, infrared radiation, microwaves, and radiowaves
Size or height of a wave
The number of complete cycles of the wave passing a given point per second
Designated as v and measured in units of 1/s = 1 hertz (Hz)
Observations of photoelectric effect
Electrons are NOT ejected below a frequency of light (aka, the threshold frequency, νo) that is specific to the metal.
-Above the threshold frequency, the number of electrons ejected is independent of frequency but dependent on intensity of light.
- Above the threshold frequency, the kinetic energy of ejected electrons is linearly dependent on frequency and independent of intensity.
-The particular pattern of wavelengths absorbed and emitted by any element.
=All elements display discrete spectra.
few states with very specific energies.
What conclusions can be drawn from the fact that atoms of all elements emit only certain wavelengths of light as opposed to continuous spectra?
The fact that only a few wavelengths of light are emitted from a particular atom is direct evidence that the atom can exist in only a
Quantum mechanical model
Replaced the Bohr model of the atom.
The energy of the electron still is quantized.
Depicts electrons as waves spread through a region of space (delocalized) called an orbital.
Similar to how light (traditionally treated as a wave) can be treated as a particle per the photoelectric effect.
A function that describes the periodic behavior of a wave (a mathematical description of a wave).
- For a two-dimensional wave, we simply could use sin x and cos x.
-In three dimensions, wavefunctions are more complicated, but the idea is the same.
The solutions to this deceptively simple equation provide the wavefunctions that describe orbitals
H is an operator—it designates a complicated series of mathematical operations to be carried out.
E is energy
Psi is the wave function for an electron
In lab, you will examine the [blank] the series of transitions that result in visible light.
The 2s and 3s orbitals are larger than a [blank] and also contain nodes, places where there is a 0 probability of finding an electron.
places where there is a 0 probability of finding an electron.
nodes in orbitals are three-dimensional analogs of the nodes on a vibrating string.
Internal energy is a [blank]:
-Its value depends only on the state of the system, not on how the system arrived at that state.
have two lobes on either side of a node.
heat capacity, C
•The heat absorbed by a system (q) and its corresponding temperature change (deltaT) are directly proportional.-The constant of proportionality between q and deltaT is
Specific heat capacity
A physical property of a material that measures how much heat is required to raise the temperature of one gram of that material by 1 degrees C
-Often called "specific heat " or "heat capacity"
-Abbreviated Cs in your textbook. May also be referred to as cs some texts.
-Typical units = J/gx degrees C
Molar heat capacity
-Describes the amount of energy needed to raise the temperature of one mole of a substance by 1 degrees C.
-May be referred to as cm or cp,m-Typical units = J/molxK
have two nodes
-The set of collective techniques used to measure flow of energy.
•Constant volume calorimetry •Constant pressure calorimetry
Thermal energy exchanged between the reaction (defined as the system) and the surrounding is measured by observing the change in temperature of the surroundings
have more lobes and nodes than s, p,and d orbitals.
-Are used to measure ∆E for combustion reactions.
-This takes place at constant volume.
All calorimeters must be [blank] before the actual measurement can be made.
-In calibration, the calorimeter constant, Ccalorimeter, is determined by dividing the known amount of heat released in the calorimeter by the temperature change of the calorimeter.
-A new thermodynamic function
-Heat flow at constant pressure
-A system that releases heat
-Change in enthalpy is less than zero (ΔH < 0)
-A system that absorbs heat
-Change in enthalpy is greater than zero (ΔH > 0)
Just like two-dimensional waves, orbitals have .
for an atom shows the particular orbitals that electrons occupy for that atom.
The enthalpy change for any process is independent of the particular way the process is carried out.
This works because enthlapy is a state function
Ground state electron configurations
represent the lowest energy state for electrons.
When two electrons occupy the same orbital, we say that they are .
(the masking of nuclear charge by by other electrons)
Noble gas configuration or inner electron configuration
Short hand based on the idea that any atom contains all of the same electrons possessed by the noble gas immediately preceding it.
The distance between corresponding points on adjacent waves
Designated as λ and measured most frequently in m or nm-
The inner electrons, which lie closer to the nucleus
Represented by the noble gas with the same electronic configuration
Bonding atomic radius
is one-half the distance between two of the atoms bonded together (nonmetals) or between two of the atoms next to each other in a crystal of a metal
When waves pass through a slit comparable to their size, they bend around it ([blank]).
•Particles simply pass through.
Phenomenon in which light striking the surface of metal causes electrons to be ejected.
The experiments were carried out by Heinrich Hertz.
The explanations were provided by Albert Einstein.
-When light strikes the surface of a metal, electrons are ejected if the light has greater than a specific frequency.
-Ejected electrons are collected by the anode.
-Movement of electrons = current, so intensity can be measured.
are always smaller than their parent atoms.
Loss of electron reduces electron/electron repulsion, leaving the remaining electrons more tightly bound to the nucleus.
Einstein proposed that light could be described as a collection of packets of energy called
-Bright light has many photons, and dim light has few photons.
-The energy of a single photon of light has been shown to be proportional to its frequency and inversely proportional to its wavelength.
are always larger than their parent atoms.
Gain of electron increases electron/electron repulsion, increasing electron/electron repulsion among valence electrons.
are held more tightly from left to right across a period, resulting in smaller ions.
The first four energy levels
for the hydrogen atom.
The lowest energy state is labeled E1, and the higher states are labeled E2, E3, and E4, respectively.
-Transitions from specific energy levels to others result in emission of photons with different wavelengths that can be calculated using ΔE = hc/λ.
-Note that while the absolute energy of each level gets larger, the change in energy between levels gets smaller (correlating to longer λ) as energy increases.
Bohr model of the atom
-The energy of an electron is an atom is quantized
-The first model of the atom that explained why negatively charged electrons didn't simply collapse into the positively charged nuclues
It turns out that the Bohr model of the atom works only for hydrogen.
The Bohr model fails for multi-electron atoms.
We now know that the model is incorrect, but it was a significant contribution.
-Electrons exhibit wave-like behavior!
-The diffraction pattern is NOT the result of two electrons interfering with each other, but rather of individual electrons interfering with themselves.
Carrying out the double-slit experiment with electrons yields very interesting results:
All particles demonstrate wave behavior, and [blank] relation allows us to calculate their wavelengths:
Heisenberg's uncertainty principle
-The product of ∆x and m∆v must be greater than or equal to a finite number (h/4pi).-The more accurately we know the position of an electron, the less accurately we can know its velocity.
decrease in size from left to right across the periodic table.
are three-dimensional waves.
-They represent probability densities—in other words, the shape of the orbital represents the area in which an electron is most likely to be located.
-The maximum number of electrons that can be found in any orbital is two.
decrease in size from left to right across the periodic table.
All Group 5A elements invert first ionization energy order with the [blank] element in their row.
-The names attached to atomic orbitals that come from the functions that solve the wave equations. In other words, the solutions to Schrodinger's equation.
Principal quantum number, n
Secondary quantum number (angular momentum quantum number),l
Magnetic quantum number, ml
Principal quantum number (n)
Describes the main energy level (shell) that the electron occupies.
Larger values of n: -Higher energy orbitals
-Orbitals are farther away from the nucleus
May have any positive integer value (1, 2, 3, ...)
Angular momentum quantum number
(l): Designates the sublevel (also known as a subshell) that the electron occupies.
Tells us the shape of the orbital by indicating the number of nodes or nodal planes:
-Planes in which there is zero probability of finding an electron.
-The region on either side of the node is called a lobe.
Magnetic quantum number (ml)
Designates the specific orbital within a subshell that the electron occupies.
Tells us the orientation in space of the orbital (along which axis is it oriented).
May have any integer value from -l to +l.
When l = 0, ml = 0
When l = 1, ml = -1, 0, or +
When l = 2, ml = -2, -1, 0, +1, or +2
spin quantum number (ms)
The specifies the orientation of the spin of the electron.
-Electrons don't really "spin." It is more correct to say that they have inherent angular momentum.
-Electron spin is a fundamental property of an electron (like its negative charge).
-All electrons have the same amount of spin, and it is quantized: ms = +1/2 or -1/2
Excitation and Radiation
When an atom absorbs energy, an electron can be excited from an orbital in a lower energy level to an orbital in a higher energy level.
-The electron in this "excited state" is unstable and relaxes to a lower energy level, releasing energy in the form of electromagnetic radiation.
All Group 5A elements invert electron affinity order with the [blank] element in their row.
three-dimensional probability densities—the probability of finding an electron in a given location.
-Orbitals also can be represented as geometric shapes that encompass the volume where the electron is likely to be found most frequently.
-s, p, d, and f orbitals have characteristic shapes
form because they lower the potential energy between the charged particles that compose atom.
The two electrons in a helium atom are referred to as a
Accounts for contributions of all attractions and repulsions
Trends in lattice energies
Small highly charged ions form ionic compounds with large lattice energies.
Large ions with small charges for ionic compounds with small lattice energies.
Results when two bonding pairs are shared.
Results when three bonding pairs are shared.
Strength of the covalent bond increases as the number of bonding pairs increases.
The attraction of an atom for the shared electrons in a covalent bond.
Values are established by considering factors such as atomic size, electron configuration, electron affinity, and ionization energy.
Often, the Pauling scale is used.
The two points of positive and negative charge constitute a
polar covalent bonds
Molecules with dipoles are called [blank] and they orient themselves in an electric field.
Pauli exclusion principle
-No two electrons in an atom may have the same set of four quantum numbers.
-Two electrons in the same orbital have the same values for n, l, and ml, so they must have different ms values.
-There are only two possible values of ms (+1/2 and -1/2)
- This means no orbital can hold more than two electrons.
The average of all the valid Lewis structures for a compound that represents the real structure of the compound.
The double headed arrow does not mean that the structure alternates between the two possibilities
-Coulomb's law•Like charges repel ( positive potential energy) and opposite charges attract (negative potential energy)
For multi-electron atoms, we need to consider three key concepts associated with the energy of an electron in the vicinity of a nucleus:
is the energy required to break one mole of the bond in the gas phase.
Stronger bonds are more stable
effective nuclear charge
, the charge "felt" by electrons in comparison to the actual nuclear charge
shielding and effective nuclear charge
The relative energies or orbitals are based on [blank] (the masking of nuclear charge by by other electrons) and [blank], the charge "felt" by electrons in comparison to the actual nuclear charge.
Auf bau principle
The "building up" principle.
Orbitals should be "filled" starting with the lowest energy orbitals and proceeding to higher energy orbitals.
Each degenerate orbital will have one electron in it before any electrons are paired.
In addition, the spins in singly occupied orbitals remain parallel until the second electron enters the orbital.
Bond energies are reported as positive values because bond breaking is
"Outer" electrons Valence electrons are shown explicitly when a noble gas shorthand is used to write electronic configurations.
Valence electrons determine reactivity
screening and effective nuclear charge
Periodic trends are based largely on
Ne has more valence electrons (8) than O (6) or Li (1).
Ne's core electrons cannot effectively shield all 8 valence electrons.
The attraction between the nucleus and valence electrons is very strong.
The sea of electrons model explains four key properties of metals:
Form when an atom loses one or more electrons.
The most common cations of s- and p-block elements have noble gas electron configuration.
highest energy electrons
•All atoms lose their [blank] first to form cations.
-For s- and p-block elements, these are the s- and p-electrons with the highest value of n.
-For transition metals, this means that the ns electrons are lost before the (n-1)d electrons.
Example: The 4s and 3d orbitals are extremely close in energy, and the 3d orbitals are higher in energy initially. However, they are so close in energy that adding and removing electrons causes their energy levels to invert. For the purposes of cation formation, the 4s electrons are higher in energy and must be removed first.
Formed when an atom gains one or more electrons.
The most common anions of s- and p-block elements have noble gas electron configuration after ionization.
Heteronuclear diatomic molecules are
Valence bond theory
Describes how covalent bonding takes place.
This involves the overlap of atomic orbitals to form mixed or hybridized orbitals.
Forces of attraction between molecules in a sample
Originate from the interactions among charges, partial charges, and temporary charges on molecules
Polar molecules have a
instantaneous fleeting dipole moment
Non-polar molecules may have an
First ionization energy
-The minimum energy needed to remove an electron from a neutral atom in the gas phase.
-The energy released when an electron is added to a gas-phase atom.
A graph showing the variation in the temperature of a sample as it is heated at constant rate and constant pressure.
Simple cubic unit cell:
(8 corner atoms)((1/8 atoms per cell) = 1 atom per cell
Body-centered cubic unit cell:
(8 corner atoms)((1/8 atom per cell) + 1 body centered atom = 2 atoms per cell
-Help keep track of valence electrons, especially for main group elements, allowing prediction of bonding in molecules.
To draw a Lewis dot symbol, the valence electrons are represented by dots and are placed around the element symbol.
Face centered cubic unit cell
(8 corner atoms)((1/8 atoms per cell) + (6 faces)(1/2 atoms per cell)) = 4 atoms per cell
-An atom will form covalent bonds to achieve a complement of eight valence electrons.
-As we will see later, this is more of a suggestion than a rule.
The valence shell electronic configuration is ns²np⁶ for a total of eight electrons
For the n = 1 shell, so hydrogen and helium violate the octet rule and share only two electrons
Band theory (quantitative)
Based on molecular orbital theory
Sea of electrons model (qualitative)
•Valence electrons of metals are delocalized (not tied to any specific atom) and move freely throughout the solid
Bonding pair (of electrons)
-A pair of electrons shared by two atoms
Lone pair (of electrons)
-Paired electrons that are associated with a single atom.
-Also referred to as nonbonding electrons.
Sea of electrons model(qualitative) and Band theory (quantitative)
•We will use two models to explain the qualitative and quantitative properties of metals:
bonding orbital .
When the wave functions combine constructively, the resulting molecular orbital has less energy than the original atomic orbitals; it is called a
When the wave functions combine destructively, the resulting molecular orbital has more energy than the original atomic orbitals; it is called an
N molecular orbitals
N atomic orbitals combine to form
is based the idea that a crystalline solid is very like a large molecule, and its valence electrons occupy the MOs formed from atomic orbitals of each atom in the crystal.
The band of molecular orbitals that is occupied by electrons
The band of molecular orbitals that is unoccupied.
Electrons carry a current by moving through a material, and this motion can be thought of as electrons moving from one orbital to another.
Electrons in a filled band cannot move readily to conduct electricity
To conduct electricity, electrons must move from the valence to the conduction band.
In a conductor, there is no energy gap between the valence band and the conduction band.
In semiconductors there is a small energy gap.
In insulators there is a large energy gap.
A pure substance is doped with an element that has fewer valence electrons than it does
• The number of electrons in the valence band is decreased so that conductivity can be achieved without needing to promote electrons across the band gap.
• An acceptor level is created.
Which kind of material (n- or p-type) would result if pure germanium was doped with:
The ratio of the change in concentration to the elapsed time.
Units = M/s
The change in molar concentration of a reactant divided by the time interval.
Valence shell electron pair repulsion (VSEPR) theory
Molecules assume a shape that allows them to minimize the repulsions between electron pairs in the valence shell of the central atom.
-A simplistic view of electrons in molecules (both as bonding and lone pairs) that leads to surprisingly accurate predictions of molecular shapes.
Order of repulsion: bonding pair to bonding pair < bonding pair to lone pair < lone pair to lone pair
(also called molecular shape or shape) takes into account the number of bonding pairs versus the number of lone pair electrons.
- Each electron geometry has subcategories for molecular geometry
-If there are no lone pairs of electrons on the central atoms, then the electron geometry and molecular geometry are the same.
•Refers to the rate at a single moment, and it is given by the slope of a line tangent to the curve defined by the change in concentration versus time.
•We almost always use initial rates in kinetics
Orders of reaction
•Represented by m and n in the previous rate law example.
• Typically, they are integers or half-integers, and they must be experimentally determined.
-Created by a linear combination of atomic orbitals, producing an equal number of hybrid orbitals (remember that orbitals are mathematical in nature).
-Two atomic orbitals combine, two hybrid orbitals are generated.
-The new hybrid orbitals have different shapes than the atomic orbitals from which they came
Integrated rate law
Equation that quantifies the relationship between concentration and time.
•Integrated rates laws are derived by taking the integral of the rate law.
• Since zero, first, and second order rate laws are all a little different, we end up with different versions of the integrated rate law for each order of reaction.
•The time it takes for its concentration to fall to one-half its original value.
•In other words, the half-life is the time required for [A] to become equal to ½[A]0.
Activated complex or transition state
At the moment of an effective collision, both bond breaking and bond formation are occurring simultaneously.
•For an instant, as bond rearrangement is occurring, this high-energy, unstable species exists in the reaction mixture.
The greater the energy of the activated complex, the higher the activation energy and the slower the reaction (at a given temperature).
elementary steps or fundamental steps.
•Reaction mechanisms are composed of a series of single events or steps.
•A physical or chemical process in which the rate of the forward and reverse reactions occur at the same rate.
raised to a power equal to that factor.
When the coefficients in a balanced chemical equation are all multiplied by a constant factor:•the corresponding equilibrium constant is
Chemical equilibria with reactants and products that are all in the same phase
•Equilibria systems having more than one phase
•Heterogeneous equilibrium expressions do not contain terms for solids and liquids
•The concentration of a solid or liquid does not change because these substances are pure.
Reaction quotient (Q)
The ratio of the product and reactant concentrations when a reaction is not at equilibrium
•All equilibrium constants are reaction quotients, but not all reaction quotients are equilibrium constants
less moles of gas
Compression = Reducing volume = increasing pressure
•Causes a reaction to shift in the direction that has [blank]
Molecular orbital theory
has proved to be the most successful theory of the chemical bond.-Electrons occupy molecular orbitalsthat spread throughout the entire molecule.
•The addition of wavefunctions is called "forming a linear combination" and the molecular orbital described is called a linear combination of atomic orbitals.
•Orbitals are no longer localized on a single atom.
more moles of gas
Expansion = Increasing volume = decreasing pressure•Causes a reaction to shift in the direction that has more [blank]
The first theory of acids and bases was developed by
•A substance that contains hydrogen and produces H+ (a proton) in aqueous solution.
Molecular orbital theory summary:
-Molecular orbitals (MOs) can be approximated by linear combination of atomic orbitals (AOs). Natomic orbitals = N molecular orbitals.
-When two AOs combine to form two MOs, one MO is lower in energy (bonding MO)and the other is higher in energy (antibonding MO).
-When assigning the electrons of a molecule to MOs, we fill the lowest energy MOs first with a maximum of two spin-paired electrons per orbital.
•A substance that contains the hydroxyl group (OH) and produces hydroxide ions (OH-) in aqueous solution.
A proton donor
A proton acceptor
The Bronsted-Lowry theory is has less restrictive definitions for acids and bases.
two conjugate acid-base pairs
The reaction of a Bronsted-Lowry acid with a Bronsted-Lowry base contains
The band gap for a [blank] is intermediate between that of conductor and an insulator.
•Some electrons have enough thermal energy to reach the conduction band even at room temperature, so they have at least minimal conductivity.
• If the temperature is increased, more electrons should occupy the conduction band.
Thus, the conductivity of semi-conductors increases at higher temperatures.
There are relatively few elements that can be classified as semiconductors, but silicon is the most common.
•No element really is suitable for widespread use in its pure form.
-Instead, the semiconductor relies on doping.
A pure substance is doped with an element that has more valence electrons than it does•
The "extra" electron cannot go into the full valence band and creates a new donor level that is close in energy to the conduction band.
• Conductivity increases slightly.
Protons cannot exist on their own in aqueous solution.
They instead attach to a water molecule to form H₃O⁺
Deprotonate completely in solution.
•The reaction is strongly product favored.
•Deprotonate incompletely in solution •The reaction is reactant favored
Substances that can act as acids or bases, depending on circumstances.
-A mathematical equation that summarizes the dependence of reaction rate on concentration.-There are two useful forms of the rate law:
•Differential rate law (usually referred to simply as the rate law)
•Integrated rate law
Rate constant, k
-A proportionality constant that relates rate and concentration of reaction species
• The value of k is characteristic of the reaction and the temperature at which it takes place.
Reactions in which one molecule transfers a proton to another molecule of the same kind is called
We can express hydroxide ion concentrations using the
•Protonate completely in solution.
•The reaction is strongly product favored.
Rates of chemical reactions depend on [blank].
-Qualitatively, we can say that most reactions go faster as temperature increases.
-This is why we cook food—heating accelerates reactions that lead to breakdown of cell walls and the decomposition of proteins.
-It's also why we refrigerate food—to slow down natural chemical reactions that lead to decomposition
-The equation that quantifies the temperature dependence of the rate constant, k.
Chemical reactions involve breaking and making bonds.
•Forming new bonds releases energy, but at least some energy input is needed to start breaking the bonds in the reactants.
•Regardless of whether a reaction is ultimately endothermic or exothermic, the reaction must first overcome an energy threshold called the activation energy.
•Regardless of whether a reaction is ultimately endothermic or exothermic, the reaction must first overcome an energy threshold called the [blank]
•Protonate incompletely in solution
•The reaction is reactant favored
Frequency factor or pre-exponential factor (A)
-This is a proportionality constant, and it arises from considerations at the molecular level.
-It represents the number of times the reactants approach the activation barrier per unit time.
-Most approaches to the activation barrier do not have enough total energy to make it over.
-Dependent on both activation energy and temperature
-A number between 0 and 1 that represents the fraction of molecules that have enough energy to make it over the activation barrier on a given approach
-A useful model for describing the motion of molecules, and it will help us investigate the role of temperature in kinetics.
•For two molecules to react, they must first collide. But not all collisions are effective.
They must occur with:
•sufficient kinetic energy
-A collection of one or more molecular steps that account for the way reactants become products.
-It is the sequence of single steps by which we believe reactants are converted to products.
•Reaction mechanisms are composed of a series of single events or steps.
-These are called elementary steps or fundamental steps.
melting, evaporation, and sublimation
Entropy increases during
-Form in one step of a mechanism and are consumed in another
-NOT found in the balanced equation for the overall reaction but plays a key role in the mechanism
Each step in the reaction mechanism has its own rate, and the overall reaction cannot go any faster than the [blank] of the steps in the mechanism.
-The slow step in the mechanism is called the rate-limiting step or rate-determining step.
rate-limiting step or rate-determining step.
Each step in the reaction mechanism has its own rate, and the overall reaction cannot go any faster than the slowest of the steps in the mechanism.
-The slow step in the mechanism is called
the number of steps in a mechanism, the rate-limiting step, and whether an overall reaction is endothermic or exothermic.
Energy diagrams can be used to determine
work by providing an alternative pathway (a different reaction mechanism) between reactants and products).
•The new pathway has lower activation energy than the original.
freezing, condensation, and deposition
Entropy decreases during
-The stage in a chemical reaction when there is no further tendency for the composition of the reaction mixture to change.
-All chemical equilibria are dynamic equilibria.
-The forward reaction rate is exactly equal to the reverse reaction rate. Rateforward = Ratereverse
•Neither reaction stops at equilibrium. They both continue with no net change in concentration of products or reactants.
-The ratio at equilibrium of the concentrations of the products raised to their stoichiometric coefficients divided by the concentrations of the reactants raised to their stoichiometric coefficients
Factors affecting entropy:
•Number of particles
A chemical reaction that generates more moles of gas than initially were present will increase the entropy of a sample.
•Increasing temperature increases kinetic energy.
•The number of ways to distribute the kinetic energy of the sample increases, resulting in an increase in disorder.
solute and solvent phases.
Solutions can be composed of a variety of
•It depends on the temperature.•Enthalpy, entropy, and temperature all effect spontaneity.
Does ice melt spontaneously?
When a system at equilibrium is disturbed, the system shifts in a direction that minimizes the disturbance.
LeChatelier's principle is based on the comparison of Q to K.
•Typically, when a stress is applied, we change the value of K.• The system shifts to re-achieve a state in which Q = K.
Addition of either reactant or product shifts the reaction away from the addition.
•Removal of reactant or product shifts the reaction toward the substance removed.
Effect of a change in pressure.
-The composition will tend to change in a way that minimizes the resulting increase in pressure.
-An increase in number of particles increases pressure while a decrease in particles decreases pressure.
• ΔH > 0(disfavors spontaneity)
• ΔS > 0 (favors spontaneity)
What are the signs for ∆H and ∆S for the melting of ice?
Gibbs Free Energy
The spontaneity of some processes and reactions is temperature dependent.
•Since the reaction or process is spontaneous at some temperatures and not at others, then the sign for
ΔG goes from negative to positive at a specific temperature.•ΔG is 0 at exactly that temperature.
position of equilibrium
The increase in pressure by the addition of an inert gas does NOT affect the [blank]
-Reacting gases continue to occupy the same volume, so their individual molar concentrations and partial pressures remain unchanged.
-Provided that gases can be regarded as ideal, the equilibrium composition is unaffected despite the change in total pressure.
Processes that occur spontaneously only at lower temperatures are said to be [blank] since the negative DH favors spontaneity.
Processes that occur spontaneously only at higher temperatures are said to be [blank] since the positive ∆S favors spontaneity.
proton transfer reaction
A substance can act as an acid only if a base is present to accept its acidic proton.
-This is called a
-Loss of a proton; results in formation of the conjugate base
-Gain of a proton; results in formation of the conjugate acid
Acidity constant or acid ionization constant, Ka
-The equilibrium constant for the deprotonation reaction of a weak acid.
Water can act as an acid or a base per the Bronsted-Lowry theory, depending on the identity of the other reactant. It is [blank]:
-Substances that can act as acids or bases, depending on circumstances.
Autoprotolysis constant for water (Kw)
-Also called the ion product constant for water
-Relates Ka and Kb for a conjugate acid-base pair
-Reactions in which one molecule transfers a proton to another molecule of the same kind is called autoprotolysis
The [blank] is a convenient means of conveying the hydronium ion concentration.
Most solutions have pH ranging from 0 to 14, but values outside this range are possible.
• A note on significant figures:
•The number of digits following the decimal point in a pH value is equal to the number of significant figures in the corresponding molar concentration.21
Basicity constant or base ionization constant, Kb
-The equilibrium constant for the protonation reaction of a weak base.
-A process that occurs without continuous intervention.
-Spontaneity does not relate to time.
•Some processes are spontaneous in the thermodynamic sense, but they occur so slowly that they cannot be observed even over several lifetimes.
Recall that thermodynamics deals with relative chemical potentials of reactants and products while kinetics deals with reaction speed.
Both endo- and exothermic processes can be spontaneous, but if you list many spontaneous processes you'll see that a majority are exothermic.
•Clearly, a relationship is implied between enthalpy and spontaneity, but it is not an exclusive relationship. •A new thermodynamic property is needed.
Which do you think is more likely to happen spontaneously: an endothermic reaction or an exothermic reaction? Why?
-A thermodynamic state function increases with the number of energetically equivalent ways to arrange the components of a system to achieve a particular state.
Systems tend to move to a state of of maximum randomness because a random arrangement of particles is more [blank] than an ordered arrangement.
-Random arrangements can be achieved in more ways.
-The probability of all the coins in the box coming up heads (perfectly ordered) is much less probable than a random mixture of heads and tails.
The probabilities of the ordered and random arrangements are proportional to the number of ways that the arrangements can be achieved.
-Two coins: Each can come up in one of two ways, so there are 2² = 4 possible arrangements.
-Twenty coins: Each can come up in one of two ways, so there are 2²⁰ = 1,048, 576 possible arrangements.
-Think about how many molecules are in a mole of substance.
S = k ln W
The entropy of a particular state is related to the number of was that the state can be achieved:
S = entropy
•k = Boltzmann's constant = R/NA = 1.38 x 10⁻²³ J/K
•W = number of was that the state can be achieved
also explains why a gas expands into a vacuum.
-If the two bulbs have equal volume, the each molecule has a 1 in 2 chance of being in bulb A and a 1 in 2 chance of being in bulb B when the stopcock is open.
-It's extremely unlikely that all molecules will be in one bulb because there is only one way to achieve that state.
-The state in which one mole of molecules can be randomly distributed between the two bulbs is 2⁶.⁰²²³, giving a higher entropy.
There are a number of different ways that the sample can achieve the same total energy with different molecules moving at different speeds.
relates to entropy.
Solution formation also
-The two ideal gases below mix spontaneously with one other because the mixture of gases has greater entropy than the separated components.'
-Solution formation results from a increase in entropy upon mixing.
-Solution formation is prohibited when there is a decrease in entropy upon mixing.
When two ideal gases mix, it is in the absence of intermolecular forces.
•When the solvent and/or solute are solids or liquids, [blank] play a big role in determining whether one substance is soluble (or miscible) in another
Second law of thermodynamics
•The natural progression of a system and its surroundings (which together make up the universe) is:
- from order to disorder
- organized to random
- from low entropy to high entropy
The second law of thermodynamics does NOT mean that all processes result in increases in entropy
-As long as the increase in entropy of he surroundings is larger than the decrease in the entropy of the system, the overall change in entropy for the entire process can still be positive.
Gibbs free energy, G
-A new state function that predicts the spontaneity of a given reaction or process.
•Changes in this function can predict whether or not a process is spontaneous under conditions of constant pressure and temperature.
•The signs of ΔG are opposite the sign of ΔSuniverse
• ΔG > 0 = nonspontaneous
• ΔG < 0 = spontaneous
The decrease in Gibbs free energy represents the maximum amount of energy available for a system to do work.
Third law of thermodynamics-
The entropy of a perfect crystal of any pure substance approaches zero as the temperature approaches absolute zero.
Standard molar entropy, S°
-The entropy of one mole of any given chemical substance under standard conditions.
-It is measured by determining the change in entropy from near 0 K to 298 K at 1 atm.
Factor impacting standard molar entropy:
The standard (Gibbs) free energy of formation (∆G°f)
is the standard Gibbs free energy of reaction per mole for the formation of a compound from its elements in their most stable form.
Oxidation-reduction reactions (or redoxreactions)
are those in which electrons are transferred from one reactant to another.
loss of electrons
Gain of electrons
reactions involving an oxidation and a reduction are called redox reactions
Assigned by imagining that each atoms in a formula unit is present in its ionic form (which it might not be) and determining the "charge" on this ion
The oxidation OR reduction part of a reaction considered alone.
-The oxidation half-reaction shows the removal of electrons from a species that is being oxidized.
Cu(S) --> Cu²⁺ (aq) + 2e⁻
-The reduction half-reaction shows the gain of electrons by the species that is being reduced
Ag⁺ (aq) + e⁻ --> Ag(s)
Neither half-reaction in a pair can occur on its own because oxidation and reduction must take place in concert with one another.
-Half-reactions can be added to give an overall reaction, but the number of electrons gained must equal the number lost.
•The species that causes oxidation
•It is the species that is reduced during the reaction.
•The species that causes reduction
•It is the species that is oxidized during the reaction.
-Device in which an electric current (flow of electrons through a circuit) is produced by a spontaneous chemical reaction or used to bring about a nonspontaneous reaction.
•An electrochemical cell in which a spontaneous chemical reaction is used to generate an electric current.
•A battery is a collection of galvanic cells joined in series.
Galvanic cell summary
:-Oxidation occurs at the anode (one half-cell)
-Reduction occurs at the cathode (one half-cell)
-Electrons flow from the anode to the cathode
-The overall redox reaction is spontaneous, and chemical energy is converted to electrical energy (lighting the light bulb)
As electrons flow from the anode to the cathode, there is a build-up of negative charge in the cathode half-cell and build-up of positive charge in the anode half-cell.
-This prevents further flow of electrons.
-The salt bridge allows cations to flow into the cathode and anions to flow into the anode to offset the build-up of charge resulting from electron flow.
-Oxidation occurs at the anode (one half-cell)
-Reduction occurs at the cathode (one half-cell)
-Electrons flow from the anode to the cathode
-The overall redox reaction is spontaneous, and chemical energy is converted to electrical energy (lighting the light bulb)
-Anions flow from the salt bridge into the anode.
-Cations flow from the salt bridge into the cathode.
electromotive force (EMF) or cell potential (Ecell).
Though we offset the build-up of charge at the anode and cathode with the salt bridge, the build-up is important because it represents the potential for work.
-We call this potential the
-Cell potential is a measure of the potential energy difference (per unit charge) between the two electrodes (anode and cathode)
Standard cell potential (Eo)
•The characteristic contribution to the cell potential every electrode makes under standard conditions
•Because they are always written for reduction half-reactions, they are also called standard reduction potentials.
•The standard hydrogen electrode (SHE) is used to assign a cell potential to individual half reactions.
The SHE arbitrarily is assigned a standard reduction potential of 0 V.
• When paired with another half-cell, whatever reading appears on the voltmeter is assigned to that half-cell as its standard reduction potential.
Standard Electrode Potentials
In a galvanic cell, the half-reaction with the largest positive potential will run as written
•This will be a reduction reaction.
•It occurs at the cathode.
•The other half reaction will be forced to run in reverse.
•This will be an oxidation reaction.
•It occurs at the anode.